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SECONDARY GENERAL SCIENCE

CHEMISTRY

NEW SYLLABUS

FORM 3 AND 4

2008

Abraham Kobia Kathali

Abraham Kobia is a veteran chemistry teacher with a long experience in teaching chemistry at

both local and international level. He holds a bachelor of education degree in science education.

He is credited with the development of the Exam Vista assessment computer software for

teachers of science and mathematics in Kenyan Secondary Schools.

KINETIC THEORY AND GAS LAWS

Properties of matter

Matter is made up of particles that are in constant motion. The motion depends on the

temperature of the substance. The higher the temperature, the faster the particles move (more

energy). Increase in temperature increase weakens inter-particle forces, causing particles to

spread apart and increase in volume/size (i.e. Expansion). Gases have greatest average energy

while solids have smallest average energy.

There are three physical states of matter.

Solid

Liquid

Gas

Properties of matter

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Property Solid Liquid Gas

Can they be

compressed

(squashed)?

Solids cannot be

compressed

Liquids cannot be

compressed

Gases can be

compressed easily

Do they flow?

Solids do not flow

Liquids do flow

Gases do flow

Do they keep their

shape?

The shape of a solid

stays the same

A liquid takes the

shape of the bottom of

the container

A gas takes the shape

of the entire container

Do they keep their

volume?

The volume of a solid

stays the same

The volume of a liquid

stays the same

A gas has no definite

volume

How are the particles

arranged?

Very close together

Regular arrangement

Close together

Random arrangement

Far apart

Random arrangement

Kinetic Particle Theory

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According to the kinetic theory, matter is made up of particles (atoms, molecules or ions) which

are in constant motion because they have energy at all temperatures above zero Kelvin (absolute

zero, 0K) or -2730C. This energy is in the form of kinetic energy.

Assumptions of the kinetic theory

Matter is made up of particles that are in constant motion

The higher the temperature, the faster the particles move (more energy)

Increase in temperature increase weakens interparticle forces, causing particles to spread

apart and increase in volume/size (i.e. Expansion)

Gases have greatest average energy while solids have smallest average energy

Consider heating a solid;

When ice is heated, the particles vibrate more strongly as they gain kinetic energy and the

particle attractive forces are weakened. Eventually, at the melting point, the attractive forces are

too weak to hold the particles in the structure together in an ordered way and so ice melts. The

particles become free to move around and lose their ordered arrangement. Energy is needed to

overcome the attractive forces and give the particles increased kinetic energy of vibration. On

heating further, the particles gain more kinetic energy and move faster. In evaporation and boiling

the particles with the highest kinetic energy can ‘escape’ from the attractive forces of the other

liquid particles. The particles lose any order and become completely free to form a gas or vapour.

Boiling is rapid evaporation anywhere in the bulk liquid and at a fixed temperature called the

boiling point and requires continuous addition of heat.

Cooling and Heating Curves

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Table 2.1: Analysis of the stages

A to B Ice warms up. Temperature rises

from -10 to 00C. No change of state

B to C Temperature stays constant.

Change of state occurs. Ice

changes to liquid water.

C to D Water warms up. Temperature

changes from 00C to 1000C

D to E Temperature remains constant.

Change of state occurs. Boils to

steam at 1000C

Due to the kinetic theory we begin to understand why gases exert pressure. The molecules of a

gas are far apart and in continuous random motion, colliding with each other and with the walls of

the vessel in which the gas is held. The molecules have mass, so they have energy hence they

exert force on each collision and hence pressure.

If the temperature of the gas is increased at constant volume, the molecules gain more energy

and move faster, hitting the walls with more force and exerting greater pressure.

If the volume of the gas is increased at constant temperature, the molecules have more space in

which to move. The frequency of collisions decreases reducing the pressure.

Gas laws

Gases and the Kinetic Theory

(a) Causes of gas pressure

All the molecules in a gas are in rapid motion, with a wide range of speeds, and repeatedly hit

the walls of the container in huge numbers per second. The average force and hence the

pressure they exert on the walls is constant since pressure is force on unit area.

(b) Boyle’s law

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If the volume of a fixed mass of gas is halved by halving the volume of the container, the number

of molecules per cm3 will be doubled. There will be twice as many collisions per second with the

walls, i.e. the pressure is doubled. This is Boyle’s law.

(c) Temperature

When a gas is heated and its temperature rises, the average speed of its molecules increases.

If the volume of the gas is to remain constant, its pressure increases due to more frequent and

more violent collisions of the molecules with the walls. If the pressure of the gas is to remain

constant, the volume must increase so that the frequency of collisions does not.

Boyle’s law

Boyle's law is one of several gas laws and a special case of the ideal gas law. Boyle's law

describes the inversely proportional relationship between the absolute pressure and volume of a

gas, if the temperature is kept constant within a closed system. The law was named after chemist

and physicist Robert Boyle, who published the original law in 1662.

Robert Boyle

The law states,

The pressure of a fixed mass of gas is inversely proportional to its volume if its temperature is

kept constant.

For a fixed amount of an ideal gas kept at a fixed temperature, P [pressure] and V [volume] are

inversely proportional (while one increases, the other decreases).

Consider a gas trapped in a container as shown. The mass, hence number of moles are constant

and do not change during the course of the investigation. The piston is frictionless and moves

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smoothly without allowing the gas to escape. When the pressure changes, the volume of the gas

changes as shown.

P1 = 1 atm P2 = 2 atm P3 = 4 atm

V1 = 12dm3 V2 = 6 dm3 V3 = 3dm3

12 x 1= 12 6 x 2 = 12 3 x 4 = 12

P1V1 = P2V2 = P3V3

∴ PV = constant at constant temperature and mass.

Graphical representation of Boyle’s law

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Graph between P & V at constant temperature is a smooth curve known as "parabola"

Graph between 1/P & V at constant temperature is a straight line.

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If pressure, p is doubled, the volume is halved. That is, p is inversely proportional to V. In

symbols

P ∝ 1/V or p = constant x 1/V

PV = constant

p1V1 = p2V2 = constant. This is Boyle’s law

EXAMPLE 1

(a) State Boyle’s law

(b) Sketch a graph of the relationship between 1/P & V at constant temperature.

(c) Explain why a gas in a container exerts pressure.

SOLUTION

(a) The pressure of a fixed mass of gas is inversely proportional to its volume if its

temperature is kept constant.

(b)

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(c) All the molecules in a gas are in rapid motion with a wide range of speeds, and repeatedly hit

the walls of the container in huge numbers per second. The average force and hence the

pressure they exert on the walls is constant since pressure is force on unit area.

EXAMPLE 2

A certain mass of ammonia occupies 600 ml at a certain pressure. When the pressure is changed

to 4 atmospheres it occupies a volume of 2.4 litres, temperature remaining constant. What was

the initial pressure?

SOLUTION

P1V1= P2 V2

Substituting the values

P1x 600 = 4 x 2400

Initial pressure =12160 mm of Hg

Charles’ Law

Charles's law (also known as the law of volumes) is an experimental gas law which describes

how a fixed mass of a gas tends to expand when heated at constant pressure.

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The volume of a fixed mass of gas is directly proportional to its absolute temperature if the

pressure is kept constant.

P = 1 atm P = 1 atm P = 1 atm

We can then say that the volume V is directly proportional to the absolute temperature T, i.e.

doubling T doubles V, etc. Therefore

V ∝ T or V = constant x T

Or V/T = Constant

Volume V1 = 1dm3 V2 = 2dm3 V3= 3dm3

Temperature (OC) 0 ºC 273 ºC 546 ºC

Temp (K) 273 K 546K 819K

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Volume/Temp 1/273 2/546 3/819

Equation ∴ V1/T1 = V2/T2 = V3/T3

∴ V/T = constant

Graphical representation

Graph between Volume and absolute temperature of a gas at constant pressure is a "straight

line"

Absolute scale of temperature or absolute zero

If the graph between V and T is extrapolated, it intersects T-axis at -273.16 0C .At -273.16 0C

volume of any gas theoretically becomes zero as indicated by the graph.

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The absolute temperature

Practically volume of a gas can never become zero. Actually no gas can achieve the lowest

possible temperature and before -273.16 0C all gases are condensed to liquid. This temperature

is referred to as absolute scale or absolute zero. At -273.16 0C all molecular motions are ceased.

This temperature is called Absolute Zero.

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Degrees on this scale are called Kelvin’s and are denoted by K while θ stands for a Celsius scale

temperature. They are exactly the same size as Celsius degrees. Since –273 0C = 0K,

conversions from 0C to K are made by adding 273. T = 273 + θ

0 0C = 273K

15 0C = 273 + 15 = 288K

The letter T represents Kelvin or absolute temperatures and θ stands for a Celsius scale

temperature.

EXAMPLE 1

A given mass of a gas is at a temperature of 3oC. When the gas is heated to 95oC at a constant

pressure, at occupies a volume of 460 ml. What is the initial volume of the gas?

SOLUTION

T1 = 3oC + 273 = 276 K T2 = 95oC + 273 = 368 K

Substituting the values

Initial volume = 345 ml.

Pressure law

The pressure of a fixed mass of gas is directly proportional to its absolute temperature if the

volume is kept constant.

P ∝ T or P = constant X T

or P/T = Constant

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The three equations can be combined giving

PV = constant

T

For cases in which p, V and T all change from say p1, V1 and T1 to p2, V2 and T2, then

P1V1 = P2V2

T1 T2

EXAMPLE 1

A certain volume of a gas at 0oC is at a pressure of 70 cm of mercury. If the gas is compressed

into a cylinder of 24 litres capacity under a pressure of 1750 mm and a temperature of 136.5oC,

what is the original volume of the gas?

SOLUTION

T1 = 0 + 273 = 273 K T2 = 136.5oC + 273 = 409.5 K

Original volume = 40 lt.

DIFFUSION AND KINETIC ENERGY

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If a bottle of perfume is opened in one corner of a room, it spreads in the whole room by diffusion.

The natural rapid and random movement of the particles means that gases readily ‘spread’ or

diffuse.

Diffusion is the movement of gas or solid particles from a region of high concentration to a region

of low concentration.

Diffusion is fastest in gases where there is more space for them to move. The rate of diffusion

increases with increase in temperature as the particles gain kinetic energy and move faster.

DIFFUSION OF BROMINE GAS

A coloured gas, heavier than air (greater density), is put into the bottom gas jar and a second gas

jar of lower density colourless air is placed over it separated with a glass cover.

If the glass cover is removed then (i) the colourless air gases diffuses down into the coloured

brown gas and (ii) bromine diffuses up into the air. The particle movement leading to mixing

cannot be due to convection because the more dense gas starts at the bottom!

No 'shaking' or other means of mixing is required. The random movement of both lots of particles

is enough to ensure that both gases eventually become completely mixed by diffusion.

This is clear evidence for diffusion due to the random continuous movement of all the gas

particles and, initially, the net movement of one type of particle from a higher to a lower

concentration ('down a diffusion gradient'). When fully mixed, no further colour change distribution

is observed BUT the random particle movement continues!

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DIFFUSION OF AMMONIA AND HYDROGEN CHLORIDE

The following experiment is set up. One filter soaked in a solution of ammonia solution and the

other soaked in a solution of concentrated hydrochloric acid are placed on the end of along glass

tubing as shown.

When colourless NH3 and HCl fumes meet, dense white smoke (fumes) of ammonium

chloride are observed.

NH3 (ag) + HCl (ag) NH4 Cl(s)

Ammonia is diffused more rapidly than the hydrogen chloride because the gas traveled a longer

distance in the same amount of time. Gases with greater R.M.M have higher densities than gases

which have small molecules e.g. hydrogen chloride are heavier than ammonia molecules. If the

concentration of hydrochloric acid and that of ammonia were increased in a separate experiment,

the rate of diffusion would be faster.

Gas

Mr.

NH3 17

HCl 35.5

Large heavy molecules move more slowly than small, light molecules. Therefore, dense gases

diffuse more slowly than gases of low density. The rate of diffusion depends on the molecular

mass/density of gas. Rate of diffusion is inversely proportional to mass of a gas.

RATE OF DIFFUSION;

Increases with temperature

Decreases with increasing R.M.M or R.A.M

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Increases with concentration

Graham’s Law of Diffusion

Thomas Graham

Graham’s law of diffusion relates the rate of diffusion of a gas to its density.

It states that the rate of diffusion of a gas at constant temperature and pressure is inversely

proportional to the square root of its density.

The density of a gas is proportional to the molecular mass MM

The law can be summarized as

And

These two can be combined to give;

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Diffusion in Liquids

If you drop a little ink in a beaker of water it will spread by itself in the beaker of water and the

color spreads uniformly. Let us study diffusion of liquids with the help of the following

experiments:

EXPERIMENT 1

Place a layer of deep blue CuSO4 solution in a tall beaker and allow it to settle. CuSO4 is a salt

which ionizes completely when placed in water into Cu2+ and SO42- ions. Then carefully transfer a

layer of water on to the CuSO4 solution layer using a pipette. Cover the beaker with a lid and

leave it undisturbed for several hours as shown in figure.

Initially water floats on top of the CuSO4 solution as it is less dense. The boundary line very soon

disappears and after a few hours the solution becomes a homogeneous pale blue color as shown

in figure. The intermingling can be explained by describing the motion of molecules of water,

copper and sulphate ions moving randomly in all directions.

EXPERIMENT 2

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Diffusion of a Potassium Permanganate in Hot and Cold Water

Purpose:

The purpose of this experiment is to demonstrate the principle of diffusion. Diffusion is the net

movement of particles (atoms, ions, or molecules) from an area of greater concentration to areas

of lesser concentration.

Materials and Methods:

For this experiment, the student will require two petri dishes, two 150 ml beakers, hot plate,

potassium permanganate (purple dye), and a spatula. Obtain about 50 ml of cold water in one

beaker and about 50 ml or boiling water in the other.

Procedure

Place the petri dishes on the table and add cold water to one dish and boiling water to the other.

Allow the water in the dishes to become still. Using the spatula, add a crystal of potassium

permanganate to the center of each dish.

Observe the experiment for several minutes and note what happens to the purple dye in each

Petri dish.

Conclusion: Do you observe diffusion in each?

How does the rate of diffusion differ?

Why is there a greater rate of diffusion in one?

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It will be observed that the rate of diffusion increases with increasing temperature.

Evaporation, boiling and kinetic theory

On heating particles gain kinetic energy and move faster. In evaporation and boiling the highest

kinetic energy molecules can ‘escape’ from the attractive forces of the other liquid particles. The

particles lose any order and become completely free to form a gas or vapour. Energy is needed

to overcome the attractive forces in the liquid and is taken in from the surroundings. This means

heat is taken in, so evaporation or boiling are endothermic (require energy input) processes. If the

temperature is high enough boiling takes place. Boiling is rapid evaporation anywhere in the bulk

liquid and at a fixed temperature called the boiling point and requires continuous addition of heat.

The rate of boiling is limited by the rate of heat transfer into the liquid. Evaporation takes place

more slowly at any temperature between the melting point and boiling point, and only from the

surface, and results in the liquid becoming cooler due to loss of higher kinetic energy particles.

Differences between evaporation and boiling

Evaporation Boiling

Takes place on the surface of the liquid Takes place throughout the liquid

Causes cooling Does not cause cooling

Occurs at any temperature although the rate

increases with increasing temperature

Occurs only at the boiling point and no

temperature change can occur during boiling

Only a few molecules have enough energy to

escape

All molecules have the same energy and can

escape

Factors that affect the rate of evaporation

Evaporation occurs at all temperatures at the surface of the liquid. It happens more rapidly

when:

i) The temperature is higher, since then more molecules in the liquid are moving fast

enough to escape from the surface,

ii) The surface area of the liquid is large so giving more molecules a changes to escape

because more are near the surface, and

iii) Wind or draught is blowing over the surface carrying vapour molecules away from

the surface thus stopping them from returning to the liquid and making it easier for

more liquid molecules to break free.

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EXAMPLE

(a) Use the kinetic particle theory of matter to explain why energy is needed to melt a solid, at its

melting point, to form a liquid.

(b) A student puts a drop of coloured ink into water. The ink slowly spreads throughout the water.

Use the kinetic particle theory of matter to explain this observation.

SOLUTION

(a) Increase the potential energy of the molecules

OR do work in separating the molecules

against intermolecular forces/bonds

(b) Molecules are moving around randomly

spread in all directions)

Exercise

1) (a) State Charles’Law

(b) Sketch a graph of temperature (OC) on the x-axis, against volume of an ideal gas. Clearly

show the x- intercept

(c) What name is given to this intercept and what does it stand for?

(d) Explain why the pressure of a gas increases when temperature is increased

2) (a) State Boyle’s law

(b) Sketch a graph of the relationship between 1/P & V at constant temperature.

(c) Explain why a gas in a container exerts pressure

3) (a) Define the term diffusion

(b) State grahams law of diffusion

4) The volume of a given mass of gas, at 150oC is 400 ml. At what temperature, will it occupy a

volume of 600 ml at the same pressure?

5) A certain mass of carbon dioxide occupies a volume of 480 litres at 1 atmosphere pressure.

What pressure must be applied to confine it to a cylinder of 12 litre capacity, temperature

remaining constant?

6) (a) Use the kinetic particle theory of matter to explain why energy is needed to melt a solid, at

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its melting point, to form a liquid.

(b) A student puts a drop of coloured ink into water. The ink slowly spreads throughout the water.

Use the kinetic particle theory of matter to explain this observation.

THE MOLE:

What is a Mole?

A mole is a word which represents a number. Like the word "dozen" represents the number 12,

so "mole" represents the number 6 x 1023. This number is also called Avogadro's number. It is a

very very big number (6 followed by 23 zeros). In the same way that you can have a dozen

atoms, or cars, or apples, so you can also have a mole of atoms, or cars, or apples.

1 Dozen = 12

1 Mole = 6 x 1023

Needless to say, chemists are concerned with atoms, ions, molecules and compounds. A mole is

defined as the number of atoms in exactly 12 grams of 12C (carbon twelve). The number is called

Avogadro's number. In the above definition, 12 is the mass number of carbon. So, one mole of

carbon atoms has a mass of 12 grams. The relative atomic mass, which can be written as Ar or

RAM, is the number just above the element in the periodic table

Amedeo Avogadro (1776-1856)

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Relative atomic mass (r.a.m) and moles

Every atom has its own unique atomic mass based on a standard comparison or relative scale

e.g. it has been based on hydrogen H = 1, oxygen O = 16, Carbon is 12.

The relative atomic mass was earlier defined as the number of times an atom of an element was

heavier than one atom of hydrogen.

However there are complications due to isotopes and so very accurate atomic masses are not

whole numbers. Isotopes are atoms of the same element with different masses due to different

numbers of neutrons. The very accurate atomic mass scale is based on a specific isotope of

carbon, carbon-12, 12C = 12.0000 units exactly, for most purposes C = 12 is used for simplicity.

Avogadro later compared the values of R.A.M to the mole.

He concluded that the value of R.AM of an element expressed in grams contains 1 mole of

atoms.

Just as one mole of carbon atoms has a mass of 12 g,

So the mass of one mole of the atoms of any element

is its "relative atomic mass" expressed in grams.

For example, look up the relative atomic mass of sodium (Na), (the larger number above it in the

periodic table).

One mole of sodium has a mass of 23 g.

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One mole of helium has a mass of 4 g,

One mole of neon has a mass of 20 g,

One mole of magnesium has a mass of 24 g,

One mole of calcium has a mass of 40 g

MASS, MOLES AND R.A.M

The mass of an element can be calculated if the number of moles is known. This can be done

using the equation and the memory triangle shown.

Mass (g) = moles x R.A.M

Example 1

How many moles are in 1.15 g of sodium?

Solution

We first need to find the R.A.M of sodium. This is 23 from the periodic table.

Moles = Mass (g)/ R.A.M

= 1.15g/23

=0.05 mol

Example 2

Calculate the mass of 2.4 moles of magnesium.

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Solution

R.F.M of magnesium = 24

Mass (g) = moles x R.F.M

= 2.4 x 24

= 57.6g

This is easy for elements which exist as atoms.

As discussed earlier, the number of protons added to the number of neutrons is known as the

relative atomic mass. This is the mass of 1 mol of an atom relative to the mass of 1 mol of C

atoms that have 6 protons and 6 neutrons, which is taken to be 12.00 g. However there are

complications due to isotopes and so very accurate atomic masses are not whole numbers.

Isotopes are atoms of the same element with different masses due to different numbers of

neutrons. The very accurate atomic mass scale is based on a specific isotope of carbon, carbon-

12, 12C = 12.0000 units exactly, for most purposes C = 12 is used for simplicity.

Atoms of elements having the same atomic number with different mass numbers are called

isotopes

The strict definition of relative atomic mass (Ar) is that it equals average mass of all the isotopic

atoms present in the element compared to 1/12th the mass of a carbon-12 atom.

Examples of relative atomic mass calculations

Example1 : chlorine consists of 75% chlorine-35 and 25% chlorine-37.

Solution

Think of the data based on 100 atoms, so 75 have a mass of 35 and 25 atoms have

a mass of 37.

The average mass = [ (75 x 35) + (25 x 37) ] / 100 = 35.5

So the relative atomic mass of chlorine is 35.5 or Ar(Cl) = 35.5

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Example 2: Bromine consists of 50% 79Br and 50% 81Br, calculate the Ar of bromine.

Solution

Ar = [ (50 x 79) + (50 x 81) ] /100 = 80

So the relative atomic mass of bromine is 80 or Ar(Br) = 80

What about elements which exist as molecules or compounds?

Relative formula mass of a compound (R.F.M):

To calculate the mass of one mole of a compound, the number of each type of atom in the

compound is multiplied by that atoms relative atomic mass and all those numbers added

together.

This value is called the relative formula mass (or relative molecular mass or molar mass) of a

compound.

Notice that r.f.m, r.m.m, or Mr have no units because they are ratios. The molar mass is obtained

from r.f.m, r.m.m, or Mr by simply adding g (grams)

If all the individual atomic masses of all the atoms in a formula are added together you have

calculated the relative formula mass (for ionic compounds) or molecular mass (for covalent

elements or compounds), Mr. can be used for any element or compound. Whereas relative atomic

mass (above) only applies to a single atom, anything with at least two atoms requires the term

relative formula/molecular mass.

Calculating R.F.M

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(a) For example, the mass of one mole of carbon dioxide (CO2) is

(1 x RAM of carbon) + (2 x RAM of oxygen)

= (1 x 12) + (2 x 16) = 44 g.

So, one mole of carbon dioxide has a mass of 44 g.

The Relative Formula Mass of carbon dioxide is 44.

This may also be called the Relative Molecular Mass (RMM),

since carbon dioxide is a molecule.

(b) The mass of one mole of calcium carbonate (CaCO3) is

(1 x RAM of calcium) + (1 x RAM of carbon) + (3 x RAM of oxygen)

= (1 x 40) + (1 x 12) + (3 x 16) = 100 g.

So, one mole of calcium carbonate has a mass of 100 g.

The Relative Formula Mass of calcium carbonate is 100.

(c) The mass of one mole of magnesium oxide (MgO) is

(1 x RAM of magnesium) + (1 x RAM of oxygen)

= (1 x 24) + (1 x 16) = 40 g.

So, one mole of magnesium oxide has a mass of 40 g.

The Relative Formula Mass of magnesium oxide is 40.

Mass from amount:

The key mathematical equation needed here is;

Mass (g) = relative formula mass (g mol-1) x amount (mol)

Using the triangular relationship from above if the mass section is covered over then the amount

multiplied by the relative formula mass gives the mass.

Example

(i) What is the mass of 0.25 mol of NaCl?

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0.25 mol of NaCl = 58.5 g mol-1 × 0.25 mol

= 14.63 g

(ii) What is the mass of 3 mol of Al2(SO4)3?

3 mol of Al2(SO4)3 = 342 g mol-1 × 3 mol

= 1026 g or 1.026 kg

Moles from mass:

The key mathematical equation needed here is ;

Amount (mol) = mass (g) / relative formula mass (g mol-1)

Using the triangular relationship from above if the amount section is covered over then the mass

divided by the relative formula mass gives the amount -

Example

(i) What amount is 117 g of NaCl?

117 g of NaCl = 117 g / 58.5 g mol-1

= 2 mol

(ii) What amount is 68.4 g of Al2(SO4)3?

68.4 g of Al2(SO4)3 = 68.4 g / 342 g mol-1

= 0.2 mol

Molar mass from mass and amount:

The key mathematical equation needed here is;

Relative formula mass (g mol-1) = mass (g) / amount (mol)

Using the triangular relationship from above if the molar mass section is covered over then the

mass divided by the amount gives the relative formula mass.

Example

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What is the molar mass of a compound for which 0.2 mol of it has a mass of 42 g?

molar mass of compound = 42 g / 0.2 mol

= 210 g mol-1

R.F.M = 210

Every mole of any substance contains the same number of the defined species.

The actual particle number is known and is called the Avogadro Constant and is equal to 6.023 x

1023 'defined species' per mole. This means there are that many atoms in 12g of carbon (C = 12)

or that many molecules of water in 18g (H2O = 1+1+16 = 18, H = 1; O = 16)

Note. Relative is just a number based on the carbon-12 relative atomic mass scale.

Molar mass is a term used to describe the mass of one mole i.e. the relative

atomic/formula/molecular mass in grams (g).

Example 1:

1 mole of ammonia, NH3, Consists of 1 mole of nitrogen atoms combined with 3 moles of

hydrogen atoms. Or you could say 2 moles of ammonia is formed from 1 mole of nitrogen

molecules (N2) and 3 moles of hydrogen molecules (H2).

Example 2: 1 mole of aluminium oxide,

Al2O3, consists of 2 moles of aluminium atoms combined with 3 moles of oxygen

atoms

(Or 1.5 moles of O2 molecules).

For calculation purposes learn the following formula for 'Z' and use a triangle if necessary.

(1) mole of Z = g of Z / atomic or formula mass of Z,

(2) or g of Z = mole of Z x atomic or formula mass of Z

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(3) or atomic or formula mass of Z = g of Z / mole of Z

where Z represents atoms, molecules or formula of the particular element or compound defined

in the question.

Example 1:

How many moles of potassium ions and bromide ions in 0.25 moles of potassium bromide?

1 mole of KBr contains 1 mole of potassium ions (K+) and 1 mole of bromide ions (Br-

).

So there will be 0.25 moles of each ion.

Example 2:

How many moles of calcium ions and chloride ions in 2.5 moles of calcium chloride?

1 mole of CaCl2 consists of 1 mole of calcium ions (Ca2+) and 2 moles of chloride ion

(Cl-).

So there will be 2.5 x 1 = 2.5 moles of calcium ions and 2.5 x 2 = 5 moles chloride

ions.

mass of NaCl formed = moles x formula mass = 0.

4 x 58.5 = 23.4g NaCl

Using the Avogadro Constant, you can actually calculate the number of particles in known

quantity of material.

Example 3:

How many water molecules are there in 1g of water, H2O?

Formula mass of water = (2 x 1) + 16 = 18

Every mole of a substance contains 6 x 1023 particles of 'it' (the Avogadro Constant).

Moles water = 1 / 18 = 0.0556

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Molecules of water = 0.0556 x 6 x 1023 = 3.34 x 1022

PERCENTAGE COMPOSITION IN A COMPOUND:

Calculating % Composition (from masses of each element)

Divide the mass of each element by the total mass of the compound and multiply by 100

Calculating % Composition (from formula)

1. Calculate formula mass

2. Divide the total atomic mass of each element by the formula mass and

multiply by 100

Example 1.

What is the percentage of oxygen in carbon dioxide gas?

mass of oxygen in one mole of carbon dioxide gas = 2×16 g

= 32 g

mass of one mole of carbon dioxide = 12 g + (2×16 g)

= 44 g

percentage of oxygen = (32 g/44 g) × 100

= 72.7 %

Example 2.

