1. C5a Moles and Empirical Formulae
• The relative atomic mass of an element is the average mass of an atom of the
element compared to the mass of _______ of an atom of _____________.
• It can be found on the periodic table. From the relative atomic mass and atomic
number we can deduce the number of ______ and ______ in the nucleus of an atom.
Relative Formula Mass Calculations – revision
Calculate the relative formula mass of the following compounds (show calculations).
1. NaOH =
2. CaSO4 =
3. Fe2O3 =
4. Na2O =
5. Ca(NO3)2 =
Note that the relative formula mass has no units
Molar Mass Calculations
• The unit for the amount of a substance is the mole.
You don’t need to know anything about moles at the moment. It is just an easy
way for chemists to handle very large numbers of atoms.
A pair of shoes = __ shoes
A trio of apples = __ apples
A dozen apples = __
A mole of apples = _________ i.e. 602 000 000 000 000 000 000 000 apples
• The molar mass of a substance is its relative formula mass in grams.
(it is the same number as the relative formula mass but it has units)
Relative Formula Mass Molar Mass
H2O 18 18 g/mol
NaOH
CaSO4
Fe2O3
Na2O

2. Higher
It is important for chemists to know exactly how much of a chemical to use in a
reaction for two reasons:
 Cost - wasting chemicals costs money
 Safety – some chemicals are more hazardous than others and we must make
sure that they react completely
Consider the following reaction:
We want to react 10 grams of lithium with iodine to make lithium iodide
lithium + iodine lithium iodide
2Li I2 2LiI
10g
Two Problems:
 Lithium reacts vigorously with water and oxygen so we need to make sure that
there is enough iodine to completely react all of the lithium.
 Iodine is hazardous so we don’t want to have too much left over.
1. Work out how many moles (atoms) there are in 10g of Li
The relative atomic mass (from the periodic table) tells us that each lithium atom has a
molar mass of ___ g/mole
moles = mass
Therefore in 10g there must be molar mass
2. Look at the symbol equation above
2 atoms of lithium react with ___ molecule of iodine
2 moles Li __ mole I2
2.5 moles
3. Mass of Iodine needed to react with lithium
The relative atomic mass of iodine is ____
So the molar mass = _____ g/mol mass = moles x molar mass
So 12.5 moles =
= g of iodine
So, to make sure that all 10 grams of lithium reacts, we must have a minimum of 159g
of iodine
4. Supposing we wanted to add a 2% excess of iodine to make certain the lithium has
reacted. How much iodine in total would we need?

3. Practice Questions
1. Determine the number of moles of an element from the mass of that element
moles = mass / molar mass
a) 2 g of hydrogen, H2 f) 184g of sodium
b) 1 g of hydrogen g) 4.6g of sodium
c) 10g of hydrogen h) 32g of oxygen
d) 56g of iron i) 4g of oxygen
e) 14g of iron j) 10.8g of Al
2. determine the number of moles of a compound from the mass of that compound
a) 10g of NaOH
b) 272g of CaSO4
c) 32g Fe2O3
d) 31g Na2O
e) 16.4g Ca(NO3)2
f) 58.5g of NaCl
g) 75g CaCO3
h) 5.1g Al2O3
3. Determine the masses of the different elements present in a given number of
moles of a compound

