Relative masses of atoms and molecules — Combined Science 0653 - Chemistry Revision Notes & Topical Questions

Section 1 of 21Relative atomic mass, Ar

Ar

• = average mass of the isotopes of an element compared with 1/12th of the mass of an atom of carbon-12.

Standard

• Standard carbon-12 is the standard atom with a fixed mass of 12 units.

Values

• Values magnesium has Ar = 24, so it is twice as heavy as carbon;
  • hydrogen has Ar = 1, one-twelfth of a carbon-12 atom.

Units

• Ar has no units because it is a ratio.

Section 2 of 21Finding Ar on the Periodic Table

Position

• Position the relative atomic mass is shown on the atomic symbol along with the atomic number.

Size

• Ar is always larger than the atomic number except for hydrogen, where they are the same.

Key

• Key use the key on the table to identify which number is Ar.

Section 3 of 21Relative molecular mass and relative formula mass

Mr

• = sum of the relative atomic masses of all the atoms in a molecule.

Relative formula mass

• Relative formula mass is the term used for the total mass of an ionic compound.

Section 4 of 21Calculating Mr

Calculating Mr

Write the formula and count the atoms of each element.

Multiply each Ar by the number of atoms.

Add the results.

• Brackets
  • multiply everything inside brackets by the number outside.

Section 5 of 21Mr of common substances

Substance	Formula	Calculation	Mr
Hydrogen	H2	2 × 1	2
Water	H2O	(2 × 1) + 16	18
Potassium carbonate	K2CO3	(2 × 39) + 12 + (3 × 16)	138
Calcium hydroxide	Ca(OH)2	40 + (2 × 16) + (2 × 1)	74
Ammonium sulfate	(NH4)2SO4	(2 × 14) + (8 × 1) + 32 + (4 × 16)	132

Section 6 of 21Reacting masses in simple proportions

Conservation of mass

• Conservation of mass in a chemical reaction the total mass of reactants equals the total mass of products.

Fixed ratio

• Fixed ratio the ratio of the masses that react is always the same, whatever the units.
• 2Ca + O₂ → 2CaO gives 80 : 32 : 112, so 80 g of calcium reacts with 32 g of oxygen to give 112 g of calcium oxide.
• The same ratio works for kilograms and tonnes.

Section 7 of 21More from the specification

More from the specification

Relative molecular mass determination

• Relative molecular mass determination summing standard relative atomic masses of all formula atoms provides a means to calculate molecular weight.

Chloride

• Salt of hydrochloric acid.

Nitrate

• Salt of nitric acid.

Empirical formula

• The simplest whole-number ratio of atoms or ions.

Carbon dioxide

• A gas that turns limewater cloudy.

Sodium chloride

• Compound of sodium and chlorine; common salt.

Simple molecule

• Small group of atoms joined by covalent bonds, with weak forces between molecules.

Ionic formula

• Ionic formula ions combine in the ratio that makes the total charge zero, so Al³⁺ and Cl⁻ give AlCl₃.

Balanced equation

• Balanced equation has the same number of each type of atom on both sides, which is why total mass is conserved.

Section 8 of 21Why equations are balanced

1Law of conservation of mass

• Law of conservation of mass atoms are neither created nor destroyed, so each element has the same number of atoms on both sides.

2Ratio

• Ratio the balancing numbers give the correct ratio of reactants to products.

Section 9 of 21Balancing method

1

Write the correct formulae for all reactants and products.

2

Compare the numbers of atoms of each element on both sides.

3

Change the coefficients in front of formulae one at a time.

4

Recount all elements after every change.

5

Balance elements that appear on their own last.

Balancing method

Subscripts

• Subscripts never change the small numbers, because that changes the substance.

Groups

• Groups treat an unchanged group such as NO₃ as one unit.