What is the percentage of water of crystallization in hydrated copper (II) sulphate, CuSO4·5H2O?

mass of water present in one mole of

hydrated copper(II) sulphate =(5×18 g)

= 90 g

mass of one mole of hydrated copper(II)

sulphate = 64 g + 32 g + (4×16 g) + (5×18 g)

= 250 g

percentage of water of crystallization present = (90 g/250 g) × 100

= 36 %

32

REACTING MASS CALCULATIONS

You can use the ideas of relative atomic, molecular or formula mass and the law of conservation

of mass to do quantitative calculations in chemistry. Underneath an equation you can add the

appropriate atomic or formula masses. This enables you to see what mass of what, reacts with

what mass of other reactants. It also allows you to predict what mass of products are formed (or

to predict what is needed to make so much of a particular product). You must take into account

the balancing numbers in the equation (e.g. 2Mg), as well of course, the numbers in the formula

(e.g. O2).

NOTE: The symbol equation must be correctly balanced to get the right answer!

Example 1:

In a copper smelter, how many tonne of carbon (charcoal, coke) is needed to make 16 tonne of

copper? (b) How many tones of copper can be made from 640 tones of copper oxide ore?

(a) 2CuO(s) + C(s) 2Cu(s) + CO2(g)

(atomic masses Cu=64, O=16, C=12)

Formula Mass ratio is 2 x (64+16) + (12) ==> 2 x (64) + (12 + 2x16)

= Reacting mass ratio 160 + 12 ==> 128 + 44

(In the calculation, impurities are ignored)

12 of C makes 128 of Cu

Scaling down numerically: mass of carbon needed

= 12 x 16 / 128 = 1.5 tonne of C

(b) 160 of CuO make 128 of Cu (or direct from formula 80 CuO 64 Cu)

Scaling up numerically: mass copper formed

33

= 128 x 640 / 160 = 512 tones Cu

Example 2:

What mass of carbon is required to reduce 20 tonne of iron(II) oxide ore if carbon

monoxide is formed in the process as well as iron?

(Atomic masses: Fe = 56, O = 16)

Reaction equation: Fe2O3 + 3C 2Fe + 3CO

Formula mass Fe2O3 = (2x56) + (3x16) = 160

160 mass units of iron oxide reacts with 3 x 12 = 36 mass units of carbon

So the reacting mass ratio is 160: 36

So the ratio to solve is 20: x, scaling down,

x = 36 x 20/160 = 4.5 tones carbon needed.

Note: Fe2O3 + 3CO 2Fe + 3CO2 is the other most likely reaction that reduces

the iron ore to iron.

MOLARITY

It is very useful to be known exactly how much of a dissolved substance is present in a solution of

particular concentration or volume of a solution. So we need a standard way of comparing the

concentrations of solutions.

The concentration of an aqueous solution is usually expressed in terms of moles of dissolved

substance per cubic decimetre, mol dm-3; this is called molarity, sometimes denoted in shorthand

as M.

Note: 1dm3 = 1 litre = 1000ml = 1000 cm3, so dividing cm3/1000 gives dm3, which is handy to

know since most volumetric laboratory apparatus is calibrated in cm3 (or ml), but solution

concentrations are usually quoted in molarity, that is mol/dm3 (mol/litre).

Equal volumes of solution of the same molar concentration contain the same number of moles of

34

solute i.e. the same number of particles as given by the chemical formula.

You need to be able to calculate:

The number of moles or mass of substance in an aqueous solution of given

volume and concentration

The concentration of an aqueous solution given the amount of substance and

volume of water.

You should recall and be able to use each of the following relationships.

(1) molarity of Z = moles of Z / volume in dm3

(2) molarity x formula mass of solute = solute concentration in g/dm3,

dividing this by 1000 gives the concentration in g/cm3

(3) (concentration in g/dm3) / formula mass = molarity in mol/dm3,

(4) moles Z = mass Z / formula mass of Z

(5) 1 mole = formula mass in grams

Example 1

What mass of sodium hydroxide (NaOH) is needed to make up 500 cm3 (0.5 dm3) of a 0.5M

solution? [Ar's: Na = 23, O = 16, H = 1]

1 mole of NaOH = 23 + 16 + 1 = 40g

35

for 1000 cm3 (1 dm3) of 0.5M you would need 0.5 moles NaOH

which is 0.5 x 40 = 20g

however only 500 cm3 of solution is needed compared to 1000 cm3

so scaling down: mass NaOH required = 20 x 500/1000 = 10g

Example 2

How many moles of H2SO4 are there in 250cm3 of a 0.8M sulphuric acid solution? What mass of

acid is in this solution? [Ar's: H = 1, S = 32, O = 16]

Formula mass of sulphuric acid = 2 + 32 + (4x16) = 98, so 1 mole = 98g

if there was 1000 cm3 of the solution, there would be 0.8 moles H2SO4

but there is only 250cm3 of solution, so scaling down ...

moles H2SO4 = 0.8 x (250/1000) = 0.2 mol

mass = moles x formula mass, which is 0.2 x 98 = 19.6g of H2SO4

Dilution

Concentrated solutions can be mixed with solvent to make dilute solutions. This is the kind of

thing people do everyday with consumer products like fruit juice. Some concentrated solutions

are used as "stock" solutions. Dilute solutions are typically used but the concentrated solutions

require less storage space. In recent years accidents have occurred in the health care

professions when dilutions were done incorrectly. Molarity and dilutions

In dilutions, the amount of solvent is increased, but the amount of solute is kept constant. The

result is a decreased concentration, but a greater volume.

The idea is that the volumes may change but the number of moles does not. This means that the

original number of moles and the final number of moles are the same.

Number moles original = number of moles final

Moriginal x Voriginal = Mfinal x Vfinal

36

Where M is molarity, and V is volume,.

This equation can be used to figure out any one of the terms if given the other three. What is also

true is that the difference between the original volume and the final volume tells how much

solvent is needed to make the desired "dilution"

Vadded = Vfinal - Vinitial

TITRATION: ACID AND ALKALI

Titrations can be used to find the concentration of an acid or alkali from the relative volumes used

and the concentration of one of the two reactants.

You should be able to carry out calculations involving neutralisation reactions in aqueous solution

given the balanced equation or from your own practical results.

1. Note again: 1dm3 = 1 litre = 1000ml = 1000 cm3, so dividing cm3/1000 gives dm3.

2. and other useful formulae or relationships are:

moles = molarity (mol/dm3) x volume (dm3=cm3/1000),

molarity (mol/dm3) = mol / volume (dm3=cm3/1000),

1 mole = formula mass in grams.

In most volumetric calculations of this type, you first calculate the known moles of one reactant

from a volume and molarity. Then, from the equation, you relate this to the number of moles of

the other reactant, and then with the volume of the unknown concentration, you work out its

molarity.

Example 1:

25 cm3 of a sodium hydroxide solution was pipetted into a conical flask and titrated with 0.2M

hydrochloric acid. Using a suitable indicator it was found that 15 cm3 of acid was required to

neutralise the alkali. Calculate the molarity of the sodium hydroxide and concentration in g/dm3.

Equation NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)

moles HCl = (15/1000) x 0.2 = 0.003 mol

moles HCl = moles NaOH (1 : 1 in equation)

37

so there is 0.003 mol NaOH in 25 cm3

scaling up to 1000 cm3 (1 dm3), there are ...

0.003 x (1000/25) = 0.12 mol NaOH in 1 dm3

molarity of NaOH is 0.12M or mol dm-3

since mass = moles x formula mass, and Mr(NaOH) = 23 + 16 + 1 = 40

concentration in g/dm3 is 0.12 x 40 = 4.41g/dm3

Example 2:

20 cm3 of a sulphuric acid solution was titrated with 0.05M potassium hydroxide. If the acid

required 36 cm3 of the alkali KOH for neutralisation what was the concentration of the acid?

equation 2KOH(aq) + H2SO4(aq) K2SO4 + 2H2O(l)

mol KOH = 0.05 x (36/1000) = 0.0018 mol

mol H2SO4 = mol KOH / 2 (because of 1 : 2 ratio in equation above)

mol H2SO4 = 0.0018/2 = 0.0009 (in 20 cm3)

scaling up to 1000 cm3 of solution = 0.0009 x (1000/20) = 0.045 mol

mol H2SO4 in 1 dm3 = 0.045, so molarity of H2SO4 = 0.045M or mol dm-3

since mass = moles x formula mass, and Mr(H2SO4) = 2 + 32 + (4x16) = 98

Concentration in g/dm3 is 0.045 x 98 = 4.41g/dm3

Exercise

1) (a) Define the term relative atomic mass

38

(b) How many moles of calcium ions and chloride ions in 2.5 moles of calcium chloride?

2) 3.1g of XCO3 are heated to constant mass, 2.0g of the metal oxide formed. When heated in a

stream of hydrogen, the oxide is reduced to 1.6g of pure metal.

a)Write down the equation for the two reactions.

b)Determine;

i.The relative atomic mass of x

ii.The volume of hydrogen used at s.t.p.

iii.The volume of carbon dioxide produced on decomposition of XCO3 at s.t.p.

3) 100cm3 of natural gas contains 20% propane (C3H8), 60% Methane (CH4), and the rest is

carbon monoxide.

a) Write balanced equation for the combustion of each of the gases in the mixture

b)Calculate the volume of oxygen needed for complete combustion of this mixture.

(All gas volumes measured at the same conditions of temperature and pressure)

4) The concentration of a solution of sodium hydroxide was found by titrating the solution with

0.200 mol dm-3

sulphuric acid. 25.0 cm3 of the sodium hydroxide solution required 31.5 cm

3 of

the sulphuric acid for complete reaction.The equation for the reaction is

2NaOH + H2SO4 Na2SO4 + 2H2O

(a)(i)Explain why Universal indicator is not a suitable indicator for use in titrations.

(ii) Name a suitable indicator for this titration.

(b)(i)Calculate the concentration, in mol dm-3

, of sodium hydroxide in the solution.

(ii)Calculate the concentration, in g dm-3

, of sodium hydroxide in the solution.

(Relative atomic masses: H = 1.0; O = 16; Na = 23)

(c)Sodium hydroxide solution is used to test for copper (II) ions in solution.

(i) Describe what you would see in this test.

ORGANIC CHEMISTRY

Products from Oil

Coal, Oil and Natural Gas Formation - Fossil Fuels.

39

Just as coal has formed by the action of heat and pressure on the remains of trees and plants on

land over millions of years, so oil and natural gas have formed by the action of heat and pressure

on the remains of sea plants and animals over millions of years.

The remains were buried in sediments which excluded the air (kept out oxygen) and stopped

them decaying. More sediment buried the remains deeper and deeper until pressure and heat

eventually turned them into coal, oil and natural gas. They are called fossil fuels because they are

buried underground (from Latin fossilis - dug up).

Fossil fuels are a finite resource and non-renewable.

The oil deposits are formed in porous rock sediments. Porous rock has pores in it. Pores are

small holes (see for example sandstone). The small holes allow the oil and natural gas to pass

through the rock and rise until they are stopped by a layer of non-porous rock. Non-porous rock

(for example shale) has no holes, and acts as a barrier to prevent the oil and natural gas rising.

The oil and natural gas become trapped underground.

The oil is called crude oil (or petroleum, from Latin - rock oil), and has natural gas in it or in a

pocket above it trapped by non-porous rock. Drilling through the rock allows the oil and gas to

escape to the surface. Natural gas is mostly methane (CH4). Crude oil is a mixture of substances

(mostly hydrocarbons).

So what are Hydrocarbons?

40

Crude oil is a mixture of substances which are mostly hydrocarbons.

A hydrocarbon is a compound containing hydrogen and carbon only.

Since crude oil is a mixture of different hydrocarbon compounds, the different hydrocarbons will

have different boiling points. A sample of crude oil will therefore have a range of boiling points,

and the mixture can be separated by fractional distillation.

Fractional Distillation of Crude Oil.

Naming the fractions

The hydrocarbon fractions are mainly alkanes.

Name Number of

Carbon Atoms

Boiling Point

(°C) Uses

Refinery Gas 3 or 4 below 30 Bottled Gas

(propane or butane).

41

Petrol 7 to 9 100 to 150 Fuel for car

engines.

Naphtha 6 to 11 70 to 200 Solvents

and used in petrol.

Kerosene (paraffin) 11 to 18 200 to 300 Fuel for aircraft

and stoves.

Diesel Oil 11 to 18 200 to 300

Fuel for road

vehicles

and trains.

Lubricating Oil 18 to 25 300 to 400 Lubricant for engines

and machines.

Fuel Oil 20 to 27 350 to 450 Fuel for ships

and heating.

Greases and Wax 25 to 30 400 to 500 Lubricants

and candles.

Bitumen above 35 above 500 Road surface

and roofing.

Crude oil is heated until it boils and then the hydrocarbon gases are entered into the bottom of

the fractionating column. As the gases go up the column the temperature decreases.

The hydrocarbon gases condense back into liquids and the fractions are removed from the sides

of the column. The smaller the hydrocarbon molecule, the further it rises up the column before

condensing. The fractionating column operates continuously.

The temperatures shown are approximate. A sample of crude oil may be separated in the

laboratory by fractional distillation. The collection vessel is changed as the temperature rises to

collect the different fractions.

Naming hydrocarbons.

Hydrocarbons are named according to the number of carbon atoms in the molecule.

Number of

Carbon Atoms Name Prefix Alkane Alkene

42

1 meth methane -

2 eth ethane ethene

3 prop propane propene

4 but butane butene

5 pent pentane pentene

6 hex hexane hexene

7 hept heptane heptene

8 oct octane octene

The hydrocarbon fractions are mainly alkanes.

Properties of Different Fractions.

The different hydrocarbon fractions obtained from crude oil condense at different temperatures.

The larger the hydrocarbon molecule (the more carbon atoms it has)

1) The higher the condensing temperature (the higher the boiling point).

2) The more viscous it is (it takes longer to flow - like syrup).

3) The less volatile it is (it evaporates less quickly).

4) The less flammable it is (it does not set fire so easily).

Gases from volatile hydrocarbons are denser than air and pose a fire hazard at ground level. This

is why ignition sources (such as smoking) are not allowed at petrol stations.

Families of organic compounds

Homologous series

Organic compounds belong to different families, though all are based on carbon C, hydrogen H,

and other elements such as oxygen O and nitrogen N etc. The compounds in each family have a

similar chemical structure and a similar chemical formula. Each family of organic compounds

forms what is called a homologous series. Different families arise because carbon atoms readily

join together in chains (catenation) and strongly bond with other atoms such as hydrogen, oxygen

and nitrogen. The result is a huge variety of 'organic compounds'. The name comes from the fact

that most of the original organic compounds studied by chemists came from plants or animals.

43

A homologous series is a family of compounds which have a general formula and have similar

chemical properties because they have the same functional group of atoms (e.g. C=C alkene, C-

OH alcohol or -COOH carboxylic acid).

Members of a homologous series have similar physical properties such as appearance,

melting/boiling points, solubility etc. but show trends in them e.g. steady increase in

melting/boiling point with increase in carbon number or molecular mass.

The molecular formula represents a summary of all the atoms in the molecule.

The structural or displayed formula shows the full structure of the molecule with all the individual

bonds and atoms.

What are alkanes?

These are obtained directly from crude oil by fractional distillation. They are saturated

hydrocarbons and they form an homologous series called alkanes with a general formula CnH2n+2

Saturated hydrocarbons have no C=C double bonds, only carbon-carbon single bonds, and so

has combined with the maximum number of hydrogen atoms. i.e. no more atoms can add to it.

Alkanes are the first homologous series. Examples of alkanes are:

The gases Methane CH4, ethane C2H6, propane C3H8,

butane C4H10

Liquids Pentane C5H12, hexane C6H14 C7H16etc

The Names of all alkanes end in ...ane

Names.

Alkanes are the simplest homologous series of compounds and their names follow this pattern,

CH4 - methane

C2H6 - ethane

C3H8 - propane

C4H10 - butane

44

C5H12 - pentane

I.e. they have a prefix (meth-, eth-, prop-, but-, etc.), which depends on the number of carbon

atoms in the molecule and a common suffix (-ane).

The general chemical form la for an alkane is CnH2n+2

Structural formulae

As well as using a normal type of molecular formula to describe an organic molecule, they can be

represented by drawing out their structure i.e. by showing how the atoms are connected, or

bonded to each other.

In order to do this a few rules have to be followed;

(i) Carbon atoms must be bonded four times

(iii) Hydrogen atoms must bond only once.

Name Molecular formula Structural formula

Methane CH4

Ethane C2H6

Propane C3H8

ISOMERISM

Isomerism occurs when two or more compounds have the same chemical formula but have

different structures, e.g. for the molecular formula C4H10 there are two possibilities - one 'linear'

and one with carbon chain 'branching’. Butane is linear while its branched isomer is methyl

propane.

45

1. Butane

2. 2-Methylbutane

3. 2,2-Dimethylpropane

As the number of carbon atoms increases, the number of possible isomers increases rapidly. All

families or homologous series exhibit isomerism.

Formula Number of structural isomers

C5H12 3

C6H14 5

C7H16 9

CHEMICAL REACTIONS OF ALKANES

Substitutional reactions of alkanes

46

Alkanes are most inert of all homologous series. They are not very reactive unless burned. But

they will react with strong oxidising chemicals like chlorine when heated or subjected to u.v light.

A substitution reaction occurs and a chloro-alkane is formed e.g. a hydrogen atom is swapped for

a chlorine atom and the hydrogen combines with a chlorine atom forming hydrogen chloride. This

process is called halogenation.

The UV light causes the formation of free radical halogen atoms by providing enough energy for

the bond between the two halogen atoms to break.

A halogen atom attacks the alkane, substituting itself for a hydrogen atom. This substitution may

occur many times in an alkane before the reaction is finished.

2. Combustion

Alkanes, along with all other types of hydrocarbon, will burn in an excess of oxygen to give

carbon dioxide and water only as the products,

e.g.CH4 (g) + 2O2(g) CO2(g) + 2H2O(g)

If there is not enough oxygen present then instead of carbon dioxide, carbon monoxide, CO, is

produced. Carbon monoxide is particularly toxic and absorbed into blood, through respiration,

very easily. For domestic heating systems it is particularly important that enough air can get to the

flame to avoid carbon monoxide being generated in the home. Car engines also require a lot of

air and there is a lot of research going on to make the internal combustion engine more efficient,

and so put out less carbon monoxide.

3. Reactivity

Alkanes are saturated hydrocarbons. Molecules of saturated hydrocarbons contain only single

bonds between all carbon atoms in the series. Hence their reactivity with other chemicals is

relatively low.

PHYSICAL PROPERTIES OF ALKANES

47

Alkanes experience inter-molecular van der Waals forces. Stronger inter-molecular van der

Waals forces give rise to greater boiling points of alkanes.

There are two determinants for the strength of the van der Waals forces:

the number of electrons surrounding the molecule, which increases with the alkane's

molecular weight

the surface area of the molecule

Boiling point

Under standard conditions, from CH4 to C4H10 alkanes are gaseous; from C5H12 to C17H36 they

are liquids; and after C18H38 they are solids. As the boiling point of alkanes is primarily determined

by weight, it should not be a surprise that the boiling point has almost a linear relationship with

the size (molecular mass) of the molecule. As a rule of thumb, the boiling point rises 20 - 30 °C

for each carbon added to the chain; this rule applies to other homologous series. A straight-chain

alkane will have a boiling point higher than a branched-chain alkane due to the greater surface

area in contact, thus the greater van der Waals forces, between adjacent molecules.

Melting point

The melting points of the alkanes follow a similar trend to boiling points for the same reason as

outlined above. That is, (all other things being equal) the larger the molecule the higher the

melting point.

Solubility and density

Alkanes, like all other organic chemicals are insoluble in water. They are however soluble in

organic liquids. Alkanes are non-polar and are hence soluble in other non-polar liquids and not in

water, as water is a polar molecule. Alkanes are said to be hydrophobic in that they repel water.

For this reason they do not form hydrogen bonds and are insoluble water.

The density of the alkanes usually increases with increasing number of carbon atoms, but

remains less than that of water. Hence, alkanes form the upper layer in an alkane-water mixture.

SUMMARY

Name Number of carbon

atoms

Melting Point Boiling Point Density (g/cm3)

48

Methane 1 -182 -161 Gas

Ethane 2 -183 -88 Gas

Propane 3 -188 -42 Gas

Butane 4 -138 0 Gas

Pentane 5 -130 36 0.626

Hexane 6 -95 69 0.659

Heptane 7 -90 99 0.684

Octane 8 -57 126 0.703

Nonane 9 -53 151 0.718

Decane 10 -29 151 0.730

What are alkenes?

Hydrocarbons, which contain two hydrogen atoms less than the corresponding alkanes, are

called alkenes. They have one double bond and are unsaturated carbon compounds. Alkenes

cannot be obtained directly from crude oil. They can only be obtained by cracking of alkanes.

Cracking

In industry the fractions obtained from the fractional distillation of crude oil are heated at high

pressure in the presence of a catalyst to produce shorter chain alkanes and alkenes.

49

E.g. C10H22 C5H12 + C5H10

PREPARATION OF AN ALKENE FROM AN ALKANE

Laboratory preparation of ethene gas by cracking

In the lab ethene is prepared by cracking kerosene or candle wax.

Kerosene is poured over sand and this is kept at the bottom of a hard glass test tube. A few

pieces of pumice stone or porcelain is kept a little distance away. The sand is slowly heated. After

a while the porcelain portion of the test tube is heated. This is done alternately. The heated

kerosene first vaporizes and then cracks. When the vapours pass over the hot porcelain, they

crack again into smaller and smaller molecules. The gases are then passed over water. Ethene is

collected by downward displacement of water. It can be understood that this method for collecting

ethene gas does not give pure ethene gas. This is because from cracking, we get many types of

molecules. All those, which are lighter than water and insoluble in water, will be collected.

PROPERTIES OF ALKENES

They are unsaturated hydrocarbons with a general formula CnH2n. Unsaturated means the

molecule has a C=C double bond to which atoms or groups can add after breaking the double

bond.

Alkenes all have a C=C double bond in their structure and their names follow this pattern. Their

names end in ...ene

50

C2H4 - ethene

C3H6 - propene

C4H8 - butene

C5H10 - pentene

The general chemical formula for alkenes is CnH2n

Name Molecular formula Structural formula

Ethene C2H4

Propene C3H6

Butene C4H8

TRENDS IN PHYSICAL PROPERTIES

Some general physical properties of alkenes are:

Physical state

The first lower member like ethene, propene and butene are colorless gases. Alkenes with five to

fifteen carbon atoms are liquids and higher ones are solids at ordinary temperatures. Alkenes

have characteristic smell.

Density

Alkenes are lighter than water.

51

Solubility

Alkenes are insoluble in water and soluble in organic solvents such as benzene, ether etc.

Boiling point

The boiling points of alkenes gradually increase with an increase in the molecular mass (or chain

length). This indicates that inter-molecular attractions become stronger with an increase in the

size of the molecule. Branched chain alkenes have lower boiling points than the corresponding

straight chain isomers.

Melting point

The melting points of alkenes increase with an increase in the molecular mass.

Hydrogenation

Alkenes can be converted back to alkanes by hydrogenation in the presence of a catalyst

This reaction is also called saturation of the double bond. In ethene the carbon atoms are said to

be unsaturated. In ethane the carbon atoms have the maximum number of hydrogen atoms

bonded to them, and are said to be saturated.

TESTING FOR ALKENES

You can tell the difference between an alkane and alkene by adding bromine water to the

substance and giving the tube a shake.

When bromine water is added to an alkane the solution will turn red-brown (the colour of bromine

water) and stay that way on shaking. Alkenes however will turn red-brown, but then the colour

disappears on shaking, this is because the bromine has reacted with the alkene.

52

Ethene by dehydration of alcohols

To obtain pure ethene gas, another method is followed. This is from a chemical reaction with

ethanol and concentrated sulphuric acid.

The temperature of the mixture of ethanol or ethyl alcohol and concentrated sulphuric acid is

increased to 160°C. The acid acts as a dehydrating agent and picks up a water molecule from the

ethanol molecule, leaving the reaction product as ethene gas.

The laboratory equipment to produce ethene gas is shown below. About 20 to 25 ml of ethanol is

taken in a round bottomed flask.

53

Concentrated sulphuric acid is added to it from a thistle funnel slowly. Heat is supplied from a

Bunsen burner and the temperature of the flask is raised to 160°C. Ethene gas starts evolving

and it can be collected over water by downward displacement of water.

Uses of ethene

Ethene is used for manufacturing organic compounds such as ethyl alcohol and ethylene

glycol. Ethylene glycol is used for making artificial fibbers like polyesters.

Ethene is used for manufacture of plastics. These plastics are made from polymerization

of ethene into polythene. Polythenes are used for making bags, electrical insulation, etc.

Ethene is used artificial ripening of fruits such as mangoes, bananas, etc.

Alcohols

The alcohols form a homologous series with the functional group C-OH. It is the presence of this

functional group that gives alcohols their characteristic properties. The simplest homologous

series of alcohols have the general formula CnH2n+1OH. As seen earlier a functional group is an

atom or a group of atoms which defines the function or the mode of activity of a given carbon

compound. The functional group determines all the properties of the compound. Organic

compounds with –OH group is also called an alcoholic group. Methanol, CH3OH is the simplest of

the alcohols. The next higher alcohol is ethanol, C2H5OH.

Alcohols all have an -OH group and their names follow this pattern,

CH3OH - methanol

C2H5OH - ethanol

C3H7OH - propanol

C4H9OH - butanol

C5H11OH - pentanol

The general chemical formula for an alcohol is CnH2n+1OH.

An alcohol is produced by replacing one hydrogen from an alkane and with an hydroxyl group.

The two equations below show how methanol and ethanol are made.

54

STRUCTURES OF ALCOHOLS

Name Molecular formula Structural formula

Methanol CH3OH

Ethanol CH3CH2OH

Propanol CH3CH2CH2OH

Preparation of ethanol

Ethanol is an alcohol with the molecular formula C2H5OH. Ethanol is not a hydrocarbon because

the molecule contains oxygen as well as hydrogen and carbon. Ethanol is a colourless liquid

which boils at 78 °C. It is completely miscible with water. This means that it mixes with water in

any quantity.

Ethanol is a neutral liquid.

Preparation of ethanol

Hydration of alkenes

55

Adding the components of water across the double bond of ethene is called the hydration of

ethene. The reaction is carried out at a high temperature and a high pressure.

ethene + steam ethanol

C2H4(g) + H2O(l) C2H5OH(l)

Alcohol by fermentation.

Ethanol can be made by the fermentation of a sugar solution (batch process) or buy the reaction

of ethene with steam in the presence of phosphoric acid (continuous process). A continuous

process means that reactants are constantly being put in at one end of the reaction vessel and

products are constantly being taken out at the other end of the reaction vessel.

Yeast is a microorganism containing an enzyme which will convert a sugar (glucose) solution into

carbon dioxide and alcohol (ethanol).

Glucose + yeast carbon dioxide + ethanol.

Carbon dioxide gas bubbles out of the solution into the air leaving a mixture of ethanol and water.

Ethanol can be separated from the mixture by fractional distillation. Fermentation must be carried

out in the absence of air to make alcohol. If air is present, ethanoic acid is made instead of

alcohol.

Different alcoholic drinks contain different amounts of alcohol. Some people drink alcohol for

enjoyment, some drink to excess and some people become addicted to alcohol. The harmful

effects (physical and social) of drinking excess alcohol are widespread and reach all parts of

society.

Yeast is also used in the baking of bread. The carbon dioxide produced causes the bread to rise

and fills the bread full of bubbles. The alcohol evaporates during the baking process.

Complete Combustion of Ethanol.

