IGCSE Chemistry

Separating Solid/Liquid Mixtures:

Solute – the solid which dissolves in a solvent

Solvent – the liquid that the solute dissolves in

Solution – formed when a solute dissolves into another solvent: solute + solvent solution

Saturated solution – a solution which contains as much dissolved

solute as it can at a particular temperature

Soluble – when the solute can dissolve in a solvent

Insoluble – when the solute cannot dissolve in a solvent

Filtration – the process of separating a solid from liquid using a fine

filter paper which does not allow the solid to pass through. The

solid (residue) will stay in the filter and the liquid (filtrate) will be in

the container under the filter.

Decanting – the process of separating a liquid from solid (which has settled) or an

immiscible heavier liquid by pouring the solution into another container. The solid

or the immiscible heavier liquid will stay at the bottom while the liquid will pour out.

Centrifuging – the separation of the

components of a mixture by rapid spinning. The denser particles

are flung to the bottom of the containing tubes. The liquid can

then be decanted off.

Evaporation – the separation of a liquid and a dissolved solid by

heating the solution. The liquid will evaporate completely

leaving the solid behind.

Crystallisation – the process of forming crystals from a

liquid. This occurs when a solution is saturated the salt

begins to crystallise and can be removed with large

scoops.

Simple Distillation – the process of boiling a liquid and

then condensing the vapour produced back into a liquid.

It is used to purify liquids and to separate mixtures of

liquids.

Separating Liquid/Liquid Mixtures:

Miscible – description of liquids that form a homogeneous layer

when two are mixed together

Immiscible – description of liquids that form two layers when

two are mixed together

Separating Funnel – funnel that allows the layers in immiscible

liquids to separate

Fractional Distillation - process used to separate miscible liquids

into liquids that have different boiling points. When the mixture

is heated, liquids with low boiling points evaporate and turn to

vapour and can then be separated as liquids. Those with high

boiling points remain liquids

Separating Solid/Solid Mixtures:

Sublimation – heating of substances, in which one will sublime (I2 and CO2 will sublime)

Magnetism – takes out those things which are attracted to a magnet (Like Fe, Co and Ni).

Chromatography – substances dissolved in water (or other solvents) travel along chromatography

paper at different speeds. This difference in properties is used to separate some chemicals in

analytical laboratories. The substances move at different speeds due to their different solubilities in

the solvent.

Locating Agent – used in chromatography to make the spots show when the substance is not visible

Rf Values – used to identify which spot is which item

Solvent Extraction – when substances are extracted from a mixture by using a solvent which

dissolves only those substances required

How the purity of a substance can be shown:

Pure Impure

Melting Point Sharp Melting Point; usually high Range of temperatures

Boiling Point Sharp Boiling Point; usually low Range of temperatures

Chromatography One well-defined Spot on chromatogram Several spots on chromatogram

Element - the simplest building blocks of the physical world. There are 92 naturally occurring

elements. The periodic table is a list of elements in order of atomic number.

Atom – the defining structure of an atom which typically includes a nucleus of protons and neutrons

with electrons orbiting the nucleus

Molecule - Two or more atoms joined together discreetly. (Usually non-metals; can be elements or

compounds)

Compound - two or more elements (different types) chemically bonded. With compounds you can

always give formula.

Mixture - two or more substances not chemically bonded

Charge Mass

Proton + 1

Electron - 0

Neutron NO CHARGE 1

Ion - Charged particles that are formed when an atom loses or gains an electron

Cat-ion - A positive ion

An-ion - A Negative ion

Valence Electrons - Electrons on the outer shell of an element

The Atomic Number is the number of Protons and Electrons unless it is an ion. (Protons = Electrons)

The Mass Number is the number of Protons + Number of Neutrons. To work out the number of

neutrons you must calculate: Mass Number - Atomic Number = Number of Neutrons

Isotope - An element that occupies the same place in periodic table but has a different number of

neutrons. The number of protons and electrons are the same.

Alloy – a mixture of a metal and another element (usually a metal)

Bohr Diagram Example

9 P

10 N

Lewis Diagram Example

F

Magnesium

12

Mg 24

Atomic

Number

Mass

Number

Element:

Fluorine

Electron

Configuration:

2 , 7

Bonding

Metal + Electrons = Metallic Bonding

Metal + Non-Metal = Ionic Bonding

Non-Metal + Non-Metal = Covalent Bonding

Metallic Bonding:

Metallic Bond is the attraction between the metal ions and the delocalised electrons

Too many atoms to count, unlike small molecules like H2O

Number of delocalised electrons = number of electrons in out shell of element

Metals conduct electricity because the electrons can move

Metals are easily shaped (malleable) because cat-ions are in compact layers and can move

Metallic Lattice – the regular arrangement in metal ions in solid metals

Ionic Bonding:

Ionic Bond is the electrostatic attraction between positive and negative ions

Only show valence electrons in diagrams.

Show transfer of electrons with arrow

Write the correct formula

Both atoms should end up with full valence shells

Polyatomic ions – ions containing more than one atom. Brackets must be used to write formulae

involving more than one of these ions. E.g. Al2(SO4)3; SO4 is a polyatomic ion

Covalent Bonding:

Covalent bond is the sharing between atoms

to gain full valence shells

Only show valence electrons on diagrams

No arrows

Atoms will be connected

+ + + + + +

A

+ + + + + +

A

+ + + + + +

A

Cat-Ions

Free (delocalised)

electrons, they can

move anywhere

Na Cl Na Cl

+ -

H H

O

Polar Bonds

Non-Polar Bonds – covalent bonds that involve exactly equal sharing of the bonded pair(s) of

electrons (e.g. Cl with Cl, O with O) or close enough to equal sharing (e.g. C with H) that the shared

pair of electrons is equidistant between the two atoms and thus the electronic charge is evenly

balanced around all atoms

Polar Bonds – covalent bonds involve uneven sharing of the electron pair or pairs, with one of the

atoms (e.g. F, O or Cl) having a slightly stronger attraction for the shared pair of electrons in the

bond than the other atom (e.g. C or H). As a result the covalent bonds are closer to that atom with

stronger attraction. This gives that atom a slightly negative charge (-) and the other atom a slightly

positive charge (+)

Polar Molecules – molecules with at least one polar bond and an asymmetrical shape so the dipoles

do not cancel

Non-Polar Molecules – molecules where all bonds are non-polar or with a symmetrical shape causing

the dipoles to cancel out

H H

O

+

+

-

-

Oxygen atom has stronger attractions

and hydrogen has weaker attractions.

Covalently bonded electrons are

closer to oxygen making it slightly

negative and the hydrogen atoms

slightly positive creating dipoles

Molecule in asymmetrical shape

therefore dipoles do not cancel

Two polar bonds

Water is Polar

H H

C H H

Carbon and Hydrogen have the same

attraction

Covalently bonded electrons are

shared almost equally creating no

dipoles

All bonds are non-polar

Molecule is in symmetrical shape so

if there were polar bonds they would

cancel out

Methane is Non-Polar

Shapes of Molecules

Molecule Example Shape Bond Angle

CO2 Linear O-C-O = 180°

H2O Bent/V-Shaped H-O-H = 105°

BF3 Triangular Planar F-B-F = 120°

NH3 Pyramidal H-N-H = 107°

CH4 Tetrahedral H-C-H = 109.5°

Radioactivity – when the nucleus in a radioactive atom is unstable and so it emits particles/waves to

form a more stable atom

Radioisotopes – a radioactive isotope; one having an unstable nucleus and emitting radiation during

its decay to a stable form

Radiation – the particles/waves emitted by radioactive substances

Radioisotopes and their Uses

Radioisotopes Uses

Carbon – 14 Carbon Dating; When an organism dies it stops taking in new carbon atoms so the amount of carbon-14 slowly drops as the atoms decay. By measuring the radiation

from the carbon-14 atoms the age of the remains can be determined.

