Relative Masses of Atoms and Molecules — Cambridge O Level Chemistry 5070 Revision Notes & Topical Questions

Section 1 of 12Relative atomic mass

Symbol

• Ar, shown beneath the symbol in the Periodic Table and larger than the atomic number.
  • (except hydrogen where they are equal)

Why needed

• Why needed atoms are too small to weigh, so masses are compared.
• Standard
  • carbon-12 has a fixed mass of 12 units.
• Definition
  • Ar is the average mass of the isotopes of an element compared with 1/12 of the mass of a carbon-12 atom.
• Examples
  • carbon 12, magnesium 24 (twice as heavy), hydrogen 1 (one twelfth).

Section 2 of 12Relative molecular and formula mass

Mr

• Mr the total of the relative atomic masses of all atoms in the formula.

Molecules

• Molecules called relative molecular mass.
• Ionic compounds
  • called relative formula mass.

Method

• Method multiply each Ar by the number of atoms, then add.

Water

• H₂O is 2 × 1 + 16 = 18.
• Calcium hydroxide
  • Ca(OH)₂ is 40 + 2 × (16 + 1) = 74.

Section 3 of 12Conservation of mass

Law

• Law mass cannot be created or destroyed, so reactant mass equals product mass.

Use

• Use with Ar and Mr values the balanced equation gives reacting masses.
• Ratio
  • the mass ratio of reactants and products is always the same whatever the units.

Section 4 of 12Reacting masses method

Reacting masses method

Write the balanced equation.

Calculate the relative mass of each substance multiplied by its coefficient.

Scale the ratio to the mass given.

Example 2Ca + O₂ → 2CaO

• 2 × 40 = 80 units of calcium react with 32 units of oxygen to give 112 units of calcium oxide.

Grams

• 80 g Ca + 32 g O₂ gives 112 g CaO.
• Tonnes
  • 40 tonnes of calcium gives 56 tonnes of calcium oxide.

Section 5 of 12Worked pattern: methane

Equation

• CH₄ + 2O₂ → CO₂ + 2H₂O.

Relative masses

• 16 + 64 = 44 + 36.
• Ratio
  • 16 g methane gives 44 g carbon dioxide.
• Scaling
  • 32 g methane gives 44 × 2 = 88 g carbon dioxide.

Section 6 of 12Worked method for Mr

1

Write out the formula.

2

Multiply the Ar of each element by the number of atoms of it.

3

Where brackets appear, multiply everything inside by the number outside.

4

Add the totals.

Carbon dioxide

• CO₂ is 12 + 2 × 16 = 44.

Sodium carbonate

• Na₂CO₃ is 2 × 23 + 12 + 3 × 16 = 106.
• Magnesium nitrate
  • Mg(NO₃)₂ is 24 + 2 × (14 + 48) = 148.

Section 7 of 12Why the Periodic Table value is not always whole

Cause

• Ar is an average of isotopes weighted by abundance.

Chlorine

• 35.5 because natural chlorine is a mixture of Cl-35 and Cl-37.
• Exam use
  • round to the values given in the data sheet and use them as stated.

Section 8 of 12Using reacting masses in practice

Scaling

• Scaling the mass ratio holds for any quantity, so grams can be scaled up to kilograms or tonnes.

Example

• Example if 80 g of calcium needs 32 g of oxygen, 8 g of calcium needs 3.2 g.
• Industry
  • chemical firms use these ratios to decide how much raw material to buy.

Checking

• Checking the total mass of reactants should equal the total mass of products.

Gases

• Gases a product that escapes explains an apparent loss of mass.
• Example
  • heating calcium carbonate loses mass as carbon dioxide leaves.

Section 9 of 12Typical exam questions

Define

• Define state Ar and Mr using the carbon-12 reference.

Calculate

• Calculate add up atoms carefully, including brackets.

Apply

• Apply use the ratio from the balanced equation to find an unknown mass, as in finding how much oxide forms from a given mass of metal.

Layout

• Layout show the equation, the relative masses, the ratio and the final scaled answer.

Section 10 of 12Worked examples

1. Calculate the Mr of ammonium sulfate, (NH₄)2SO₄. Ar: N = 14, H = 1, S = 32, O = 16. [2]2 marks

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2. Calculate the mass of carbon dioxide produced when 32 g of methane burns completely in excess oxygen. [3]3 marks

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3. What mass of calcium oxide forms from 20 g of calcium in excess oxygen? [3]3 marks

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4. Magnesium has an Ar of 24. How many times heavier than a carbon atom is it? [1]1 mark

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Section 11 of 12Key terms

Term	Meaning
Relative atomic mass (Ar)	Average mass of isotopes of an element compared with 1/12 of carbon-12.
Relative molecular mass (Mr)	Sum of the Ar values of all atoms in a molecule.
Relative formula mass	Sum of the Ar values in the formula of an ionic compound.
Conservation of mass	Total mass of reactants equals total mass of products.
Reacting masses	The masses of substances that react or form in the ratio of the balanced equation.

Section 12 of 12Exam technique

1

Saying Ar has units — it is a ratio with no units.

2

Forgetting to multiply by coefficients when using the equation.

3

Not multiplying the atoms inside brackets, such as the 2 in Ca(OH)₂.

4

Using mass numbers instead of Ar for chlorine, which is 35.5.