AQA GCSE Chemistry 8462: A simple model of the atom, symbols, relative atomic mass, electronic charge and isotopes – Study Guide

This guide teaches section 4.1.1 A simple model of the atom, symbols, relative atomic mass, electronic charge and isotopes of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 onwards and exams in 2018 onwards (version 1.1). It covers every statement from 4.1.1.1 to 4.1.1.7. This content is examined on Paper 1 at Foundation and Higher Tier, and Paper 2 can draw on it too. One statement is marked (HT only) in the specification; it is labelled Higher tier only below. No required practical sits in 4.1.1.

For quick recall, use the revision notes for this section. To test yourself, use the practice questions for this section. For the rest of topic 4.1, see the topic overview, topic revision notes and topic practice questions. The course hub is AQA GCSE Chemistry, and the printable checklist lists every specification point.

What this section covers

Spec point	What you must be able to do	Tier
4.1.1.1 Atoms, elements and compounds	Use names and symbols; name compounds from formulae; write word equations, formulae and balanced equations	Both
4.1.1.1 (HT only)	Write balanced half equations and ionic equations	Higher tier only
4.1.1.2 Mixtures	Describe, explain and give examples of separation processes; suggest a technique from given information	Both
4.1.1.3 Development of the model of the atom	Describe why scattering evidence changed the model; compare plum pudding and nuclear models	Both
4.1.1.4 Relative electrical charges	Use the nuclear model to describe atoms	Both
4.1.1.5 Size and mass of atoms	Use nm and standard form; find protons, neutrons and electrons in atoms and ions; relate atomic scale to everyday objects	Both
4.1.1.6 Relative atomic mass	Calculate relative atomic mass from percentage abundances	Both
4.1.1.7 Electronic structure	Show the electronic structure of the first 20 elements as numbers and as a diagram	Both

4.1.1.1 Atoms, elements and compounds

All substances are made of atoms. An atom is the smallest part of an element that can exist. Each element has a chemical symbol: O is an atom of oxygen, Na is an atom of sodium. There are about 100 elements, and they are shown in the periodic table. You are given a periodic table in the exam, but must know the names and symbols of the first 20 elements, Groups 1 and 7, and other elements in the specification.

A compound contains two or more elements chemically combined in fixed proportions. Compounds are formed from elements by chemical reactions, and they can only be separated back into elements by chemical reactions. A reaction always forms a new substance and often has a detectable energy change.

Naming compounds from formulae

• Metal first, non-metal second.
• A non-metal on its own changes its ending to -ide: chloride, oxide, sulfide.
• Common groups keep their own names: CO₃ is carbonate, SO₄ is sulfate, NO₃ is nitrate, OH is hydroxide.

Worked example 1: name and count

Name Na₂CO₃ and Mg(NO₃)₂, and count the atoms in one formula unit of each.

Na2CO3:     sodium carbonate
            Na = 2, C = 1, O = 3          total = 6 atoms

Mg(NO3)2:   magnesium nitrate
            the bracket is doubled:
            Mg = 1, N = 1 x 2 = 2, O = 3 x 2 = 6
                                          total = 9 atoms

A number after a bracket multiplies everything inside it.

Writing and balancing equations

1. Write the word equation.
2. Replace each name with its correct formula. Never change a formula to make it balance.
3. Balance by putting numbers in front of formulae until each element has the same number of atoms on both sides.

Worked example 2: balancing

Aluminium burns in oxygen to form aluminium oxide, Al₂O₃.

word:        aluminium + oxygen -> aluminium oxide
unbalanced:  Al + O2 -> Al2O3
balance O:   O2 has 2, Al2O3 has 3; lowest common multiple 6
             Al + 3O2 -> 2Al2O3
balance Al:  right side now has 4 Al
             4Al + 3O2 -> 2Al2O3
check:       Al 4 = 4,  O 6 = 6

Ionic and half equations (Higher tier only)

An ionic equation shows only the particles that change. Ions that are the same on both sides are spectator ions; leave them out. A half equation shows one species gaining or losing electrons, written e⁻. Both kinds must balance in atoms and in charge.

