Sec 3 Chemistry Atomic Structure Notes

Overview

What you’ll learn

State the relative charge and relative mass of protons, neutrons and electrons.
Use proton number and nucleon number to count the sub-atomic particles in an atom or ion.
Define isotopes and explain why they share chemical properties but differ in mass.
Calculate relative atomic mass from percentage abundances and explain why it is often not a whole number.
Write electronic configurations for the first 20 elements and for simple ions.
Describe how cations and anions form, with electron configurations as evidence.

In 30 seconds

1An atom has a tiny dense nucleus of protons and neutrons, with electrons arranged in shells around it.
2Proton: charge +1, mass 1. Neutron: charge 0, mass 1. Electron: charge −1, mass negligible (1/1840).
3Proton number (Z) = the number of protons — it fixes the element’s identity. Nucleon number (A) = protons + neutrons.
4In any neutral atom, electrons = protons. For an ion, electrons = protons − (charge): a 2+ ion has two fewer electrons.
5Isotopes have the same proton number but different numbers of neutrons — same chemistry, different mass.
6Relative atomic mass (Ar) = weighted average of isotope masses by abundance — that is why chlorine is 35.5, not 35 or 37.
7Electrons fill shells 2, 8, 8 for the first 20 elements: sodium is 2.8.1, calcium is 2.8.8.2.
8Metals lose electrons to form positive cations; non-metals gain electrons to form negative anions — both aim for a full outer shell.

Notes · 1 of 5

The three sub-atomic particles

Protons and neutrons sit in the nucleus; electrons occupy shells far outside it.

Almost all the atom’s mass is crammed into the nucleus, about 100 000 times smaller than the atom itself. Around it the electrons occupy shells — energy levels — which is where all the chemistry happens.

The sub-atomic particles
Particle	Relative charge	Relative mass	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	−1	1/1840 (negligible)	Shells around the nucleus

Because proton charges (+1) and electron charges (−1) are equal and opposite, equal numbers of each make a neutral atom. Losing electrons leaves a positive charge; gaining them leaves a negative one — the particle itself is never charged into fractions, so ion charges are always whole numbers.

Notes · 2 of 5

Counting particles: proton and nucleon numbers

The two-number label on every nuclide tells you everything about its particles.

Every nuclide is written with a proton number (Z) — the number of protons — and a nucleon number (A) — protons plus neutrons — as A over Z beside the symbol. Carbon-14 is written with A = 14, Z = 6: so it has 6 protons, 6 electrons, and 14 − 6 = 8 neutrons.

1Read Z
The proton number fixes the element: any atom with 11 protons is sodium, whatever else it carries.
2Read A
Neutrons = A − Z. Sodium-23: 23 − 11 = 12 neutrons.
3Neutral or ion?
Neutral: electrons = protons. Charged: subtract the charge. Na⁺ has 11 − 1 = 10 electrons; S²⁻ has 16 + 2 = 18 electrons.

Notes · 3 of 5

Isotopes and relative atomic mass

Same protons, different neutrons: identical chemistry, averaged mass.

Isotopes of an element have the same proton number but different nucleon numbers — they differ only in neutron count. Chlorine-35 and chlorine-37 both have 17 protons, so both react as chlorine; but one has 18 neutrons and the other 19.

What isotopes share and what they change
Property	Across isotopes
Chemical reactions	Identical — chemistry depends on electron arrangement, and isotopes have the same configuration
Physical properties depending on mass	Differ slightly — density, rate of diffusion (mass-dependent)
Charge and proton number	Identical
Neutrons and nucleon number	Different

The same method works for three isotopes: multiply each mass number by its percentage, add the products, and divide by 100. Neon (90.9% Ne-20, 0.3% Ne-21, 8.8% Ne-22) gives Ar ≈ 20.2 — again, a weighted average, never a count of any single atom.

Notes · 4 of 5

Electron shells and configurations

Filling order 2, 8, 8 — the configuration decides the chemistry.

Electrons occupy shells that hold 2, 8 and 8 electrons for the first twenty elements. Fill from the innermost shell outwards. The outer-shell electron count is the chemically important number: Group I elements have 1 outer electron, Group VII have 7, Group 0 (Noble gases) have full shells — the stable arrangement.

Configurations worth memorising (Z = proton number)
Element (Z)	Configuration	Outer electrons
Carbon (6)	2.4	4
Oxygen (8)	2.6	6
Neon (10)	2.8	8 (full shell)
Sodium (11)	2.8.1	1
Chlorine (17)	2.8.7	7
Calcium (20)	2.8.8.2	2

Notes · 5 of 5

Ions: losing and gaining electrons

Cations from metals, anions from non-metals — the nucleus is never involved.

Ion formation
Particle	Made by	Example	Electrons after
Cation (positive)	Metal loses electrons	Na → Na⁺ + e⁻	11 → 10
Anion (negative)	Non-metal gains electrons	Cl + e⁻ → Cl⁻	17 → 18

The driving force is the full outer shell: Noble-gas configurations are unusually stable, so atoms gain or lose the fewest electrons needed to reach one. Sodium only needs to lose 1; chlorine only needs to gain 1 — a perfect match, which is why sodium chloride forms so readily.

Revise

Key terms

The words examiners expect you to use precisely.

Proton number (Z): The number of protons in a nucleus; it identifies the element.
Nucleon number (A): The total number of protons and neutrons in a nucleus.
Neutron number: A − Z, the number of neutrons in the nucleus.
Isotope: Atoms of the same element (same proton number) with different numbers of neutrons.
Relative atomic mass (Ar): The weighted average mass of an element’s isotopes compared with 1/12 of a carbon-12 atom.
Electronic configuration: The arrangement of electrons in shells, written 2.8.8-style for the first 20 elements.