The combustion of alcohols in air occurs very easily. Methanol for example burns with a blue

flame. Alcohols burn in air to give carbon dioxide and water. A large amount of heat is also

generated during this chemical reaction. In fact, spirit lamps in laboratories are examples of

burning of alcohol in air.

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Ethanol can be used as the fuel in cars, called biofuel. Car engines can be specially designed to

burn only ethanol. It is the common fuel in Brazil where plentiful sugar cane provides a renewable

resource for fermentation and lack of oil resources means that imported oil for fuel is expensive.

Up to 30% ethanol can be blended with petrol (called gasohol) and used in ordinary car engines

to reduce the use of fossil fuels. Industrial Methylated Spirits (IMS) is mainly ethanol. Other

chemicals (including methanol) have been added making industrial methylated spirits poisonous

to drink. IMS is used in industry as a solvent

and in domestic burners for heating and cooking.

Physical Properties of alcohols

1. Physical State: Alcohols are colourless liquids at ordinary room temperatures.

2. Odour: Lower members of the alcoholic group have a characteristic fruity smell.

3. Density: Alcohols are lighter than water.

4. Solubility: Alcohols are completely soluble in water. Methanol and ethanol are completely

miscible in water. Higher members of alcohols tend to be less soluble in water.

5. Acidic nature: alcohols are neutral liquids and have no effect on litmus or other acid tests.

6. Conductivity: Alcohols are covalently bonded compounds and are hence non-ionic and non-

conductors of electricity.

Oxidation of ethanol

Oxidation can be defined as the addition of oxygen to a substance. Oxidation can be done by

heating the ethanol with a mixture of sulphuric acid and potassium dichromate (VI) solution. The

mixture turns from orange to green. The ethanol is oxidised to form ethanoic acid which is a

useful organic chemical. BUT it is this oxidation of ethanol that results in alcoholic drinks turning

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sour (e.g. cider, wine) when exposed to air. Ethanoic acid (old name 'acetic acid') is the basis of

vinegar and is also used in making esters.

Ethanol, from a solution made from fermented sugar cane, can be concentrated by fractional

distillation. In some countries it is mixed with petrol to give an alternative motor vehicle fuel.

EXERCISE

1) (i) Explain what you understand by each of the following terms.

(a) hydrocarbon

(b) Homologous series

(ii) The table below shows some of the fractions of fractional distillation of crude oil.

Name Carbon Atoms B.P. (°C) Uses

Refinery Gas 3 or 4 below 30 Bottled Gas (propane or

butane).

Petrol 7 to 9 100 to 150 Fuel for car engines.

Naphtha 6 to 11 70 to 200 Solvents and used in petrol.

Kerosene (paraffin) 11 to 18 200 to 300

11 to 18 200 to 300 Fuel for road vehicles and

trains.

Lubricating Oil 18 to 25 300 to 400 Lubricant for engines

Fuel Oil 20 to 27 350 to 450 Fuel for ships

25 to 30 400 to 500 Lubricants and candles.

Bitumen above 35 above 500

(a) Complete the blank spaces in the table.

+ O2 + H2O

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(b) How does the number of carbon atom affect the boiling point of a fraction?

2) In catalytic cracking, large saturated hydrocarbon molecules are broken down into simpler

ones. Some of these simpler molecules are unsaturated.

Describe the difference between a saturated and an unsaturated hydrocarbon.

3) Methanol is made from carbon monoxide.

CO(g) + 2H2(g) CH3OH(g) the forward reaction is exothermic

(i) Describe how hydrogen is obtained from alkanes.

(ii) Suggest a method of making carbon monoxide from methane.

(iii) Which condition, high or low pressure, would give the maximum yield of methanol?

Give a reason for your choice.

(d) For each of the following predict the name of the organic product.

(i) reaction between methanol and ethanoic acid

(ii) oxidation of propan-1-ol by potassium dichromate(VI)

(iii) removal of H2O from ethanol (dehydration)

Soaps

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Soaps and detergents are used by all of us as cleaning agents in our everyday lives. These

products are sodium salts of organic acids. They are special organic acids that have long chain

hydrocarbons with hydroxyl groups. These types of acids are also called fatty acids. We have

seen earlier that the name fatty acid comes from the fact that these acids, such as palmitic acid

(C15H31COOH), oleic acid (C17H33COOH) and stearic acid (C17H35COOH) are obtained from

hydrolysis of fats.

Saponification

Soapiness in a soap comes from sodium salts of fatty acids like stearic acid, oleic acid and

palmitic acid. These chemicals are sodium stearate C17H35COO- Na+, sodium oleate C17H33COO-

Na+, sodium palmitate C15H31COO- Na+. The naturally occurring fatty acids such as palmitic acid

(C15H31COOH), oleic acid (C17H33COOH) or stearic acid (C17H35COOH) is heated with sodium

hydroxide (NaOH), soapy sodium salt along with an alcohol is formed. This type of reaction we

have seen while studying hydrolysis of ester. The alcohol called glycerol is formed when higher

fatty acids are used in the reaction.

Preparation of soap

The following raw materials are required for the manufacturing of soaps :

Animal fats or vegetable oils

NaOH

NaCl

Heating fat along with NaOH makes soap. Sodium salt of a fatty acid and an alcohol is formed.

This process is called saponification. One such saponification reaction is shown below.

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In saponification, a water molecule is removed from a fatty acid or an ester. When sodium

stearate, sodium oleate or sodium palmitate is formed, it is precipitated from the solution by

addition of NaCl. The addition of sodium chloride makes the sodium salt of fatty acid partially

insoluble in water. This product then separates out of the solution and starts floating on top. The

soap flakes are removed and pressed into moulds after addition of perfumes and colours. The

solution remaining has glycerol and sodium chloride. Glycerol is recovered as it finds important

use in many industries like cosmetics, paints, etc. Soap can be made from many fatty ester or oils

such as vegetable oil, olive oil, coconut oil and other oils.

Experiment

MAKE SOAP FROM FATS

Soap can be made from many oils and fats. The reaction is a double displacement involving a

strong base such as sodium hydroxide and fats.

(a) Obtain animal fat from a butcher. Boil this fat in water and the oil will separate on the surface.

When cold, the fat will solidify and it can be separated from the water. Melt the fat again and

strain through several layers of cloth.

(b) Weigh this fat and then weigh out about one third as much sodium hydroxide pellets. Take

care not to touch either the solid sodium hydroxide or the solution, because it is very caustic.

Heat the fat in an iron saucepan or dish and, when it is molten, slowly add the sodium hydroxide

solution with continuous stirring. Heat with a small flame to avoid boiling over. Allow the fat and

the sodium hydroxide to boil for 30 minutes. Stir the mixture frequently.

(c) The next stage is to weigh out common salt, sodium chloride; about twice the weight of

sodium hydroxide used in (b) is needed. After the 30 minutes boiling, stir this salt well into the

mixture. Then allow to cool. The soap separates as a layer at the top. Separate this soap from the

liquid below, melt and pour into matchboxes where it will solidify again as small bars of soap.

Cleansing action of soap

We all know soap is used to remove dirt and grime from substances. Generally dirt and grime get

stuck because they have an oily component which is difficult to remove by plain brushing or

washing by water. A soap molecule has two parts: the long chain organic part and the functional

group –COO- Na+. A tail and a head. It has to be remembered that this is not an ion, the atoms

are all covalently bonded, the electrical charges show how the charges get polarized in the group.

A soap molecule has a tadpole like structure shown below.

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The organic part is water insoluble but is soluble in organic solvents or in oil or grease. The ionic

part is soluble in water, as water is a polar solvent. When soap is added to water in which dirty

clothes are soaked, the two parts of the soap molecule dissolves in two different mediums. The

organic tail dissolves in the dirt, grime or grease and the ionic head dissolves in water. When the

clothes are rinsed or agitated, the dirt gets pulled out of the clothes in the water by the soap

molecule. In this way the soap does its cleaning work on dirty and grimy clothes or hands.

The soap molecules actually form a closed structure because of mutual repulsion of the positively

charged heads. The structure is called a micelle. The micelle pulls out the dirt and grime more

efficiently.

Synthetic detergents (Soapless Detergents)

Soaps have problems in use for hard water. Hard water has calcium and magnesium, which does

not create lather. Synthetic detergents overcome this problem. Synthetic detergents do not have

sodium salts of fatty acids. In spite of this the detergents have all the properties of soap. Synthetic

detergents have long chain molecules such as sodium n-dodecyl benzene sulphonate and

sodium n- dodecyl sulphate. The chemical structure of these molecules is shown below.

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We can see that the synthetic detergent molecule is similar to the soap molecule. It has a long

hydrocarbon chain (or tail) and a short ionic head. These two parts are water repellent and water

attracting respectively. The cleansing function of a synthetic detergent molecule is similar to that

of the soap.

Synthetic detergents are manufactured from long chain hydrocarbons obtained as a byproduct of

the petroleum industry. The hydrocarbons are treated with concentrated sulphuric acid and

sodium hydroxide. Neutral sodium salt is obtained, which is the synthetic detergent.

Advantages and disadvantages of synthetic detergents with respect to soaps

Soaps Detergents

Soap cannot be used in hard water. Synthetic detergents can be used in hard

water.

Soap is made from vegetable/edible oils Synthetic detergents are made from

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byproducts of petroleum industry which helps

to conserve valuable edible oils.

Soaps cannot be used in acidic medium as this

would precipitate the fatty acids.

Synthetic detergents can be used in any

medium including acidic medium.

Soaps have weak cleansing action. Synthetic detergents have strong cleansing

action.

Soaps are not very soluble in water. Synthetic detergents are highly soluble.

Soaps are biodegradable and do not cause any

pollution.

Some synthetic detergents are not

biodegradable and cause water pollution. This

in some cases can become quite serious.

EXERCISE

1) (a) Equal volumes of different samples of water were shaken with equal

volumes of soap solution in three separate test tubes.

The height of the lather in each test tube was measured.

sample of water height of lather (cm)

A 0

B 6

C 6

(i) What is formed in the mixture of A and soap solution instead of lather?

(ii) What is the name of the type of water that does not form a lather with soap solution?

(iii) How could you treat another sample of A so that it would form a lather when shaken

with soap solution?

(iv) Which of the samples could be pure water?

(b) Describe and explain what you would see when the following solutions are shaken with

soap solution.

(i) Calcium nitrate solution

(ii) Lithium nitrate solution

2) (a) what is hard water?

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(b) What causes ‘temporary’ and ‘permanent’ hardness in water?

(c) Explain, using suitable equations how each type of hardness may be removed from water

3) The structures below represent two cleansing agents

R-COO-Na+

R-OSO3-Na+

Give one advantage and one disadvantage of using each one of them.

POLYMERS

Polymerization

Is the process of formation of large molecules or macromolecules with high molecular weight from

small molecules.

"Poly" means "many" and "mer" means "parts." So "polymer" means "many parts." The parts are

usually the same part used repeatedly in a chain-like manner. Polymers are also referred to as

plastics because they are easily molded.

Polymers are not just a human invention. Nature has many examples of polymers. Proteins and

carbohydrates are natural polymers. Synthetic fibers, resins, plastics, and synthetic rubber are all

synthetic or manmade polymers.

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POLYMERISATION

The simple compounds from which these macromolecules are made are called monomers. Many

monomers combine together to form a high molecular weight substance called polymer.

Monomers

It's often easy to see if something could become a polymer. From the shape of these parts, you

can probably guess how they might stick together. These parts are called monomers.

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Polymer

The rounded end will connect to the open end. This approach is used by many toys that let you

build things. It is also used in chemistry to build things. This model represents a polymer

There are two types of polymers. These are additional and condensation polymers.

Additional polymers

Alkenes such as ethene contain a double bond between carbon atoms. These molecules can

take part in additional reactions where the double bond is broken and other atoms attach to the

carbons. The only product of additional polymerisation is the polymer. Ethene monomers produce

polyethene; propene produces polypropene and so on.

Various conditions can be used to produce different types of additional polymers. Generally a

high pressure, a temperature above room temperature and a catalyst are needed.

Addition polymerisation:

All alkenes will react with free radical initiators to form polymers by a free radical addition

reaction.

Some definitions -

monomer - a single unit e.g. an alkene.

The alkene monomer has the general formula:

Where R is any group of atoms, e.g. R=CH3 for propene.

The reaction progresses by the separate units joining up to form giant, long chains -

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Polymer- a material produced from many separate single monomer units joined up together.

An addition polymer is simply named after the monomer alkene that it is prepared from.

Alkene Additional polymer

Ethene polyethene

Propene poly propene

Phenylethene phenylethene

Chloroethene poly chloroethene

The structure above shows just 4 separate monomer units joined together. In a real polymer,

however, there could be 1000's of units joined up to form the chains. This would be extremely

difficult to draw out and so the structure is often shortened to a repeat unit.

There are 3 stages to think about when drawing a repeat unit for an addition polymer -

1) Draw the structure of the desired monomer.

2) Change the double bond into a single bond and draw bonds going left and right from the

carbon atoms.

3) Place large brackets around the structure and a subscript n and there is the repeat unit.

Additional Polymers Uses

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Plastic bags, bowls,

packaging

Insulation and pipes

Crates, boxes, plastic

ropes

Non-stick frying pans

What is natural rubber?

Natural rubber is a soft and sticky solid. Rubber is a long chain polymer made out of isoprene (2 –

methyl butadiene) monomer. The long molecule forms a coil like structure.

The unique property of rubber is that it is elastic. When rubber is stretched, the molecular bonds

can be extended out. When released, the molecules coil back to their original shape.

Natural rubber looses its rubber-like properties at temperatures above 60C. Also its wear

resistance and tensile strengths are low.

What is happening inside rubber during vulcanization?

At the vulcanizing temperature (practically, 140-200°C), a complex chemical reaction starts to

occur and leads to the formation of sulphur crosslink as shown schematically below.

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a, b and c denote the number of sulphur molecules which range from 1-6.

Why rubber has to be vulcanized?

The process of vulcanization can improve the quality of rubber. Raw rubber is heated with sulphur

during vulcanization. This makes the rubber hard, more elastic and strong. During the process of

vulcanization, the sulphur atoms attach themselves to extra loose bonds in the rubber molecule

and also cross-link the molecules. The cross-linking locks the molecules in place and prevents

slipping. Thus making the vulcanized rubber more strong. Vulcanized rubber is non-sticky and

has higher elasticity. It does not loose its properties easily and can be used in a temperature

range of – 40C to 100C.

Condensation polymers

Condensation polymerisation involves linking lots of small monomer molecules together by

eliminating a small molecule. This is often water from two different monomers, a H from one

monomer, and an OH from the other, the 'spare bonds' then link up to form the polymer chain.

Classes of condensational polymers

There are two categories of condensational polymer. These are synthetic and Natural polymers.

Synthetic polymers: These are man-made polymers e.g. nylon and polyesters (Terylene)

Natural polymers: These are naturally occurring polymers e.g. carbohydrates and proteins. Nylon

and proteins have a similar link, commonly known as peptide or amide link.

Polyesters (Terylene) and carbohydrates have a similar link commonly known as ester link.

Link Found in

Peptide or amide Nylon and proteins

Ester Polyesters (terylene) and carbohydrates

Nylon and proteins

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Nylon

Nylon is a synthetic condensation polymer formed when di-carboxylic acid and di-amine with

elimination of water. Hence the name condensation.

Nylon is a thermoplastic polymer, and can be spun, cast, machined and moulded, and thus finds

a wide variety of uses, such as nylon fabrics, clothing and footwear, machine parts, fishing line,

ropes, guitar strings, and all sorts of sporting equipment.

Proteins

Amino acids are carboxylic acids (like ethanoic acid) but one of the hydrogen atoms of the 2nd

carbon atom is substituted with an amino group (nitrogen + two hydrogens gives -NH2). Hydrogen

on the same 2nd carbon can be substituted with other groups of atoms (R) to give a variety of

amino acids.

is another amino acid called 2-aminopropanoic acid

or The simplest is aminoethanoic acid

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All amino acids have the general structure H2N-CH(R)-COOH.

They polymerise together, by condensation polymerisation, forming proteins or polypeptides. The

peptide linkage is formed by elimination of water between two amino acids.

HNH-CH(R)-COOH + HNH-CH(R)-COOH H2N-CH(R)-CO-HN-CH(R)-COOH +

H2O

+ nH20

Polyesters (Terylene) and carbohydrates

Polyesters (terylene)

Terylene (a polyester) is formed by condensation polymerisation and the structure of Terylene

represented as.

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This is the same kind of 'ester linkage' (-COOC-) found in fats which are combination of long

chain fatty carboxylic acids and glycerol (alcohol with 3 -OH groups, a 'triol').

Terylene and nylon are good for making 'artificial' or 'man-made' fibres used in the clothing and

rope industries.

Properties of Polyester

Polyester and nylon have similar properties. Polyester is

Very strong. So garments made from polyester fibres last longer than the ones made

from natural fibres

Very resistant to wrinkles. Polyester garments retain their crease

Not a good absorbent of water. Therefore polyester clothes dry faster

Highly abrasion and moth resistant

Resistant to ordinary chemicals and biological agents

Uses of Polyester Fibres

Polyester is used in making:

Textiles, sarees, dress materials, and curtains

Blended textiles with natural fibres. For example, when terylene is mixed with cotton, it is

called terrycot and with wool, it is called terrywool. These blended fibres are again used

in making a whole range of textiles and garments

Sails for sailboats that are light, strong, don't stretch and don't rot in contact with water

Water hoses for fire-fighting operations

Conveyor belts

Carbohydrates

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Carbohydrates are a whole series naturally occurring molecules based on the elements carbon,

hydrogen and oxygen. Their formation can be described in terms of a large number of sugar units

joined together by condensation polymerisation

n C6H12O6 ==> (C5H10O5)n + nH2O (where n is a very large number to form the natural

polymer)

The [rectangles] below, represent the rest of the carbon chains in each unit

Plus many H2O

molecules etc.

Acid hydrolysis of complex carbohydrates (e.g. starch) gives simple sugars. This can be

brought about by e.g. warming starch with hydrochloric acid solution to form glucose. The

hydrolysis products from polysaccharides can be analysed with paper chromatography.

(C5H10O5)n + nH2O n C6H12O6 (where n is a very large number)

CONSEQUENCES OF MAKING & USING PLASTICS

The favorite properties of plastics are that they are inert and won't react with what is stored in

them. They also are durable and won't easily decay, dissolve, or break apart. These are great

qualities for things you keep, but when you throw them away, they won't decompose.

n

+ n H20

Carbohydrate

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EXERCISE

1) More than 60 000 plastic materials, or polymers, are in use. The table gives some

information about five important polymers.

(a) Which polymer would be most suited for making a pipe to carry lubricating oil at

100 °C? Give two reasons for your answer.

(b) State one use for poly (ethene).

(c) Describe some of the problems of the disposal of waste polymers.

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(d) Poly (propene) is made from the monomer propene. Draw the structure of poly

(propene).

2) Propanoic acid, C2H5CO2H, is a weak acid.

(a) Explain what is meant by the term weak acid.

(b) Propanoic acid reacts with sodium carbonate. Write the equation for this reaction.

(c) Magnesium reacts with propanoic acid to form magnesium propanoate and hydrogen.

A student added 4.80 g of magnesium to 30.0 g of propanoic acid.

(i) Which one of these reactants, magnesium or propanoic acid, is in excess?

Explain your answer.

(ii) Calculate both the number of moles of hydrogen and the volume of hydrogen formed at r.t.p.

(d) Terylene has the simplified structure shown.

A student added 4.80 g of magnesium to 30.0 g of propanoic acid.

(i) Which one of these reactants, magnesium or propanoic acid, is in excess?

Explain your answer.

(ii) Calculate both the number of moles of hydrogen and the volume of hydrogen formed at r.t.p.

(d) Terylene has the simplified structure shown.

(i) State the functional groups on the monomers used to make Terylene.

(ii) State the type of polymerisation that occurs when Terylene is made.

(iii) State one large scale use of Terylene.

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(e) Many problems are caused by the disposal of plastics.

Describe one method of disposal of a plastic and a problem caused by this method.

3) Glucose and starch are carbohydrates.

(i) The chemical formula of glucose is C6H12O6. State the total number of atoms which are

combined in one molecule of glucose.

(ii) Starch is a polymer which has been formed from glucose. Explain the meaning of this

statement.

CHEMISTRY OF NON-METALS

Forms of carbon and sulphur

Allotropy and Allotropes of Carbon

The existence of an element in more than one form, in the same physical state, is called allotropy.

The various forms are called allotropes. Allotropy is the result of different arrangement of the

atoms in the crystal structure of the element. As a result, the allotropes exhibit very different

physical properties but have similar chemical properties. Thus diamond and graphite are the

allotropes of carbon. Many other elements like sulphur also exhibit this allotropic property.

As stated before, carbon exists in nature in two crystalline forms - diamond and graphite. Though

both of these are pure carbon, they have different physical properties. The flow chart below

shows the allotropes of carbon.

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DIAMOND

Carbon has an electronic arrangement of 2, 4. In diamond, each carbon shares electrons with

four other carbon atoms - forming four single bonds.

Diamond is the purest form of natural carbon. It occurs as small crystals embedded in rocks.

These are supposed to have been formed by the crystallization of carbon under extreme pressure

and temperature in the interior of the earth. Nowadays, synthetic industrial diamonds are being

manufactured by subjecting graphite to very high temperatures and pressures. Carbon atoms in

diamond have tetrahedral structure. Each atom of carbon is surrounded by four other atoms that

together forms the tetrahedral structure, as shown in the figure below.

In the diagram some carbon atoms only seem to be forming two bonds (or even one bond), but

that's not really the case. We are only showing a small bit of the whole structure.

This is a giant covalent structure - it continues on and on in three dimensions. It is not a molecule,

because the number of atoms joined up in a real diamond is completely variable - depending on

the size of the crystal.

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The physical properties of diamond

Diamond:

Has a very high melting point (almost 4000°C). Very strong carbon-carbon covalent

bonds have to be broken throughout the structure before melting occurs.

Is very hard. This is again due to the need to break very strong covalent bonds operating

in 3-dimensions.

Doesn’t conduct electricity. All the electrons are held tightly between the atoms, and

aren't free to move.

Is insoluble in water and organic solvents. There are no possible attractions which could

occur between solvent molecules and carbon atoms which could outweigh the attractions

between the covalently bound carbon atoms.

Chemical Properties of Diamond

1. Diamond is chemically very inert. It does not react with any substance at ordinary

temperatures.

2. When heated in oxygen to about 800oC, it completely burns to form carbon dioxide. This shows

that diamond is pure form of carbon.

3. When heated in the absence of air to 1500oC, the atoms get rearranged to form graphite.

4. Diamond is affected slowly by molten sodium carbonate, forming carbon monoxide.

5. When heated with concentrated sulphuric acid and potassium dichromate it gets oxidised to

carbon dioxide.

Uses of Diamond

1. Diamond is used as a gem (except the black variety) due to its brilliance.

2. Black variety of diamond is use for cutting glass, as drilling bits for industrial drills, for polishing

other diamonds etc.

The giant covalent structure of graphite

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Unlike the tetrahedral arrangement of atoms in diamond, the carbon atoms in graphite are

arranged in the form of hexagonal rings in layers. Each carbon atom uses three of its electrons to

form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding

level. These "spare" electrons in each carbon atom become delocalised over the whole of the

sheet of atoms in one layer. They are no longer associated directly with any particular atom or

pair of atoms, but are free to wander throughout the whole sheet. The important thing is that the

delocalised electrons are free to move anywhere within the sheet - each electron is no longer

fixed to a particular carbon atom. There is, however, no direct contact between the delocalised

electrons in one sheet and those in the neighbouring sheets. Graphite has a layer structure which

is quite difficult to draw convincingly in three dimensions.

The atoms within a sheet are held together by strong covalent bonds So what holds the sheets

together?

In graphite you have the ultimate example of Van der Waals dispersion forces. As the delocalised

electrons move around in the sheet, very large temporary dipoles can be set up which will induce

opposite dipoles in the sheets above and below - and so on throughout the whole graphite

crystal. Because of this property, it is also used as a lubricant.

The physical properties of graphite

Graphite

Has a high melting point, similar to that of diamond. In order to melt graphite, it isn't

enough to loosen one sheet from another. You have to break the covalent bonding

throughout the whole structure.

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Has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks.

You can think of graphite rather like a pack of cards - each card is strong, but the cards

will slide over each other, or even fall off the pack altogether. When you use a pencil,

sheets are rubbed off and stick to the paper.

Has a lower density than diamond. This is because of the relatively large amount of

space that is "wasted" between the sheets.

Is insoluble in water and organic solvents - for the same reason that diamond is insoluble.

Attractions between solvent molecules and carbon atoms will never be strong enough to

overcome the strong covalent bonds in graphite.

Conducts electricity. The delocalised electrons are free to move throughout the sheets. If

a piece of graphite is connected into a circuit, electrons can fall off one end of the sheet

and be replaced with new ones at the other end.

Uses of Graphite

1. Graphite is used in making the 'lead' of pencils.

2. It is used in the production of refractory crucibles, which can withstand very high temperature.

3. Graphite being a conductor of electricity finds application in making electrodes.

4. It is used in making polishes and paints.

5. Graphite is used as lubricant in machines, which have to be operated at high temperatures. All

such machines cannot be lubricated with oils, grease, etc. as they vaporize immediately at the

high temperature. As a lubricant it is used as dry powder or mixed with water or oil. When mixed

with water, it is called 'aqua-dag' and when mixed with oil, it is called 'oil dag'.

6. It is used for making electrotypes for printing, in the following manner. Wax impressions are

made and then a thin layer of graphite powder is applied. Copper is deposited on this thin layer.

The layer of graphite is used to give the negative electrical connection for carrying out electrolysis

to deposit copper on it . After coating the required thickness, the wax can be melted out by

dipping it in hot water.

7. Graphite is extensively used in nuclear reactors, to absorb neutrons. This helps in moderating

nuclear reactions.

Amorphous carbon

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Apart from diamond and graphite, which are crystalline forms of carbon, all other forms of carbon

are amorphous (shapeless) allotropes of carbon.

1. Coke

Coke is the amorphous allotrope of carbon, which is derived from coal. When coal undergoes

destructive distillation, it yields two allotropes of carbon, namely coke and gas carbon. Destructive

distillation is a chemical process, which involves is the breaking up of a complex substance by

heating it in the absence of air.

Uses of coke

It is a very good fuel and when ignited it burns almost with no smoke. It is a non-conductor of heat

and electricity. It acts as a good reducing agent and is extensively used in the production of

producer gas, water gas and hydrogen.

2. Sugar Charcoal

Sugar charcoal can be obtained by dehydrating cane sugar, either by treating it with concentrated

sulphuric acid or by heating it in the absence of air.

It is the purest form of the amorphous variety of carbon. It is used in the preparation of artificial

diamonds.

3. Wood Charcoal

Wood charcoal is obtained by the destructive distillation of wood. The chief products formed are

wood charcoal.

Properties.

Wood charcoal is black, porous, brittle and soft. Though denser than water it can float on water,

as it contains plenty of air bubbles trapped in the pores. Wood Charcoal is not a conductor of

electricity.

Uses

Wood charcoal is mostly used as a fuel as it catches fire easily.

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It is used in gas masks as an adsorbent.

It is also used as a decolourising agent for sugar, oils, alcohol, petroleum products, etc.

It is used in gun powder which is a mixture of charcoal powder, potassium nitrate and

sulphur.

4. Animal Charcoal

It is prepared by the destructive distillation of bones of animals. It is porous and can adsorb

colouring matter. It is mostly used in sugar industry to decolourise sugar.

SULPHUR

It takes the form of a yellow solid naturally and can be found in this state near volcanoes. Sulphur

is also present in a number of metal ores, for example zinc blende (zinc sulphide, ZnS ).

Sulphur has chemical symbol S. It has 16 protons and 16 neutrons. An atom of S is represented

as 3216S. Sulphur is a non-metal and exists in the earth’s crust either as pure sulphur or as a

metal-sulphide.