Cobalt - 60 Cancer Treatment; weak beams of radiation will kill cancer cells more easily than healthy cells. Cobalt-60 is also used to kill germs and bacteria.

Krypton – 81 Tracers; a small amount of krypton is breathed in, it decays in the lungs and the radiation can be detected and viewed on a screen where it shows up as bright spots.

Dark patches show where the lungs are not working properly.

Cesium - 137 Kill Germs and Bacteria; a low dose of gamma radiation is used to kill bacteria in food that causes it to decay. They also kill germs and bacteria on surgical equipment.

States of Matter

Solid Liquid Gas

Fixed shape and volume No fixed shape, fixed volume No fixed shape or volume

Particles are held together by relatively strong forces

Particles have weaker forces so are further apart

Particles are very far apart

Incompressible Slight compressibility Compressible

Particles do not have free movement but can vibrate

around fixed positions

Particles can move throughout bulk of liquid

Particles are very spread out and move in random fashion

Matter – all the substances and materials from which the physical universe is composed

Kinetic Particle Theory – a theory which accounts for the bulk properties of matter in terms of the

constituent particles. It states that:

All matter is made up of tiny, moving particles, invisible to the naked eye. Different

substances have different types of particles (atoms, molecules or ions).

The particles move all the time. The higher the temperature, the faster they move and the

forces of attraction weaken

Heavier particles move more slowly than lighter particles at a given temperature

Diffusion – the process by which different substances mix as a result of the random motions of their

particles. Particles with smaller Mr diffuses faster.

Intimate mixing – when diffusion takes place between a liquid and a gas

Brownian Motion – random motion of particles caused by smaller and faster moving water particles

constantly colliding with them and moving them around

When the object is melting or vaporising, heat energy is

being added but temperature is not changing. The average

kinetic energy stays the same. Energy changes to potential

energy by separating.

Absolute Zero – the theoretical temperature (which can never be reached) at which all particle

motion stops. Absolute zero is -273°C or 0 Kelvin. To calculate Kelvin = °C + 273. A 1K change equals

a 1°C change.

Pressure of a Gas

The free moving particles of a gas will spread evenly within a container and collide with the

walls. This will exert a force on the wall when it bounces off.

When this happens on a large scale (billions of particles) there is an average force exerted on

the wall. This creates a pressure due to Pressure =

Vapour Pressure – particles that gain enough energy to become gaseous at the top of a liquid

Melting Evaporation

Sublimation

Solidification

Deposition

Solid

Melting

Liquid

Vaporising Gas

Condensation

Ionic Compounds

Solid ionic compounds have no moving charged particles, they do not conduct electricity.

Liquid and aqueous ionic compounds have free moving charged particles (ions) in solution

which can carry charge under the influence of an electric filed.

Ionic Bond Covalent Bond

Bonding - Results from the attraction between positive and negative

ions - Occurs when a metal reacts

with a non-metal

- Formed by the sharing of electrons between atoms

- Occurs when two non-metals react

Properties - Does form ions so it will conduct electricity as liquid or

aqueous - Forms a crystal lattice

structure - Has a high melting point

- Does not form ions so does not conduct electricity

- Forms a shared electron structure

- Has a low melting and boiling point

Intermolecular – these are attractions between molecules when they are close together and are

broken when substances melt or boil.

Intramolecular – this refers to the covalent bonding. These bonds are only broken during a chemical

reactions and never when melting or boiling. There are three varying degrees of strength of these

bonds that depend on the type of molecules:

1. Temporary dipole attractions – the weakest attraction between non-polar molecules

2. Permanent dipole attractions – the next strongest attraction between polar molecules

3. Hydrogen bonding – strongest attraction between polar molecules and extremely reactive non-

metals (like F, O, or N)

Physical Changes – when the appearance/form of the substance changes but the actual identity and

characteristics of the substance remain the same

Chemical Change/Reaction – when substances chemically combine and alter one another forming

new substances with different properties and characteristics.

Conductors

To conduct, charged particles must be present and these charged particles must be free to

move.

There are two types of conductors:

o Elements which conduct in both solid and liquid because their outer shell electrons

are mobile e.g. metals

o Electrolytes conduct because they contain positive and negative ions. In electrolytes,

the mobile ions carry the current under the influence of an electric field, and the

electrolyte is decomposed/discharged as the ions gain or lose electrons at the

electrodes e.g. Sodium Chloride solution

Allotropes – different forms of the same element.

Allotropy – when an element can exist in more than one physical form in the same state

Giant covalent structures – structures with a network of covalent bonds throughout it. They take a

lot of energy to break and have high boiling and melting points.

Diamond Graphite Silicon Dioxide (Silica)

Stru

ctu

re

Bo

nd

ing

Each carbon atom has four covalent bonds with other

carbon atoms

Each carbon atom has three covalent bonds with other

carbon atoms. Van der Waal’s forces hold the layers

together.

Each silicon atom has four covalent bonds with oxygen

atoms.

Arr

ange

men

t Carbon atoms link together to form a giant lattice

structure

Arranged in hexagons and are arranged in layers on top of each other. Electrons move

throughout layers

Atoms link together to form giant lattice structure

Pro

p. Does not conduct electricity,

high melting point, insoluble Conducts Electricity, high melting point, insoluble

Does not conduct electricity, high melting point, insoluble

Loo

k

A hard, colourless, transparent crystal which

sparkles in light

A soft dark grey, shiny solid with a slippery feel

A hard, colourless, transparent crystal which sparkles in light (Quartz)

Use

Jewellery, Glass Cutters, Polishers

Pencils, Electrodes, Lubricants

Cement

Calculating Ar (Relative Atomic Mass/RAM):

The Ar of an element is defined as the average mass of its isotopes compared with one-

twelth of the mass of one atom of Carbon 12

Carbon 12 has a mass of 12 therefore, one-twelth of the mass of one atom of Carbon 12 is 1.

To calculate Ar you use the equation: Ar =

.

Average mass of isotopes of the elements =

Calculating Mr (Relative Formula Mass/Relative Molecular Mass):

To calculate the Mr of a compound, you add the mass number of each of the elements

e.g.

NaCl = 23 + 35 = 58 or MgBr2 = 24 + 80 + 80 = 184

The Mole Concept

Avogadro’s Constant – equal volumes of all gases measured under the same conditions of

temperature and pressure contain equal numbers of molecules

Moles – the amount of a substance which contains 6x1023 atoms, ions or molecules. This number is

called the Avogadro’s constant. One mole of atoms has a mass equal to the relative atomic mass (Ar)

in grams. One mole of molecules has a mass equal to the relative molecular mass (Mr) in grams.

The calculation of no. of moles is moles =

The calculation of Molar Mass is Molar Mass =

The calculation of mass is mass = number of moles x Molar Mass

The calculation of concentration is concentration =

The calculation of Volume is Volume =

The calculation of no. of moles is moles = concentration x Volume

The calculation of no. of moles is moles =

One mole of any gas at rtp = 24

The calculation of Volume = Volume x 24

Empirical Formula

Percentage composition by mass gives a ratio by mass of the elements contained in a

compound

This ratio by mass can be converted to a ratio by moles if the % figures (or the known mass)

are divided by the respective relative atomic masses

Ratios of moles will seldom be whole number ratios

To calculate the ratio of numbers or atoms in the empirical formula divide all the mole ratios

by the smallest calculated mole value

The amount of water associated with a particular salt (water of crystallisation) can be

calculated using this method as well.

Step Calcium Bromine

1 - % By Mass (or Mass) 20% 80%

2 – Mole Ratio 20/40 = 0.5 80/80 = 1

3 – Ratio by atoms 0.5/0.5 = 1 1/0.5 = 2

4- Empirical Formula Ca = 1 Br = 2

n M

m

c V

n

n 24

V

Percentage Yield – the percentage of the reactants that are converted to products. In some reactions

this will be 100% but others it is much less than 100%.