Section 10 of 21Typical balanced equations

Type	Equation
Metal oxide + acid	MgO(s) + 2HNO3(aq) → Mg(NO3)2(aq) + H2O(l)
Displacement	2Al(s) + 3CuO(s) → Al2O3(s) + 3Cu(s)
Combustion of methane	CH4 + 2O2 → CO2 + 2H2O

Section 11 of 21Checking an equation

Count

• Count list each element on both sides, including atoms inside brackets.
• In Mg(NO₃)₂ there are 1 Mg, 2 N and 6 O.

Order

• Order check state symbols last.

Section 12 of 21Linking the equation to relative masses

Balancing numbers

• Balancing numbers multiply each Mr by its coefficient to compare masses.
• In 2Al + 3CuO → Al₂O₃ + 3Cu the Mr values of CuO (79.5) and Al₂O₃ (102) are multiplied by 3 and 1.

Use

• Use balance the equation first, before using relative masses.

Section 13 of 21More from the specification

1Stoichiometric equation balance

• Stoichiometric equation balance conserving atomic counts on both reactant and product sides provides a fundamental means to write valid chemical equations.

2Relative atomic mass (Ar) standard

• Relative atomic mass is the weighted mean mass of naturally occurring isotopes of an element relative to 1/12th of the mass of an atom of carbon-12.

3Relative isotopic mass and mass spectrometry

• Mass spectra indicate relative abundance and precise m/z values for individual isotopes, used to compute average relative atomic masses.

4Avogadro's constant and molar mass

• One mole of any substance contains exactly 6.022 x 10²³ elementary entities, with molar mass M in g/mol numerically equal to Mr.
  • (Avogadro's constant, N_A)

5Balancing ionic equations

• Balancing ionic equations ionic compounds separate into aqueous ions whose positive and negative charges balance to give neutral overall formula units.

6Combustion of hydrocarbons

• Combustion of hydrocarbons hydrocarbon fuels containing carbon and hydrogen react with oxygen gas,
  • forming carbon dioxide and water vapour as products.

7Sodium chloride synthesis

• Sodium chloride synthesis neutral sodium chloride forms when reactive sodium metal reacts with toxic chlorine gas, satisfying atomic conservation of mass.

8Nitrate

• Salt of nitric acid.

9Ionic compound

• A compound made of positive and negative ions held together by electrostatic attraction.

10Empirical formula

• The simplest whole-number ratio of atoms or ions.

11Carbon dioxide

• A gas that turns limewater cloudy.

12Sodium chloride

• Compound of sodium and chlorine; common salt.

13Simple molecule

• Small group of atoms joined by covalent bonds, with weak forces between molecules.

14Hydrogen

• A gas that gives a squeaky pop with a lighted splint.

15Oxygen

• A gas that relights a glowing splint.

16Compound

• Two or more elements chemically combined in a fixed ratio.

17Chloride

• Salt of hydrochloric acid.

18Charge

• Q = It, measured in coulombs.
  What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

19Sulfate

• Salt of sulfuric acid.

20Hydroxide

• A compound containing the OH⁻ ion, formed by Group I metals with water.

Section 14 of 21Relative atomic mass and relative molecular mass

Ar

• Ar is the relative atomic mass shown on the Periodic Table, larger than the atomic number except for hydrogen.

Mr

• = sum of the Ar of all atoms in a molecule.

Relative formula mass

• Relative formula mass is used for ionic compounds.

Section 15 of 21Conservation of mass in calculations

Rule

• Rule the total mass of reactants equals the total mass of products.

Use

• Use with relative masses it gives the quantities of reactants or products.

Section 16 of 21Method for a reacting mass calculation

1

Write the balanced symbol equation.

2

Calculate the Mr of each substance involved.

3

Multiply each Mr by its coefficient.

4

Scale the ratio up or down to the mass given.

5

Check the masses of reactants add up to the masses of products.

Section 17 of 21Illustration with methane

Substance	CH4	2O2	CO2	2H2O
Mass units	16	64	44	36

Totals

• 16 + 64 = 80 on the reactant side and 44 + 36 = 80 on the product side.

Scaling

• Scaling doubling the mass of methane doubles the mass of carbon dioxide formed.