Since S has 16 protons, it also has 16 electrons; the electronic configuration of S is 2, 8, 6. S is

placed in Group VI A of the periodic table, just after phosphorus, and below oxygen. The reaction

of S is similar to oxygen.

Sulphur is found as a free element or in combined state in nature. Free sulphur is found in at a

large depth below the earth’s surface. Metal sulphides such as Zn, Fe, Ag, Ca, Pb, Cu are found

in abundant quantities. Mineral ores containing S are:

Copper pyrites : CuFeS2

Galena : PbS

Zinc blende : ZnS

Iron pyrite : FeS2

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Sulphur is found as H2S gas in petroleum gas, coal gas. H2S is the familiar pungent smell of

onions. It is present in hair, eggs, many proteins and wool.

Extraction of pure sulphur : Frasch process

Since sulphur in free state is found at depths of more than 150 to 300 meters below the earth’s

surface, the method of extraction of sulphur differs from other metal or non-metal extractions.

Sulphur’s relatively low melting point (115C) is utilized in this process. This is known as the

Frasch process. Here compressed super heated water (at 170C) is pressed into a pipe which

reaches up to the sulphur deposits. The sulphur here melts. Introducing hot compressed air

through another pipe brings it up. The molten sulphur and water mixture is forced up and is

collected in a settling tank. The sulphur is cooled and water is evaporated. The sulphur extracted

in this way is more than 99% pure.

The sulphur obtained by Frasch process is a yellow and brittle solid or powder.

Physical properties of sulphur:

Since S has 6 electrons in its outermost shell, it needs 2 more electrons to complete its shell. But

S combines with 7 other atoms to make a sulphur molecule that has a total of 8 sulphur atoms.

Thus each S atom shares 2 electrons with its neighboring atom. The bonds are covalent in

nature. A molecule of sulphur is represented as S8. It is a ringed molecule. The structure is shown

below.

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These large molecules each have many electrons , so the Van der Waals forces are quite strong

and the melting point is quite high (119oC). There are two ways of packing the sulphur rings, so

solid sulphur exists in two crystalline forms, called rhombic and monoclinic.

Allotropes of sulphur

Rhombic sulphur Monoclinic sulphur

Yellow transparent crystals Amber coloured needles

Melt pt. 113 oC Melt pt. 119 oC

Obtained when sulphur crystallises Obtained when sulphur

from solution. solidifies above 95.6 oC.

Sulphur is a yellow crystalline solid. It is tasteless and odourless. The melting point of S is 115C.

Sulphur is an insulator and is a poor conductor of heat and electricity. S is insoluble in water but

is soluble in CS2. Sulphur forms covalent bonds and shows allotropic forms. The allotropes have

different crystalline shapes such as rhombic and monoclinic. There is another allotrope which has

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no shape and is called plastic sulphur. Vapours of sulphur are pungent and although not

poisonous, they can cause health problems to humans.

Effect of heat on sulphur:

A sulphur molecule consists of 8 atoms in a ring form. When heated, S melts at 115C and a pale

yellow liquid is formed. The S8 ringed molecules are connected to other molecules in a long

chain. On heating, the long chain breaks up. The individual molecules can slip over each other

when melted. On further heating, the liquid becomes dark brown and viscous. When the

temperature rises beyond 160C, the intra-molecular bonds break. Sulphur boils at 444C. At this

temperature the large molecule breaks up into pieces of S2 molecule. This molecule is pale

yellowish-brown in colour. The vapours of S contains S2 molecules.

The Properties Of Sulphur:

Non-metallic yellow solid at room temperature

Brittle

Insoluble in water

Soluble in organic solvents, for example methylbenzene

Non-conductor of electricity whether solid, molten of dissolved.

Relatively low melting point and boiling point.

Uses of sulphur :

S is used to make H2SO4 acid, which is used in the manufacture of many compounds

such as detergents, plastics, explosives, etc.

S is used for making CS2 molecule, gun powder, matches etc.

S is used for manufacture of fire works.

S is used in the rubber industry for vulcanization of rubber.

S is used for making germicides, fungicides.

S is used in many medicines.

S is used in photographic development (sodium thiosulphate or hypo).

S is used for making bleaching agents.

S is used in making artificial hair colours or dyes.

Carbon (IV) oxide:

Carbon dioxide is easily prepared by the action of dilute hydrochloric acids on metal carbonates

(normally calcium carbonate or marble). Vigorous effervescence occurs as bubbles of carbon

dioxide are liberated.

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CaCO3(s) + 2HCl (aq) CaCl2 + CO2 (g) + H2O (l)

Since carbon dioxide is 1.53 times as heavy as air, it usually collected by upward displacement of

air. In this case, it is collected by downward displacement of water.

Properties of carbon (IV) oxide

Physical properties

Carbon (IV) oxide is a colourless gas with a faint pungent smell. It does not burn or support

combustion, except in extreme cases, and is not poisonous (it is the gas in fizzy cool drinks). It

can however cause death by suffocation, when it is present in sufficient concentrations.

Carbon (IV) oxide is a linear molecule. The gas condenses to a liquid at 0 ºC under a pressure of

35 atm. At normal pressure, Carbon (IV) oxide condenses directly to a solid at -78.5 ºC. This

solid, known as dry ice, is widely used as cooling agent. Solid Carbon (IV) oxide does not melt

under conditions of normal atmospheric pressure, but passes directly into the gas phase, a

process known as sublimation.

Chemical properties

1. Reaction with water

When carbon dioxide is dissolved in water, carbonic acid, H2CO3, is produced in small quantities:

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CO2 (g) + H2O (l) H2CO3 (aq)

Carbonic acid is a weak diprotic ( forms two hydrogen ions) acid which gives rise to salts known

as carbonates, which contain the carbonate anion CO32-.

2. Reaction with limewater

Test for carbon dioxide

Limewater is a clear colourless solution of calcium hydroxide (slaked lime). Calcium carbonate is

precipitated when carbon dioxide is passed through a clear solution of calcium hydroxide in

water.

The lime water turns milky serving as a test for carbon dioxide liberation.

On passing excess of CO2 gas the milky ness disappears due to the formation of calcium

hydrogen carbonate, which is soluble in water.

However, if this colourless solution of calcium hydrogen carbonate is boiled, it decomposes

forming the insoluble calcium carbonate, water and carbon dioxide. As a result, the milkiness

reappears.

4. Reaction with alkalis

Carbon dioxide is readily absorbed by alkalis such as sodium hydroxide and potassium

hydroxide, to form their respective carbonates.

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The above reaction of carbon dioxide with potassium hydroxide is used to purify air. When air is

passed through a solution of potassium hydroxide, it absorbs carbon dioxide.

Hence, air, which is first passed through aqueous potassium hydroxide, and then passed through

limewater, does not turn the latter milky, as all the carbon dioxide is absorbed by the potassium

hydroxide

5. Nature

Carbon dioxide is slightly acidic. It turns blue litmus paper red. During rainy season, blue litmus

paper kept open in the laboratory slowly turns red, due to the presence of carbon dioxide in air.

6. Combustibility

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Carbon dioxide is neither combustible, not a supporter of combustion. A burning splinter or a

burning candle gets put off, but metals like potassium, sodium, magnesium etc. continue to burn

in carbon dioxide.

Burning magnesium in carbon dioxide

Ignite a ribbon of magnesium, and introduce it in the jar of carbon dioxide. The magnesium ribbon

continues to burn in carbon dioxide. Deposits of carbon can be seen on the inner sides of the jar.

At the temperature of ignition, these metals are able to reduce the carbon dioxide to carbon by

taking away the oxygen i.e., tiny black solid particles of carbon get deposited inside the jar.

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Action on heated carbon

When carbon dioxide is passed over red-hot carbon in the form of coke charcoal, the carbon

dioxide loses one of its two atoms of oxygen. As a result, carbon dioxide gets reduced and

becomes carbon monoxide. At the same time the hot carbon also gets converted to carbon

monoxide.

Uses of Carbon Dioxide

1. In nature

a) Carbon dioxide is used in photosynthesis by green plants to produce carbohydrates.

b) To induce natural breathing.

2. To extinguish fires

Soda-acid fire extinguishers produce carbon dioxide to put out fires.

3. As a refrigerant

Solid carbon dioxide called "Dry ice" can provide temperatures as low as -109.3o F. It is superior

to ordinary ice, for the following reasons:

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i) It provides much lower temperature than ice.

ii) It lasts longer.

iii) It freezes faster.

iv) It does not wet the food being chilled, as it sublimes directly into gaseous state.

4. Manufacture of fertilizer

Carbon dioxide is used extensively in the manufacture of urea, an important in nitrogenous

fertilizer.

5 In the baking industry

Baking powder is used in all the food preparations. The addition of baking powder during baking

produces carbon dioxide which makes the dough "rise". The small pores in a loaf of bread are the

spaces in which carbon dioxide was formed. Yeast is also used in baking instead of baking

powder. It produces carbon dioxide by anaerobic respiration.

Baking powder contains starch, sodium hydrogen carbonate and an acid forming ingredient, such

as tartaric acid of calcium hydrogen phosphate [Ca(H2PO4)2] or alum [Na2SO4.Al2(SO4)3.24 H2O].

6. In medicine

A mixture of 97% oxygen and 3% carbon dioxide, called carbogen is used to revive persons

affected by carbon monoxide poisoning, pneumonia, asphyxiation etc.

7. Manufacture of aerated drinks

Carbon dioxide is extensively used in aerated drinks. Increasing the pressure increases the

solubility of the gas. The fizz in the drink is due to carbon dioxide being liberated when the

pressure is reduced.

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Carbon (II) oxide (Carbon monoxide):

Carbon(II) oxide is an odourless, tasteless and colourless gas, which is insoluble in water. It is

extremely poisonous. Under no circumstances must the gas be inhaled or smelled. It is not

usually prepared in a school laboratory. If need be, the gas should be prepared in a fume

chamber.

Preparation of Carbon (II) oxide

1. By dehydrating oxalic acid with hot concentrated sulphuric acid

Carbon (II) oxide is prepared with the help of oxalic acid and concentrated sulphuric acid as

shown below.

Sulphuric acid reacts with oxalic acid and removes from it one molecule of water (both the

hydrogen atoms, along with an oxygen atom). The product left behind due to this reaction, is a

molecule of carbon (IV) oxide and a molecule of Carbon (II) oxide. The carbon dioxide can be

removed by passing it through a concentrated solution of potassium hydroxide.

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2. Preparation of Carbon (II) oxide by dehydrating Methanoic acid

Methanoic acid (also known as formic acid) has the formula HCOOH. Methanoic acid can also be

dehydrated in a similar way by hot concentrated sulphuric acid. Sulphuric acid removes two

atoms of hydrogen and one atom of oxygen as a molecule of water from it, and leaves behind

one molecule of Carbon (II) oxide.

Carbon (II) oxide is produced whenever organic matter is burnt in a limited quantity of oxygen.

For this reason, it is found in the exhaust gases of motor vehicles, as well as in cigarette smoke.

Indoor fireplaces can be a serious hazard if ventilation is poor.

Industrial preparation

It is prepared industrially (mixed with hydrogen) by passing steam over coke at temperatures

above 900 ºC. The resulting gas mixture is known as water gas, and it is used as a fuel:

Physical Properties of Carbon (II) oxide

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Property description

Nature

Carbon (II) oxide is colorless, almost odorless and tasteless gas.

Density

It is very slightly lighter than air.

Solubility

Carbon (II) oxide is only very slightly soluble in water.

Poisonous nature

This is a highly poisonous gas. Air containing even less than 1% of

Carbon (II) oxide, can be fatal, if breathed in for about 10 to 15 minutes.

Chemical Properties of Carbon (II) oxide

a) Nature

Carbon (II) oxide is a neutral oxide. It is neither acidic nor basic.

b) Stability

It is very stable and cannot be decomposed by heat.

c) Combustibility

It is a combustible gas. It burns well in air or oxygen to form carbon (IV) dioxide. The formation of

carbon (IV) dioxide is tested by passing it through a solution of lime water.

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This is a highly exothermic reaction. Hence it is a very good fuel.

d) Reducing property

Carbon (II) oxide is a powerful reducing agent. When CO is passed over heated metallic oxides, it

takes away the oxygen to form carbon (IV) dioxide and reduces the oxides to their respective

metals.

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SULPHUR (IV) OXIDE

Sulphur (IV) dioxide is a colourless gas, about 2.5 times as heavy as air, with a suffocating smell,

faint sweetish odour.

Occurrence

Sulphur (IV) dioxide occurs in volcanic gases and thus traces are present in the atmosphere.

Other sources of sulphur (IV) oxide are the combustion of the iron pyrites which are contained in

coal. Sulphur (IV) dioxide also results from various metallurgical and chemical processes.

Preparation of Sulphur (IV) oxide

Sulphur (IV) dioxide is prepared by burning sulphur in oxygen or air.

S + O2 SO2

Sulphur (IV) dioxide is usually made in the laboratory by heating concentrated sulphuric acid with

copper turnings.

Cu + 2 H2SO4 CuSO4 + SO2 + 2 H2O

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Sulphur (IV) dioxide is released by the action of acids on sulphites or acid sulphites (e.g. by

dropping concentrated sulphuric acid into a concentrated solution of sodium hydrogen sulphite).

NaHSO3 + H2SO4 NaHSO4 + SO2 + H2O

Properties of Sulphur Dioxide

Sulphur (IV) dioxide is a colourless liquid or pungent gas, which is the product of the combustion

of sulphur on air. Its melting point is -72.7 0C, its boiling point is -100C and its relative density is

1.43.

Sulphur (IV) dioxide is an acidic oxide which reacts with water to give sulphurous acid.

SO2 + H2O H2SO3

Sulphur (IV) dioxide is a good reducing and oxidising agent.

Summary

Physical Properties

Colour Colourless

Odour Pungent odour

Density compared to

air (heavier or

lighter)

Heavier than air

Chemical Properties

Solubility in water Soluble. It reacts with water to

form a strong acid.

Burning Does not support combustion

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Moist pH paper Acidic reaction

Red rose petals Are bleached and lose their colour

Specific test None

Uses of Sulphur (IV) Dioxide

a. Sulphur (IV) dioxide is a reducing agent and is used for bleaching and as a fumigant and

food preservative.

b. Large quantities of Sulphur (IV) dioxide are used in the contact process for the

manufacture of Sulphuric (VI) acid.

c. Sulphur (IV) dioxide is used in bleaching wool or straw, and as a disinfectant.

d. Liquid Sulphur (IV) dioxide has been used in purifying petroleum products

e. It is used as a bleaching agent.

REACTION AND PROPERTIES OF SULPHURIC (VI) ACID

Sulphuric (VI) acid is a dense, oily liquid once known as oil of vitriol. Pure Sulphuric (VI) acid is

almost twice as dense as water (1.98 g cm-2). As water is added the density drops. Car batteries

contain concentrated Sulphuric (VI) acid. As the battery is discharged, the concentration of the

acid falls. By measuring the density of the acid the driver can check whether the battery is flat or

not.

Action as an oxidising agent

It behaves as an oxidising agent only when hot and concentrated:

Cu + 2H2SO4 CuSO4 + H2O + SO2

The Sulphuric (VI) acid is reduced to sulphur (IV) oxide.

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Action as a dehydrating agent

Concentrated Sulphuric (VI) acid has a great affinity for water. (It is important when diluting the

concentrated acid to add the acid to water and NEVER water to acid.) The reaction is highly

exothermic.

So great is its affinity for water that it can dehydrate compounds containing hydrogen and oxygen:

C12H22O11 + nH2SO4 12C + (11H2O + nH2SO4)

sucrose + sulphuric Carbon + (water + sulphuric acid)

acid

CH3CH2OH + nH2SO4 CH2 = CH2 + (H2O + nH2SO4)

ethanol + sulphuric ethene + (water + sulphuric acid)

acid

It is used for drying gases, especially SO2 and HCl, but cannot be used to dry a reducing gas

such as H2S or an alkaline gas such as NH3.

Action as a dehydrating agent

The properties of acids are due to the hydrogen ions in solution. Concentrated Sulphuric (VI) acid

contains molecules, rather than ions. Since it contains very few hydrogen ions it does not react

significantly with metals and can safely be stored in steel containers. A piece of magnesium

ribbon does not dissolve in concentrated Sulphuric (VI) acid.

Diluted with water, Sulphuric (VI) acid behaves as a typical acid:

it reacts with metals to form sulphates plus hydrogen gas

it reacts with metal carbonates to form metal sulphates plus carbon dioxide plus water

It neutralises bases to form sulphates plus water.

Industrial uses

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THE CONTACT PROCESS

This page describes the Contact Process for the manufacture of Sulphuric (VI) acid, and then

goes on to explain the reasons for the conditions used in the process. It looks at the effect of

proportions, temperature, pressure and catalyst on the composition of the equilibrium mixture, the

rate of the reaction and the economics of the process.

A brief summary of the Contact Process

The Contact Process:

Makes sulphur(IV) dioxide;

Converts the sulphur dioxide into sulphur trioxide (the reversible reaction at the heart of

the process);

Converts the sulphur trioxide into concentrated sulphuric acid.

Making the sulphur (IV) oxide

This can either be made by burning sulphur in an excess of air:

101

or by heating sulphide ores like pyrite in an excess of air:

In either case, an excess of air is used so that the sulphur (IV) oxide produced is already mixed

with oxygen for the next stage.

Converting the sulphur dioxide into sulphur trioxide

This is a reversible reaction, and the formation of the sulphur trioxide (sulphur (VI) oxide is

exothermic.

A flow scheme for this part of the process looks like this:

Converting the sulphur trioxide into Sulphuric (VI) acid

This can't be done by simply adding water to the sulphur trioxide - the reaction is so

uncontrollable that it creates a fog of Sulphuric (VI) acid. Instead, the sulphur trioxide is first

dissolved in concentrated Sulphuric (VI) acid:

102

The product is known as fuming Sulphuric (VI) acid or oleum.

This can then be reacted safely with water to produce concentrated Sulphuric (VI) acid - twice as

much as you originally used to make the fuming sulphuric acid.

Summary

Explaining the conditions

The proportions of sulphur dioxide and oxygen

The mixture of sulphur (IV) oxide and oxygen going into the reactor is in equal proportions by

volume.

Avogadro's Law says that equal volumes of gases at the same temperature and pressure contain

equal numbers of molecules. That means that the gases are going into the reactor in the ratio of 1

molecule of sulphur dioxide to 1 of oxygen.

That is an excess of oxygen relative to the proportions demanded by the equation.

According to Le Chatelier's Principle, Increasing the concentration of oxygen in the mixture

causes the position of equilibrium to shift towards the right. Since the oxygen comes from the air,

this is a very cheap way of increasing the conversion of sulphur dioxide into sulphur trioxide.

Why not use an even higher proportion of oxygen? This is easy to see if you take an extreme

case. Suppose you have a million molecules of oxygen to every molecule of sulphur dioxide.

The equilibrium is going to be tipped very strongly towards sulphur trioxide - virtually every

molecule of sulphur dioxide will be converted into sulphur trioxide. Great! But you aren't going to

produce much sulphur trioxide every day. The vast majority of what you are passing over the

catalyst is oxygen which has nothing to react with.

103

By increasing the proportion of oxygen you can increase the percentage of the sulphur dioxide

converted, but at the same time decrease the total amount of sulphur trioxide made each day.

The 1: 1 mixture turns out to give you the best possible overall yield of sulphur trioxide.

The temperature

Equilibrium considerations

You need to shift the position of the equilibrium as far as possible to the right in order to produce

the maximum possible amount of sulphur trioxide in the equilibrium mixture.

The forward reaction (the production of sulphur trioxide) is exothermic.

According to Le Chatelier's Principle, this will be favoured if you lower the temperature. The

system will respond by moving the position of equilibrium to counteract this - in other words by

producing more heat.

In order to get as much sulphur trioxide as possible in the equilibrium mixture, you need as low a

temperature as possible. However, 400 - 450°C isn't a low temperature!

Rate considerations

The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to

produce as much sulphur trioxide as possible per day. It makes no sense to try to achieve an

equilibrium mixture which contains a very high proportion of sulphur trioxide if it takes several

years for the reaction to reach that equilibrium.

You need the gases to reach equilibrium within the very short time that they will be in contact with

the catalyst in the reactor.

The compromise

400 - 450°C is a compromise temperature producing a fairly high proportion of sulphur trioxide in

the equilibrium mixture, but in a very short time.

104

The pressure

Equilibrium considerations

Notice that there are 3 molecules on the left-hand side of the equation, but only 2 on the right.

According to Le Chatelier's Principle, if you increase the pressure the system will respond by

favouring the reaction which produces fewer molecules. That will cause the pressure to fall again.

In order to get as much sulphur trioxide as possible in the equilibrium mixture, you need as high a

pressure as possible. High pressures also increase the rate of the reaction. However, the reaction

is done at pressures close to atmospheric pressure!

Economic considerations

Even at these relatively low pressures, there is a 99.5% conversion of sulphur dioxide into

sulphur trioxide. The very small improvement that you could achieve by increasing the pressure

isn't worth the expense of producing those high pressures.

The catalyst

Equilibrium considerations

The catalyst has no effect whatsoever on the position of the equilibrium. Adding a catalyst doesn't

produce any greater percentage of sulphur trioxide in the equilibrium mixture. Its only function is

to speed up the reaction.

Rate considerations

In the absence of a catalyst the reaction is so slow that virtually no reaction happens in any

sensible time. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to

be set up within the very short time that the gases are actually in the reactor.

Properties of Sulphuric (VI) acid

Chemical Formula H2SO4

105

Molar mass 98 g mol-1

Melting point 10oC

Boiling point 340oC

Density 1.83g cm-3

COMPOUNDS OF NITROGEN

Nitrogen (IV) oxide:

Nitrogen (IV) oxide, NO2 is a deep red-brown gas, which may be prepared by the action of

concentrated nitric acid on copper:

Cu + 4HNO3 Cu (NO3)2 + 2H2O + 2NO2

Or by the decomposition of heavy-metal nitrates, such as lead nitrate:

2Pb(NO3)2 2PbO + 4NO2 + O2

At temperatures below 100 ºC, it forms dinitrogen tetroxide, N2O4:

2NO2 N2O4 (REVERSIBLE ARROW)

Nitrogen dioxide will support combustion, as shown by the fact that a glowing splint of wood will

ignite in this gas.

106

AMMONIA

Ammonia is a colorless gas. It has a characteristic pungent odor. It is bitter to taste. Its vapor

density is 8.5. Hence it is lighter than air. When cooled under pressure ammonia condenses to a

colorless liquid, which boils at -33.4oC. When further cooled, it freezes to a white crystalline snow-

like solid, which melts at -77.7oC. Ammonia is one of the most soluble gases in water. At 0oC and

760 mm of Hg pressure one volume of water can dissolve nearly 1200 volumes of ammonia. This

high solubility of ammonia can be demonstrated by the fountain experiment.

Preparation of Ammonia

By Heating any Ammonium Salt with an Alkali

In the laboratory, ammonia is usually prepared by heating a mixture of ammonium chloride and

slaked lime in the ratio of 2 : 3 by mass. The arrangement of the apparatus is shown in the figure

6.2. As ammonia is lighter than air, it is collected by the downward displacement of air.

107

Drying of Ammonia

The drying agent used for ammonia is quick lime. Other drying agents such as concentrated

sulphuric acid or phosphorus (V) oxide or fused calcium chloride cannot dry an alkaline gas like

ammonia. Sulphuric acid and phosphorus (V) oxide are both acidic. They react with ammonia,

forming their respective ammonium salt.

Industrial Preparation of Ammonia

Haber's Process

Ammonia is manufactured by Haber's process using nitrogen and hydrogen.

Raw materials

a. Nitrogen is obtained in large scale from air. This is done by fractonal distillation of liquid

air.Air free from dust and carbon dioxide is cooled under high pressure and low

temperature to about 200oC and then allowed to warm. As nitrogen has lower boiling

point (-169oC) as compared to oxygen (-183oC) it turns to gas leaving oxygen in liquid

state. Nitrogen can also be obtained by heating a ammonium nitrite (in small amounts)

b. Hydrogen is obtained from cracking of petroleum

108

c. Catalyst is finely divided iron

Reactants: Nitrogen gas -1 volume and hydrogen gas -3 volumes

Conditions

Temperature Pressure Catalyst

450 – 500OC 250 at (optimum) Finely divided iron

The reaction in Haber’s process is exothermic and so external heating is not required once the

reaction starts. Lowering the temperature to 450o - 500oC favours the reaction, but lowering the

temperature below 450o - 500oC brings down the yield.

Yield analysis at different temperatures

109

Chemical Properties of Ammonia

Combustibility

Ammonia is neither combustible in air nor does it support combustion. However it burns in oxygen

with a greenish-yellowish flame producing water and nitrogen.

a) Burning of Ammonia in Oxygen

Activity

Set the apparatus as shown in figure 6.7.

110

Firstly, when ammonia is passed through the longer tube and is made to ignite, it does not catch

fire. Then oxygen is sent through the shorter tube. Now when ammonia is ignited, it catches fire

and the following reaction takes place:

Although the products formed in the above reaction are insignificant, it is an extremely important

reaction from viewpoint of industry. This is because in the presence of platinum, catalytic

oxidation of ammonia can take place to give various important products.

b) Catalytic Oxidation of Ammonia

The platinum coil is heated at 800oC in a combustion tube till it becomes white hot. Then

ammonia and oxygen are passed through the tube. Under these conditions and in the presence

of the catalyst, ammonia combines with free oxygen or oxygen of the air, to form nitric oxide and

water vapour.

As colourless nitric oxide comes out into the air, it cools down and combines with the oxygen of

the air to form reddish brown fumes of nitrogen dioxide.

111

The formed nitrogen dioxide is converted into nitric acid in the presence of water and oxygen.

The importance of the above reactions lies in the production of nitric acid, which is a very

important industrial product.

Basic Nature

Absolutely dry ammonia or pure liquefied ammonia is neutral. In the presence of water however,

it forms ammonium hydroxide, which yields hydroxyl ions. As a result of this reaction, it exhibits

basic nature. It is a weak base and is perhaps the only gas that is alkaline in nature.

Ammonia is an alkaline gas. When damp red litmus paper is introduced into the gas, it turns blue

due to the presence of hydroxide ions as shown in the equation above.

112

Colour changes with other indicators

Test for ammonia gas

TEST METHOD OBSERVATIONS TEST CHEMISTRY

Strong pungent odour

tested with damp red

litmus

Litmus turns blue Only common alkaline

gas

conc. hydrochloric

acid

White clouds with

HCl fumes.

forms fine ammonium

chloride crystals with

HCl

Ammonia as a reducing agent

Heated dry ammonia gas can reduce copper (II) oxide to pure copper. This reaction can be used

to prepare nitrogen.

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NH3 (g) + CuO(s) N2 (g) + H2O (g)

The gas passes through a U- tube surrounded by cold water which contains some melting ice.

This helps to condense the vapour produced to liquid water. Nitrogen is finally collected by

downward displacement of water.

Solubility of ammonia.

Fountain experiment

Fill a clean dry round-bottomed flask with dry ammonia, close it by a one holed stopper, through

which a long jet tube is introduced. The free end of the tube is dipped into a trough of water as

shown.

Add two or three drops of an acid and a small quantity of phenolphthalein to the water in the

trough. This water is colorless. Pour a small quantity of spirit or ether on a layer of cotton and

place it over the inverted flask.

114

Due to the cooling effect produced by the process of evaporation of spirit or ether, the ammonia

gas contracts a little and as a result, small quantity of the water gets sucked up. As soon as this

water enters the flask, the ammonia dissolves in it, forming a partial vacuum. As a result of it,

water rushes in and comes out of the tube as a jet of fountain. The color of the water turns deep

pink.

