Edexcel A-Level Chemistry: Atomic Structure and Periodicity — Complete Revision Guide (9CH0)

Atomic structure is the very first piece of chemistry on the Edexcel 9CH0 specification, and it is also the topic that quietly underpins almost every later one. Bonding theory rests on electron configurations. Energetics depends on ionisation and electron affinity. The colour and catalytic behaviour of transition metals comes from d-orbital structure. The periodicity of Period 3 oxides and chlorides flows directly from electronegativity and atomic radius. If your atomic structure is shaky, every later topic gets harder. If it is fluent, the rest of the course makes sense.

This guide is a topic-by-topic walkthrough of the atomic structure and periodicity content in the 9CH0 specification. It covers atomic models from Dalton through to the modern shell model, electron configuration including s, p, d and f sublevels and the special cases of chromium and copper, first and successive ionisation energies and what they reveal about structure, time-of-flight mass spectrometry, periodic trends in atomic radius and electronegativity, and the chemistry of Period 3 elements and their oxides and chlorides. For each topic you will see the core ideas, common pitfalls, a worked example or short numerical illustration, and a link into the LearningBro Atomic Structure and Periodicity course.

What the Edexcel 9CH0 Specification Covers

Edexcel A-Level Chemistry (9CH0) is assessed through three written papers. Paper 1 (Advanced Inorganic and Physical Chemistry) is one hour 45 minutes and worth 90 marks. Paper 2 (Advanced Organic and Physical Chemistry) is also one hour 45 minutes and worth 90 marks. Paper 3 (General and Practical Principles in Chemistry) is two hours 30 minutes and worth 120 marks, and it draws on the whole specification including practical work. There is no coursework. Practical skills are assessed through written questions on the core practicals (CPAC).

Atomic structure and periodicity sits in Topics 1 and 2 of the specification (refer to the official specification document for exact wording). It is examined most heavily on Paper 1, but the foundational ideas — particularly electron configuration, ionisation energy and periodicity — appear synoptically on every paper. The table below maps the main sub-topics to a typical paper weighting.

Sub-topic	Spec area	Typical Paper 1 weight
Atomic models and the modern atom	Topic 1	2-4 marks
Electron configuration (s, p, d, f)	Topic 1	4-6 marks
Ionisation energies and trends	Topic 1	6-8 marks
Time-of-flight mass spectrometry	Topic 1	4-6 marks
Periodicity (atomic radius, electronegativity)	Topic 2	4-6 marks
Period 3 elements and their reactions	Topic 2	4-6 marks
Period 3 oxides and chlorides	Topic 2	4-6 marks

These weightings are estimates, modelled on the Edexcel 9CH0 paper format rather than guarantees for a single sitting. What is reliable is that ionisation-energy explanations and electron configuration appear on virtually every paper, and that Period 3 trends are a frequent extended-response topic.

Atomic Models and the Modern Atom

The history of atomic models is short on the specification but worth getting right because it informs how we describe sub-atomic particles today. Dalton proposed indivisible atoms; J. J. Thomson's plum-pudding model placed electrons inside a positive sphere; Rutherford's gold-foil scattering showed the atom is mostly empty space with a small dense positive nucleus; Bohr added quantised electron orbits to explain emission spectra; modern quantum mechanics replaces orbits with orbitals — three-dimensional regions of probability.

You should be able to state the relative mass and charge of the proton, neutron and electron (1, +1; 1, 0; 1/1836 ≈ 0, -1), define mass number (protons + neutrons) and atomic number (protons), and define an isotope as atoms with the same atomic number but different mass numbers. Relative atomic mass (Ar) is the weighted mean mass of the isotopes of an element relative to one twelfth of the mass of a carbon-12 atom.

A common pitfall is conflating mass number with relative atomic mass. Mass number is an integer for a single isotope; relative atomic mass is a weighted average across the natural abundance of all isotopes and is rarely a whole number. Another pitfall is forgetting that ions have the same number of protons and neutrons as the parent atom — only electron count changes.