Section 18 of 21Ions in solution

• Dissociation
  • ionic compounds in solution separate into their ions.
  • HCl → H⁺ + Cl^-
  • KOH → K⁺ + OH^-

Section 19 of 21More from the specification

More from the specification

Relative formula mass calculation

• Relative formula mass calculation summing the unified atomic weights of all constituent elements provides an exact means to evaluate molecular formula mass.

Empirical Formula Determination

• The simplest whole-number ratio of atoms in a chemical compound, calculated by dividing the mass of each element by its relative atomic mass (Ar).

Balancing Ionic Formula Charges

• The chemical formula of an ionic compound must have an overall net charge of zero, balancing positive cations with negative anions.
  • (like Na⁺ or Cu²⁺)

Formulas of Polyatomic Ions

• Common compound anions including sulfate (SO₄²⁻), nitrate (NO₃⁻),
• and hydroxide (OH⁻) require brackets in formulas when more than one group is present.

Relative Formula Mass (Mr)

• Calculated by adding together the relative atomic masses of all atoms indicated in the formula unit (e.g. copper(II) chloride CuCl₂ or sodium sulfate Na₂SO₄).

Empirical Formula Determination

• The simplest whole-number ratio of atoms in a chemical compound, calculated by dividing the mass of each element by its relative atomic mass (Ar).

Balancing Ionic Formula Charges

• The chemical formula of an ionic compound must have an overall net charge of zero, balancing positive cations with negative anions.
  • (like Na⁺ or Cu²⁺)

Formulas of Polyatomic Ions

• Common compound anions including sulfate (SO₄²⁻), nitrate (NO₃⁻),
• and hydroxide (OH⁻) require brackets in formulas when more than one group is present.

Relative Formula Mass (Mr)

• Calculated by adding together the relative atomic masses of all atoms indicated in the formula unit (e.g. copper(II) chloride CuCl₂ or sodium sulfate Na₂SO₄).

Empirical formula

• The simplest whole-number ratio of atoms or ions.

Sodium chloride

• Compound of sodium and chlorine; common salt.

Simple molecule

• Small group of atoms joined by covalent bonds, with weak forces between molecules.

Oxygen

• A gas that relights a glowing splint.

Chloride

• Salt of hydrochloric acid.

Charge

• Q = It, measured in coulombs.
  What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

Sulfate

• Salt of sulfuric acid.

Hydroxide

• A compound containing the OH⁻ ion, formed by Group I metals with water.

Nitrate

• Salt of nitric acid.

Element

• A substance made of only one type of atom.

Section 20 of 21Key terms

Term	Meaning
Relative atomic mass	Average mass of an element's isotopes compared with 1/12 carbon-12.
Relative molecular mass	Sum of the Ar of all atoms in a molecule.
Relative formula mass	Sum of Ar for an ionic compound's formula.
Isotope	Atom of the same element with a different number of neutrons.
Conservation of mass	Total mass is unchanged in a reaction.

Term	Meaning
Balanced equation	Equation with equal numbers of each atom on both sides.
Coefficient	Number in front of a formula.
Subscript	Small number in a formula.
Relative formula mass	Sum of Ar of atoms in the formula.
Conservation of mass	Total mass of reactants equals total mass of products.

Term	Meaning
Ar	Relative atomic mass.
Mr	Relative molecular mass.
Relative formula mass	Mr of an ionic compound.
Reacting mass	Mass of reactant that reacts or product formed.
Dissociate	Separate into ions in solution.

Section 21 of 21Exam technique

1

Adding atomic numbers instead of Ar.

2

Forgetting to multiply by the subscript or the bracket number.

3

Giving Ar or Mr units.

4

Using O instead of O_{2} for oxygen molecules.

1

Changing small subscript numbers to balance.

2

Using relative masses before balancing the equation.

3

Writing the products first.

4

Forgetting to balance elements that appear alone last.

1

Adding atomic numbers instead of Ar.

2

Forgetting to multiply by the subscript.

3

Using the mass of one O atom instead of O_{2}.

4

Mixing up Ar and Mr.