Theoretical Yield – the amount of a substance that should be produced through a chemical reaction

Actual Yield – the amount of a substance that is actually produced through a chemical reaction

Percentage Yield =

e.g.

(30.7g) CaCO3 (11.7g) CO2 + CaO; 100 = Mr of CaCO3, 44 = Mr of CO2

n(CaCO3) =

= 0.307 moles

44 x 0.307 = 13.508

= 86.6%

Types of Reactions

Single Displacement - one element displaced

Most Metals + Water Metal Hydroxide + Hydrogen Gas

Most Metals + Acid Metal Salt + Hydrogen Gas e.g.

Cu + HCl CuCl2 + 2H

A + BC A + B

Double Displacement - two elements displaced

Acid + Base Salt + Water e.g.

2KI + Pb(NO3)2 PbI2 + 2KNO3

AB + CD CB + AD

Synthesis - two elements combine to form one compound

Most Metals + Oxygen Gas Metal Oxides e.g.

2Mg + O2 2MgO

A + B AB

Decomposition - one compound becomes two (breaks apart/decompose)

Most Metal Hydroxides Metal Oxide + Water Liquid

Most Metal Nitrates Metal Oxide + Nitrogen Dioxide Gas + Oxygen Gas

Group One Metal Nitrates Group 1 Metal Nitrate + Oxygen Gas

Most Metal Carbonates Metal Oxide + Carbon Dioxide Gas e.g.

CuCO3 CuO + CO2

AB A + B

Combustion – oxygen combines with another compound to form water and CO2

Organic Molecule + excess Oxygen Gas Carbon Dioxide + Water e.g.

C10H8 + 12 O2 10 CO2 + 4 H2O

AB + C AC + BC

Redox Reactions

Oxidation – the gain of oxygen; the loss of hydrogen; the loss of electrons; increase in the oxidation

number

Reduction – the loss oxygen; the gain of hydrogen; the gain of electrons; decrease in the oxidation

number

Oxidising Agents (Oxidants) – the molecule that is reduced (lost oxygen, gained hydrogen etc.)

Reducing Agents (Reduced) – the molecule that is oxidised (gained oxygen, lost hydrogen etc.)

Oxidation Number – numerical bookkeeping system to help keep track of electron movements

between substances. The oxidation number of ions is there electric charge. Elements have an

oxidation number of 0.

Photochemical Reactions

Photosynthesis – the conversion of water and carbon dioxide into glucose and oxygen for energy.

The glucose produced is used to make sugars and starch as carbohydrates.

Carbon Dioxide + Water Glucose + Oxygen Gas

6 CO2 + 6 H2O C6H12O6 + 6 O2

Respiration – reverse of photosynthesis; combustion reaction

Glucose + Oxygen Gas Carbon Dioxide + Water

C6H12O6 + 6 O2 6 CO2 + 6 H2O

Photography – this photochemical reaction is used as the basis of black and white photography. The

photographic film is made of flexible plastic, coated with a layer of gelatine with millions of particles

of silver bromide spread through it which is changed to silver when light falls on the exposed parts of

the film.

Silver Bromide Silver + Bromide

2AgBr 2Ag + Br2

Hydrogen as a fuel:

Hydrogen is considered to be the fuel of the future and is being trialled by motor

manufacturers as an alternative to fossil fuels such as petrol.

Hydrogen is non toxic, produces more energy per gram than any other fuel and burns cleanly

to form water so there is no exhaust pollution. It has a lower flammability than fossil fuels.

Hydrogen is obtained from the electrolysis of water which is plentiful. So far this is not a cost

efficient alternative to using fossil fuels and non-renewable if produced using fossil fuels or

nuclear energy.

Hydrogen is difficult to transport and store. Because it is too light to liquefy easily, a large

fuel tank would be needed.

It is explosive in correct proportions with air.

Energy

Exothermic Reactions – when energy is lost to the surroundings during a reaction.

Energy Profile for an Exothermic Reaction – the graph shows the variation in energy during the

course of a chemical reaction where heat is released.

Endothermic Reactions – when energy is absorbed by the products from the surroundings during a

reaction.

Activation Energy (Ea) – the initial energy that is required for a reaction to begin

Calorimeter – determines the amount of heat generated in a chemical reaction by the rise in

temperature of the reaction chamber and the water jacket around the reaction vessel.

Bond Energy – the amount of energy needed or released to break or form a bond. Bond breaking is

endothermic. Bond forming is exothermic.

ΔH = Total bond energy of all bonds broken – Total bond energy of all bonds formed

Equilibrium Reactions

Irreversible Reactions – reactions that has products that cannot turn back into their reactants.

Reversible Reactions – reactions that has products that can react back into the original reactants.

Dynamic Equilibrium – when there is no overall change in the amount of products and reactants

even though the reaction is ongoing. Dynamic Equilibrium can only take place in a closed system.

The position of dynamic equilibrium is not always at a half-way point, as in it may be at a position

where there are more products than reactants.

Le Chatelier’s principle – if a closed system at equilibrium is subject to a change then the system will

adjust in such a way as to minimise the effect of the change.

Energy Level

Reaction Progress

Reactants

Products

Reaction Progress

Products

Reactants

Energy Level

Energy Lost

(ΔH)

Ea

Ea

Energy Absorbed

(ΔH)

Factors affecting Equilibrium

Factor Increase of Factor Decrease of Factor

Temperature

Equilibrium shifts to decrease the temperature so it shifts to

the endothermic direction

Equilibrium shifts to increase the temperature so it shifts to

the exothermic direction

Concentration

Equilibrium shifts to decrease the concentration

Equilibrium shifts to increase the concentration

Pressure

Equilibrium shifts to decrease the pressure so it shifts in the

direction of the least molecules

Equilibrium shifts to increase the pressure so it shifts in the

direction of the most molecules

Catalyst

Speeds up the time it takes to reach equilibrium but does not

change the position

-

Haber Ammonia Process

Haber Process – the process by which ammonia is made from nitrogen and hydrogen. Nitrogen is

obtained from air and hydrogen is obtained from methane. It follows the following equation:

N2 + 3 H2 2 NH3 ΔH = -92 kJ/mol

Increasing the temperature will produce less ammonia because this will use up the added

heat. Lowering the temperature will produce a greater yield of ammonia but will decrease

the rate of the overall reaction.

Increasing pressure should move the equilibrium to the right to produce more ammonia.

However this will increase the cost because of the thickness of the walls of the plant needed

to contain the reaction and it means the temperature will increase and its disadvantages.

Conditions of the Haber Process

Pressure of 200 atm and Temperature between 380 and 450 °C

Ground Iron catalyst to increase the rate of reaching equilibrium at the lower temperature

The equilibrium mixture is cooled, allowing ammonia to liquefy and be removed.

Unused Nitrogen and Hydrogen is continuously recycled back into the system.

Rates of Reaction

Rates of reactions can be measured by the:

Time for a solid to dissolve or form

Loss in Mass (gas given off) over time

Volume of Gas collected per time

Time for a colour to appear or disappear

Collision Theory - in order for a chemical reaction to occur the particles must collide with each other

and have sufficient energy to react. The rate of reaction depends on the number of successful

collisions there are in a given time. When particles move faster, they have more kinetic energy.

Rates of reaction are affected by:

1. Concentration - adds more

particles so they can collide

with each other. At the

beginning of the reaction, the

concentration is at its

highest, so the initial rate is

the fastest as there are a

large number of reactant

particles per unit volume and

more collisions will occur. As

the reactants decrease in

concentration, there will be less

collisions and the rate slows

2. Particle Size/Surface Area -

larger area means there is more

room for the particles to roam

and will collide easier. The

amount of product remains the

same but the surface area is

different

3. Gas Pressure – an increase in pressure forces the particles to come closer together and

increases the chance of successful collisions.

4. Temperature – a higher temperature gives particles more energy for collisions and makes

the particles move faster so they are more likely to collide.