Dissolving Ammonia in Water

Due to its high solubility, ammonia cannot be passed through water like many other gases.

Ammonia is dissolved in water, as shown below.

This arrangement is called funnel arrangement and its principle is the same as that discussed for

HCl gas. The funnel arrangement prevents back suction of water, which can cause damage to the

apparatus used. It provides larger surface area for dissolution of ammonia. A very strong solution

of ammonia in water is called liquor ammonia. Ammonia can be obtained from it by boiling.

The properties of ammonia

Colourless gas

Distinctive pungent smell

Less dense than air

115

Very soluble in water to give an alkaline solution

Action with Acids

Ammonia reacts with the acids to form their respective ammonium salts. The ammonium salts

appear as white fumes

Ammonia gas + acid ammonium salt

Uses of Ammonia

Following are the chief uses of ammonia:

1) Ammonia is used in the industrial preparation of nitric acid by Ostwald's process.

116

2) Fertilisers, such as ammonium sulphate, ammonium nitrate, ammonium phosphate, urea etc.

are manufactured with the help of ammonia.

3) It is used in the manufacture of other ammonium salts, such as ammonium chloride,

ammonium carbonate, ammonium nitrite etc.

4) It finds use in the manufacture of nitrogen compounds such as sodium cynamide, plastics,

rayon, nylon, dyes etc.

5) It is used in the manufacture of sodium carbonate by Solvay's process. (Ammonia and carbon

dioxide are treated with aqueous sodium chloride, crystals of sodium hydrogen carbonate are

formed. They are heated to yield sodium carbonate)

6) Ammonia acts as refrigerant in ice plants. Evaporation of a liquid needs heat energy. About

17g of liquid ammonia absorb 5700 calories of heat from the surrounding water. This cools the

water and ultimately freezes it to ice.

7) Ammonia is used to transport hydrogen. Hydrogen is dangerous to transport, as it is highly

combustible. So it is converted to ammonia, liquefied, transported and then catalytically treated to

obtain hydrogen.

8) Many ammonium salts are used in medicines. Inhaling the fumes produced by rubbing

ammonium carbonate in the hands can revive people who have fainted.

9) It is used as a cleansing agent. Ammonia solution emulsifies fats, grease etc. so it can be used

to clean oils, fats, body grease etc. from clothes. It is also used to clean glassware, porcelain,

floors etc.

10) It is used as laboratory reagent.

CHLORINE

Halogen is elements in group (vii) of the periodic table. Chlorine is a halogen as well as fluorine,

bromine, iodine and astatine.

Chlorine has a symbol 35.5 Cl because it is made up of two isotopes 37 Cl and 35 Cl. It has an

electronic arrangement of 2:8:7, hence justifying its position in group (vii).

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Laboratory preparation

In order to convert hydrogen chloride to chlorine, it is necessary to remove hydrogen. Removal of

hydrogen is oxidation. A powerful oxidizing agent such as manganese (IV) oxide converts

hydrogen chloride (HCl) to chloride (Cl2)

The most common laboratory method for preparation of Chlorine is to heat of Manganese Dioxide

with concentrated Hydrochloric Acid.

The gas is bubbled through water to remove any traces of hydrochloric gas that may be present

and then it is dried by bubbling it through concentrated sulphuric acid.

Chlorine may also be prepared by dropping cold concentrated Hydrochloric Acid on crystals of

Potassium Permanganate.

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2 KMnO4 + 16 HCl 2 MnCl2 + 2 KCl + 8 H2O + 5 Cl2

The gas is also bubbled through water to remove any traces of Hydrochloric Acid gas that may be

present and then it is dried by bubbling it through concentrated Sulphuric Acid.

Manufacture OF CHLORINE

Membrane cell

Chlorine is manufactured industrially as a by-product in the manufacture of Caustic Soda by the

electrolysis of brine.

2 NaCl + 2 H2O Cl2 + H2 + 2 NaOH

The membrane cell has titanium anode and a nickel cathode. Titanium is chosen because it is not

attacked by chlorine. The anode and the cathode compartments are separated by an ion

exchange membrane. The membrane is selective; it allows Na+ ions, H+ and OH- ions to flow but

not Cl- ions. These ions cannot flow backwards, so products are kept separate and cannot react

with each other

At anode, the Cl- ions are discharged more readily than OH- ions because they are in higher

concentration and are hence preferred.

2 Cl- Cl2 + 2e-

A pale green gas is seen coming off at the anode

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At cathode, it is the H+ ions that accept electrons, as sodium is more reactive than hydrogen.

2 H+ + 2e- H2

Bubbles of hydrogen are seen at the cathode. The remaining ions of Na+ and OH- join up and

come off as sodium hydroxide, NaOH.

Products and uses

Product Uses

Chlorine Poisonous greenish yellow gas

Used for making; PVC, solvents for dry

cleaning, paints and dye stuffs, bleaches, weed

killers, pesticides, killing bacteria in swimming

pools and in domestic water treatment.

Hydrogen Colourless flammable gas

Used for making:

Margarine, nylon, hydrochloric acid

Sodium hydroxide Alkaline and corrosive substance

Used for making; Soap, detergents, textiles,

textiles and paper.

Properties of chlorine

Test for Chlorine Gas, Cl2(g).

1) Will turn moist litmus or universal indicator paper red,

and then bleach it white.

(2) Is green-yellow in colour.

(3) Has a pungent choking smell and is poisonous. It is twice as dense as air.

(4) Will put out a lit splint.

120

Collecting Chlorine.

Chlorine is denser than air and can be collected by downward delivery or using a gas syringe.

Reactions

1. Chlorine is a highly reactive element, and undergoes reaction with a wide variety of other

elements and compounds.

2. Chlorine is a good bleaching agent, due to its oxidising properties.

3. Chlorine is soluble in water (which solution is called Chlorine Water) and this loses its yellow

colour on standing in sunlight, due to the formation of a mixture of Hypochlorous Acid and

Hydrochloric Acid.

Cl2 + H2O HOCl + HCl

4. Chlorine gas supports the vigorous combustion of many elements to form their chlorides. For

example, Sulphur and Phosphorus burn in the gas.

Cl2 + S SCl2

Cl2 + P PCl3 + PCl5

Bleaching Action

If chlorine is passed through water, it forms two acids, hydrochloric acid.

Cl2 (g) + H2O (l) HCl (aq) + HOCl (aq)

Hypochlorous acid (the second acid) is the source of oxygen and is responsible for the bleaching

of chlorine.

HClO (aq) + Dye HCl (aq) + oxidized Dye

(Coloured) (Colourless)

It is important to wash bleached clothes thoroughly to remove hydrochloric acid formed after the

process.

121

Reaction of Chlorine with Hydrogen

A mixture of Chlorine and Hydrogen explodes when exposed to sunlight to give Hydrogen

Chloride. In the dark, no reaction occurs, so activation of the reaction by light energy is required.

Hydrogen and chlorine gas also combine directly in presence of sunlight. A jar of hydrogen is

inverted and placed on a jar containing chlorine in the sun. Soon hydrogen chloride is formed

In diffused sunlight, the reaction slows down, and in the dark it is very slow.

Cl2 + H2 2 HCl

Hydrogen chloride is highly soluble in water. It dissolves to form hydrochloric acid. This reaction

can be used to produce hydrochloric when the hydrogen chloride gas produced is dissolved in

water as shown.

122

Bleaching Action of Chlorine

Chlorine bleaches (removes the color) organic colors by the process of oxidation in presence of

moisture. The bleaching action takes place in few steps:

a) Chlorine first dissolves in water to give a mixture of hydrochloric acid and hypochlorous acid.

b) As mentioned earlier hypochlorous acid is very unstable and decomposes to give hydrochloric

acid and oxygen.

c) The nascent oxygen oxidises the coloring matter to colorless matter thereby bleaching them.

Bleaching by chlorine is permanent. Chlorine bleaches cotton fabrics, wood pulp litmus etc.

However, it is not used to bleach delicate articles such as silk, wool etc., as it is a strong

bleaching and oxidizing agent. This dual action will damage the base material.

You must have noticed that in the very first stage, presence of water is essential to produce

hypochlorous acid.

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The dry coloured fabric does not bleach. The coloured fabric soaked in water is bleached. This

shows that chlorine only bleaches in the presence of water. If water is absent no bleaching can

take place. Dry chlorine therefore does not bleach.

Tests for Chlorine

1. Chlorine is a greenish yellow gas with a pungent and irritating odor.

2. It turns wet blue litmus paper red and then bleaches it.

3. Colored petals and the leaves of plants can be bleached by it.

4. Chlorine turns wet starch potassium iodide paper blue by displacing the iodine from the

potassium iodide, and causing iodine to turn the starch blue.

Pollution

EXERCISE.

1) (a) Sulphuric acid is manufactured by the Contact process.

Use words from the box to complete the paragraph below. Each word may be used once,

more than once or not at all.

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Air Sulphur sulphur dioxide

sulphur trioxide Sulphuric acid water

The raw materials for the Contact process are..........................

and.......................................................They are heated together to form

.......................................................More air and ..................................................... are then

heated and passed over the catalyst. The catalysed reaction produces

.......................................................

(b) State two uses of sulphuric acid.

2) Diamond and graphite are both allotropes of carbon

(a) What is Allotropes?

(b) Explain why;

(i) Graphite Conduct Electricity while diamond does not.

(ii) Diamond is hard while graphite is soft

3) Nitrogen (I) oxide (dinitrogen oxide), N2O, is prepared by gentle heating of ammonium

nitrate as shown.

(a) Write the equation for the reaction taking place in the flask

(b) Suggest a use for nitrogen (I) oxide

125

4) Chlorine, hydrogen and sodium hydroxide are made by the electrolysis of concentrated

aqueous sodium chloride. Aqueous sodium chloride contains the following ions, Na+, H+,

OH- and Cl- Concentrated aqueous sodium chloride can be electrolysed using inert electrodes.

The electrode reactions are represented below.

Cathode 2H+ + 2e- → H2

Anode 2Cl- → Cl2 + 2e-

(i) Explain why hydrogen, not sodium, is formed at the cathode.

(ii) Suggest why, as the electrolysis proceeds, the concentration of Sodium hydroxide in the

electrolyte increases.

(iii) Give one use of chlorine

5. Sulphur dioxide, SO2, and nitrogen dioxide, NO2, are both atmospheric pollutants formed

during the combustion of coal at a power station.

(a) (i) State another source of sulphur dioxide as an atmospheric pollutant.

(ii) State another source of nitrogen dioxide as an atmospheric pollutant.

(b) Nitrogen dioxide and sulphur dioxide both cause acid rain. They are removed from the flue

gases released from the power station by reaction with moist calcium carbonate in a process

called flue gas desulphurisation. Calcium carbonate reacts with sulphur dioxide to make a solid

called calcium sulphite and a gas.

(i) What is the name of this gas?

126

(ii) Nitrogen dioxide reacts with calcium carbonate to make a solid. Suggest the name

of this solid.

(iii) Describe one environmental effect of acid rain.

FORM 4

ACIDS, BASES AND INDICATORS

Indicators

An indicator is a substance or mixture of substances that when added to the solution gives a

different colour depending on the pH of the solution. Universal indicator solution or paper, is

prepared from mixing several indicators to give a variety of colours to match the pH. It is a very

handy indicator for showing whether the solution is very weakly/strongly acidic (pH <7) or alkaline

(pH > 7) or neutral (pH = 7) and gives the pH to the nearest pH unit.

pH can be very accurately measured with a special instrument called a pH meter using a glass

electrode which is calibrated with buffer solutions of accurately known pH.

Acidic solutions have a pH of less than 7, and the lower the number, the stronger the acid it. The

colour can range from orange-yellow (pH 3-6) for partially ionised weak acids like ethanoic acid

(vinegar) to carbonated water. Strong acids like hydrochloric, sulphuric and nitric are fully ionised

and give a pH 1 or less! and a red colour with universal indicator or litmus paper.

Neutral solutions have a pH of 7. These are quite often solutions of salts, which are themselves

formed from neutralising acids and bases. The 'opposite' of an acid is called a base. Some bases

are soluble in water to give alkaline solutions - these are known as alkalis. Alkaline solutions

have a pH of over 7 and the higher the pH the stronger is the alkali. Weak alkalis (soluble bases)

like ammonia give a pH of 10-11 but strong alkalis (soluble bases) like sodium hydroxide give a

pH of 13-14. They give blue-purple-violet colour with universal indicator or litmus paper.

Table 3.3: Other common indicators used in the laboratory

Indicator colour in acid pH<7 colour in neutral pH=7 colour in alkali pH >7

Litmus red 'purple' blue

Phenolphthalein colourless colourless >9 pink

Methyl orange <3.5 red, orange about yellow yellow

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pH 5, > 6 yellow

Methyl red <5 red, orange, >6

yellow yellow yellow

Bromothymol blue <6 yellow

MAKING CABBAGE INDICATOR

Acid base indicators are chemicals that change colour in the presence of different pH levels.

These are usually larger organic molecules. Some, like that in purple cabbage, are natural.

You will be making an acid base indicator from purple cabbage. This indicator is a very good one

with good color changes.

RED CABBAGE

The color of red cabbage is affected by the concentration of protons (acid level) in the plant. The

red of this leaves is affected by the acidity level in the petals. Change the acid levels, and the

color of the rose will change.

Red Cabbage

Finding out:

Preparing a red cabbage indicator

MATERIALS NEEDED:

Tea strainer

2 Glass quart jars with lids

1 Quart distilled water

uncooked purple cabbage

Hotplate and pan

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PROCEDURE:

a. Fill one jar with cabbage leaves that have been crushed into small pieces.

b. Heat the distilled water to boiling, and fill the jar containing the pieces of cabbage

with the hot water.

c. Allow the jar to stand until the water cools to room temperature.

d. Poor the cooled cabbage solution through the tea strainer into the second quart

jars. Discard the cabbage leaves.

e. Store the cabbage indicator in a cool place until needed.

RESULTS:

The hot water dissolved the colored chemicals in the cabbage. These colored chemicals

turn red when mixed with an acid, and green when mixed with a base. Red cabbage has

a particular wide range of color change depending on the pH. This indicator can be used

to test for the presence of either an acid or a base.

Table 3.5: Comparing indicators

pH Universal indicator Red Cabbage indicator

4.0 Red Red

5.0 Orange-Red Purple

6.0 Yellow-Orange Purple

7.0 Dark Green Purple

8.0 Light Green Blue

9.0 Blue Blue-Green

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10.0 Reddish-Violet Green

11.0 Violet Green

12.0 Violet Green

13.0 Violet Green-Yellow

14.0 Violet Yellow

The pH values of some common solutions are shown in the table below.

Table: Ph of common substances

Substances PH

Strongly

acidic

Hydrochloric acid (HCl) 0.0

Gastric juices 1.0

Lemon juice 2.5

Vinegar 3.0

Wine 3.5

Tomato juice 4.2

Black coffee 5.0

Acid rain 5.6

Urine 6.0

Weakly

Acidic

Neutral

Weakly

Alkaline

Rain water 6.5

Milk 6.5

Pure water 7.0

Blood 7.4

Baking soda solution 8.5

Toothpaste 9.0

Borax solution 9.2

Milk of magnesia 10.5

Lime water 11.0

Strongly

Alkaline

Household ammonia 12.0

Sodium hydroxide

(NaOH)

14.0

Introduction to Acids and Bases

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Humans have been familiar with acids and bases for thousands of years. They didn’t know the

exact chemistry but they knew a lot about them.

Vinegar means sour wine. Vinegar is acetic acid in water. Acetum is a historical name for vinegar

and comes from the word acere` meaning to be sour. Citrus fruits also have a sour or tart taste

due to a different acid called citric acid. Sweet-Tarts get their sour or tart taste from citric acid.

What is an acid?

When a substance dissolves in water, the solution may be acidic, neutral or alkaline. An acid is

any substance which produces H+ ions or H3O+ ions when dissolved in water. H+ ions are called

hydrogen ions; H3O+ ions are called hydroxonium ions. You will mostly see acids in reactions as

forming H+ ions. In reality, H+ is a single proton, and does not exist on its own. It always attaches

to something; in water it joins to H2O to form H3O+ ions. All acids taste sour and are mostly

derived from oxides of non-metals dissolved in water.

DID YOU KNOW

The word acid comes from Latin word acidus meaning sour, sharp, or tart taste.

EXAMPLES OF ACIDS

1. Sulphuric acid nitric acid

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ETHANOIC ACID

Inorganic acids are usually stronger than organic acids.

Table : dissociation of acids

Name of acid Nature Chemical Formula Dissociation in water

Hydrochloric acid In-organic HCl H+ + Cl-

Nitric Acid In-organic HNO3 H+ + NO3-

Sulphuric acid In-organic H2SO4 2H+ + SO4-

Ethanoic acid organic CH3COOH H+ + CH3COO-

Measure of acidity (pH).

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A discussion of acids and bases is not complete without an explanation of "pH". Everyone has

heard of pH but few actually know what that means. First of all "pH" gets its name from "potential

for Hydrogen" more specifically, the hydrogen ion (H+). So it's a scale that reflects the

concentration of H+.

PH is a measure of how acidic or how alkaline a solution in water is. The pH scale goes from 1 to

14, with 1 being very strongly acidic, and 14 being very strongly alkaline. A pH of 7 is neutral.

Each colour stands for a particular pH as shown below

ph scale and universal indicator

You can measure the pH of a solution using universal indicator. Just as litmus paper will be red

for an acid and blue for an alkali, so universal indicator is a mixture of indicators which will give

different colours for a different pH.

An acid is a substance that dissolves in water to produce hydrogen ions (H+). This gives a

solution which:

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Contains an excess of H+ ions,

Turns litmus red.

Has a pH lower than 7.

Turn blue litmius paper red

A strong acid (HCl or H2SO4 or HNO3 ) will have a pH of 1 (red).

Acid strength increases with decreasing PH

Acids can be neutralised (which means cancelled out) by alkalis

Strong acids are corrosive (which means that they dissolve things)

Weak acids include lemon juice and vinegar

Strong acids include battery acid, stomach and laboratory acids

A weak acid will have a pH of 3 to 4 (orange). Examples of weak acids are ethanoic acid

(vinegar), citric acid (lemon juice) and rain water. Rain water has a natural pH of 5·5 (see

carbonic acid). Water and salts are neutral, pH 7 (green). A weak alkali (ammonia) will have a pH

of 11 to 12 (blue). A strong alkali (Ca(OH)2 or NaOH) will have a pH of 14 (purple).

Diluting Concentrated Acids

Dilution of concentrated acid should always be done in a fume cupboard. It is important only to

pour acid into water, not the other way around, especially with concentrated acids. Acids may

quickly absorb water, creating a lot of heat in the process. When acid is poured into water, the

134

heat can quickly become evenly distributed in the water. If water is poured into acid, the water

may quickly boil, spraying acid everywhere.

What is the difference between strong acid and concentrated acid?

A strong acid is one that is completely dissociated or ionised in water, when the concentration is

not too high. Examples are HCl, H2SO4 and HNO3

A weak acid is one that is only partly dissociated, no matter what the concentration. For example,

acetic acid (the active ingredient in vinegar) is only about 1% dissociated in water, even in dilute

solutions. An example of a weak acid is ethanoic acid

A concentrated acid is an acid which is either pure (no solvent) or has a high concentration.

Sulphuric acid (about 98% by weight, the other 2% is water) is a concentrated strong acid.

Acids are found in:

citrus fruits (le mon juice, orange juice)

vinegar (ethanoic acid)

Car batteries (sulphuric acid)

your stomach (hydrochloric acid)

Rainwater is a little acidic, but pollution (e.g. sulphur dioxide) from burning fossil fuels may make

it even more acidic, forming acid rain. When acids are present in food, they usually taste sour

(think of the taste of lemon juice or vinegar).

Did you Know?

Bees and wasps are both insects which use a sting as part of their defense.

The pain we feel if we are stung is because of the change in our blood pH. The pH values

of their stings are shown on the diagrams.

135

Properties of Acids.

They have a pH less than 7, see pH. They will turn blue litmus paper red. Hydrochloric acid and

sulphuric acid will react with.

Any alkali or base, see neutralisation.

Any metal above hydrogen in the reactivity series. The metal will fizz, giving off hydrogen

gas, and leaving the metal salt in solution. It is not safe to put a metal into an acid which is

above magnesium in the reactivity series. Any chloride or sulphate can be safely made by

reacting the appropriate metal (from lead to magnesium in the reactivity series) with

hydrochloric acid to make the chloride or sulphuric acid to make the sulphate.

Any metal carbonate or metal hydrogen carbonate. The metal carbonate or metal hydrogen

carbonate will bubble giving off carbon dioxide gas, leaving the metal salt and water. Any

chloride or sulphate can be made by reacting the appropriate metal carbonate or hydrogen

carbonate with hydrochloric acid to make the chloride or sulphuric acid to make the sulphate.

Strong and Weak Acids - Strength and Concentration.

Acids and alkalis can be described as strong or weak. This does not mean the same as

concentrated or dilute. The strength of an acid or alkali depends on how ionised it is in water. A

strong acid or alkali is completely (100%) ionised. For hydrochloric acid

hydrogen chloride (in water) hydrogen ion + chloride ion

HCl(aq) H+(aq) + Cl-(aq)

All of the hydrogen chloride molecules become hydrogen ions and chloride ions in water

(see examples for other strong acids).

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For sodium hydroxide

NaOH(aq) Na+(aq) + OH-(aq)

Sodium hydroxide exists as ions both in water and in the solid.

(See examples for other strong alkalis). A weak acid or alkali is only partly (less than 100%)

ionised. For ethanoic acid

Ethnic acid (in water) hydrogen ion + ethanoic ion

CH3CO2H(aq) H+(aq) + CH3CO2-(aq)

Some of the ethanoic acid molecules become ions in water but most of them stay as molecules.

The reaction is reversible (shown by the arrow).

Ammonia

Ammonia + water ammonium ion + hydroxide ion

NH3(g) + H2O(l) NH4+(aq) + OH-(aq)

Some of the ammonia molecules become ions in water but most of them stay as molecules. See

also Concentration and Differences between Strong and Weak Acids.

Common uses of Acids.

1. Steel used in construction is acid treated before painting. Dilute sulphuric or hydrochloric

acid will remove any surface rust which would otherwise spread under the painted

surface. 'Rust remover' used to repair cars is dilute phosphoric acid - H3PO4.

2. Baking powder contains tartaric acid.

3. 'Lime scale' removers contain dilute acids. Try using lemon juice or vinegar (weak acids).

Lime scale is calcium carbonate (also called furring).

4. A wasp sting is alkaline. It may be neutralised with a weak acid (lemon juice or vinegar).

5. A bee sting is acidic. It may be neutralized by an alkali.

BASES AND ALKALIS

Bases are substances that react with acids and neutralise them. They are usually metal oxides,

metal hydroxides, metal carbonates or metal hydrogen carbonates. Many bases are insoluble -

they do not dissolve in water. If a base is soluble and does dissolve in water, we call it an alkali.

All alkalis are also bases, but not all bases are alkalis.

137

Here are two examples:

Copper oxide is a base because it will react with acids and neutralise them, but it is not

an alkali because it does not dissolve in water.

Sodium hydroxide is a base and also an alkali, because not only will it react with acids

and neutralise them, but it dissolves in water.

DID YOU KNOW

Al Kali is the Arabic name for the plant we call the saltwort plant. What was discovered a couple

thousand years ago was that the ashes from the saltwort plant had the ability to neutralize the

power of acids.

Which base is an alkali?

Alkali is pronounced like alcohol, with 'lie' at the end instead of 'hol'. An alkali is any substance

which produces OH- ions in water. OH- ions are called hydroxide ions. A strong base is one which

is fully ionized in aqueous solution. A weak base is one which is partially ionized in aqueous

solution. Both strong and weak bases are alkalis because they are soluble in water. Some bases

are insoluble in water. They are called insoluble bases.

138

Bases and Alkalis

The outer circle encloses all bases, while the inner circle selects those which are alkalis or

soluble bases.

1. SODIUM HYDROXIDE

Sodium hydroxide (NaOH), also known as caustic soda, is a caustic metallic base. Sodium

hydroxide forms a strong alkaline solution when dissolved in a solvent such as water, however,

only the hydroxide ion is basic. It is used in many industries, mostly as a strong chemical base in

the manufacture of pulp and paper, textiles, drinking water, soaps and detergents and as a drain

cleaner.

Pure sodium hydroxide is a white solid; available in pellets, flakes, granules and as a 50%

saturated solution. It is hygroscopic and readily absorbs carbon dioxide from the air, so it should

be stored in an airtight container. It is very soluble in water with liberation of heat.

139

Chemical properties

Sodium hydroxide is completely ionic, containing sodium cations and hydroxide anions. The

hydroxide anion makes sodium hydroxide a strong base which reacts with acids to form water

and the corresponding salts, e.g., with hydrochloric acid, sodium chloride is formed:

NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)

This type of reaction with a strong acid, releases heat, and hence is referred to as exothermic.

Such acid-base reactions can also be used for titrations, which is a common method to determine

the concentration of acids. Another type of reaction that sodium hydroxide is involved in is with

acidic oxides. The reaction of carbon dioxide has already been mentioned, but other acidic oxides

such as sulphur dioxide (SO2) also react completely. Such reactions are often used to "scrub"

harmful acidic gases (like SO2 and H2S) and prevent their release into the atmosphere.

2NaOH + CO2 → Na2CO3 + H2O

Concentrated alkalis are corrosive. They can attack metals and destroy skin if spilled. They are

just as dangerous as concentrated acids, but many people do not realise this. Strong alkaline

substances are just as dangerous as strong acidic substances, causing very serious burns if they

come into contact with your skin.

2. POTASSIUM HYDROXIDE

140

Potassium hydroxide is the inorganic compound with the formula KOH. Along with sodium

hydroxide, this colourless solid is a prototypical "strong base".

Potassium hydroxide is usually sold as translucent pellets, which will become tacky in air because

KOH is hygroscopic. Consequently, KOH characteristically contains varying amounts of water. Its

dissolution in water is strongly exothermic, meaning the process gives off significant heat.

Strength of Alkalis.

A substance which will neutralize an acid, but does not dissolve in water, is called a base.

For example, copper (II) oxide, iron (II) oxide and zinc carbonate are bases. They do not dissolve

in water.

Alkalis are the bases that dissolve in water. The three common alkalis you will find in the

laboratory are;

1) Sodium Hydroxide solution - NaOH (aq).

2) Calcium Hydroxide solution - Ca (OH)2(aq), (lime water)

(Lime water is used in the test for carbon dioxide).

3) Ammonia solution - NH3 (aq).

1 and 2 are strong alkalis, 3 is a weak alkali - see pH.

They all ionise in water to form hydroxide ions (OH- ions).

1) NaOH (aq) Na+(aq) + OH-(aq)

2) Ca (OH) 2(aq) Ca2+(aq) + 2OH-(aq)

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3) NH3 (aq) + H2O(l) NH4+(aq) + OH-(aq)

NH4+ is an ammonium ion.

If there are excess of (OH) - ions when a compound is dissolved in water, the solution is called a

base or an alkaline solution. A base is generally a metal hydroxide solution. Strong bases have a

PH value of 14

Table below the dissociation of alkalis in water

Table : Strength of Alkalis

Name of alkali Chemical Formula Dissociation in water Strength in water

Sodium Hydroxide NaOH Na+ + (OH)- Strong

Potassium Hydroxide KOH H+ + NO3- Strong

Ammonium Hydroxide NH4OH NH4+ + (OH)- Weak

Any metal oxide or hydroxide is a base. If the base dissolves in water it is called an alkali.

Alkalis are found in:

1. oven cleaner (sodium hydroxide)

2. soap

3. cleaning fluid e.g. spray-and-wipe (ammonia)

Notice the connection between these substances? Alkalis are often found in substances for

cleaning.

Properties of Alkalis.