Worked example 3 (Higher tier only)

(a) Hydrochloric acid reacts with sodium hydroxide solution. Write the ionic equation.

H+ Cl-  +  Na+ OH-  ->  Na+ Cl-  +  H2O
Na+ and Cl- are unchanged, so remove them:
H+ + OH- -> H2O        charge: (+1) + (-1) = 0 on each side

(b) Potassium reacts with chlorine to form potassium chloride. Write the two half equations.

K -> K+ + e-           (each potassium atom loses one electron)
Cl2 + 2e- -> 2Cl-      (each chlorine molecule gains two electrons)

Lost electrons go on the right; gained electrons go on the left. Topic 4.4 develops this; see the chemical changes study guide.

4.1.1.2 Mixtures

A mixture consists of two or more elements or compounds not chemically combined. Each substance keeps its own chemical properties. Mixtures are separated by physical processes. These involve no chemical reaction and make no new substances.

Process	What it separates	Property used
Filtration	Insoluble solid from a liquid	The solid cannot pass through the filter paper
Crystallisation	Dissolved solid from its solution	The solvent evaporates; the solute does not
Simple distillation	Solvent from a solution	The solvent boils at a much lower temperature than the solute
Fractional distillation	Liquids that mix, with different boiling points	Each liquid condenses at its own boiling point up the column
Chromatography	Dissolved substances, such as dyes	Different attraction to the paper and solubility in the solvent

Crystallisation. Heat to evaporate some solvent until the solution is saturated, leave it to cool so crystals form, then filter and dry them. Simple distillation. The solvent boils, condenses in a condenser and is collected; the solute stays in the flask.

Worked example 4: suggesting a technique

A sample of muddy sea water contains sand, dissolved salt and water. Suggest how to obtain pure water.

1. Filter first to remove the sand. Sand is insoluble, so it stays on the filter paper.
2. Distil the filtrate (simple distillation). Water boils off and is condensed; salt stays in the flask.

Name the technique and the property that makes it work. Chromatography returns in topic 4.8 (required practical 6); see the chemical analysis study guide.

4.1.1.3 The development of the model of the atom

New experimental evidence may lead to a scientific model being changed or replaced.

1. Before the electron was discovered, atoms were thought to be tiny spheres that could not be divided.
2. The discovery of the electron led to the plum pudding model: a ball of positive charge with negative electrons embedded in it.
3. The alpha particle scattering experiment led to the conclusion that the mass of an atom is concentrated at the centre (the nucleus) and that the nucleus is charged. This nuclear model replaced the plum pudding model.
4. Niels Bohr adapted the nuclear model: electrons orbit the nucleus at specific distances. His theoretical calculations agreed with experimental observations.
5. Later experiments showed that the positive charge of a nucleus could be subdivided into a whole number of smaller particles, each with the same positive charge. These were named protons.
6. James Chadwick’s experimental work gave evidence for neutrons in the nucleus, about 20 years after the nucleus became an accepted idea.

Details of Bohr’s and Chadwick’s experimental work are not required.

Plum pudding versus nuclear model

Plum pudding model	Nuclear model
Positive charge spread through the whole atom	Positive charge concentrated in a tiny nucleus
Mass spread through the whole atom	Almost all the mass in the nucleus
Electrons embedded in the positive ball	Electrons outside the nucleus
No empty space	Mostly empty space

Worked example 5: why the model changed

Explain why the results of the scattering experiment led to the plum pudding model being replaced.

• In the plum pudding model, charge and mass are spread thinly, so nothing could strongly deflect an alpha particle.
• Some were deflected through large angles. That needs concentrated positive charge and mass: a nucleus.
• The old model could not explain this, so it was replaced.