5. Catalysts – provides particles an alternate way of reacting if the activation energy of the

particle is too low, without itself being consumed. E.g. Iron in Haber process, Enzymes in

human body

Acids and Bases

Acid – a substance that acts as a donor of hydrogen ions

Base – a substance that acts as an acceptor of hydrogen ions

Alkali – soluble bases

Acid Base

Sour Taste Bitter Taste

pH less than 7 pH greater than 7

In solution, contains hydronium ions (H3O+) In solution, contains hydroxide ions (OH-)

Turns blue litmus red Turns red litmus blue

Turns phenolphthalein colourless Turns colourless phenolphthalein pink

Corrosive Soapy feel

Reacts with metals to produce salt and hydrogen Cannot react with metals

Examples of Acids Examples of Bases

Hydrochloric Acid HCl Sodium Hydroxide NaOH

Nitric Acid HNO3 Potassium Hydroxide KOH

Sulphuric Acid H2SO4 Calcium Hydroxide Ca(OH)2

Ethanoic Acid CH3COOH Ammonia Solution NH3 (aq)

Hydronium Ion – same as a single proton because when a hydrogen atom loses an electron, only a

proton remains. H+ is irresistibly attractive to water molecules and therefore it would form H3O+.

Dissociation – breaking apart

Strong Acids – in aqueous solutions, strong acids donate all their protons to water molecules.

Weak Acids – there is only a slight tendency to donate protons to water molecules, therefore an

aqueous solution of a weak acid contains mainly undissociated molecules and a low concentration of

H3O+.

Strong Acids Weak Acids

Dissociation in Aqueous Solution Completely dissociate Partially dissociate

Equilibrium None (forward only) Equilibrium reaction

Electrolyte Good Poor

Electrical conductivity Good Poor

[H3O+] Higher Lower

pH value Lower Higher

Examples HCl, HNO3, H2SO4 CH3COOH, NH4+

Amphiprotic – substances can act as both an acid and a base e.g. H2O, HCO3-, HSO4

-

Amphoteric – substances will undergo chemical reactions with both acids and bases

Neutralisation – an alkali or base can neutralise an acid by removing the H+ ions and converting them

to water. Neutralisation always produces a salt.

Concentration – a measure of the amount of acid per dm3, refers to the proportion or ratio of acid to

water in the solution

Concentrated Acids – high proportion of acid to water

Dilute Acids – low proportion of acid to water

Monoprotic – having one transferrable proton

Diprotic – having two transferrable protons

Titration – an indicator shows when the acid properties are just destroyed by the alkali. The salt can

then be recovered by evaporating the water away allowing the salt to crystallise. This method is

used when the base, acid and salt are all soluble.

Oxides

Oxides of metals are bases (they will react with acids to form salts)

Oxides of non-metals are acids (they will react with acids and bases)

Some metal oxides are amphoteric (they will react with acids and bases)

Some non-metal oxides are neutral

Oxide ions immediately react with water and then dissolve to form hydroxide ions. Although

potassium hydroxide solution exists, potassium oxide solution does not exist

Metal Oxides – compounds of metal cations and the oxide anion O2-. Few metal oxides react or

dissolve in water. The main metal oxides which are considered soluble are potassium and sodium

oxides, as well as barium, calcium, and magnesium oxides in decreasing amounts. Metal oxides are

either basic or amphoteric. The basic oxides will only react with acids, while the amphoteric oxides

will react with both acids and bases.

Non-metal Oxides – covalently bonded compounds of a non-metal with oxygen. They are either

acidic or neutral oxides. The acidic oxides react with water immediately and dissociate to form acid

solutions while the neutral oxides do nothing when placed in water. The acidic oxides will react only

with bases, while the neutral oxides are unreactive with both acids and bases.

Solubility of Ionic Compounds in Water

Precipitate – a solid formed in a solution.

Sparingly Soluble (SpSol) – materials have very low solubilities

Hydrolyse (Hyd) – reacts with water

Always Soluble Usually Soluble Usually Insoluble

All NO3- All SO4

2- EXCEPT Ba, Pb, Ag, Ca All CO32- EXCEPT Group 1

All NH4+ All Cl- EXCEPT Ag, Pb All O2- EXCEPT Group 1 and 2

Group 1 All I- EXCEPT Ag, Pb All OH- EXCEPT Group 1; Ca and Ba are slightly soluble All Br- EXCEPT Ag, Pb

NO3- Cl- Br- SO4

2- CO32- OH- O2- I-

NH4+ White

Soluble White

Soluble White

Soluble White

Soluble White

Soluble - - White

Soluble

Na+ White Soluble

White Soluble

White Soluble

White Soluble

White Soluble

White Soluble

White Hyd

White Soluble

K+ White Soluble

White Soluble

White Soluble

White Soluble

White Soluble

White Soluble

White Hyd

White Soluble

Al2+ White Soluble

White Soluble

White Soluble

White Soluble

- White Insoluble

White Insoluble

White Hyd

Zn2+ White Soluble

White Soluble

White Soluble

White Soluble

White Insoluble

White Insoluble

White Insoluble

White Soluble

Ca2+ White Soluble

White Soluble

White Soluble

White Insoluble

White Insoluble

White SpSol

White Hyd

White Soluble

Cu2+ Blue Soluble

Blue Soluble

Black Soluble

Blue Soluble

Green Insoluble

Blue Insoluble

Black Insoluble

-

Ag+ White Soluble

White Insoluble

White Insoluble

White Insoluble

White Insoluble

- Brown Insoluble

Yellow Insoluble

Fe2+ Green Soluble

Yellow Soluble

Green Soluble

Green Soluble

Grey Insoluble

Green Insoluble

Black Insoluble

Grey Soluble

Fe3+ Violet Soluble

Brown Hyd

Red Hyd

Yellow Soluble

- Brown Insoluble

Red Insoluble

-

Identification of Cations

Add a few drops of NaOH Add excess NaOH Add Ammonia to fresh sample

NH4+ No precipitate formed When warm = litmus blue -

Cu2+ Blue precipitate formed - -

Fe2+ Green precipitate formed - -

Fe3+ Orange/Brown precipitate - -

Al3+ White precipitate Precipitate dissolves Insoluble white precipitate

Ca2+ White precipitate Precipitate remains -

Zn2+ White precipitate Precipitate dissolves Soluble white precipitate

Identification of Anions

Test with red litmus Add HNO3 and Ba(NO3)2

Add dilute HNO3 and AgNO3 to fresh sample

Add Al and NaOH to fresh sample; warm

CO3- Litmus turns blue - CO2 released (HNO3 only) -

SO42- No change Precipitate forms - -

Cl- No change No precipitate White precipitate -

I- No change No precipitate Yellow precipitate -

NO3- No change No precipitate No precipitate NH3 produced

Identification of Gases

Ammonia Turns damp litmus blue; forms white smoke when in contact with HCl fumes

Carbon Dioxide Turns limewater milky

Chlorine Bleaches damp litmus paper

Hydrogen ‘Pops’ with lighted splint

Oxygen Relights a glowing splint

Physical Properties of Metals, Non-Metals and Metalloids

Metals Non-Metals Metalloids

Lustre (shiny) No lustre Can be shiny or dull

Good conductor of heat Poor conductor of heat Fair conductor of heat

Good conductor of electricity Poor conductor of electricity Fair conductor of electricity

Malleable Not Malleable Malleable

Ductile Not Ductile Ductile

High Density Low Density Solids

High Melting Point Low Melting Point

Chemical Properties of Metals and Non-Metals

Metals Non-Metals

Easily loses electrons Tends to gain electrons

Oxides generally basic and amphoteric Oxides generally neutral

Corrodes easily

Alkali Metals

Group One Metals

Very low density and therefore floats on water. The densities increase down the group.

Silvery and shiny when freshly cut, however they quickly tarnish

Low melting point

Low boiling point

The reactivity increases down the group. Since the valence electron is further from the

nucleus, the attractive force holding it is weaker and therefore other stronger forces can

easily remove it.