Show 4 more key terms

Cation: A positively charged ion, formed when an atom loses electrons.
Anion: A negatively charged ion, formed when an atom gains electrons.
Nuclide: A specific combination of protons and neutrons, labelled A and Z.
Noble-gas configuration: A full outer shell of electrons — the stable arrangement atoms approach when they ionise.

Revise

Common mistakes

Where students most often lose marks in this chapter — and what to do instead.

Mistake: Saying an ion’s charge changes its number of protons.
Ions form by losing or gaining electrons only; the nucleus is untouched.
Instead: Charge = protons − electrons. Na⁺ still has 11 protons, just 10 electrons.
Mistake: Counting neutrons as "A − charge".
The charge involves electrons, not neutrons.
Instead: Neutrons = nucleon number − proton number (A − Z).

Show 10 more mistakes

Mistake: Writing electrons as mass 1.
Electrons are about 1/1840 of a proton mass — negligible next to protons and neutrons.
Instead: Proton and neutron: mass 1 each; electron: negligible mass, charge −1.
Mistake: Believing isotopes have different chemical properties.
Chemistry is decided by electron configuration, which isotopes share.
Instead: Isotopes react identically; only mass-dependent physical properties (density, diffusion rate) differ.
Mistake: Averaging isotope masses without weighting by abundance, e.g. (35 + 37)/2 = 36 for chlorine.
The isotopes are not equally abundant — 75% is chlorine-35.
Instead: Ar = Σ(mass × %abundance) ÷ 100, giving 35.5 for chlorine.
Mistake: Expecting Ar always to be a whole number, and "correcting" 35.5 to 36.
Ar is a weighted average across a mixture of isotopes, so it is usually not whole.
Instead: Report the value on the Periodic Table (Cl = 35.5) in calculations.
Mistake: Writing potassium (Z = 19) as 2.8.9.
The third shell holds at most 8 for the first 20 elements.
Instead: Potassium is 2.8.8.1 — the fourth shell starts once the third holds 8.
Mistake: Giving O²⁻ the configuration 2.6.
2.6 is the neutral atom; the 2− ion has gained two electrons.
Instead: O²⁻ is 2.8 — a full outer shell, the same as neon.
Mistake: Saying ions form "to become more stable" without saying how.
Vague wording earns nothing; the mechanism is the full outer shell.
Instead: State the gain/loss of electrons and the Noble-gas configuration reached.
Mistake: Mixing up mass number and relative atomic mass.
Mass number counts one nuclide’s particles; Ar averages a natural sample.
Instead: A is a whole-number count per atom; Ar is the sample average (often non-whole, e.g. chlorine’s 35.5).
Mistake: Reading a negative charge as "lost electrons", e.g. saying Cl⁻ has 16 electrons.
Negative charge means extra electrons.
Instead: Cl⁻: 17 + 1 = 18 electrons.
Mistake: Claiming the nucleus contains electrons.
The Rutherford scattering evidence puts only protons and neutrons in the nucleus.
Instead: Nucleus = protons + neutrons; electrons live in shells.

Check yourself

Quick quiz

20 questions with instant explanations. About 5 minutes.

Question 1 of 3

Question 1: current question
Question 2: not answered
Question 3: not answered

Recall sub-atomic particles

Which particle carries a charge of +1 and a relative mass of 1?

Practise

Topical questions

Exam-style practice in three tiers, from recall to data questions.

4Foundation
5Standard
3Challenging

12 questions worth 49 marks, each with a model answer, mark scheme and examiner tips.

Practise topical questions

Exam practice

Practice paper

A 30-mark paper set out like a school paper, with a separate mark scheme. Give yourself 40 min.

Question paper

Chapter 3 Practice Paper

30 marks · 40 min · 4 pages

5 multiple-choice questions and structured questions with answer lines. Chapter 3 (sub-atomic particles, isotopes, relative atomic mass, configurations, ions).

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Mark scheme

PRO

Chapter 3 Practice Paper — mark scheme

Answers · Marking points · 3 pages

Model answers with a mark for each creditworthy point — for marking at home or in class.

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Where this chapter is tested

Each school sets its own scope — check your assessment outline.

Questions

Frequently asked questions

What does Sec 3 Chemistry Chapter 3 cover?

It covers sub-topic 2.2 of O-Level Pure Chemistry 6092: the three sub-atomic particles with their charges and masses, proton and nucleon numbers, isotopes and relative atomic mass from abundances, electronic configurations for the first 20 elements, and the formation of simple ions.

Why is chlorine’s relative atomic mass 35.5?

Natural chlorine is a mixture of about 75% chlorine-35 and 25% chlorine-37. The relative atomic mass is the abundance-weighted average: (75×35 + 25×37)/100 = 35.5. No single atom weighs 35.5 — the average of the sample does.

How do I count particles in an ion?

Protons never change: read Z. Neutrons = A − Z, also unchanged. Electrons = Z − charge, so a 2+ ion has two electrons fewer and a 2− ion has two more than the neutral atom.

Why do I write potassium as 2.8.8.1 and not 2.8.9?

For the first twenty elements the third shell holds at most eight electrons. Once those eight are placed, the nineteenth electron opens the fourth shell — so potassium (Z = 19) is 2.8.8.1 and calcium (Z = 20) is 2.8.8.2.

When is this chapter usually tested in school?

Atomic structure follows the particle theory early in Sec 3, so it appears in Weighted Assessment 2 (Term 2) and the end-of-year examination — and isotope arithmetic plus electron configurations feed directly into bonding and the Periodic Table chapters later. Scope varies by school, so check your school’s assessment outline.

Sec 3 Chemistry Atomic Structure notes