Worked example. An element X has two isotopes of mass numbers 63 (69.2 percent abundance) and 65 (30.8 percent abundance). Calculate Ar. Ar = (63 × 69.2 + 65 × 30.8) / 100 = (4359.6 + 2002) / 100 = 6361.6 / 100 = 63.616, so Ar ≈ 63.6. The element is copper.

See the atomic models lesson for the development of the modern atom in detail.

Electron Configuration (s, p, d and f Sublevels)

Electrons occupy quantised energy levels (shells) labelled n = 1, 2, 3, .... Within each shell are sublevels: s (one orbital, holds 2 electrons), p (three orbitals, 6 electrons), d (five orbitals, 10 electrons) and f (seven orbitals, 14 electrons). Edexcel expects you to use spdf notation to write the configuration of any element up to krypton with confidence, and beyond for transition metal ions.

The order of filling follows the Aufbau principle (lowest energy first), Hund's rule (maximum unpaired spin within a sublevel) and the Pauli exclusion principle (no two electrons share all four quantum numbers). The filling order for the first few sublevels is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p — note that 4s fills before 3d because 4s is slightly lower in energy in the neutral atom. When 3d ionises, however, the 4s electrons are removed first. This is a frequent exam question.

Two anomalous configurations must be memorised. Chromium is [Ar] 3d^5 4s^1, not [Ar] 3d^4 4s^2; copper is [Ar] 3d^10 4s^1, not [Ar] 3d^9 4s^2. Both anomalies arise because a half-filled or fully-filled d sublevel gains extra stability that more than offsets the cost of unpairing a 4s electron.

Worked example. Write the electron configuration of Fe and Fe^3+. Iron is element 26, so Fe is 1s^2 2s^2 2p^6 3s^2 3p^6 3d^6 4s^2, or [Ar] 3d^6 4s^2. To form Fe^3+, remove three electrons — first the two 4s electrons, then one 3d electron — giving [Ar] 3d^5. The half-filled d^5 configuration is one reason Fe^3+ is a stable oxidation state.

See the electron configuration lesson for full coverage including transition metal ions.

First and Successive Ionisation Energies

The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly positive ions: X(g) → X+(g) + e-. The unit is kJ mol^-1. The state symbol (g) is essential — many candidates lose the mark by writing X(s).

Three factors determine ionisation energy: nuclear charge (more protons pull harder), atomic radius (further electrons feel a weaker pull), and shielding (inner electrons reduce the effective nuclear charge felt by the outer electron). Across a period, nuclear charge rises faster than radius shrinks, so first ionisation energy generally rises. Down a group, both radius and shielding increase, and first ionisation energy falls.

Two anomalies in Period 2 must be explained. Between beryllium and boron, ionisation energy drops slightly because boron's outer electron is in a 2p orbital, which is slightly higher in energy than the 2s orbital from which Be's outer electron is removed. Between nitrogen and oxygen, ionisation energy drops because oxygen's 2p^4 configuration has paired electrons in one 2p orbital — the repulsion within the pair makes one electron easier to remove than removing one of nitrogen's three unpaired 2p electrons. Equivalent anomalies appear between Mg/Al and P/S in Period 3.

Successive ionisation energies (1st, 2nd, 3rd, ...) always increase because each subsequent electron is removed from an increasingly positive ion. Large jumps between successive values reveal shell structure: a big jump after the second ionisation suggests the third electron came from an inner shell, indicating two outer-shell electrons and therefore Group 2.

Worked example. The first six ionisation energies of an element (kJ mol^-1) are 786, 1577, 3232, 4356, 16091, 19785. Identify the group. There is a large jump between the fourth and fifth values, indicating the fifth electron is in a deeper shell. The element therefore has four outer-shell electrons and is in Group 4 (likely silicon).

See the ionisation energies lesson for the full set of trend explanations.