1. They have a pH greater than 7, see pH.

2. They will turn red litmus paper blue.

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They will react with acids to form a salt and water, see neutralisation.

Uses of Alkalis.

Sodium hydroxide is used in the manufacture of paper, soap and ceramics.

Calcium hydroxide (called 'slaked lime', or just 'lime'), is added to soils or lakes to make them

less acidic (see acid rain).

Indigestion may be caused by too much hydrochloric acid in the stomach. Indigestion tablets

contain a base such as magnesium oxide, or calcium carbonate to neutralise the acid.

A bee sting is acidic. It may be neutralised by a weak alkali such as ammonia solution.

Table : Weak, strong acids and alkalis

strong acid ph <2 weak acid ph 2-6 neutral ph 7 weak alkali ph 8-

12

strong alkali ph

>12

H2SO4 sulphuric

acid (battery

acid) pH 1

vitamin C

(ascorbic acid)

H2O water NaHCO3 sodium

hydrogen

carbonate

('bicarb', baking

soda) pH 8

NaOH sodium

hydroxide pH 13-

14

HCl hydrochloric

acid (in the lab is

same as your

stomach!) pH 1

CH3COOH

acetic/ethanoic

acid (vinegar)

NaCl sodium

chloride (salt

water) pH 7

NH3 ammonia pH

11

KOH potassium

hydroxide pH 13-

14

HNO3 nitric acid

pH 1

Fruit juices e.g.

oranges and

lemons contain

citric acid pH 2-3

MgSO4

magnesium

sulphate (Epsom

salts) pH 6.5-7.0

Na2CO3 sodium

carbonate

(washing soda)

pH 11

oven cleaner

(may contain

NaOH) pH can

be >12

rain water (pH C6H12O6 sugar toothpaste pH 8

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5.5 normal, down

to 3.5 when

polluted)

pH 7

milk pH 6 C2H5OH ethanol

(alcohol) pH 7

bleach pH 11

wine/beer pH 6 Ca(OH)2 calcium

hydroxide

(limewater) pH

10

Mg(OH)2

magnesium

hydroxide ('milk

of magnesia') pH

10

EVERYDAY USES OF ACIDS AND ALKALIS

Food preservation

Food can be made to last longer by 'pickling'. This means storing in vinegar (ethanoic

acid)

The acid stops bacteria from multiplying and decaying the food.

Soil treatment

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Different plants prefer different soil pH's. Some are acid-loving and some are alkali-

loving.

Soil pH can be changed by adding lime or mildly acidic compounds.

Sting bites

Bee stings are acidic and should be neutralised by applying a mild antacid (washing

soda).

Wasp stings are alkaline and need to be treated with vinegar.

Hair and skin care

Shampoos tend to be alkaline so that they can remove grease. After shampooing, a

conditioner will neutralise the pH of the hair.

After using a face cleanser it’s a good idea to use a 'pH balanced' moisturiser.

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Digestion & indigestion

The human stomach produces quite a strong acid to help in digestion. If it produces too much we

get 'acid indigestion' or heartburn, and have to take a mild alkali in the form of an ANTACID.

Toothpaste

Soap

Sodium hydroxide is used in the manufacture of soap. Soaps are alkaline. However it should not

be above a pH of 10. Nine or less would be better. If your soap's above a pH of 10 it's likely to

cause a burning sensation on your skin.

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Neutral substances

Neutral substances have a pH of 7

Neutral substances are neither acidic or alkaline

Because a substance is neutral it doesn't mean it isn't harmful

Neutral substances include water

Hard and Soft Water

HARD and SOFT WATER: Many compounds dissolve in water without chemical change but may

have a variety of consequences. Water which readily gives a lather with soap is described as soft

water. Detergents usually give a good lather with any water. Some of these dissolved substances

make the water hard. This means the water does not readily give a good lather with soap and so

wastes soap as well as causing a 'scum', though it does not affect soapless detergents. The

simplest test for 'hardness' is to shake the water soap. Most hardness is due to water containing

dissolved calcium or magnesium compounds.

How water becomes hard

We all breathe our carbon dioxide gas; the internal combustion engine exhausts carbon dioxide;

aeroplanes produce carbon dioxide, and so on. Therefore, the atmosphere must contain carbon

dioxide. Carbon dioxide is soluble in water, so that rainwater contains a certain amount of

carbonic acid.

When acidified rainwater runs of these rocks, some of the calcium and magnesium go into

solution as free ions.

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1. Calcium and magnesium sulphates are washed out of rock formations on hillslopes such as

those on Mt. Kenya. These make water permanently hard.

2. Calcium and magnesium carbonates dissolve in acid rainwater to form hydrogen carbonates.

These make the water temporarily hard. Temporary hardness can be removed by boiling.

CaCO3(s) + H2O(l) + CO2(g) ==> Ca(HCO3)2(aq)

Ca(HCO3)2(aq) ==> CaCO3(s) + H2O(l) + CO2(g)

'Scum' formation.

Hard water contains dissolved compounds that react with soap to form scum. With soaps made

from the sodium salts of fatty acids, insoluble calcium or magnesium salts of the soap are formed.

This wastes soap.

CaSO4(aq) + 2C17H35COONa(aq) ==> (C17H35COO)2Ca(s for scum!) + Na2SO4(aq)

Hard water can be made soft by removing the dissolved calcium and magnesium ions.

1. If due to hydrogencarbonates it is removed by boiling

2. The addition of sodium carbonate (as 'washing soda' crystals), which dissolves

and precipitates out the calcium or magnesium ions as their insoluble carbonates

((s) formed).

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CaSO4(aq) + Na2CO3(aq) ==> CaCO3(s) + Na2SO4(aq)

3. Ion exchange polymer resin columns hold hydrogen ions or sodium ions. These

can be replaced by calcium and magnesium ions when hard water passes down

the column. The calcium or magnesium ions are held on the negatively charged

resin. The freed hydrogen or sodium ions do not form a scum with soap.

4. e.g. 2[resin]-H+(s) + Ca2+(aq) ==> [resin]-Ca2+[resin]-(s) + 2H+(aq)

Permanently hard water means the hardness cannot be removed by boiling eg when caused by

dissolved magnesium or calcium sulphate. Temporary hard water means it is softened by boiling

eg when caused by magnesium or calcium hydrogencarbonate.

EXERCISE

1) (a) What is an acid?

(b) What do you understand by pH of a solution?

(c)What is a strong acid? Give two examples of a strong acid.

(d) A form one student incorrectly wrote that a weak acid is a dilute acid.

Explain why this was incorrect.

(e) Write down an example of a weak acid, and write an equation to show how it ionizes in water.

(f) Suggest a test other than the ph test, that would be used to show that a given solution is an

acid, and give the results of the positive test.

2) In many countries supplies of clean water for drinking are obtained from river water.

(a) State two processes that are used to convert river water into water which is safe for humans

to drink.

(b) Safe drinking water may still contain dissolved compounds which make the water hard.

(i) Name a metallic element whose compounds cause hardness in water.

(ii) Suggest a reason why some natural water supplies are hard and others are not.

(iii) Describe how a soap solution can be used to find out whether a sample of water is

hard.

(iv) Some types of water are said to contain temporary hardness. Describe one way in

which temporary hardness may be removed from water.

(c) Some types of salt used to flavour food are mixtures of sodium chloride and potassium

chloride. Sodium chloride and potassium chloride are both ionic compounds.

Describe and explain the difference between a sodium atom and a sodium ion.

149

3) In many countries supplies of clean water for drinking are obtained from river water.

(a) State two processes that are used to convert river water into water which is safe for humans

to drink.

(b) Safe drinking water may still contain dissolved compounds which make the water hard.

(i) Name a metallic element whose compounds cause hardness in water.

(ii) Suggest a reason why some natural water supplies are hard and others are not.

(iii) Describe how a soap solution can be used to find out whether a sample of water is

hard.

(iv) Some types of water are said to contain temporary hardness. Describe one way in

which temporary hardness may be removed from water.

Energy changes in physical and chemical processes

Energy - Endothermic and Exothermic.

A chemical reaction always has a change in energy. In a reaction going from reactants to

products, either:

1) Heat is given out - called Exothermic

Or

2) Heat is taken in - called Endothermic.

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The large majority of reactions are exothermic, they give heat out.

Some physical processes are associated with a change in energy. Melting and boiling are

endothermic. Freezing and condensing are exothermic. Breaking bonds (overcoming the force of

attraction) requires energy, you have to put heat in - it is endothermic. This is why melting and

boiling are endothermic.

Making bonds gives out energy - it is exothermic. This is why freezing and condensing are

exothermic. In a chemical reaction you need to put energy in to break bonds in the reactants, you

get energy out when new bonds are formed to make the products.

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A table of bond enthalpies for some covalent bonds

Bond Bond energy ( kJ/mol)

H-H 436

C-H 435

O-H 464

C-C 497

C=C 347

O=O 497

C=O 803

Cl-Cl 242

H-Cl 431

If you get out more energy than you have to put in, then overall the reaction is exothermic. This is

what normally occurs.

Energy Level Diagrams.

The change in energy can be plotted against the progress of a reaction, as the reactants turn into

products.

Going from reactants to the top of the curve, you are going up the energy scale. Energy (heat) is

being put in to break bonds in the reactants.

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At the top of the curve, the bonds in the reactants have been broken.

The amount of energy put in to break these bonds is called the activation energy.

The activation energy is the minimum amount of energy needed for the reaction to occur. A

catalyst may work by lowering the activation energy for a reaction.

Going from the top of the curve to the products, you are going down the energy scale. Energy

(heat) is given out as bonds form in the products.

The reactants are higher up the energy scale than are the products. The amount of energy (heat)

you need to put in (the activation energy) is less than the amount of energy (heat) you get out.

This is a typical exothermic reaction.

The difference in energy levels between the reactants and the products is given the symbol H

(pronounced 'delta H'). This is the amount of heat given out (or taken in) during the reaction. For

an exothermic reaction, H is negative. For an endothermic reaction, H is positive.

Example. Calculate the enthalpy change for the reaction:

H2 (g) + Cl2 (g) 2HCl (g)

H-H + Cl-Cl H-Cl + H-Cl

Enthalpy of bonds broken( KJ) Enthalpy of bonds made( kJ)

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H-H 1 X 436 = 436 2 H-Cl = 2 x 431

Cl-Cl 1 x 242 = 242

Total = 678 Total = 862

Heat of reaction = energy difference

H = Energy supplied to break bonds – Energy released when bonds form

H = 678 – 862

H = - 184 kJ to form two moles of HCl

This is the balance of energy given out to the surroundings

Energy released on bond formation = 431 + 431 = 862 kJ given out

92 kJ released to form 1 mole of HCl formed)

H = -92 kJ/mol

If you have to put in more energy than you get out, then the reaction is endothermic, but except in

rare circumstances, the reaction doesn't happen.

Experiment

REACTIONS THAT GIVE OUT HEAT ENERGY

Be careful! The reaction is vigorous so do not do the experiment in a stoppered bottle!

(i) Put white anhydrous copper sulphate powder to a depth of about 1 cm in a test-tube. Hold a

thermometer with the bulb in the powder. Add water drop by drop. Record the changes of the

thermometer reading.

(ii) Put about 10 mL of strong aqueous copper sulphate solution into a wide test-tube or small

beaker. Support a thermometer with the bulb in the solution. Add magnesium powder, or ribbon, a

little at a time until the blue colour disappears. Note any changes in the thermometer reading.

(iii) To a little water in a wide test-tube, add concentrated sulphuric acid, drop by drop, down the

side of the tube. Stir gently with a thermometer after the addition of each drop. Note any changes

in the thermometer reading.

Experiment

HEAT OF A NEUTRALIZATION REACTION

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a. Dissolve 40 g of sodium hydroxide pellets in water and make up to 500 mL. This is a 2M

solution.

b. Prepare 500 mL of a 2M hydrochloric acid solution.

c. Leave the solutions to cool to room temperature. Note the actual temperature of the

solutions when cool.

d. Then add the acid to the base quite rapidly and stir with a thermometer. Note the

maximum temperature reached.

e. The increase of temperature should be about 13oC.

(iii) Since the volume of water has been doubled by adding one solution to the other, the

final solution contains 1 mole of OH- (aq) ions which reacted with 1 mole of H+(aq) ions to

form 1 mole of water molecules. We must assume that the specific heat of this

moderately weak solution is the same as that of water.

Examples of Endothermic and Exothermic Reactions.

Most reactions are exothermic. Examples are

1) Combustion of methane (natural gas).

2) Neutralisation.

3) Adding concentrated sulphuric acid to water is highly exothermic

(gives out a lot of heat).

4) Adding water to anhydrous copper (II) sulphate is exothermic. Anhydrous copper (II) sulphate

is white. Anhydrous means "without water". Anhydrous copper (II) sulphate is copper (II) sulphate

which is completely dry. When water is added to it, it turns into the familiar hydrated blue crystals.

Hydrated means 'with water'. This is used as a test for water.

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If blue (hydrated) copper (II) sulphate crystals are heated, an endothermic reaction occurs, they

lose their water, turn white and become anhydrous copper (II) sulphate crystals again. The

reaction is reversible. Blue copper (II) sulphate crystals will also turn white in the presence

of concentrated sulphuric acid.

Experiment

Reactions that take in heat energy (endothermic)

a. Put 10 mL of water in a test-tube. Read the temperature of the water.

b. Dissolve about 2 g of potassium nitrate in the water. Note the final temperature. The

temperature should fall through 90oC.

c. This means that, in the process of dissolving in the water, the particles have absorbed

energy. This energy has been taken from the surrounding water in the form of heat. A

similar result can be obtained by using potassium chloride instead.

Few reactions are endothermic. Examples are:

1) Photosynthesis. Chlorophyll is a very clever catalyst, which allows plants to make sugar from

carbon dioxide in the air. The energy needed for the reaction comes from sunlight.

2) Thermal Decomposition of compounds.

3) The reaction of ethanoic acid with sodium carbonate.

4) Dissolving some salts in water is an endothermic process. Potassium chloride and ammonium

nitrate both take in energy when they dissolve.

If you put a thermometer in these salts as they are dissolving, you will see that the temperature

drops.

5) Melting, boiling and evaporation are all endothermic processes (not reactions).

Definitions of heats of reaction

Standard conditions: This means the reactants/products start/finish at a specified temperature,

pressure and concentration whatever the 'temporary' temperature change in the reaction - which

is required to calculate the enthalpy change. The net energy change is based on the products

returning to the same temperature and pressure that the reactants started at. The most frequently

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used standard conditions are a temperature of 298 K/25oC (K = 273 + oC) and a pressure of 1

atm/101 kPa and a concentration of 1.00 mol dm-3. The use of standard conditions enables a

database of delta H change to be assembled from which you can do theoretical calculations.

1. Standard enthalpy of formation

1. Standard Enthalpy of Reaction ΔHr/react/reaction is the enthalpy change (heat absorbed/released,

endothermic/exothermic) when molar quantities of reactants as stated in an equation react under

standard conditions

Examples

NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(aq) (exothermic)

ΔHθr,298 = -57.1 kJ mol-1

2. Standard Enthalpy of Combustion ΔHc/comb/combustion is the enthalpy change when 1 mole of a

fuel (or any combustible material) is completely burned in oxygen (or air containing oxygen). You

should ensure just 1 mole of fuel appears in the equation to accompany the delta H value which

is always negative i.e. always exothermic.

Examples

C3H8(g) + 5O2(g) ==> 3CO2(g) + 4H2O(l) ΔHθc,298K(propane) = -2219 kJ mol-1

CH3COOH(l) + 2O2(g) ==> 2CO2(g) + 2H2O(l) ΔHθc,298K(ethanoic acid) = -876 kJ mol-1

3. Standard enthalpy of neutralisation

Standard enthalpy of neutralisation is the energy released when unit molar quantities of acids and

alkalis completely neutralise each other at 298K.

NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l) ΔHθneutralisation = -57.1 kJ mol-1

4. Enthalpy of fusion.

( Hfus) heat of fusion; molar heat of fusion; molar enthalpy of fusion. The change in enthalpy

when one mole of solid melts to form one mole of liquid. Enthalpies of fusion are always positive

because melting involves overcoming some of the intermolecular attractions in the solid.

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5. Enthalpy of vaporization. ( Hvap) heat of vaporization.

The change in enthalpy when one mole of liquid evaporates to form one mole of gas. Enthalpies

of vaporization are always positive because vaporization involves overcoming most of the

intermolecular attractions in the liquid.

Enthalpy of solution (DHsol)

When a solute is dissolved in a solvent a solution is formed. During dissolution of a solute in any

solvent, a certain amount of heat is either absorbed or evolved. Such heat changes under

constant pressure conditions are known as the enthalpy of solution. 'The change in enthalpy

when one mole of a solute is dissolved in a specified quantity of a solvent at a given temperature

is called enthalpy of solution'.

To avoid the amount of solvent, heat of solution is usually defined for an infinite dilute solution.

Thus, heat of solution at infinite dilution is the heat change when one mole of a substance is

dissolved in such a large quantity of solvent so that further dilution does not give any further heat

change.

For example, dissolution of sodium chloride

Here 'aq' represents aqueous meaning a large excess of water.

For substances, which dissolve with the absorption of heat (endothermic), the enthalpy of solution

is positive while for the substances which dissolve by liberating heat (exothermic), the enthalpy of

solution is negative.

For example, when KCl is dissolved in water, heat is absorbed. Thus, the enthalpy of solution of

KCl is positive. For a 200 times dilution

(Water: KCl = 200: 1), the enthalpy change during the process,

So, the enthalpy of solution of KCl at a dilution of 200 is 18.6 kJ mol-1. The dissolution of CaCl2(s)

in water is an exothermic process. So, the enthalpy of solution of calcium chloride (CaCl2) is

negative. At a dilution of 400, the enthalpy change for the reaction,

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So, the enthalpy of solution of CaCl2(s) at a dilution of 400 is -75.3 kJ mol-1.

RATE OF REACTION

Collision Theory.

A chemical reaction can only occur between particles when they collide (hit each other). Particles

may be atoms, ions or molecules. There is a minimum amount of energy

which colliding particles need in order to react with each other. If the colliding particles have less

than this minimum energy then they just bounce off each other and no reaction occurs. This

minimum energy is called the activation energy.

The faster the particles are going, the more energy they have. Fast moving particles are more

likely to react when they collide. You can make particles move more quickly by heating them up

(raising the temperature). The rate of a reaction may be measured by following the loss of a

reactant, or the formation of a product.

There are three reactions which may be studied to show how the rate can be changed. They are;

(1) The reaction between calcium carbonate and dilute hydrochloric acid.

(2) The reaction between sodium thiosulphate solution and hydrochloric acid.

(3) The decomposition of hydrogen peroxide solution.

Changing the Rate of a Reaction.

There are 5 ways to increase the rate of a chemical reaction. They are all understood in terms of

collision theory. The rate of a chemical reaction may be increased by

1) Raising the temperature.

2) Increasing the concentration (in solution).

3) Increasing the pressure (in gases).

4) Increasing the surface area of a solid.

5) Use a catalyst.

The opposite of 1, 2, 3 and 4 will decrease the rate of a reaction. A catalyst (strictly speaking) will

change the rate of a reaction. In practice a catalyst is mainly used to make a reaction go faster.

The graph shows how most of the reactions progress to completion.

159

The rate of reaction is fastest where the gradient (or slope) of the graph is highest and least

where the graph is least steep as shown. Where the graph is parallel to the time axis (or where

there is no change in amount of product) means the reaction is over.

How can a reaction end?

A reaction can terminate due to either of the following:

a. All the reactants are used up

b. One of the reactants is used up even if the other is in excess.

CONCENTRATION AND RATE OF REACTION

The reaction between sodium thiosulphate and hydrochloric acid can take a noticeable time.

Sulphur is produced during the reaction making the solution cloudy. The rate of reaction can be

found by finding the time taken to reach a certain degree of cloudiness in the solution. The

degree of cloudiness in this case may be defined as the point at which a black cross marked

below the reaction vessel can no longer be seen by looking through the solution from above. In

this experiment the concentration of sodium thiosulphate is made variable, whilst the

concentration of acid is kept constant.

Experiment

The reaction between sodium thiosulphate solution and dilute hydrochloric acid.

160

a. Make up 500 mL of aqueous solution containing 20 g sodium thiosulphate. 2 M

hydrochloric acid is also needed. Bench dilute acid is usually of this strength.

b. Using a measuring cylinder, put 50 mL of thiosulphate solution into a 100 mL beaker.

c. Place the beaker on a black cross marked on a sheet of paper.

d. Add 5 mL of the acid and note the time given by the second hand of a clock. Stir the acid

into the solution. Note the time when the cross is no longer visible through the sulphur in

the solution.

e. Repeat the experiment with a smaller concentration of thiosulphate. Take 40 mL of

thiosulphate solution and add 10 mL of distilled water. Stir and then add 5 mL of acid as

before. The time for the cross to become invisible should be greater than for the last

experiment.

f. Repeat the experiment using 30 mL, 20 mL and 10 mL of this sulphate mixed with 20 mL,

30 mL and 40 mL of distilled water.

g. Plot concentration of the thiosulphate solution against time taken for the reaction.

Concentration values may be taken as the volume of the original thiosulphate solution

used. Since I/time, reciprocal of time, is the measure of the rate of the reaction, plot

thiosulphate concentrations against I/time. The equation for the reaction can be written

as:

HCl(aq) + Na2S2O3(aq) --> NaCl(aq) + SO2(g) + S(s) + H2O(l)

How concentration increases rate of reaction

161

Increasing the concentration of a substance in solution means that there will be more particles

per dm3 of that substance. The more particles that there are in the same volume the closer to

each other the particles will be. This means that the particles collide more frequently with each

other and the rate of the reaction increases.

Consider the reaction between sodium thiosulphate solution and dilute hydrochloric acid above.

solid sulphur (S(s)) is formed in the flask. Increasing the concentration of sodium thiosulphate

means that the solid sulphur will be produced more quickly and there will be less time before the

cross can no longer be seen. When making different concentrations of sodium thiosulphate, take

care to use the same total volume of sodium thiosulphate plus hydrochloric acid for the

comparison to be fair.

HOW DOES TEMPERATURE AFFECT THE RATE OF A REACTION?

Experiment

Use the reaction in experiment above to investigate the effect of temperature.

a. Put 10 mL of sodium thiosulphate solution into the 100 mL beaker and stir in 40 mL of

water.

b. Use this concentration for the series of experiments with the temperature of the solution

as the variable.

c. Add 5 mL of acid as before and record the initial time and the temperature of the solution.

d. Record the final time when the black cross below the beaker is no longer visible. Repeat

the experiment, each time warming the thiosulphate solution to just over 30oC

Why increasing temperature increases rate of reaction.

Raising the temperature has the same effect on all reactions. Raising the temperature makes the

particles move faster. This means that the particles collide more frequently with each other and

the rate of the reaction increases.

Also, the faster the particles are traveling, the greater is the proportion of them which will have the

required activation energy for the reaction to occur.

Consider a reaction of marble chips with lime stone

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The variation rates of reaction with temperature can also investigated using marble chips and

dilute acid.

Apparatus

a. Conical flask

b. Electronic balances

c. Stop watch

d. Cotton

Chemicals

a. Dilute hydrochloric acid

b. Marble chips

Method;

a. Measure accurately 40cm3 of 2M hydrochloric acid and transfer it into a conical flask.

b. Weigh accutately 10g of marble chips

c. Place the conical flask and its contents onto an electronic balance and note the mass.

d. Carefully slide the marble chips into the acid and cover the mouth of the conical flask wit

cotton wool to prevent splashing. At the same time, start the stop watch.

163

e. Record the mass every 2 minutes in a suitable table

f. Plot a graph of loss in mass against time.

Analysis

In this graph, A is the curve for the rate at low temperature, while B is the graph for the reaction at

higher temperature. As a general guide, raising the temperature of a reaction by 10 °C will double

the rate of a reaction. The gradient of the plot will be twice as steep.

Experiment: Catalyst and rate of reaction

How does a catalyst affect the rate of a reaction?

Apparatus

a. Burette

b. Weighing balance

c. Flask

d. Stop watch

Chemicals

a. Hydrogen peroxide

b. Copper (II) oxide, nickel oxide, manganese (lV) oxide and zinc oxide.

164

Procedure

The variable in this reaction is the substance used as a catalyst in the decomposition of an

aqueous solution of hydrogen peroxide.

a. Set up the burette filled with water as in a standard water displacement experiment. 2 mL

of 20-volume hydrogen peroxide will give enough oxygen almost to fill the burette.

b. Weigh out 1 g each of copper (II) oxide, nickel oxide, manganese (lV) oxide and zinc

oxide.

c. Put 50 mL of water in the flask and add 2 mL of hydrogen peroxide solution. Add the 1 g

of copper oxide. Immediately insert the bung with the delivery tube into the flask.

d. Time the volume of oxygen given off at intervals of 15 seconds. Plot the volume of

oxygen produced every 15 seconds against the time of the reaction. Repeat the

experiment using the other oxides as catalysts. Plot a graph for each experiment.

e. Manganese (IV) oxide is usually used as a catalyst in this reaction. The catalyst is not

used up during the reaction.

What is a Catalyst?

A catalyst is a substance that will change the rate of a reaction. A catalyst is often used to make a

reaction go faster. The catalyst itself does not take part in the reaction as a reactant. It is not

changed by the reaction; it is not used up during the reaction. It is still there in the same form

when the reaction is complete.

A catalyst is usually a transition metal, a transition metal oxide (see uses of transition metals) or

an enzyme in living cells. An exception is aluminium oxide, used in the cracking of hydrocarbons.

A substance which works well as a catalyst for one reaction might not work well as a catalyst for a

different reaction.

How does a catalyst work?

A catalyst works by providing a convenient surface for the reaction to occur. The reacting

particles gather on the catalyst surface and:

1) collide more frequently with each other

2) more of the collisions result in a reaction between particles because the catalyst can lower the

activation energy for the reaction. A catalyst is often used as a powder, so that it has a bigger

surface area per gram.

Increasing the Pressure (in gases).

165

Increasing the pressure of a reaction where the reactant is a gas is similar to increasing the

concentration of a reactant in a solution. The gas particles (usually molecules) will be closer

together when the pressure increases. This means that the particles collide more frequently with

each other and the rate of the reaction increases. If the reaction is reversible, then increasing the

pressure will favour the side of the reaction which has the smaller volume.

Increasing the Surface Area of a Solid.

Experiment

SIZE OF PARTICLES AND RATE OF REACTION

Marble chips can be broken up with a hammer and graded into 3 or 4 sizes:

(a) Coarse powder;

(b) Pieces about half the size of a rice grain;

(c) Pieces as large as rice grains;

(d) The original lumps of marble chips.

a. Place four 100 X 16 mm test tubes in a stand. Weigh approximately 2 g of each grade,

size, of marble chips and put the four grades separately into each of the four tubes.

b. Obtain four balloons and blow them up several times to stretch them.

c. Put 5 mL of bench hydrochloric acid into each of the four balloons and slip the mouth of

the balloon over the top of the tube without letting any acid into the tube. When each

balloon is in place, tip the acid into each test-tube at the same time and observe which

balloon is the fastest and the slowest to be blown up.

d. The smallest should give the carbon dioxide in the shortest time. Instead of using marble

chips you can use granulated zinc.

e. Do not use metals in powder form because the reaction may be too vigorous and even

cause an explosion. Be careful! This reaction produces hydrogen gas!

f. Instead of collecting the gas in a balloon or plastic bag, a more accurate method would

be to collect the gas in a gas syringe and compare the volume of gas given off in unit

time for each grade of marble chips.

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g. Another accurate method is to stand a conical flask containing the marble chips and acid

on a balance and record the loss in mass every half minute. Carbon dioxide is a heavy

gas and most balances will enable the loss in mass to be found as the gas escapes.