4. The Law of Conseravtion of Mass
• Matter cannot be created or destroyed.
• Atoms of elements can rearrange to form new substances but atoms cannot
disappear completely.
• In a chemical reaction mass is conserved
• This means that the total mass of the reactants must be equal to the total mass of
the products.
mass of reactants = mass of products
e.g. lithium + iodine lithium iodide
2Li I2 2LiI
10g 159g ____ g
1. Calcium carbonate, CaCO3, decomposes on heating to produce calcium oxide, CaO and
carbon dioxide, CO2.
a) Write a balanced symbol equation for the reaction:
b) If 100g of CaCO3 produces 56g of CaO, how many grams of CO2 are produced?
c) If 100 tonnes of CaCO3 produces 56 tonnes of CaO, how many tonnes of CO2 are
produced?
d) If 300g of CaCO3 produces 168g of CaO, how many grams of CO2 are produced?
e) If 25g of CaCO3 produces 14g of CaO, how many grams of CO2 are produced?
Higher
f) On a separate sheet of paper, calculate out how many moles of CaCO3, CaO and CO2
are produced in b) to e). Do you notice any patterns in terms of the number of moles?
2. 64 g of copper reacts with 16 g of oxygen. What mass of copper oxide forms?
3. 160g of anhydrous copper sulfate, CuSO4, is dissolved in water. When it crystallises
it forms hydrated copper sulfate crystals, CuSO4.5H2O, with a mass of 250g. What
mass of water has been used to make the crystals?

5. Extension: for question 2 a) calculate how many moles of each substance there are and
b) work out the molecular formula for copper oxide.

6. HIGHER – Practice calculations on reacting quantities
1. If you add 5.3g of sodium carbonate, Na2CO3 to sulphuric acid what mass of
sodium sulfate Na2SO4 would be made?
Na2CO3 + H2SO4 Na2SO4 + CO2 + H2O
2. 200 tonnes of calcium carbonate is thermally decomposed to produce calcium
oxide, CaO, and carbon dioxide, CO2. Write a balanced symbol equation and then
calculate how many tonnes of calcium oxide can be made.
CaCO3 CaO + CO2
3. Ammonium nitrate fertiliser is made in the following reaction:
NH3 + HNO3 NH4NO3
How many tonnes of ammonia are needed to make 2400 tonnes of ammonium nitrate?
NH4NO3 is made from NH3
4. How much carbon is needed to reduce 223g of lead oxide?
2PbO + C 2Pb + CO2
5. Calculate the mass of water produced when 1g of hydrogen burns completely in air.
2H2 + O2 2H2O
6. What mass of magnesium is needed to produce 20g of magnesium oxide?
2Mg + O2 2MgO

7. 7. Zinc oxide is heated with carbon to form zinc and carbon monoxide:
zinc oxide + carbon zinc + carbon monoxide
ZnO + C Zn + CO
a) How much zinc oxide do you need to make 130g of zinc?
b) How many moles of ZnO is used?
c) What mass of carbon is needed to react with the zinc oxide?
d) What mass of carbon monoxide will be released?
e) The carbon monoxide reacts with oxygen in the air to make carbon dioxide:
CO + O2 CO2
Using the information above, calculate what mass of carbon dioxide will be
produced. There are 2 ways you could do this, see if you can find both
methods.

8. 8. A student adds 4.8g of magnesium to excess dilute hydrochloric acid. Mg
+ 2HCl MgCl2 + H2
a) What mass of magnesium chloride will be made?
b) How many moles of magnesium are being used?
c) How many moles of hydrochloric acid are needed? (look at the balanced equation)
d) How many moles of hydrogen will be produced?

9. Teacher Copy
Empirical Formulae
An empirical formula gives the simplest whole number ratio of each type of atom in a
compound.
molecular formula empirical formula
e.g. C2H6
C2H4
C6H6
CH4
C6H12O6
Calculating the empirical formula of a compound
1. Write the symbols of all the elements in a line
2. Write the mass of % of each element under their symbol
3. Divide by the molar mass (from periodic table)
4. Divide all the numbers by the smallest one to get a ratio
Example 1: Find the empirical formula of a compound that contains 13.2% lithium,
26.4% nitrogen and 60.4% oxygen
1. Write the symbols of all the elements in a line Li N O
2. Write the mass of % of each element under their symbol
3. Divide by the molar mass
4. Divide all the numbers by the smallest to get a ratio
Empirical formula =
Example 2: Analysis of a sample of sodium oxide found it contained 2.3g of sodium
and 0.8g of oxygen. Find the empirical formula of sodium oxide
1. Write the symbols of all the elements in a line Na O
2. Write the mass of % of each element under their symbol
3. Divide by the molar mass
4. Divide all the numbers by the smallest to get a ratio
Empirical formula =

