1 of 45© Boardworks Ltd 2009. 2 of 45© Boardworks Ltd 2009

1 of 45 © Boardworks Ltd 2009



How many particles?

Scientists often need to know how many particles of a substance they are dealing with.

Instead of counting, scientists use the relative masses of different atoms and molecules to measure how many are present in a given sample.

…thirty-four million and

one…

Unfortunately, atoms and molecules are tiny and it is impossible to count them directly.


How many particles?

For example:

One carbon atom has a relative atomic mass of 12.

The quantity of atoms in 24 g of pure graphite can be calculated as 24

12= 2

So, 24 g of carbon contains twice as many atoms as 24 g of magnesium, because magnesium atoms are twice as heavy.The unit for these results is called a mole.

One magnesium atom has a relative atomic mass of 24.

The quantity of atoms in 24 g of pure magnesium can be calculated as 24

24= 1


What are moles?

For example, the relative atomic mass of carbon is 12, so one mole of carbon atoms weighs 12 g.

24 g of pure graphite therefore contains 2 moles of carbon atoms.

The term mole is used to describe how many atoms or molecules of a substance are present in a sample.

relative atomic mass

One mole of a substance is equal to its relative atomic mass, or relative formula mass, in grams.


How many moles?


How much is in a mole?

Scientists have calculated that one mole of any substance contains 602,000,000,000,000, 000,000,000 (6.02 × 1023) particles.

One mole of carbon weighs 12 g, so 12 g of carbon contains 6.02 × 1023 carbon atoms.

How many carbon atoms are in 6 g of carbon?

One mole of sodium weighs 23 g, so 23 g of sodium contains 6.02 × 1023 carbon atoms.

One mole of water weighs 18 g, so 18 g of water contains 6.02 × 1023 water molecules.

3.01 × 1023


Avogadro’s number

The number of particles in one mole is called Avogadro’s number.

If you collected 6.02 × 1023 marbles and spread them over the surface of the Earth, they would form a layer of marbles 50 miles thick!

It is named after Amedeo Avogadro, an Italian scientist working in the early 19th century.

6.02 × 1023 really is a staggeringly large number of particles.

crustmantle

marbles


Calculating moles

The number of moles in a sample is calculated using this equation:

mass of substance is measured in grams.

some elements exist as molecules (for example, O2). You will need to use the molecular rather than atomic mass for these.

for compounds, the molar mass is the sum of the relative atomic masses of all the atoms in the formula.

number of moles = mass of substancemolar mass

Remember:


Calculating molar mass


Calculating moles: questions

1. How many moles of iron are in 28 g of pure iron?

Number of moles = = 0.5 moles of iron56

28

2. How many moles of CO2 are in 11 g of carbon dioxide?

Relative molecular mass = 12 + 16 + 16 = 44

Number of moles = = 0.25 moles of CO244

11

number of moles = mass of substancemolar mass


Using moles in calculations

Moles are very useful when reacting substances together. They provide us with a quick way of calculating how many particles of each substance are present.

Chemists use moles to calculate:

how much of each reactant they will need

how much of a substance they are likely to make in the end.


Using moles: example 1

If you reacted 56 g of iron with excess copper sulfate solution, what mass of copper would you expect to make?

mass = moles × RAM

Step 1: Write a balanced equation for the reaction:

Fe + CuSO4 Cu + FeSO4

Step 2: Turn the mass of iron into moles of iron:

moles = = = 1 mole of ironmass

molar mass5656

Step 3: Use the balanced equation to work out how many moles of copper you should make:

1 mole of Fe makes 1 mole of Cu

Step 4: Turn moles of copper into mass of copper:

= 63.5 g of Cu= 1 × 63.5


Using moles: example 2

What mass of iron oxide would you need to react with excess aluminium to make 1120 g of iron?

mass = moles × RAM


Fe2O3 + 2Al Al2O3 + 2Fe

Step 2: Turn mass of iron into moles of iron:

moles = = = 20 moles of ironmassRAM

1,12056

Step 3: Use the balanced equation to work out how many moles of iron oxide you would need:

1 mole of iron oxide makes 2 moles of iron, so 10 moles of iron oxide are needed to make 20 moles of iron.

Step 4: Turn moles of iron oxide into mass of iron oxide:= 1,600 g of Fe2O3= 10 × 160


Using moles in calculations



Amedeo Avogadro


Avogadro’s law

Avogadro’s law states that:

What would happen to the volume of a mole of gas if you:

Equal volumes of gases, at the same temperature and pressure, contain the

same number of molecules.

This means that one mole of any gas at a given temperature and pressure will always have the same volume.

At room temperature and pressure (RTP), one mole of any gas has a volume of 24 dm3 (24 litres).

increased the temperature? increased the pressure?