4.1.1.4 Relative electrical charges of subatomic particles

Particle	Relative charge
Proton	+1
Neutron	0
Electron	-1

An atom has equal numbers of electrons and protons, so it has no overall charge. The number of protons is the atomic number: the same for every atom of one element, different for different elements.

4.1.1.5 Size and mass of atoms

• Atomic radius: about 0.1 nm = 1 × 10⁻¹⁰ m (1 nm = 1 × 10⁻⁹ m).
• Nuclear radius: less than 1/10 000 of the atom’s radius, about 1 × 10⁻¹⁴ m.
• Almost all the mass of an atom is in the nucleus.

Particle	Relative mass
Proton	1
Neutron	1
Electron	Very small

The mass number is the number of protons plus neutrons. Isotopes are atoms of the same element with different numbers of neutrons. An atom is written with the mass number top left and the atomic number bottom left, for example ²³₁₁Na.

protons   = atomic number
neutrons  = mass number - atomic number
electrons = protons - charge on the ion   (atoms: charge 0)

Worked example 6: protons, neutrons and electrons

Particle	Protons	Neutrons	Electrons
³¹₁₅P	15	31 - 15 = 16	15
⁷⁹₃₅Br⁻	35	79 - 35 = 44	35 + 1 = 36
⁵⁶₂₆Fe³⁺	26	56 - 26 = 30	26 - 3 = 23

A negative ion has gained electrons; a positive ion has lost them. The nucleus never changes.

Worked example 7: size and scale

(a) A line of atoms, each 0.2 nm across (twice the radius), is 1 mm long. How many atoms are in the line?

diameter = 2 x 1 x 10^-10 m = 2 x 10^-10 m
1 mm     = 1 x 10^-3 m
number   = (1 x 10^-3) / (2 x 10^-10) = 5 x 10^6 atoms

(b) A model atom is built with a radius of 10 m. How big should its nucleus be?

nuclear radius = 10 m / 10 000 = 0.001 m = 1 mm

4.1.1.6 Relative atomic mass

The relative atomic mass (Ar) of an element is an average that takes account of the abundance of its isotopes.

Ar = (mass 1 x % 1 + mass 2 x % 2 + ...) / 100

Worked example 8

A sample of copper is 69% copper-63 and 31% copper-65. Calculate the relative atomic mass to 1 decimal place.

Ar = (63 x 69 + 65 x 31) / 100
   = (4347 + 2015) / 100
   = 6362 / 100
   = 63.62  ->  63.6

Check: the answer is nearer 63, the more abundant isotope.

4.1.1.7 Electronic structure

Electrons occupy the lowest available energy levels (innermost available shells). For the first 20 elements, the first shell holds 2, the second 8, the third 8, and the fourth takes the rest. You can answer in terms of shells or energy levels.

Worked example 9

Give the electronic structure of silicon (atomic number 14) and potassium (19), and describe the diagram for silicon.

silicon:    14 electrons -> 2, 8, 4
potassium:  19 electrons -> 2, 8, 8, 1

Silicon’s diagram: a nucleus with three circles round it, carrying 2, 8 and 4 crosses from the inside out.

Common errors

• Changing a small number inside a formula to balance an equation. Only add numbers in front.
• Counting atoms in Mg(NO₃)₂ as 1 + 1 + 3: the 2 multiplies the whole bracket.
• Half equations that balance atoms but not charge.
• Saying the scattering experiment showed electrons orbit in shells. That was Bohr’s later change.
• Adding electrons for a positive ion. A 3+ ion has three fewer electrons than protons.
• Forgetting to divide by 100 in a relative atomic mass calculation.
• Putting 8 electrons in the first shell, or 3 in the first shell of lithium (2,1 is correct).

Official syllabus

AQA GCSE Chemistry (8462) specification, for teaching from September 2016, exams from June 2018 (version 1.1), published by AQA: section 4.1.1, A simple model of the atom, symbols, relative atomic mass, electronic charge and isotopes.