Time-of-Flight Mass Spectrometry

Time-of-flight (TOF) mass spectrometry is the standard technique for measuring isotopic masses and molecular masses on the specification. The four stages are ionisation, acceleration, drift (or flight) and detection. In ionisation, the sample is ionised either by electron impact (a high-energy electron knocks an electron off, forming X+) or by electrospray (the sample is dissolved, sprayed through a charged needle, and gains a proton to form XH+). Ions are accelerated through a potential difference, giving each the same kinetic energy. They drift through a field-free region; lighter ions travel faster, heavier ions slower. The detector records the time of flight, and the mass-to-charge ratio (m/z) is calculated.

Because every ion gains the same kinetic energy, ½mv^2 = constant, so v ∝ 1/√m. The time to traverse a fixed drift length is t = d/v, so t ∝ √m. Doubling the mass increases flight time by a factor of √2. This proportionality is examined directly.

Worked example. Two isotopes have m/z values of 24 and 26. The lighter ion takes 1.20 × 10^-5 s to traverse the drift region. How long does the heavier ion take? t_2 / t_1 = √(m_2/m_1) = √(26/24) = √1.0833 ≈ 1.0408. So t_2 = 1.20 × 10^-5 × 1.0408 ≈ 1.249 × 10^-5 s.

The mass spectrum of an element shows isotope peaks at each isotopic mass, with peak heights proportional to abundance. Ar can be calculated from the spectrum exactly as in the earlier worked example. For molecules, the highest m/z peak (the molecular ion peak, M+) gives Mr.

See the mass spectrometry lesson for sample spectra and Ar calculations.

Periodicity: Atomic Radius and Electronegativity

Periodicity is the repeating pattern of properties as you move across the table. Two trends dominate at A-Level. Atomic radius decreases across a period because nuclear charge increases while shielding stays roughly constant — outer electrons are pulled in tighter. Down a group, atomic radius increases because each new period adds a shell of shielding. Electronegativity, the ability of an atom in a covalent bond to attract the bonding electrons, increases across a period (greater nuclear charge, smaller radius) and decreases down a group (more shielding, larger radius). Fluorine is the most electronegative element on the Pauling scale.

Trend	Across period	Down group
Atomic radius	Decreases	Increases
First ionisation energy	Generally increases	Decreases
Electronegativity	Increases	Decreases
Metallic character	Decreases	Increases
Melting point	Variable (peaks at giant covalent)	Group-specific

Many candidates lose marks by stating the trend without giving the explanation — the examiner wants the underlying reason in terms of nuclear charge, radius and shielding, not just the direction. A clean answer pattern is "as you move across the period, nuclear charge increases and atomic radius decreases while shielding remains approximately constant; therefore the outer electron is held more tightly".

See the periodicity lesson for trend explanations and graph-reading practice.

Period 3 Elements and Their Reactions

Period 3 — sodium through to argon — is the standard set of elements for examining periodic trends in reactions. You need the reactions with oxygen, water and chlorine for the metals and selected non-metals.

Sodium burns in oxygen with a yellow flame to give sodium oxide (Na2O); magnesium burns brightly to give MgO; aluminium reacts slowly because of its protective Al2O3 layer; silicon, phosphorus and sulfur all burn to give covalent oxides (SiO2, P4O10, SO2 → SO3). Sodium reacts vigorously with cold water to give NaOH and H2; magnesium reacts very slowly with cold water but readily with steam to give MgO and H2; aluminium and silicon do not react. Sodium burns in chlorine to give NaCl; magnesium gives MgCl2; aluminium gives Al2Cl6 (a covalent dimer in the gas phase); silicon gives SiCl4; phosphorus gives PCl3 or PCl5; sulfur gives a mixture of chlorides.

A common pitfall is writing the wrong stoichiometry — particularly forgetting that Mg reacts as 2Mg + O2 → 2MgO and that aluminium oxide is Al2O3 (not AlO). Another is omitting state symbols in equations involving aqueous and gaseous species.

See the Period 3 reactions lesson for full equations and observations.

Period 3 Oxides and Chlorides

The oxides and chlorides of Period 3 show a clean trend in bonding type and behaviour with water that examiners love. The metallic oxides Na2O, MgO and Al2O3 are giant ionic (Al2O3 has some covalent character because Al^3+ is small and highly polarising). The non-metallic oxides SiO2, P4O10, SO2/SO3 are giant covalent (SiO2) or simple molecular (the rest). With water, basic oxides give alkaline solutions, acidic oxides give acidic solutions, and amphoteric Al2O3 reacts with both acids and bases.