A solid in a solution can only react when particles collide with the surface. The bigger the area of

the solid surface, the more particles can collide with it per second, and the faster the reaction rate

is. You can increase the surface area of a solid by breaking it up into smaller pieces. A powder

has the largest surface area and will have the fastest reaction rate. This is why catalysts are often

used as powders.

In the reaction between calcium carbonate and dilute hydrochloric acid

HCl (aq) + CaCO3(s) --> CaCl2 (aq) + CO2(g) + H2O(l)

Calcium carbonate may be used in the form of marble chips. The following apparatus may be

used to measure the loss in mass with time. Cotton wool is used to prevent loss of solid or liquid

due to splashing. It only allows the gas to escape.

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The reaction rates can be compared using large marble chips, and the same mass of small

marble chips. The reaction can be followed by plotting the loss of mass against time.

The reaction rate is faster (the slope is steeper) for the reaction with small marble chips (greater

surface area). Note that the final loss of mass is the same for both reactions. This is because the

same mass of calcium carbonate (marble chips) will give the same mass of carbon dioxide

whether the chips are large or small. The smaller chips will just do it more quickly.

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Does light affect the rate of a chemical reaction?

Chemical reactions often involve an energy change, and this is often shows itself as heat being

produced, but it may show itself in the production of light, as in the example of a flame. You

already know that the application of energy in the form of heat will often will often start or

accelerate a reaction, so it is natural to ask whether application of energy in the form of light will

do the same. For majority of reactions, the presence of light does not affect the rate at all, but

there are cases where light has considerable effect, and these reactions are very important.

Light energy (uv or visible radiation) can initiate or catalyse particular chemical reactions.

As well as acting as an electromagnetic wave, light can be considered as an energy 'bullets'

called photons and they have sufficient 'impact' to break chemical bonds, that is, enough energy

to overcome the activation energy.

The greater the intensity of light (visible or ultra-violet) the more reactant molecules are likely to

gain the energy react, so the reaction speed increases.

Examples:

Silver salts are converted to silver in the chemistry of photographic exposure of the film.

Silver chloride (AgCl), silver bromide (AgBr) and silver iodide (AgI) are all sensitive to light

('photosensitive'), and all three are used in the production of various types of photographic film to

detect visible light and beta and gamma radiation from radioactive materials.

Each silver halide salt has a different sensitivity to light. When radiation hits the film the silver ions

in the salt are reduced by electron gain to silver

Ag+ + e- ==> Ag (X = halogen atom, Cl, Br or I)

and the halide ion is oxidised to the halogen molecule by electron loss

2X- ==> X2 + 2e-

so overall the change via light energy is: 2AgX ==> 2Ag + X2

AgI is the least sensitive and used in X-ray radiography, AgCl is the most sensitive and used in

'fast' film for cameras.

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Photosynthesis in green plants:

The conversion of water + carbon dioxide ==> glucose + oxygen

6H2O(l) + 6CO2(g) ==> C6H12O6(aq) + 6O2(g)

requires the input of sunlight energy and the green chlorophyll molecules absorb the photon

energy packets of light and initiate the chemical changes summarised above.

EXERCISE

1) When an acid is added to bleach, chlorine gas is produced.

The diagram shows apparatus for carrying out this reaction and measuring the

volume of chlorine gas produced.

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gas collected

bleachand acid

The reaction was carried out using an excess of acid (10 cm3). The chlorine was

collected and its volume recorded every half minute. The results are shown in the

table.

time (minutes) 0.0 0.5 1.0 1.5 2.0 2.5

volume of chlorine (cm3)

0 10 17 22 26 29

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(a) (i) Draw a graph of these results on the grid.

40

30

20

10

0

volumeof chlorinein cm

0 1 2 3 4 5

time in minutes

3

(ii) The reaction had not finished after 2.5 minutes.

Explain how you know this is true.

(iii)The reaction finished in less than 5 minutes. Continue your graph to show this. Use your

graph to estimate the total time for the reaction.

(b) The reaction was repeated under the same conditions but using 20 cm3 instead of 10

cm3 of acid.

Explain how, if at all, the rate of reaction would change compared to the first experiment.

2) Hydrochloric acid and magnesium ribbon react to produce hydrogen gas.

2HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)

You are to plan the details of an experiment, based on the volume of gas produced in

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the reaction, to investigate how the rate of reaction depends on the concentration of the

hydrochloric acid.

(a) Using scientific knowledge linking molecular or ionic collisions to rates of reaction, predict the

relationship between

(i) the rate of formation of hydrogen gas and the concentration of the hydrochloric acid

(ii) the rate of formation of hydrogen gas and the temperature of the reaction.

(b) In an experiment to determine the rate of reaction with respect to HCl identify the independent

variable.

(c) Identify one variable, other than temperature, that must be controlled in the experiment.

3. Hydrogen reacts with oxygen according to the equation given below

H2(g) + ½ O2(g) → H2O(l) ΔH° = –286 kJ mol–1

Calculate the volume of hydrogen that would produce 7.15 KJ of enrgy when completely burnt

(Molar volume of gas = 24 litres

4. Graphite burns in oxygen according to the equation

C (graphite) + O2 (g) → CO2 (g) ΔH° = –394 kJ mol–1

The enthalpy change when 3 grams of graphite burn is -98.5KJ

Calculate the standard enthalpy (ΔH°) of combustion of graphite under standard condition.

REVERSIBLE REACTIONS.

When a reaction is reversible, it means that it can go both forwards and backwards. The forward

reaction is the one we want, in which reactants are converted into products.

A reversible reaction is shown by the sign , a half-arrow to the right (forward reaction),

and, a half-arrow to the left (backward reaction).

Most reactions are not reversible (irreversible) and have the usual complete arrow

only pointing to the right.

The backward reaction is where the products become changed back into the original reactants.

Reactions in both directions occur at the same time. In a closed system, after awhile an

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equilibrium mixture is reached, where a certain proportion of the mixture exists as reactants and

the rest exists as products. A closed system means that none of the products or reactants can

escape to the outside environment. It is like being inside a bottle with a cork in the top.

When equilibrium has been reached, it does not mean that the reactions have stopped. It means

that the forward reaction is making products at the same rate that the backward reaction is

making reactants. It is said to be a dynamic equilibrium. Dynamic means moving or changing, to

remind you that the reaction has not stopped.

The thermal decomposition of ammonium chloride

On heating strongly above 340oC, the white solid ammonium chloride thermally decomposes into

a mixture of two colourless gases ammonia and hydrogen chloride. On cooling the reaction is

reversed and solid ammonium chloride reforms.

This is an example of sublimation but here it involves both physical and chemical changes. When

a substance sublimes it changes directly from a solid into a gas without melting and on cooling

reforms the solid without condensing to form a liquid.

Ammonium chloride + heat ammonia + hydrogen chloride

NH4Cl(s) NH3(g) + HCl(g)

Note:

Reversing the reaction conditions reverses the direction of chemical change,

typical of a reversible reaction.

Thermal decomposition means using 'heat' to 'break down' a molecule into

smaller ones.

The decomposition is endothermic (heat absorbed or heat taken in) and the

formation of ammonium chloride is exothermic (heat released or heat given out).

This means if the direction of chemical change is reversed, the energy change is

also reversed.

Similarly, ammonium sulphate also sublimes when heated above 235oC and thermally

decomposes into ammonia gas and sulphuric acid vapour.

(NH4)2SO4(s) NH3(g) + H2SO4(g)

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Amount of Product from Reversible Reactions.

The amount of product (called the yield) which you get from a reversible reaction depends on the

conditions (use of temperature, pressure and catalyst). This is a useful summary for any

reversible reaction including the Contact Process and the Haber process.

1. Increasing the temperature favours the endothermic reaction.

2. Increasing the pressure favours the smaller volume.

3. Using a catalyst gives the equilibrium conditions more quickly.

Le Chatelier's Principle.

For a reversible reaction, Le Chatelier's Principle states that

"The equilibrium position will respond to oppose a change in the reaction conditions".

What this means in practice is;

1) If you remove a product, the equilibrium mixture changes to make more products.

It tries to get back to the composition it had before the product was removed. You can carry on

removing product until all the reactants have turned into product (useful!).

The reverse of 1 is also true. If you remove a reactant, the equilibrium changes to make more

reactant.

2) Heat may be treated as a reactant (for an endothermic reaction) or as a product (for an

exothermic reaction). If you remove heat from an exothermic reaction (cool it down), the

equilibrium will change to produce more products. This will not only produce more heat, but also

produce more of the chemical product that you want in the equilibrium mixture. If you add heat to

an exothermic reaction (raise its temperature), the reverse will happen, and you will get less

product in the equilibrium mixture. The opposite of the above is true for an endothermic reaction.

3) For a reversible reaction involving gases, increasing the pressure will change the equilibrium to

make more of the side of the reaction which has the smaller volume.

Decreasing the pressure will change the equilibrium to make more of the side of the reaction

which has the larger volume.

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Reversible reactions and Equilibrium

When a reversible reaction occurs in closed system equilibrium is formed, in which the original

reactants and products formed coexist.

In equilibrium, there is a state of balance between the concentrations of the reactants and

products and once a state of chemical equilibrium is reached there is no further change in

concentrations BUT the reactions don't stop!

At equilibrium the rate at which the reactants reactants change into products is exactly equal to

the rate at which the products change back to the original reactants.

However the final relative equilibrium amounts of the reactants and products depend on the

reaction conditions e.g. the temperature and pressure.

For industrial processes, it is important to maximise the concentration of the desired products and

minimise the 'leftover' reactants. A set of rules can be used to predict the best reaction conditions

to give the highest possible yield of product.

Rule 1a: If the forward reaction forming the product is endothermic, raising the temperature

favours its formation increasing the yield of product (lowering the temperature decreases the

yield).

Rule 1b: If the forward reaction forming the product is exothermic, decreasing the temperature

favours its formation (increasing temperature decreases the yield).

Rule 2a: Increasing the pressure favours the side of the equilibrium with the least number of

gaseous molecules as shown by the balanced symbol equation.

Rule 2b: Decreasing the pressure favours the side of the equilibrium with the most number of

gaseous molecules as shown by the balanced symbol equation.

Rules 1 above, and rule 3, below, apply to any reaction, BUT rule 2 above, ONLY applies to a

reaction with one or gaseous reactants or products.

Rule 3a: If the concentration of a reactant (on the left) is increased, then some of it must change

to the products (on the right) to maintain a balanced equilibrium position.

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Rule 3b: If the concentration of a reactant (on the left) is decreased, then some of the products

(on the right) must change back to reactants to maintain a balanced equilibrium position.

e.g. nitrogen + hydrogen ammonia

or N2(g) + 3H2(g) 2NH3(g)

If the nitrogen or hydrogen concentration was increased, some of this extra gas would change to

ammonia.

If the nitrogen or hydrogen concentration was decreased, some of ammonia would change back

to nitrogen and hydrogen.

So in terms of enforced change ==> system response:

a. Increasing nitrogen ==> decreases hydrogen and increases

ammonia.

b. Increasing hydrogen ==> decreases nitrogen and increases

ammonia.

c. Increasing ammonia ==> increases nitrogen and hydrogen.

d. Decreasing ammonia ==> decreases nitrogen and hydrogen.

e. Decreasing nitrogen ==> increases hydrogen and decreases

ammonia.

f. Decreasing hydrogen ==> increases nitrogen and decreases

ammonia.

Rule 4: A catalyst does NOT affect the position of equilibrium, you just get there faster! A catalyst

usually speeds up both the forward and reverse reaction but there is no way it can influence the

final 'balanced' concentrations. However, the importance of a catalyst lies with economics e.g.

(i) Bringing about reactions with high activation energies at lower temperatures and so

saving energy.

(ii) Saving time is saving money.

EXERCISE

1) A dynamic equilibrium is established when hydrogen and carbon dioxide

React as shown below.

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H2(g) + CO2(g) H2O(g) + CO(g)

(a) State three features of a dynamic equilibrium

(b) The equilibrium yield of carbon monoxide increases as the temperature is increased.

Determine whether the reaction as written is exothermic or endothermic, giving a reason

for your answer.

(c)Explain the effect of an increase in pressure on:

(i) the equilibrium position

(ii) the reaction rate

Give your reasoning in each case

2) In the Ostwald process for the production of nitric acid, ammonia is oxidised at 900 °C

over a platinum/ rhodium alloy catalyst according to the equation:

4NH3 (g) + 5O2 (g) 4NO (g) + 6H2O (g)

The reaction is very exothermic.

(a) (i) Why do the concentrations of the substances in an equilibrium mixture

remain constant?

(ii) State with a reason the effect of an increase in pressure on this

equilibrium system.

(b) The mixture obtained from the catalyst chamber contains excess oxygen.

(i) Write the equation for the further reaction that occurs on cooling this

mixture.

(ii) Show, by means of an equation, how the product in (b) (i) is used to

make nitric acid.

METALS

Introduction

The Earth's crust contains many different rocks. Rocks are a mixture of minerals and from some

we can make useful substances. A mineral can be a solid metallic or non-metallic element or a

compound found naturally in the Earth's crust.

Table 3.0: Elements in the earths crust

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Element % composition in

Earths crust

Oxygen 46

Silicon 28

Aluminium 8

Iron 6

Calcium 4

Sodium 2.4

Magnesium 2.3

Potassium 2.1

Titanium 0.6

Hydrogen 0.14

A metal ore is a mineral or mixture of minerals from which economically viable amounts of metal

can be extracted. Ores are often oxides, carbonates or sulphides.

Metals and their ores

Metal Name of the ore Compound present

Aluminium Bauxite Aluminium oxide

Al2O3

Copper Copper pyrates Copper iron sulphide

CuFeS2

Iron Haematite Iron III Oxide Fe2O3

Sodium Rock salt Sodium chloride Nacl

Zinc Zinc-blende Zinc Sulphide (Zns)

Lead Galena Lead II Sulphide

Pbs

In order to extract a metal, the ore or compound of the metal must undergo a process called

reduction to free the metal (i.e. the positive metal ion gains negative electrons to form the neutral

metal atom, or the oxide loses oxygen, to form the free metallic atoms).

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The chemical that removes the oxygen from an oxide is called the reducing agent i.e. carbon,

carbon monoxide or sometimes hydrogen. Generally speaking the method of extraction depends

on the metals position in the reactivity series.

The reactivity series of metals can be presented to include two non-metals, carbon and hydrogen,

to help predict which method could be used to extract the metal.

Reactivity series Reaction with

Air Water Dilute acid

Sodium

Calcium

Magnesium

Burn very strongly in

air to form oxide

React with cold water

to evolve hydrogen

React very strong to

evolve hydrogen

Aluminium

Zinc

Iron

Burn less strongly in

air, to form oxides

React with steam

when heated to give

hydrogen

React less strongly to

give hydrogen

Lead

Copper

React slowly to form

an oxide layer

Do not react Do not react

Silver

Gold

Do not react Do not react Do not react

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RULE: Any element higher in the series can displace any other lower element from its solution in

water.

NOTES

1. Metals above zinc and carbon in the reactivity series cannot usually be extracted with carbon

or carbon monoxide. They are usually extracted by electrolysis of the purified molten ore or other

suitable compound e.g. aluminium from molten aluminium oxide or sodium from molten sodium

chloride. The ore or compound must be molten or dissolved in a solution in an electrolysis cell to

allow free movement of ions (electrical current).

2. Metals below carbon can be extracted by heating the oxide with carbon or carbon monoxide.

The non-metallic elements carbon will displace the less reactive metals in a smelter or blast

furnace e.g. iron or zinc and metals lower in the series.

Metals below hydrogen will not displace hydrogen from acids. Their oxides are easily reduced to

the metal by heating in a stream of hydrogen, though this is an extraction method rarely used in

industry. In fact most metal oxides below carbon can be reduced when heated in hydrogen, even

if the metal reacts with acid.

Some metals are so unreactive that they do not readily combine with oxygen in the air or any

other element present in the Earth's crust, and so can be found as the metal itself. For example

gold (and sometimes copper and silver) and no chemical separation or extraction is needed. In

fact all the metals below hydrogen can be found as the 'free' or 'native' element.

Other methods are used in special cases using the displacement rule. A more reactive metal can

be used to displace and extract a less reactive metal but these are costly processes since the

more reactive metal also has to be produced in the first place!

Sometimes electrolysis is used to purify less reactive metals which have previously been

extracted using carbon or hydrogen (e.g. copper and zinc). Electrolysis is also used to plate one

metal with another.

The demand for raw materials does have social, economic and environmental implications e.g.

conservation of mineral resources by recycling metals, minimising pollution etc.

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Historically as technology and science have developed the methods of extraction have improved

to the point were all metals can be produced. The reactivity is a measure of the ease of

compound formation and stability (i.e. more reactive, more readily formed stable compound, more

difficult to reduce to the metal). The least reactive metals such as gold, silver and copper have

been used for the past 10000 years because the pure metal was found naturally.

Moderately reactive metals like iron and tin have been extracted using carbon based smelting for

the past 2000-3000 years. But it is only in the last 200 years that very reactive metals like sodium

or aluminium have been extracted by electrolysis.

Methods of extraction of metals and reactivity series.

Reactivity Metal Method of extraction

Decreasing

reactivity

Potassium

Sodium

Calcium

Magnesium

Aluminium

Electrolysis of molten salt

Zinc

Iron

Tin

Lead

Reduction of oxides with carbon or carbon monoxide.

Sulphide ores are heated to produce oxides.

Copper

Silver

Gold

These occur uncombined in the ground.

Physical properties of metals

1. Metals are malleable: Metals can be hammered or beaten into thin sheets without breaking.

Malleable means that metallic bonds in the metals do not break easily. Gold (Au), silver (Ag) are

highly malleable elements. Metals can be made into thin foils because they are malleable.

2. Metals are ductile: Metals can be melted and drawn into thin wires. Because of this property,

metals are known as ductile. The ductility property follows from the malleability property. While

being drawn into wires, metals are stretched. Because of the strong metallic bonds, the metal

atoms do not separate easily. Silver (Ag), copper (Cu), aluminum (Al) are very ductile, very thin

wires can be made out of these elements.

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3. Metals are good conductors of heat and electricity: In metals, the bonds are formed by excess

or free electrons moving around large array of atoms. These electrons are able to conduct

electricity and heat. Silver (Ag), copper (Cu), aluminum (Al) are good conductors of heat and

electricity. Amongst metals, lead (Pb) is a poor conductor of electricity.

4. Metals have luster: Metals generally can be highly polished. The electronic structure of metals

is such that the metals are able to reflect incident light. This gives metal the characteristic metallic

lustre. Gold (Au), silver (Ag), copper (Cu), platinum (Pt) can be polished and give off a good

lustre.

5. Metals are highly tensile: Due to their ductility and malleability properties, metals are very

strong. Their bonds do not break easily as the electrons are shared over an array of metal atoms.

This gives metals a very high tensile strength, and metals do not break easily.

6. Physical state: Metals are solids at room temperature. Mercury (Hg) is the only exception - it is

liquid at room temperature. Density of metals is high. Only sodium (Na) and potassium (K) have

densities less than water. Thus all metals are hard materials except sodium and potassium which

are soft metals. Lead is also considered to be a soft metal.

7. Brittleness: Metals are not brittle. But Zinc (Zn) is an exception. Metals do not break easily

because of their metallic bonds.

8. Melting and boiling points: All metals, other than sodium (Na) and potassium (K), have high

melting and boiling points. The melting points of sodium (Na) and potassium (K) are below 100C.

The melting point of iron is about 1540C.

9. Solubility of metals: Pure metals are insoluble in solvents like water or in any organic solvent.

Metals can be dissolved only in acids.

10. Metals are sonorous: Metals make a characteristic sound when hit with an object. Thus

metals are sonorous. The sonorousness of metals depends on the temperature and density.

METALS AND NON-METALS

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Metals Non-metals

They are usually solids (except for mercury,

which is a liquid) at room temperature. Their

melting and boiling points are usually high.

They are solid s or gases (except for bromine,

which is a liquid) at room temperature. Their

melting and boiling points are often low.

They are usually hard and dense

Most non-metals are softer than metals (but

diamond is very hard). Their densities are often

low.

All metals are good conductors of electricity

They are poor conductors of electricity (except

graphite, a form of carbon); they tend to be

insulators.

They are good conductors of heat.

They are poor thermal conductors

Their shape can be changed by hammering

(they are malleable). They can also be pulled

out into wires (they are ductile).

Most non-metals are brittle when solid.

They are grey in colour (except gold and

copper). They can be polished.

They vary in colour. They often have a dull

surface when solid.

They make a ringing sound when struck (they

are sonorous)

They are not sonorous

ALUMINIUM

The Extraction of Aluminium

Aluminium ore is called bauxite. Bauxite contains aluminium oxide, water, iron oxide and other

impurities. The purified dry ore, called alumina, is aluminium oxide - Al2O3.The alumina must be

molten for electrolysis to work, since the ions are not free to move in the solid state.

Unfortunately, alumina has a high melting point (2040 °C) and it is not practical to do electrolysis

at such a high temperature. In the middle of the nineteenth century, it was found that alumina

dissolved in cryolite.

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Cryolite is sodium aluminium fluoride - Na3AlF6.A solution of alumina in cryolite melts at about

900 °C and electrolysis is done at about 950 °C. This reduces the cost of aluminium and saves

on energy.

Electrolysis Cell.

The steel container is coated with carbon (graphite) and this is used as the negative electrode

(cathode). Aluminium oxide (Al2O3) is an ionic compound. When it is melted the Al3+ and O2- ions

are free to move and conduct electricity. Electrolysis of the alumina/cryolite solution gives

aluminium at the cathode and oxygen at the anode.

4Al3+ + 12e- 4Al (reduction).

6O2- 3O2 + 12e- (oxidation)

Aluminium is denser than the alumina/cryolite solution and so it falls to the bottom of the cell

where it can be tapped off as pure liquid metal.

The overall reaction is:

2Al2O3 (l) 4Al (l) + 3O2 (g)

Oxygen is given off at the positive carbon anode. Carbon dioxide is also given off at the carbon

anode because hot oxygen reacts with the carbon anode to form carbon dioxide gas.

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C(s) + O2 (g) CO2 (g)

The carbon anodes slowly disappear because each molecule of carbon dioxide which is given off

takes a little piece of carbon away with it. The carbon anodes need to be replaced when they

become too small.

Recycling.

Aluminium is the third most abundant element in the Earth's crust after oxygen and silicon. It is

the most abundant metal but it is more expensive to produce than iron because of the cost of the

large amounts of electricity, which are needed for electrolysis. Extracting aluminium from recycled

scrap requires only about 5% of the energy needed to extract aluminium from its ore.

Resistance to Corrosion.

Corrosion is the name given to the process where a metal reacts with oxygen in the air to form a

metal oxide. Corrosion is the reverse of the process used to extract the metal. If the metal is iron,

corrosion is called rusting.

A rusted ship

186

In general, the more reactive a metal is the more quickly it will corrode. Aluminium is a reactive

metal. It is above carbon in the reactivity series but is resistant to corrosion because aluminium

reacts with oxygen in the air and forms a thin layer of aluminium oxide (Al2O3).

The oxide layer covers the surface of the aluminium metal and prevents any further reaction

(corrosion) from happening. Anodizing can increase the thickness of the oxide layer. The oxide

layer will react with an acid or alkali to form an aluminium salt + water. The exposed aluminium

metal will then itself react with acid or alkali to form the aluminium salt + hydrogen. It is therefore

best to avoid contact of aluminium cookware with citrus fruits or vinegar (acids), or alkali cleaners

such as caustic soda (sodium hydroxide).

Properties and Uses of Aluminium.

Properties.

Aluminium

1) is strong, malleable and has a low density.

2) Is resistant to corrosion.

3) Is a good conductor of heat and electricity?

4) Can be polished to give a highly reflective surface.

Uses of aluminium based on properties

Property Use

Low density and strength Construction of aircraft, lightweight vehicles

Easy shaping and corrosion resistance Drink cans and roofing materials

Corrosion resistance and low density Greenhouses and window frames

Good conduction of heat Boilers, cookers and cookware

Good conduction of electricity Overhead power cables hung from pylons

High reflectivity Reflectors and heat resistant clothing for fire

fighting

187

EXTRACTION OF IRON

Chemistry of the blast furnace.

Hot air is blasted into the furnace causing coke (carbon) to burn rapidly and raise the temperature

to 2000 °C.

188

Sequence of reaction processes

1. Production of heat

Carbon burns in plenty of oxygen producing heat and carbon dioxide

C(s) + O2 (g) CO2 (g)

2. Production of carbon monoxide

The carbon dioxide then reacts with hot carbon to form carbon monoxide.

CO2 (g) + C(s) 2CO (g)

3. Reduction of iron (III) oxide to iron metal

Carbon monoxide then reduces iron in the ore to iron metal.

3CO(g) + Fe2O3(s) 3CO2(g) + 2Fe(l)

189

The temperature where the reduction takes place is above 1500 °C. Iron falls to the bottom of the

furnace where the temperature is 2000 °C. Iron is liquid at this temperature and is tapped off

periodically.

4. Removal of impurities

Limestone is calcium carbonate (CaCO3) and it is added to the blast furnace to remove the

impurities in the iron ore. Calcium carbonate is decomposed by heat in the furnace

to give calcium oxide (quicklime) and carbon dioxide. This is called thermal decomposition.

CaCO3(s) CaO(s) + CO2(g)

The main impurity is silica (sand or rock) which is silicon dioxide. Silicon dioxide is solid at the

furnace temperature and the furnace would become blocked if it was not removed.

Silicon dioxide (acidic) reacts with calcium oxide (basic) to form the salt calcium silicate (called

slag) which is liquid in the furnace. Slag flows to the bottom of the furnace where it floats on the

liquid iron and is easily removed.

CaO(s) + SiO2(s) CaSiO3 (l)

The slag (CaSiO3) is allowed to cool until it becomes a solid and is used for road construction.

THE PRODUCT IRON

Iron from the blast furnace contains about 5% carbon, which comes from the coke in the furnace.

It is cast into moulds called pigs and the iron is called cast iron or pig iron.

Molten scrap iron and molten scrap steel are mixed with molten iron from the blast furnace in a

converter. This is the stage when scrap iron and steel are recycled. Carbon is removed from the

mixture by bubbling pure oxygen through it. The oxygen reacts with the carbon to form carbon

dioxide. Other non-metals in the mixture react with the oxygen to form acidic oxides. Calcium

carbonate is then added to remove the acidic oxides. These reactions produce pure iron, which is

called wrought iron. The large majority of iron from the blast furnace is made into steel for

construction, steel contains 0·1% to 1·5% carbon.

PROPERTIES AND USES OF IRON.

Iron is one of the three magnetic elements (the others are cobalt and nickel). Cast iron is very

brittle (it cracks easily) but it has a greater resistance to corrosion (see rusting) than either pure

190

iron or steel. Cast iron is used for manhole covers on roads and pavements

and as engine blocks for petrol and diesel engines. Pure iron is called wrought iron.

Wrought iron is malleable and is mainly used in ornamental work for gates.

Iron is also the catalyst in the Haber process. The large majority of iron from the blast furnace is

made into steel.

Metal Carbon content Properties Uses

Mild steel Less than 0.25% Easily worked; not

brittle

Car bodies

High carbon steel 0.45-1.5% Hard and brittle Cutting tools, razor

blades, chisels

Cast iron 2.5-4.5% Cheaper than steel,

easily moulded

Gear boxes, engine

blocks, brake discs

Stainless steel Iron 74% Tough, does not

rust

Cutlery, surgical

instruments,

kitchen sinks

Chromium 18%

Nickel 8%

Common alloy of steel is stainless steel. It does not rust. It is made by adding small amounts of

chromium and nickel to pure iron.