Transition Metals

Physical Properties (compared to Group 1) Chemical Properties (compared to Group 1)

Much harder Much less reactive

Higher tensile strength Many have excellent corrosion resistance

Higher density Show more than one valency (e.g. Fe2+ or Fe3+)

Higher melting point and boiling point Them and their compounds are useful catalysts

Many of their compounds are coloured Some are strongly magnetic

Alloys

Alloy Mixture Use

Solder 70% Tin, 30% Lead Joining wires and pipes

Brass 60-95% Copper, 5-40% Zinc Taps, hose/pipe fittings, zips, screws

Bronze 90% Copper, 10% Tin Ornaments, bells, bearings

Mild Steel 99.5% Iron, 0.5% Carbon General structural purposes, cars

Hard Steel 99% Iron, 1% Carbon Blades

Stainless Steel 74% Iron, 18% Chromium, 8% Nickel Corrosion resistance

Alnico Iron/Aluminium/Nickel/Cobalt Permanent Magnets

Metals Uses

Metal Property Uses

Aluminium Does not corrode Food containers

Low density, unreactive Containers and packaging buildings

Low density, strong, conducts Long distance wiring

Low density, strong, cheap Transport vehicles

Low density, conducts heat Car Engines

Zinc Reactive Dry cells (“batteries”)

More reactive than iron Galvanising Iron

Iron Similar expansivity Reinforcing concrete

Strong, cheap Nails

Strong and abundant Ship building

Copper Good conductor of electricity Electrical wiring

Unreactive, workable Alloys – brass and bronze

Unreactive Coinage (with Nickel)

Unreactive Hot water piping

Reactivity Series

Corrosion

Corrosion – when metals react with water and oxygen. The metal ions lose electrons to form ions.

Rusting – the corrosion of iron metal to form a red-brown compound (hydrated iron (III) oxide)

Covering with Protective Coat Preventing Oxidation of Metal

Painting Galvanising – zinc atoms react before the iron

Greasing of metal parts Sacrificial protection – a more reactive metal reacts before the metal that it is protecting Oiling of bike chains

Tin Plating – In cans Carthodic protection – an electric power source pushes electrons into the metal to prevent the loss of electrons

Plastic covering on electric wires

Galvanising – zinc coating for galvanised steel

Chromium plating of car parts

Most Reactive K Potassium

Na Sodium

Ca Calcium

Mg Magnesium

Al Aluminium

C Carbon

Zn Zinc

Fe Iron

Sn Tin

Pb Lead

H Hydrogen

Cu Copper

Ag Silver

Au Gold

Least Reactive Pt Platinum

Any metal higher on the reactivity

series will displace another lower

metal’s ions from solution.

e.g. Ca (s) + Cu2+ (aq) Ca2+ (aq) + Cu (s)

BUT Cu (s) + Ca2+ (aq) No Reaction

The more reactive metals are difficult

to extract from their ores in compound

form as they are stable.

The less reactive metals have the

greater tendency to form atoms and

therefore their compounds are less

stable.

Reduction of Metal Oxides

By Hydrogen: only the metals below Hydrogen in the reactivity series are reduced by using this

method (mainly only CuO)

CuO (s) + H2 (g) Cu (s) + H2O (l)

By Carbon: only the metals below Carbon in the reactivity series are reduced by using this method

2 PbO (s) + C (s) 2 Pb (l) + CO2 (g)

By Carbon Monoxide: only metals below Carbon in the reactivity series are reduced by using this

method

CuO (s) + CO (g) Cu(s) + CO2 (g)

Blast Furnace

Iron is extracted from Haematite or Ironsand in a Blast Furnace

A charge is a mixture of limestone, coke (carbon) and iron oxide (as well as it’s impurities,

mainly consisting of SiO2)

The charge is placed in the top of the blast furnace and hot air is blasted through at the

bottom, making the charge glow white hot.

The following reactions take place:

C (s) + O2 (g) CO2 (g) CaCO3 (s) CO2 (g) + CaO (s)

CO2 (g) + C (s) 2 CO (g) CaO (s) + SiO2 (s) CaSiO3 (l) This is known as slag.

3 CO (g) + Fe2O3 (s) 2 Fe (l) + 3 CO2 (g)

If the iron from the blast furnace solidifies, it is called

cast iron and is mostly turned into steel. Steel is

manufactured the following way:

Unwanted impurities are removed in an oxygen

furnace where the molten metal is poured into a

furnace along with some scrap iron (to recycle

it). Calcium oxide is added and a jet of oxygen is

blasted into it. The calcium oxide reacts with the

impurities forming slag that can be skimmed off.

Oxygen reacts with the excess carbon, burning

most of it away as CO2, leaving some to mix with

the iron to make the metal hard but not brittle.

Other elements are then added to gain the

desired steel properties.

Zinc from Zincblende

1. The ore zincblende (made mostly from Zinc Sulphide) is crushed and put into water through

which air is blown. Rock particles sink and the zinc sulphide floats in a froth which is

skimmed off and dried. The product of this stage is 55-75% Zinc Sulphide.

2. The Zinc Sulphide is converted to Zinc Oxide by strong heating in a furnace:

ZnS + 3 O2 2 ZnO + 2 SO2

3. Zinc Oxide is mixed with coke in a furnace and heated to 1400 °C where it is reduced to zinc:

ZnO + C Zn + CO

4. The zinc metal produced cools and the carbon monoxide is burnt, with the heat given out to

help reduce costs of the furnace.

Heating Metal Compounds

Hydroxide Nitrate Carbonate

Potassium Stable No Reaction

Decomposes 2NaNO3 2NaNO2 + O2

Stable No Reaction Sodium

Calcium

Decomposes Cu(OH)2 CuO + H2O

Decomposes 2Ca(NO3)2 2CaO +

4NO2 + O2

Decomposes MgCO3 MgO + CO2

Magnesium

Aluminium

Carbon

Zinc

Iron

Lead

Hydrogen

Copper

Halogens

Group VII elements are known as Halogens

They are non-metals

They are poisonous

Melting point and Boiling point will increase as it goes down the group because the size

increases, meaning an increase in the strength of the Van der Wall forces holding them

together, causing a higher temperature to be needed to break them

Colour goes darker as it goes down the group

Less reactive as it goes down the group because the bigger the atom, the smaller attraction

between the nucleus and incoming electron

All have similar properties because they all have seven electrons in the outer shell

Reacts with metals to form ionic compounds, containing halide ions

A more reactive halogen will displace a less reactive one from solution

Sulfur

Sulfur – mined from solid underground deposits of elemental sulphur, extracted from fossil fuels,

received from metal sulphide ores when the metal is extracted. It is used in the manufacture of

sulfuric acid.

Sulfur Dioxide – is prepared from when sulphur burns in air or oxygen (burns with a blue flame):

S(s) + O2(g) SO2(g)

Sulfur dioxide dissolves in water to form sulfurous acid (a weak acid) which can lead to the

problem of acid rain

H2O(l) + SO2(g) H2SO3(aq)

Sulfur dioxide can cause bronchiospasm in asthmatics

It is used as a bleaching agent when paper is made from wood pulp

It is used as a preservative for food by killing bacteria.

Sulphites and hydrogen sulphites are also used as preservatives because they liberate SO2 in

solution

Sulfuric Acid – a typical acid used in fertilisers, paints, pigments, dyestuffs, chemical manufacture,

soaps and detergents and fibres.

Contact Process – the industrial preparation of sulfuric acid. All reactions in it are exothermic.

1. Sulfur is burned in air:

S(s) + O2(g) SO2(g)

2. The SO2 is reacted with further oxygen over a catalyst bed (vanadium (V) oxide). The

vanadium (V) oxide is a catalyst which speeds up the reaction without being used up. It

melts at 400 °C, spreading to give a larger area. The yield of SO2 is sufficiently high for this

stage to be carried out at atmospheric pressure.