10. Example 3: Analysis of a sample of mercury oxide found it contained 2.17g of mercury
and 0.16g of oxygen. Find the empirical formula of mercury oxide
1. Write the symbols of all the elements in a line Hg O
2. Write the mass of % of each element under their symbol
3. Divide by the molar mass
4. Divide all the numbers by the smallest to get a ratio
Empirical formula =
Example 4: Find the empirical formula of a compound that contains 31.1% iron,
15.6% nitrogen and 53.3% oxygen
1. Write the symbols of all the elements in a line Fe N O
2. Write the mass of % of each element under their symbol
3. Divide by the molar mass
4. Divide all the numbers by the smallest to get a ratio
Empirical formula =
Extension: Can you suggest a name and formula for this compound?
Example 5: Analysis of a sample of gas found it contained 2.4g of carbon and 0.8g of
hydrogen. Find the empirical formula of this gas and name it.
1. Write the symbols of all the elements in a line C H
2. Write the mass of % of each element under their symbol
3. Divide by the molar mass
4. Divide all the numbers by the smallest to get a ratio
Empirical formula = Name =

11. 1. 7.2g of carbon was reacted 1.2g of hydrogen and 9.6g of oxygen. What is the
empirical formula of glucose?
2. The actual formula is C6H12O6. Calculate the mass of a molecule.
3. What percentage of glucose is carbon?
4. Write a balanced symbol equation for the complete combustion of glucose in
air.
5. a) If 9g of glucose is burnt, what mass of oxygen would be needed for
complete combustion to take place? Use your balanced equation to help you.
b) What mass of water would be produced?

12. Experiment to Measure the Increase in Mass on Complete Oxidation of
Magnesium Ribbon
lid
magnesium crucible
ribbon pipeclay
triangle
heat
heatproof
mat
1 Complete the table of results:
1 1 mass of empty crucible + lid
2 2 mass of crucible + lid + magnesium
3 3 mass of crucible + lid + contents after heating
4 4 mass of magnesium at the start (2−1)
5 mass of magnesium oxide formed (3−1)
6 mass of oxygen in the oxide (5−4)
2 Explain why the mass of the crucible and its contents increased.
3 Explain why it was necessary to raise the crucible lid from time to time.
4 Explain why the crucible lid had to be put back quickly if the magnesium flared up.
5 Calculate the empirical formula for magnesium oxide. (H)
6 Comment on how easy it was to tell if the formula for magnesium oxide is Mg2O, MgO
or MgO2. (H)
7 Look up the formula of magnesium oxide in a textbook or data book. Suggest why the
amount (in moles) of oxygen combining with the magnesium is lower than expected
from the formula.
8 What other metals, if any, could have their empirical formula found by this method?
(H)

13. Experiment to Measure the Decrease in Mass on Reduction of
Copper Oxide
burning
gas clamp here
gas from
gas tap
black copper
oxide tube must be kept
heat horizontal
1 Complete the table of results.
1 mass of empty tube
2 mass of tube + copper oxide before heating
3 mass of tube + copper after heating
4 mass of copper oxide before heating (2−1)
5 mass of copper formed (3−1)
6 mass of oxygen in the oxide (4−5)
2 State the name the brownish pink solid formed at the end of the experiment.
3 Why did the mass of the copper oxide decrease?
4 What happened to the oxygen from the copper oxide?
5 Suggest why the gas from the gas tap is kept flowing through the tube after the
Bunsen burner has been removed. (H)
6 Calculate the empirical formula for black copper oxide. (H)

14. 7 What type of chemical reaction is taking place? (H)