Gases and moles


Gases and moles


Gases and moles


Gas calculations: example 1

Hydrogen gas reacts with oxygen gas to make water vapour. If 24 dm3 of hydrogen gas is burned with excess oxygen, what volume of water vapour is produced?

volume = moles × 24


2H2 (g) + O2 (g) 2H2O (g)

Step 2: Turn volume of hydrogen into moles of hydrogen:

moles of H2 = = = 1 mole of H2volume

242424

Step 3: Use the balanced equation to work out how many moles of water vapour you should make:

1 mole of hydrogen makes 1 mole of water vapour

Step 4: Turn moles of water into volume of water vapour:= 24 dm3 of water vapour

= 1 × 24


Gas calculations: example 2

Nitrogen and hydrogen react together to make ammonia. What is the maximum volume of ammonia that can be made from 2,400 cm3 of nitrogen gas?


N2 (g) + 3H2 (g) 2NH3 (g)

Step 2: Turn volume of nitrogen into moles of nitrogen:

moles of N2 = volume24

2.424

NOTE: To convert cm3

to dm3, divide by 1,000.

= 0.1 moles of N2

=


Gas calculations: example 2 continued

Step 3: Use the balanced equation to work out how many moles of ammonia you can make:

1 mole of N2 makes 2 moles of NH3, so 0.1 moles of N2 makes 0.2 moles of NH3.

Step 4: Turn moles of ammonia into volume of ammonia:

volume = moles × 24

= 4.8 dm3 (4,800 cm3) of ammonia

= 0.2 × 24


Gas calculations



Using gases to follow reactions

Many chemical reactions produce gases.

Scientists can measure how much gas is produced and use it to follow the progress of a reaction.

The volume of gas produced can also be converted into moles and used to determine the amount of reactants present at the start of the reaction.

Which gas is produced by each of these reactions?

Magnesium + hydrochloric acid

Calcium carbonate + hydrochloric acid

Decomposition of hydrogen peroxide

hydrogen

carbon dioxide

oxygen


Measuring gases in reactions

One way of measuring how much gas is produced in a reaction is to use an upturned measuring cylinder filled with water.

What would be the problem with using this method for a gas which is very soluble in water?

As the gas bubbles into the measuring cylinder, it displaces the water.

This technique works well for gases which are less dense than air and not very soluble in water.



Another method of collecting a gas is to use a gas syringe.This technique works well for all gases.

What are the advantages and disadvantages of using a gas syringe rather than a measuring cylinder to collect the gas?



One other method of measuring how much gas is produced in a reaction is to measure the mass of the reaction mixture.

What will happen to the reading on the balance?

Why is cotton wool placed in the top of the flask?


Limiting reactants

The amount of gas produced in a reaction will depend on the limiting reactant. This is the reactant that runs out first, bringing the reaction to an end.

The other reactant(s) are said to be “in excess”. This means that there will be some left after the reaction.

This table shows hydrogen gas production during a reaction between a small amount of magnesium and excess dilute hydrochloric acid.

time (s) volume H2 (cm3)

0306090

120150180210240270300

After how many seconds was all the magnesium used up?

0150300450600700780850895910910


Reaction graphs


Following reactions using gases



Empirical and molecular formulae

The molecular formula of a compound tells you how many of each type of atom are present in one molecule.

The molecular formula of ethane is C2H6.

The empirical formula of a compound gives you the simplest whole number ratio of the types of atoms in one molecule.

The empirical formula of ethane is CH3.

Ethane: C2H6

Sometimes the empirical and the molecular formulae of a compound are the same, like for water, H2O.

Water: H2O


Empirical and molecular formulae


Why are empirical formulae useful?

Chemists sometimes make compounds whose identity they are unsure of. Finding the empirical formula of an unknown compound is an important step in identifying it.

The empirical formula can often be determined by studying the reactions of the unknown compound.

For example, an unknown hydrocarbon can be burned in oxygen to produce carbon dioxide and water vapour.

Chemists can measure the volumes of the gases produced by the reaction and use them to work out the empirical formula.


How are empirical formulae found?

oxygen in oxygen out

unknown compound

calcium chloride

potassium hydroxide solution

The unknown compound is heated in a stream of pure, dry oxygen, reacting to form steam and CO2 gas.

The steam is absorbed by a known mass of calcium chloride. The CO2 travels on and reacts with the known mass of potassium hydroxide solution.


How are empirical formulae found?

The change in mass of the calcium chloride and the potassium hydroxide can be used to calculate the relative masses of carbon and hydrogen in the original compound.

For example:If the calcium chloride increased in mass by 27 grams and the potassium hydroxide solution increased by 66 grams…

27 g of H2O was made. H2O = 1 + 1 + 16 = 18.

There are 2 g of hydrogen in every 18 g of water.

× 27 = 3 g of hydrogen were produced.

66 g of CO2 was made. CO2 = 12 + 16 + 16 = 44.

There are 12 g of carbon in every 44 g of carbon dioxide.

× 66 = 18 g of carbon were produced.

218

1244


Empirical formulae from masses


Calculating from masses or percentages



Glossary


Anagrams


Multiple-choice quiz

Documents

1 of 45© Boardworks Ltd 2009. 2 of 45© Boardworks Ltd 2009