2 SO2(g) + O2(g) 2 SO3(g)

3. The sulfur trioxide is reacted with 98% H2SO4 to form oleum, which then reacts with water to

form more sulfuric acid. The SO3 must be reacted with sulfuric acid first and not immediately

with water as it is too exothermic/violent to carry out directly.

SO3(g) + H2SO4(l) H2S2O7(l)

H2S2O7(l) + H2O(l) 2 H2SO4(l)

V2O5

Cells

Electrolyte – molten or dissolved metal compounds that conduct electricity. When electrolytes

conduct electricity, ions move.

Electrode – most electrodes are metals or graphite. When metals conduct electricity, valence

electrons move from ion to ion, from the negative to positive electrode. When graphite conducts,

the delocalised electrons between the layers can flow. The two electrodes are called the cathode

(which is negatively charged and attracts cations) and the anode (which is positively charged and

attracts anions). Reduction occurs at the cathode and oxidation occurs at the anode. The mobile ions

of the electrolyte carry the current between the electrodes. Graphite electrodes must be replaced

periodically because graphite will react with oxygen to form CO2.

Electrochemical Cells – produce electricity spontaneously via a chemical reaction (redox) between

two metals.

When two metals of different reactivity are connected electrically in a complete circuit with

a conducting wire and an electrolyte, electrons flow from the more reactive metal to the

least reactive metal.

The electron flow is called current, and the energy

transfer from the higher to lower reactivity metals is

called the voltage.

The greater the difference in reactivity between the two

metals making up the electrodes, the greater the energy

transfer and therefore greater voltage of the cell.

A dry cell uses a damp paste of ionic material (salt

bridge) between the electrodes instead of a liquid

electrolyte.

More than one cell connected together is called a

battery.

The more reactive metal is always the negative electrode.

Hydrogen Fuel Cell

Hydrogen at the anode and Oxygen at the cathode combine to form water. The reduction of

Hydrogen at the anode causes the lost electrons to form a current on their way to reducing oxygen

at the cathode.

2 H2 4 H+ + 4 e-

O2 + 4 e- 2 O2-

2 O2- + 4 H+ 2 H2O

Advantages Disadvantages

Only product is H2O (no CO2) Risk of explosions/gases takes up a lot of volume

Hydrogen is very abundant (in compounds)

It is renewable

Electrolysis

Electrolysis - the passing of a direct current through a conducting solution or liquid and the

resultant decomposition of the electrolyte. It uses electricity from a power source in order to

cause a chemical reaction (redox). It will force the oxidation or reduction of substances that are

high in reactivity that do not naturally oxidise or reduce using normal chemical processes.

+

+-

-

Direction of electron flow

Cations are attracted to the cathode and are discharged (converted to a new substance)

Anions are attracted to the anode and are discharged

The electrolyte is decomposed

The electrodes are usually made of graphite or a completely unreactive metal (e.g. platinum)

Selective Discharge Rules

The ease of discharge of an ion depends on several factors, including the nature of the electrode,

and the nature of the electrolyte (molten/aqueous, concentrated/dilute).

1. Cations always discharge at the cathode

2. Anions always discharge at the anode

3. The ions of the more reactive metals are more difficult to discharge than those of less

reactive metals

4. The sulphate and nitrate ions are never discharged (but not when altered)

5. Halide ions will be difficult to discharge

6. The more concentrated the solution, the more chance the ions will be discharged

Note: if chloride ion is present it is most likely always going to be the anode product. The following

equation represents the discharging of Hydroxide ions:

4 OH- O2 + 2 H2O + 4 e-

Production of Caustic Soda

Caustic Soda – a common name of sodium hydroxide. The electrolysis of a concentrated sodium

chloride solution (brine) produces three products: hydrogen, caustic soda and chlorine.

The Membrane Cell method – the membrane is permeable to cations so only the sodium ions can

flow through and the hydroxide ions cannot flow back to the anode. The membrade is a porous,

thin, flexible sheet.

membrane

chlorine hydrogen

35 % NaOH

water

brine

depleted brine

anode + ve

cathode - ve

The membrane cell for making NaOH

Na +

(Ti or graphite) (steel or graphite)

Anode: 2Cl- Cl2 + 2e- Cathode: 2H2O + 2e- 2OH- + H2

At the cathode, Hydroxide ions are formed (produced together with the hydrogen)

As the sodium ions move towards the cathode, a solution of sodium hydroxide is thereby

formed

The water molecules are reduced at the cathode

The sodium ions are not reduced and apart from moving through membrane do not change

Chlorine is used in PVC and in water

Sodium Hydroxide is used in pulp and paper, and soap.

Production of Aluminium

1. Mine the bauxite ore (a mixture of Al2O3 and SiO2 as well as other impurities such as Fe2O3)

2. Purify the ore by dissolving it in sodium hydroxide solution. This dissolves the alumina

(aluminium oxide) which is amphoteric but not the basic impurities e.g. Fe2O3

3. At the aluminium smelter, the alumina is dissolved in cryolite (Na3AlF6) because this gives

the mixture a much lower melting temperature (900 °C) and it conducts electricity better.

The mixture is 95% alumina and 5% cryolite.

4. Large amounts of electrical energy are passed through the mixture. The anodes are carbon

rods. The cathode lining is graphite in steel casing. The passing of electric current causes

electrolysis to occur.

Anode: Al3+ + 3 e- Al Cathode: 2 O2- O2 + 4e-

Overall: 4 Al3+ + 6 O2- 3 O2 + 4 Al

5. The molten aluminium is poured into ingots and used for many purposes.

Purification of Copper

Copper metal is readily extracted by roasting copper ores malachite (impure CuCO3) and

copper pyrites (CuFeS2) to obtain copper.

Impure copper was at the anode and pure copper was at the cathode

Aqueous copper(II) sulphate was the electrolyte

At the anode, metals more active than copper are oxidised to their cations and remain as

cations and must be removed as they accumulate.

Copper is oxidised to Copper(II), while metals less reactive than copper are not oxidised but

instead fall to the bottom of the cell and are removed through filtration of the electrolyte.

The impurities are mainly Ag and Au and are called the anode sludge.

At the cathode, copper ions are reduced to copper metal (almost 100% pure).

Cu2+ (aq) + Cu (s) Cu (s) + Cu2+ (aq)

Electroplating Metals

Electroplating – the process involving electrolysis to coat one metal with another. Often the purpose

of electroplating is to give a protective coating to the metal beneath or as a decorate coat.

To plate an object with a metal, the object to be electroplated is made the cathode in an

electrolysis cell.

The anode is made from the metal that is to be the coating

The electrolyte will be a salt solution of the metal to be electroplated

Hydrocarbons

Hydrocarbons - a substance made of Hydrogen and Carbon

Homologous Series – a series of carbon compounds differing from each other only by the addition of

more CH2 groups to increase the length of the carbon chains.

Isomers – different forms of the same molecular formula with different structural formulae

Alkanes Alkenes

Non-polar molecules

Weak intermolecular attractions

Low melting point and boiling point (but increases as size increases)

Lower density than water

Saturated Hydrocarbons

Single bonds only

Formula = CnH2n +2

Gas state between 1-4 C’s

Liquid state between 5-17 C’s

Non-polar molecules

Weak intermolecular attractions

Low melting point and boiling point (but increases as size increases)

Unsaturated Hydrocarbons

Contain a double bond

Formula = CnH2n

Extremely reactive due to double bonds breaking

Turns bromine water from red to colourless

e.g. Ethane (C2H6)

e.g. Ethene (C2H4)

Alcohols Carboxylic Acids

Polar molecules (decreases as more carbons added)

Strong intermolecular attractions

Formula = CnH2n +1OH

Colourless volatile liquids

Burn cleanly and efficiently, but with less energy from presence of oxygen

Polar molecules

Formula = CnH2n+1COOH

Strong intermolecular attractions

Weak acids

Formed through the oxidation of Alcohols

Reacts with Alcohols to form Esters

e.g. Ethanol (C2H5OH)

e.g. Ethanoic Acid (CH3COOH)

Alkane Reactions

1. Combustion – alkanes burn in oxygen to form carbon dioxide and water as long as sufficient

oxygen is present; if insufficient, carbon monoxide or carbon will be produced instead of

carbon dioxide

2. Substitution – alkanes will react with halogen molecules in a substitution reaction e.g.

C2H6 + Br2 C2H5Br + HBr

Alkene Reactions

1. Hydrogenation - Addition by hydrogen. Alkanes are formed when the H2 adds to the alkene

molecule. A catalyst of nickel or platinum is used at a temperature of about 150 °C e.g.