Oxide	Bonding	Reaction with water	pH of solution
Na2O	Ionic	Na2O + H2O → 2NaOH	~13-14
MgO	Ionic	Slight; MgO + H2O → Mg(OH)2	~9
Al2O3	Ionic with covalent character	Insoluble; amphoteric	7 (insoluble)
SiO2	Giant covalent	Insoluble	7 (insoluble)
P4O10	Simple molecular	P4O10 + 6H2O → 4H3PO4	~2
SO2	Simple molecular	SO2 + H2O → H2SO3	~3
SO3	Simple molecular	SO3 + H2O → H2SO4	~1

The chlorides show a similar bonding shift. NaCl and MgCl2 are ionic and dissolve neutrally (NaCl) or slightly acidically (MgCl2 hydrolyses a little). Al2Cl6 (or AlCl3) is covalent and hydrolyses strongly to give an acidic solution. SiCl4, PCl5 and others are covalent and hydrolyse violently in water with HCl fumes.

A common pitfall in these questions is writing "ionic" for Al2Cl6 — its small, highly charged cation polarises chloride enough that the bonding is essentially covalent. The equation for hydrolysis of SiCl4 is SiCl4 + 2H2O → SiO2 + 4HCl.

See the oxides and chlorides lesson for hydrolysis equations and pH summaries.

Common Mark-Loss Patterns

• Omitting the (g) state symbol in the ionisation energy definition.
• Stating a periodic trend without explaining it in terms of nuclear charge, radius and shielding.
• Forgetting the chromium and copper anomalous configurations in d-block questions.
• Removing 3d electrons before 4s when ionising a transition metal — the order is reversed.
• Confusing mass number (single isotope) with relative atomic mass (weighted average).
• Writing "ionic" instead of "covalent" for Al2Cl6 and AlCl3 in the gas phase.
• Mis-balancing equations for Period 3 reactions, especially Al2O3 formation.
• Misreading m/z scale on a mass spectrum or assuming the largest peak is M+ rather than the rightmost peak.

How to Revise This Topic

• Drill electron configurations. Write twenty configurations a day for a week, including transition metal ions, until you stop hesitating on Cr, Cu and Fe^3+.
• Memorise the trend explanations as templates. Write a one-sentence template for "first ionisation energy across a period" and "down a group" and reuse them. Examiners want consistent reasoning.
• Practise ionisation-energy data questions. Print a successive-IE table, redact the element, and identify the group from the jump.
• Sketch periodicity graphs from memory. First ionisation energy versus atomic number for Periods 2 and 3, atomic radius across the same range. Annotate the anomalies.
• Learn the Period 3 reactions as a sequence. Na, Mg, Al, Si, P, S — what each does with O2, water, and Cl2. Pair this with oxide and chloride hydrolysis.
• Use the LearningBro practice quizzes to test recall under timed conditions; little-and-often beats one massive session.

Linking to Other Topics

Atomic structure feeds directly into bonding (electron configurations explain why elements gain, lose or share electrons), energetics (ionisation and electron affinity are the first two terms of every Born-Haber cycle), and inorganic chemistry (the Group 2 and Group 7 trends are direct applications of periodicity, and transition-metal chemistry is unintelligible without d-orbital configurations). Mass spectrometry is also revisited in the analytical chemistry course when interpreting molecular fragmentation. Time spent on this topic returns many times across the rest of the course.

Final Word

Atomic structure and periodicity is the cheapest section of the specification to master because the content is finite and the examiner question patterns are stable. Drill electron configurations, lock in the trend explanations, and practise data-interpretation questions on ionisation energies and mass spectra. The full LearningBro Atomic Structure and Periodicity course walks you through every sub-topic with worked examples, practice questions and AI tutor feedback. Get this section right and the rest of A-Level Chemistry has a foundation it can stand on.