RUSTING OF IRON

When a metal is attacked by air, water or other surrounding substances, it is said to corrode. In

the case of iron and steel, it is said to rust. Rust is a red-brown powder consisting of hydrated iron

III oxide (FeO3.xH2O). Rusting process is increased by the presence of salt. This explains why

ships get rusty quickly. Stainless steel does not rust does not rust because of the presence of

chromium which is protected by an oxide layer just as aluminium.

RUST PREVENTION

Painting

This method is widespread for objects ranging from ships and bridges and gates. Paint only

protects the metal as long as the paint layer is unscratched.

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Oiling and greasing

The oil and greasing of moving parts of machinery forms a protective film, preventing rusting. The

treatment must be repeated to continue the protection.

Plastic coating

These are used to form a protective layer on such items as refrigerators, chairs bicycle handles

etc

Electroplating

A piece of iron can be electroplated with chromium or tin to prevent it from rusting. Tin is used as

it is unreactive and non-toxic. However if the layer of coating is scratched, the iron below will still

rust

Galvanising

A piece of iron can be coated with a more reactive metal such as zinc. This is called galvanizing.

The advantage with galvanizing is that the protection still works even when the Zinc layer is

scratched. Bodies of cars are protected from rusting by electroplating them with zinc.

Sacrificial protection

This is a form of protection in which blocks of reactive metal are attached to the iron surface. This

is a common way of protecting underground pipes. The reactive metal corrodes in preference to

iron.

Reaction of iron with air

Iron metal reacts in moist air by oxidation to give a hydrated iron oxide. This does not protect the

iron surface to further reaction since it flakes off, exposing more iron metal to oxidation. This

process is called rusting and is familiar to any car owner. Finely divided iron powder is pyrophoric,

making it a fire risk.

On heating with oxygen, O2, the result is formation of the iron oxides Fe2O3 and Fe3O4.

4Fe(s) + 3O2(g) → 2Fe2O3(s)

3Fe(s) + 2O2(g) → Fe3O4(s)

192

Reaction of iron with water

Air-free water has little effect upon iron metal. However, iron metal reacts in moist air by oxidation

to give a hydrated iron oxide. This does not protect the iron surface to further reaction since it

flakes off, exposing more iron metal to oxidation. This process is called rusting and is familiar to

any car owner.

Reaction of iron with the halogens

Iron reacts with excess of the halogens F2, Cl2, and Br2, to form ferric, that is, Fe(III), halides.

2Fe(s) + 3F2(g) → 2FeF3(s) (white)

2Fe(s) + 3Cl2(g) → 2FeCl3(s) (dark brown)

2Fe(s) + 3Br2(l) → 2FeBr3(s) (reddish brown)

This reaction is not very successful for iodine because of thermodynamic problems. The iron(III)

is too oxidizing and the iodide is too reducing. The direct reaction between iron metal and iodine

can be used to prepare iron (II) iodide, FeI2.

Fe(s) + I2(s) → FeI2(s) (grey)

Reaction of iron with acids

Iron metal dissolves readily in dilute sulphuric acid in the absence of oxygen to form solutions

containing the aquated Fe (II) ion together with hydrogen gas, H2. In practice, the Fe (II) is

present as the complex ion [Fe (OH2)6]2+.

Fe(s) + H2SO4(aq) → Fe2+(aq) + SO42-(aq) + H2(g)

If oxygen is present, some of the Fe(II) oxidizes to Fe(III).

Reaction of nitric acid with Iron

With cold and dilute nitric acid:

Iron reacts with cold and dilute nitric acid to give ferrous nitrate, water and nitric oxide.

With concentrated nitric acid (hot or cold):

193

Iron reacts with cold or hot concentrated nitric acid to give ferric nitrate, water and nitrogen

dioxide.

With highly concentrated or fuming nitric acid:

When iron reacts with highly concentrated or fuming nitric acid, a film of an insoluble oxide is

formed on it. This renders the iron passive.

CHEMICAL PROPERTIES

Aluminium

The surface of aluminium goes white when strongly heated in air/oxygen to form white solid

aluminium oxide. Theoretically its quite a reactive metal but an oxide layer is readily formed even

at room temperature and this has quite an inhibiting effect on its reactivity. Even when scratched,

the oxide layer rapidly reforms, which is why it appears to be less reactive than its position in the

reactivity series of metals would predict but the oxide layer is so thin it is transparent, so

aluminium surfaces look metallic and not a white matt surface.

Aluminium + oxygen ==> aluminium oxide

4Al(s) + 3O2 (g) ==> 2Al2O3(s)

Under 'normal circumstances' in the school laboratory aluminium has virtually no reaction with

water, not even when heated in steam due to a protective aluminium oxide layer of Al2O3.

Slow reaction with dilute hydrochloric acid to form the colourless soluble salt aluminium chloride

and hydrogen gas.

Aluminium + hydrochloric acid ==> aluminium chloride + hydrogen

2Al(s) + 6HCl (aq) ==> 2AlCl3 (aq) + 3H2 (g)

The reaction with dilute sulphuric acid is very slow to form colourless aluminium sulphate and

hydrogen.

194

Aluminium + sulphuric acid ==> aluminium sulphate + hydrogen

2Al(s) + 3H2SO4 (aq) ==> Al2 (SO4)3(aq) + 3H2 (g)

Reaction of aluminium with air

Aluminium is a silvery white metal. The surface of aluminium metal is covered with a thin layer of

oxide that helps protect the metal from attack by air. So, normally, aulumium metal does not react

with air. If the oxide layer is damaged, the aluminium metal is exposed to attack. Aluminium will

burn in oxygen with a brilliant white flame to form the trioxide aluminum (III) oxide, Al2O3.

4Al(s) + 3O2 (l) → 2Al2O3(s)

Reaction of aluminium with water

Aluminium is a silvery white metal. The surface of aluminium metal is covered with a thin layer of

oxide that helps protect the metal from attack by air. So, normally, aluminum metal does not react

with air. If the oxide layer is damaged, the aluminium metal is exposed to attack, even by water.

Reaction of aluminium with the halogens

Aluminium metal reacts vigorously with all the halogens to form aluminium halides. So, it reacts

with chlorine, Cl2, bromine, I2, and iodine, I2, to form respectively aluminium (III) chloride, AlCl3,

aluminium (III) bromide, AlBr3, and aluminium (III) iodide, AlI3.

2Al(s) + 3Cl2(l) → 2AlCl3(s)

2Al(s) + 3Br2 (l) → Al2Br6(s)

2Al(s) + 3I2 (l) → Al2I6(s)

Reaction of aluminium with acids

Aluminium metal dissolves readily in dilute sulphuric acid to form solutions containing the aquated

Al (III) ion together with hydrogen gas, H2. The corresponding reactions with dilute hydrochloric

acid also give the aquated Al (III) ion. Concentrated nitric acid passivates aluminium metal.

2Al(s) + 3H2SO4(aq) → 2Al3+(aq) + 2SO42-(aq) + 3H2(g)

2Al(s) + 6HCl(aq) → 2Al3+(aq) + 6Cl-(aq) + 3H2(g)

195

Exercise

1) Aluminium, iron, sodium and chlorine are important elements produced by the chemical

industry.

(a) State which of the elements above

(i) has atoms which are converted into ions by gaining an electron,

(ii) has atoms which contain 3 electrons in their outer shells.

(b) When chlorine gas is bubbled into a colourless solution of sodium bromide, the solution turns

orange.

Explain this observation.

(c) Fig. 7.1 shows a blast furnace which is used to convert iron(III) oxide into iron.

The balanced equations of the three main chemical reactions in the blast furnace are

shown in Fig. 7.1. Each reaction is a redox reaction.

(i) State two substances, shown in Fig. 7.1, which are reduced. Explain your answer briefly.

(ii) Use the relative atomic masses shown on the Periodic Table to calculate the

relative formula mass of iron(III) oxide. Show your working.

2) Aluminium is extracted by the electrolysis of a molten mixture that contains alumina, which

is aluminium oxide, Al2O3.

196

(a) The ore of aluminium is bauxite. This contains alumina, which is amphoteric, and iron(III)

oxide, which is basic. The ore is heated with aqueous sodium hydroxide.

Complete the following sentences.

The…………………. dissolves to give a solution of…………………………………………..

The………………… does not dissolve and can be removed by………………………[4]

(b) Complete the labelling of the diagram.

(c) The

ions that are involved in the electrolysis are Al3+ and O2-.

(i) Write an equation for the reaction at the cathode.

(ii) Explain how carbon dioxide is formed at the anode.

(d) Give an explanation for each of the following.

(i) Aluminium is used extensively in the manufacture of aircraft.

(ii) Aluminium is used to make food containers.

197

(iii) Aluminium electricity cables have a steel core.

3) Aluminium is obtained from the ore bauxite. The first stage is the purification of the ore

followed by electrolysis of the molten ore.

(i) Identify one major impurity in bauxite

(ii) Why must the ore be molten before electrolysis is done?

(iii) Explain why the anode had to be replaced at regular intervals.

ANSWERS

Gas Laws

1) (a) The volume of a fixed mass of gas is directly proportional to its absolute temperature if the

pressure is kept constant.

(b)

(c) This temperature is called Absolute Zero.

At -273.16 0C volume of any gas theoretically becomes zero as indicated by the graph. (d)

When a gas is heated and its temperature rises, the average speed of its molecules increases. If

the volume of the gas is to remain constant, its pressure increases due to more frequent and

more violent collisions of the molecules with the walls.

2) (a) The pressure of a fixed mass of gas is inversely proportional to its volume if its temperature

is kept constant.

(b)

198

(c) All the molecules in a gas are in rapid motion, with a wide range of speeds, and repeatedly hit

the walls of the container in huge numbers per second. The average force and hence the

pressure they exert on the walls is constant since pressure is force on unit area.

3) (a) Diffusion is the movement of gas or solid particles from a region of high concentration to a

region of low concentration.

(b) It states that the rate of diffusion of a gas at constant temperature and pressure is inversely

proportional to the square root of its density.

4) V1 = 400 ml V2 = 600 ml

T1 = 15oC + 273 = 288 K

Substituting the values

T2 = 432 K

T2 in degree Celsius = 432 - 273 =159oC

5) P1 V1= P2 V2

Substituting the values

199

Hence the pressure required = 30400 mm of Hg

6) (a) Increase the potential energy of the molecules

OR do work in separating the molecules

against intermolecular forces/bonds

(b) Molecules are moving around randomly

spread in all directions

Moles

1) (a) This is the mass of 1 mol of an atom relative to the mass of 1 mol of C atoms that have 6

protons and 6 neutrons, which is taken to be 12.00 g.

(b)1 mole of CaCl2 consists of 1 mole of calcium ions (Ca2+) and 2 moles of chloride ion (Cl-).So

there will be 2.5 x 1 = 2.5 moles of calcium ions[1m] and 2.5 x 2 = 5 moles chloride ions.

2) (a) XC03(s) → X0(s) + C02(g)

X0(s) + H2(g) → X + H20(g)

(b) (i) Mass of water formed = 2.0 - 1.6 = 0.4g

Mass of Carbon dioxide formed = 3.1 - 2.0g = 1.1g

Moles of Carbon dioxide = 1.1 = 0.025

44

Moles of Metal oxide (X0) = 0.0025

Moles of Metal X 0.025

R.F.M = Mass =1.6 = 64

Moles 0.025

(ii) Volume of hydrogen used = 0.025 x 22400

= 560cm3

200

(iii) Volume of Carbon dioxide = 0.025 x 22400

= 560cm3 (1m)

3) (a) C3H8 + 502 → 3C02 + 4H20

CH4 +202 → C02 + 2H20

2C0 + 02 → 2C02

(b) Combustion of C3H8

Vol = 20 x 100 = 20cm3

100

Volume of oxygen = 100cm3

Combustion of CH4

Volume = 60 x 100 = 60cm3

100

Volume of oxygen = 60cm3 x 2 = 120cm3

Combustion of C0

Volume of oxygen = 20cm3 x 0.5

= 10cm3

4) (a) (i) many/gradual colour change(s);

(ii) phenolphthalein/(screened) methyl orange etc;

(b) (i) A calculation to include:

1. moles of acid = 200

1000

531.

.

= 0.00630;

2. moles of NaOH = 0.00630 x 2 = 0.0126;

3. concentration of NaOH =1000

25

01260

.

= 0.504 (mol dm-3

);

201

(ii) A calculation to include:

1. NaOH = 40;

2. (b) (i) x 40 = 20.16 (g dm-3

);

(c) (i) A description to include:

1. blue etc;

2. precipitate;

ORGANIC CHEMISTRY

1) (i) (a) A hydrocarbon is a compound containing hydrogen and carbon only.

(b) The compounds in each series have a similar chemical structure and a similar chemical

formula. Each family of organic compounds forms what is called a homologous series.

(ii)

(a)

Name carbons B.P.(°C) Uses

Refinery Gas 3 or 4 below 30 Bottled Gas(propane or butane).

Petrol 7 to 9 100 to 150 Fuel for car engines.

Naphtha 6 to 11 70 to 200 Solvents and used in petrol.

Kerosene (paraffin) 11 to 18 200 to 300 Fuel for aircraft and stoves.

Diesel Oil 11 to 18 200 to 300 Fuel for road vehicles and trains.

Lubricating Oil 18 to 25 300 to 400 Lubricant for engines and machines.

Fuel Oil 20 to 27 350 to 450 Fuel for ships and heating.

Greases and Wax 25 to 30 400 to 500 Lubricants and candles.

Bitumen above 35 above 500 Road surface and roofing.

(b)as number of c increases, so does the B.P.

2) saturated - only single bonds(between carbons)/contains maximum

amount of hydrogen per molecule;

unsaturated - has double/multiple bonds(between carbons)/does not

contain maximum amount of hydrogen per molecule.

3) (a) (i) cracking

heat (alkane) or (alkane) and catalyst

202

NOTE thermal cracking or catalytic cracking

alkane = alkene + hydrogen

ANY TWO

OR steam reforming

CH4 + H2O = CO + 3H2

or water/steam

catalyst or heat

(ii) combustion or burning

incomplete or insufficient oxygen/air

OR ACCEPT steam reforming as above

(iii) high pressure

COND forward reaction volume decrease

or volume of reactants greater than that of products

or fewer moles of gas on the right

or fewer gas molecules on right

NOTE accept correct arguments about either reactants or products

(b) (i) methyl ethanoate

(ii) propanoic acid or propanal

(iii) ethene

SOAP

1) (a) (i) scum/calcium stearate/CaSt;

(ii) hard;

(iii) add sodium carbonate/water softener/pass through

ion-exchange/boiling/distilling;

(iv) B, C

(b) (i) A description and an explanation to include:

1. scum/precipitate;

[Ignore no lather]

2. Ca2+

ions/calcium compound/it is hard;

[Reject contains calcium]

(ii) A description and an explanation to include:

1. lather;

2. no Ca2+

/Mg2+

ions etc/soft water/Li+

ions do not cause

203

hardness;

[Reject lithium ions causes softness/ contains Li ions]

2) (a) Water which does not form lather with soap is described as hard water.

(b) Calcium and magnesium sulphates are washed out of rock formations on hillslopes. These

make water permanently hard.Calcium and magnesium carbonates dissolve in acid rainwater to

form hydrogencarbonates. These make the water temporarily hard.

(c) Permanently Hard water can be made soft by removing the dissolved calcium and magnesium

ions by ion exchange. [1m]

CaSO4(aq) + Na2CO3(aq) → CaCO3(s) + Na2SO4(aq)

Or Ion exchange polymer resin columns hold hydrogen ions or sodium ions. These can be

replaced by calcium and magnesium ions when hard water passes down the column. [1m]

e.g. 2[resin]-H+(s) + Ca2+(aq) → [resin]-Ca2+[resin]-(s) + 2H+(aq)

Temporary hardness can be removed by boiling.

CaCO3(s) + H2O(l) + CO2(g) → Ca(HCO3)2(aq)

Ca(HCO3)2(aq) → CaCO3(s) + H2O(l) + CO2(g)

3) R-COO-Na+

Advantage

Does not pollute the environment

Disadvantage

Wastage of soap

R-OSO3-Na+

No wastage of soap

204

Disadvantage

Pollutes the environment

POLYMERS

1) (a) marks only for the reasons for the choice of poly(propene)

if any other polymer chosen, {0} for the section

useable temp. is above 100 °C

insoluble in oil

(b) polythene used for cling film plastic bags etc.

(c) any two problems from

non-biodegradable litter filling landfill sites

burning gives toxic gases

(d) structure of poly(propene)

correct repeat unit

shows continuation

2) (a) acid which is only slightly or partly ionised/partly dissociated/not fully ionised

NOT: only contains a few hydrogen ions

(b) 2C2H5CO2H + Na2CO3 → 2C2H5CO2Na + CO2 + H2O

(c) (i) 24g of magnesium will need 2 x 74 g of propanoic acid to react

so 4.8g magnesium requires 29.6g acid

so acid (30g) in excess

OR

74g of propanoic acid will need ½ x 24g of Mg to react

so 30g of acid requires 4.86g Mg

so acid in excess (as only 4.8g Mg used)

OR

mol Mg = 4.8/24 = 0.2

mol acid = 30/74 = 0.405(4)/0.41 mol;

2x moles of acid required to 1 mole Mg

Mg = 0.4 x 74 = 29.6g compared with 30 g acid

205

OR

0.405/2 moles = 0.2027/0.203 moles acid compared with 0.2 moles Mg

Any two of mark for both molar masses i.e. 24 and 74 /

• use of moles i.e. 4.8/24 or 30/74

• correct understanding of the 1:2 mole ratio

(no mark for stating which reactant is in excess)

(ii) 0.2 mol H2 (allow ecf from part (i));

0.2 x 24 = 4.8 dm3 (correct unit needed)

(d) (i) alcohols and carboxylic acids are monomers (both required);

ALLOW: alkanoic acids/OH and COOH or CO2H

(ii) condensation

(iii) clothing/named clothing/sails/conveyor or fan belts/

(e) one from:

• landfill - doesn’t (bio)degrade/

• incineration/burning - harmful substances/harmful fumes/harmful gases produced

ALLOW: stated harmful gas with correct effect e.g. hydrogen chloride acid rain/

carbon dioxide global warming etc.

• recycling - difficult to sort out different polymers [1]

ALLOW: expensive/time consuming

3) (i) 24;

(ii) many glucose molecules / monomers have linked together;

to form a long chain / a polymer is a long chain molecule;

SULPHUR, CARBON, NITROGEN AND CHLORINE

1) (a) sulphur;

air;

sulphur dioxide;

sulphur dioxide;

sulphur trioxide;

(b) Any two from:

making fertilisers;

detergents;

paints;

plastics;

206

2) (a) Differnent physical forms of the same element

(b) Graphite has delocalized electrons which enable it to conduct electricity

(c) It is a giant cavalent structure where each atom forms four bonds around itself.

3) (a) NH4NO3 → N2O + 2H2O

(b) It is used as an anaesthetic, popularly called laughing gas

4) (a) (i) hydrogen is below sodium in the reactivity series

(ii) chloride ions are removed (leaving hydroxide ions)

(iii) water purification OR treatment/killing bacteria etc./

bleaching agent (for paper)/making refrigerants

/making organic chlorine compounds (named)/making

solvents /making hydrochloric acid/extraction of bromine from

seawater/other suitable use

5. (a) (i) volcanoes / treatment of sulphide ores

ALLOW: bacterial oxidation / burning natural gas

IGNORE: unqualified burning fuels / from car engines / making sulphuric acid / from

smoke / from power stations

(ii) lightning / car engines / car exhausts / high temperature furnaces / explosives

ALLOW: burning fuel in car

NOT: from cars unqualified

NOT: bacterial activity / from fertilizers

(b) (i) carbon dioxide / CO2

(ii) calcium nitrite / calcium nitrate or correct formulae

IGNORE: incorrect oxidation numbers

(iii) Any one of:

• erodes buildings / reacts with buildings or statues

ALLOW: corrodes buildings / eats away buildings

NOT: destroys buildings / damages buildings

207

• forest death / kills trees or plants / kills fish in lakes / acidifies lakes

ALLOW: damages / destroys crops

NOT: kills animals (unless in lakes / rivers)

• breathing difficulties in humans OWTTE

NOT: causes pollution / harmful (unless specified) / affects building or animals

ACIDS, BASES, SALTS AND WATER

1) (a) An acid is any substance which produces H+ ions or H3O+ ions in water.

(b) PH is a measure of how acidic or how alkaline a solution in water is.

(c) The strength of an acid or alkali depends on how ionised it is in water.

(d) A weak acid or alkali is only partly (less than 100%) ionised.

A dilute acid is any acid (weak or strong) that contains plenty of water.

(e) Any organic acid such as ethanoic acid.

CH3CO2H(aq) H+(aq) + CH3CO2-(aq)

(f) Any metal carbonate or metal hydrogen carbonate. The metal carbonate or metal

hydrogen carbonate will bubble giving off carbon dioxide gas, leaving the metal salt and

water.

2) (a) filtration;

sedimentation / treatment with aluminium sulphate;

sterilisation / boiling / treatment with chlorine / ozone;

distillation;

(b) (i) calcium / magnesium;

(ii) water (during water cycle) flows over different types of rock /

different salts dissolve from different types of rock;

(iii) water and soap mixed / shaken;

if hard scum forms / little (or no) lather / excessive soap needed for lather;

(iv) boil the water;

distillation;

use of ion exchange resin;

other correct;

(c) sodium ion has a positive charge a sodium atom is uncharged;

because sodium ion has one less electron than sodium atom;

3) (a) filtration;

sedimentation / treatment with aluminium sulphate;

208

sterilisation / boiling / treatment with chlorine / ozone;

distillation;

(b) (i) calcium / magnesium;

(ii) water (during water cycle) flows over different types of rock /

different salts dissolve from different types of rock;

(iii) water and soap mixed / shaken;

if hard scum forms / little (or no) lather / excessive soap needed for lather;

(iv) boil the water;

distillation;

use of ion exchange resin;

MOLES

1) (a) A mole is a word which represents a number. Like the word "dozen" represents the number

12, so "mole" represents the number 6 x 1023.

(b)1 mole of KBr contains 1 mole of potassium ions (K+) and 1 mole of bromide ions (Br-).So there

will be 0.25 moles of each ion.

2) (a) XC03(s) → X0(s) + C02(g)

X0(s) + H2(g) → X + H20(g)

(b)(i) Mass of water formed = 2.0 - 1.6 = 0.4 g

Mass of Carbon dioxide formed = 3.1 - 2.0g = 1.1 g

Moles of Carbon dioxide = 1.1 = 0.025

44

Moles of Metal oxide (X0) = 0.0025

Moles of Metal X 0.025

R.F.M = Mass =1.6 = 64

Moles 0.025

(ii) Volume of hydrogen used = 0.025 x 22400

= 560cm3

(iii) Volume of Carbon dioxide = 0.025 x 22400

= 560cm3

3) (a) This is the mass of 1 mol of an atom relative to the mass of 1 mol of C atoms that have 6

protons and 6 neutrons, which is taken to be 12.00 g.

209

(b)1 mole of CaCl2 consists of 1 mole of calcium ions (Ca2+) and 2 moles of chloride ion (Cl-). So

there will be 2.5 x 1 = 2.5 moles of calcium ions and 2.5 x 2 = 5 moles chloride ions.

4) (i) Pb(N03)2(s) → Pb)(s) + 2N02(g) + 02(g)

(ii) Moles of Pb(N03)2 = Mass = 6.62 = 0.02

R.F.M 331

Mass of Pb0 = 0.02 x 223 = 4.46g

(iii) Vol of N02 = (0.02x2) x22400

= 896cm3 (1m)

Vol of 02 = 0.02 x 22400

= 448cm3

5) Relative formula mass = 2x27 + 3 x (32 + 4x16) = 342

there are 4 x 3 = 12 oxygen atoms, each of relative atomic mass 16, giving a total mass of

oxygen in the formula of 12 x 16 = 192

% O = 192 x 100 / 342 = 56.1% oxygen by mass

RATES OF REACTION

1) (a) (i) correct plotting;;(tolerance +/- ½ square

-1 for each incorrect point but ignore no point at 0,0)

Allow sticks if correct length for plotting

suitable curved line passing through origin;

Ignore after 2.5 minutes

(ii) volume of gas still increasing/graph still rising/not levelled

off/not straight;

(iii) graph extended and levels off before 5 minutes;

correct time reading from graph; (when it first stops rising)

allow even if graph comes back down. Ignore lack of units but

reject wrong units. Not a range

(b) comment on rate;

correct explanation i.e (same rate) acid already in excess/acid

concentration the same;

Accept (faster rate) because acid more concentrated;

Reject: more acid(particles)

Accept (slower rate) because bleach more dilute;

210

2) (a) (i) Uses collision theory to predict that the rate of formation of

H2(g) increases as the concentration of HCl increases

(ii) Uses collision theory to explain how rate of reaction increases

with increasing temperature

(b) Concentration of HCl identified as independent variable

[HCl] is acceptable

(c) States that the (total) volume of solution must be kept constant,

Or States that the amount/size/length/mass/surface area of the

magnesium ribbon must be kept constant

3. 1 mole of hydrogen burns to produce 286KJ

7.15/286 = 0.025 Moles of hydrogen will produce 7.15KJ

Volume of hydrogen = 0.025 x 24= 0.6 litres

4. I mole of carbon = 12g

Hence 3g = 3/12 = 0.25 mol

0.25 mol of carbon = -98.5KJ

I mole of carbon = 1/0.25 x -98.5

= -394KJmol-1

REVERSIBLE REACTIONS

1) (a) - forward reaction proceeds at the same rate as backward reaction

- concentration of products and reactants remain constant

- equilibrium can only be achieved in a closed system

(b) Endothermic

With increased temperature equilibrium shifts in the endothermic

(Forward) direction)

(c) (i) Equilibrium unaffected. There are equal numbers of gaseous

moles on either side of the equilibrium

(ii) Rate increases

Increasing the pressure also increases concentration of gas

2) (a) (i) Rates of forward and back reactions the same

(ii) · Moves equilibrium to left hand side

211

· being the side with the smaller number of molecules/moles

(b) (i) 2NO + O2……. 2NO2 or N2O4

(ii) 4NO2 + 2H2O + O2 …. 4HNO3

species balance.

Answer could be in terms of a reactions that give HNO2 and HNO3 as

the first stage.

i.e. 2NO2+ H2O ….. HNO2 + HNO3

(2HNO2 + O2 ….. 2HNO3)

or 3NO2+ H2O …. 2HNO3 + NO

METALS

1) (a) (i) chlorine / Cl;

(ii) aluminium / Al;

(b) orange substance is bromine / bromine is produced;

chlorine is more reactive than bromine;

chlorine displaces bromine / chlorine reacts with bromide ;

correct reference to redox;

(c) (i) iron(III) oxide;

carbon dioxide;

because these substances lose oxygen / reduction is loss of oxygen;

oxygen;

because carbon is oxidised and so oxygen must be reduced;

(ii) (56 x 2) + (16 x 3) or 160;

2) (a) alumina or aluminium oxide

sodium aluminate

iron(III) oxide

filtration or centrifuge NOT conditional

(b) from left to right:

carbon cathode or carbon negative electrode

900 to 1000oC

aluminium

cryolite [1]

(c) (i) Al3+ + 3e → Al

not balanced

212

Al3+(aq) = 0

(ii) oxygen is formed NOT oxide

reacts with carbon anode

(d) (i) low density or light or resistant to corrosion

accept strength/weight ratio or alloys are strong

strong on its own is neutral

(ii) not attacked or corroded or unreactive

oxide layer

easily shaped or malleable or ductile

any TWO

(iii) for strength or so it does not break or does not sag or can have pylons further apart

NOT steel is a better conductor

NOT aluminium protects steel from rusting

3) (i) Iron (iii) oxide or Silicon (iv) oxide

(ii)To release free ions which conduct electricity.

Oxygen produced reacts with carbon anode to form carbon dioxide. Hence the anode wears off

and needs regular replacement.