CH2 = CH2 + H2 CH3CH3

2. Halogenation – Addition of bromine or other halogens. Halogen alkenes are formed when

halogens attach to the carbons in the double bond by covalent bonds e.g.

CH2 = CH2 + Br2 CH2BrCH2Br

3. Hydration – Addition of water. Alcohols form when water is added to alkene molecules. A

catalyst of dilute H2SO4 or H3PO4 is used.

CH2 = CH2 + H2O CH3–CH2–OH

Number of Carbons

Alkane Alkene Alcohol Carboxylic Acids

1 Methane --- Methanol Methanoic Acid

2 Ethane Ethene Ethanol Ethanoic Acid

3 Propane Propene Propanol Propanoic Acid

4 Butane Butene Butanol Butanoic Acid

5 Pentane Pentene Pentanol Pentanoic Acid

6 Hexane Hexene Hexanol Hexanoic Acid

7 Heptane Heptene Heptanol Heptanoic Acid

8 Octane Octene Octanol Octanoic Acid

9 Nonane Nonene Nonanol Nonanoic Acid

10 Decane Decene Decanol Decanoic Acid

Crude Oil

Fossil Fuel - organic matter (once living) e.g. coal (dead plant matter), oil (dead sea creature

remains).

Crude oil - made of hydrocarbons. It is the result of heat and pressure on plant and (sea) animal

remains over millions of years in the absence of air. This oil (and gas) rises up through permeable

rocks and becomes trapped under impermeable rocks, so they have to be extracted by drilling. The

oil is called crude oil because it is unrefined which makes it of little use as it is hard to transport.

Fractional Distillation - process used to separate a mixture of liquids that have different boiling

points. When the mixture is heated, liquids with low boiling points evaporate and turn to vapour and

can then be separated as liquids. Those with high boiling points remain liquids.

Cracking – allows large hydrocarbon molecules to be broken down into smaller, more useful

hydrocarbon molecules. Fractions containing large hydrocarbon molecules are vaporised and passed

over a hot catalyst. This breaks chemical bonds in the molecules, and forms smaller hydrocarbon

molecules. Cracking is an example of a decomposition reaction. e.g.

Polymers

Polymers - very large molecules made when hundreds of monomers join together to form long

chains. They have no double bonds.

Synthetic Polymers (Plastics) – man-made polymers

Monomers - a molecule that can be bonded to other identical molecules to form a polymer.

Polymerisation - the combining of monomers to form polymers

Addition Polymers – the monomer is thousands of the same alkene molecules, whose double bond is

broken to join the molecules together in one long chain e.g.

If you had n (number) of this monomer:

Then n of the monomers would join together, by breaking the double bond and connecting to the

other monomers to form a long chain:

This would then be written as a repeated unit:

Condensation Polymers – a polymer formed by a condensation reaction (one in which water is given

out).

Artificial Polymers – where a product is formed from two different types of monomers arranged

alternately and linked together. In this case the monomers usually contain a minimum of two of the

same/different functional groups at the end of their molecules. When the polymer forms, a small

molecule such as H2O or HCl is lost at each junction. Artificial polymers include:

Ester Linkage – a polymer found between di-carboxylic acids and a diols. (-COO-) e.g. Terylene

Amide Linkage - a polymer formed between di-carboxylic acids and diamines (-CONH-) e.g. Nylon

Natural Polymers – those polymers found in nature and are usually condensation polymers

Peptide (Amide) Linkage – a protein (polypeptide) formed between amino acid molecules

Polysaccharides – complex carbohydrate molecules made my polymerizing simple sugar

molecules such as glucose

C

H

H

C

H

H

O

H

H

O

C

O

C

O

O

H

H

O

+

C

H

H

C

H

H

O

O

C

O

C

O

n

n

+ nH2O

Ester Link

C

H

H

C

H

H

N

N

H

H

H

H

C

O

C

O

O

H

H

O

+n

n

C

H

H

C

H

H

N

N

H

H

C

O

C

O

+ nH2O

Amide Link

H O H

OHHOn O O OO

+

C

H

H

N

H

O

H

C

O

C

H

H

N

H

O

H

C

O

C

H

N

H

O

H

C

O

C

H

H

N

H

C

O

Amide Link

+ n H2O

+ n H2O

Proteins – polymers of amino acids formed by condensation reactions.

Amino Acids – naturally occurring organic compounds which possess both an –NH2 group and –

COOH group on adjacent carbon atoms. There are 20 naturally occurring amino acids, of which

glycine is the simplest.

Carbohydrates – a group of naturally occurring organic compounds which can be represented by the

general formula Cx(H2O)y

Sugars – any of the class of soluble, crystalline, typically sweet-tasting carbohydrates found in living

tissues and exemplified by glucose and sucrose. They are tested by warming with Benedicts or

Fehlings solutions; if the sugar is present, the colour changes from a blue solution to an orange-red

suspension or precipitate.

Starch – made up of 200-300 glucose monomers. Starch turns iodine solution from red-brown to

blue-black colour.

Cellulose – made up of about 3000 glucose monomers.

Fats and Oils

Fats – naturally occurring polyesters with the same link between ester monomers as Terylene. The

chains typically contain 12 to 20 carbon atoms. Fats and oils are rich in energy and this is their

normal function to us. They are also important in soap and detergent products. A fat molecule is

made of two components, a glycerol (the “backbone” of the molecule) and fatty acids (which are

attached to the backbone.

Soaps – long chains derivatives of fatty acids. The fatty acid is reacted with a base such as caustic

soda. This causes the formation of the sodium salt of the fatty acid which is used as soap. Soap

(sodium stearate) is an ionic compound and can remove dirt

with the covalent end (non-polar) attracting to the dirt and

the ionic end attracting to water molecules. Soap molecules

can make oils and water form a stable emulsion. However, it

forms a scum with hard water by reacting with Ca2+ or Mg2+

present.

Glycerol Fatty Acids

R is the fatty acid

chain

Formation of Triglyceride

Air

Gas Percentage in Air Percentage in Inhaled Air Percentage in Exhaled Air

Nitrogen 79 79 79

Oxygen 20 20 17

Argon 0.9 1 (with other Inert Gases) 1 (with other Inert Gases)

Carbon Dioxide 0.03 Trace 4

Others 0.07 - -

Water - Variable Saturated

Fractional Distillation of Air – the main industrial method of preparation of pure oxygen and nitrogen

Air is liquefied by compression and cooling to below the boiling point of both oxygen and

nitrogen, so that most of the “air” becomes a liquid.

The liquid “air” is allowed to warm slowly and the nitrogen (b.p. -195°C) boils off first and thus

can be extracted.

The oxygen (b.p. -183°C) boils off after the nitrogen and can then also be extracted.

Pollutants Cause Effect Solution

Sulphur Dioxide Combustion of Sulphur (found in fossil fuels)

Forms acid rain when reacted with water in

clouds

Scrubbing – converts SO2 to H2SO4

S(s) + O2(g) SO2(g) SO2(g) + H2O(l) H2SO3(aq)

Nitrogen Oxides Reaction between nitrogen and oxygen at high temp and pressure in motor vehicle engines

Contributes to Acid Rain and are a major component of

photochemical smog

Catalytic Converters

2N2(g) + 3O2(g) 2NO(g) + 2NO2(g)

Carbon Monoxide Incomplete combustion in motor vehicle engines

Strong bonds with haemoglobin,

decreasing the amount of oxygen distributed

around the body

Catalytic Converters

Lead Combustion of petrol containing lead

Brain damage, nervous system problems

Catalytic Converters, using unleaded fuel

Catalytic Converter –catalysts that convert poisonous exhaust fumes into harmless gasses in cars.

The reduction catalyst uses platinum and rhodium to help reduce the NOx emissions by ripping

the nitrogen atom out of the molecule, freeing the oxygen.

2NO N2 + O2 / 2NO2 N2 + 2O2

The oxidation catalyst removes the unburned hydrocarbons

and carbon monoxide by burning them over a platinum and

palladium catalyst.

2CO + O2 2CO2

2NO2 + 4CO 4CO2 + N2

Oxygen

Laboratory Preparation of Oxygen

1. Heating Potassium Manganate (VII)

2 KMnO4 (s) K2MnO4 (s) + MnO2 (s) + O2 (g)

purple green + black + colourless

2. Decomposing hydrogen peroxide by a catalyst of manganese (IV) oxide

2 H2O2 (l) 2 H2O (l) + O2 (g)

Industrial Uses of Oxygen

Steel making

In oxy-acetylene welding

In hospitals and ambulances for treatment of trauma patients

In rockets to combine with the fuel

In deep sea diving helium-oxygen mixtures

Compressed air is avoided in deep sea diving because the nitrogen causes ‘nitrogen narcosis’

when the diver surfaces too rapidly and the nitrogen bubbles out of the blood due to the

rapid decrease in pressure.

Nitrogen

Properties

Very slightly soluble in water

Very unreactive because of its strong intramolecular nitrogen to nitrogen triple bond

Will react with some substances under the right conditions

Has low boiling point of -196°C

Non-polar molecule with weak intermolecular attractions

Has a low melting point of -210°C

Uses

Filling spaces in food packaging and oil tanks

Liquid nitrogen is used for freezing food, gametes, and other delicate materials

Production of Ammonia in the Haber Process

Creating a non-oxidising environment for fruit storage

Ammonia

Preparation

1. Ammonia gas (NH3) can be conveniently prepared in the laboratory by heating together an

ammonium salt with a strong alkali e.g.

Ca(OH)2 (s) + 2 NH4Cl (s) CaCl2 (s) + 2 H2O (l) + 2 NH3 (g)

2. Ammonia can be collected by the downward displacement of air since it is lighter than air

3. Ammonia cannot be collected by the displacement of water because it is very soluble in

water

4. Ammonia can be collected through the Haber Process

Properties of Ammonia

Colourless

Strong choking smell

Less dense than air

Liquefies at -33°C. This makes it easy to transport and store as a liquid

Extremely soluble in water, as it is a polar molecule and can hydrogen bong with itself and

water molecules, to produce an alkaline solution

The only common alkaline gas

Uses of Ammonia

Making nitric acid

Making fertilisers such as urea, ammonium nitrate, and ammonium sulphate

Household cleaners

Dyes

Explosives

Urea Production

Urea – an important nitrogenous (nitrogen containing) fertiliser. It is a white, water-soluble solid.

Urea is less soluble than inorganic fertilisers and so releases the nitrogen slowly to plants; CO(NH2)2.

The process occurs at 190°C and 230 atm.

CO2 (g) + 2 NH3 (g) CO(NH2)2 (s) + H2O (g)

Ammonium Compounds

Ammonia – NH3; polyatomic molecule

Ammonium – NH4+; polyatomic charged ion, only found in ammonium compounds

Ammonium Salts – formed through the reaction of ammonia with the appropriate acid. They are

used as fertilisers to supply nitrogen to plants. Fertilisers are given an NPK rating (Nitrogen,

Phosphorus and Potassium are best used in fertilisers). Ammonia can be displaced from its salts

through decomposition by heating or by the action of strong bases.

Limestone

Limestone – CaCO3; various forms of lime are used to put on pastures to raise the pH because many

soils are naturally acidic. Intensive cropping also lowers the pH. The lime is basic so it neutralises the

soil and brings the pH closer to 7.

Quicklime – CaO; formed from limestone in a lime kiln (oven with extremely high temperatures).

CaCO3 (s) CaO (s) + CO2 (g)

Slaked Lime – Ca(OH)2; the solid product, in a form of white powder, of the exothermic reaction that

occurs when a minimal amount of water is added to quicklime.

CaO (s) + H2O (l) Ca(OH)2 (s)

Limewater – a solution of slaked lime in excess water. The slaked lime is only sparingly soluble but

produces an alkaline solution containing calcium and hydroxide ions.

Ca(OH)2 (s) Ca2+ (aq) + 2 OH- (aq)

Mortar – a mixture of slaked lime, sand and water and is a thick paste. It sets when it dries, then

over a long period of time becomes hard due to the formation of calcium carbonate as it absorbs

carbon dioxide from the atmosphere.

Ca(OH)2 (s) + CO2 (g) CaCO3 + H2O (l)

Cement – made from heating limestone with sand and silicates such as clay. It is a mixture of calcium

silicates and aluminates. When water is added, a complex series of reactions occur which make it

set.

Carbon Cycle

Methane – can be sourced from natural gas trapped in oil-bearing rocks, partial decomposition of

plant materials under anaerobic conditions, waste product of digestion in animals.

Water

Properties

Colourless, odourless liquid

0°C Melting Point

100°C Boiling Point

V shaped or bent; 105° bond angle

Molecules join via hydrogen bonding

Turns anhydrous copper sulphate from white to blue

CuSO4 (s) + 5 H2O (l) CuSO4.5H2O (s) + heat

Turns anhydrous cobalt chloride from blue to pink

CoCl2 (s) + 6 H2O (l) CoCl2.6H2O (s) + heat

Water is used in generating electricity, cleaning, cooling, dissolving, manufacturing products,

making concrete, and consumption in the house (e.g. toilets and showers)

Desiccator – sealed glass basins used to keep substances and papers (e.g. cobalt chloride paper) dry

The Hoffman’s Voltamater

Since water is not ionic, we cannot electrolyse pure

water as there are no charged particles to carry the

current. If we add some dilute sulphuric acid this

dissociates in the water and allows electrolysis to

occur. The apparatus is known as Hoffmann’s

Voltameter.

Cathode: 4 H2O (l) + 4 e- 2 H2 (g) + 4 OH- (aq)

Anode: 2 H2O (l) O2 (g) + 4 H+ (aq) + 4 OH- (aq)

Full: 2 H2O (l) 2 H2 (g) + O2 (g)

Water of Crystallisation

Ionic solids often have molecules of water bonded into their ionic crystal lattice. This water is

called water of crystallisation and often has a consistent, simple ratio in the formula. Formulae

are often quoted with the water of crystallisation included. E.g.Copper sulphate pentahydrate,

Sodium carbonate decahydrate.

Agents

Drying Agents – various drying agents are used to absorb water out of air or gas mixtures. Common

drying agents are:

Concentrated sulphuric acid is used as a drying agent in the preparation of some neutral or

acidic gases

Calcium oxide which is used to dry alkaline ammonia gas

Anhydrous calcium chloride is used to prepare dry hydrogen

Silica gel is used as a drying agent in equipment which is sensitive e.g. cameras

Dehydrating Agents – drying agents that are so powerful that they will remove all atoms required to

make up water from certain substances in the solution of solid form. The most common example of

a dehydrating agent is concentrated sulphuric acid which will dehydrate sucrose and hydrated

copper sulphate. It will also extract the elements of water from cloth or skin.

Purification of Water

Water is stored in dams and reservoirs. It is never completely pure and may contain

bacteria, dissolved substances and solid material which need to be removed.

Concerns over levels of pesticides in river water have led to improvements in water

purification.

Water treatment essentially involves the stages, filtration, sedimentation, (ozone

treatment) and chlorination.

Laboratory Equipment

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