AQA GCSE Chemistry 8462: The periodic table – Revision Notes

These notes condense section 4.1.2 The periodic table (4.1.2.1 to 4.1.2.6) of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 and exams from June 2018 onwards. The section is examined on Paper 1 at Foundation and Higher Tier. It contains no (HT only) statements; ionic equations, which the specification marks (HT only) in 4.1.1.1, are labelled Higher tier only below.

For full explanations and worked examples, use the periodic table study guide. Then test yourself with the periodic table practice questions. The wider topic is summarised in the Atomic structure revision notes. Course hub: AQA GCSE Chemistry. Every point is listed on the printable checklist.

4.1.2.1 Layout of the table

Term	Meaning
Atomic (proton) number	Order of the modern table
Group	Column; same number of outer electrons; similar chemical properties
Period	Row; same number of occupied shells
Periodic	Similar properties recur at regular intervals

Method: position from electronic structure (first 20 elements)

1. Count shells           -> period
2. Count outer electrons  -> group (1 to 7)
3. Full outer shell       -> Group 0

Reminder: magnesium is 2,8,2, so period 3, Group 2. Nitrogen is 2,5, so period 2, Group 5.

Working backwards: Group 3, period 3 means three shells with 3 outer electrons, 2,8,3, which is 13 electrons, so aluminium.

4.1.2.2 Development of the table

Step	What happened	Why it mattered
1	Early tables in strict order of atomic weight	Incomplete; some elements put in the wrong group
2	Mendeleev left gaps	Kept similar elements together
3	Mendeleev swapped some pairs (tellurium before iodine)	Put iodine with the halogens it resembles
4	Predicted elements found (gallium, scandium, germanium)	Successful predictions supported his table
5	Isotopes discovered	Explained why atomic-weight order is sometimes wrong

Isotope explanation in one line: relative atomic mass is an average over isotopes, so an element with a lot of heavier isotopes can have a higher relative atomic mass than the element after it, even though it has fewer protons.

Pairs to know: tellurium (52, 127.6) before iodine (53, 126.9); argon (18, 39.95) before potassium (19, 39.10).

Science idea (WS 1.1): a prediction that comes true supports a theory; one that fails would refute it.

4.1.2.3 Metals and non-metals

• Metal: reacts to form positive ions. Left and bottom of the table. Most elements.
• Non-metal: does not form positive ions. Right and top.

	Metals	Non-metals
Conduct heat and electricity	Yes	No (graphite conducts)
Melting/boiling points	Mostly high	Mostly low
Solid	Malleable, shiny	Brittle, dull
Outer electrons	Few (usually 1 to 3)	Many (4 to 7, or full)
In reactions	Lose electrons	Gain or share electrons
Oxides	Basic	Often acidic

Link to position: few outer electrons, easily lost, so left side. Further down a group, outer electrons are further from the nucleus and lost more easily, so more metallic (Group 4: carbon non-metal; tin and lead metals).

4.1.2.4 Group 0: noble gases

• Unreactive; exist as single atoms; do not easily form molecules.
• Reason: stable arrangement of electrons: 8 outer electrons (helium 2).
• Boiling point increases down the group as relative atomic mass increases.

Data to recognise (°C): He −269, Ne −246, Ar −186, Kr −152, Xe −107, Rn −62.

Method: predicting from a trend

1. Decide the direction of the trend (up or down the group).
2. Your value must lie between the neighbours.
3. If there is only one neighbour, continue the pattern of the gaps.

4.1.2.5 Group 1: alkali metals

• One outer electron, lost to form 1+ ions.
• Reactivity increases down the group.
• Explanation: more shells, outer electron further from nucleus and more shielded, weaker attraction, lost more easily.

Reacts with	Product	Observations (Li → Na → K, each more vigorous)
Oxygen	Metal oxide, M₂O	Tarnish in air; burn red (Li), yellow-orange (Na), lilac (K)
Chlorine	Metal chloride, MCl (white solid)	Vigorous when heated
Water	Metal hydroxide + hydrogen	Float and fizz; Na melts into a ball; K ignites, lilac flame

Equations (M = Li, Na or K):

4M + O₂ → 2M₂O
2M + Cl₂ → 2MCl
2M + 2H₂O → 2MOH + H₂

Solution after reaction with water: alkaline, universal indicator turns purple.

Melting points fall down the group (°C): Li 181, Na 98, K 63.

Worked reminder: predict rubidium’s melting point.

Drops: Li -> Na = 181 - 98 = 83;  Na -> K = 98 - 63 = 35
The drops are shrinking, so the next drop is smaller still (say about 20)
Prediction: about 63 - 20 = 43 °C  (any value a little below 63 °C fits)

The measured value is 39 °C, so the prediction is reasonable. A prediction above 63 °C breaks the trend and scores nothing.

Predicting reactions: caesium, near the bottom, reacts with water even more violently than potassium and forms caesium hydroxide, CsOH, and hydrogen.

4.1.2.6 Group 7: halogens

• Non-metals; diatomic molecules; seven outer electrons.
• Down the group: relative molecular mass, melting point and boiling point increase; reactivity decreases.
• Explanation (reactivity): outer shell further from nucleus, so an extra electron is attracted less strongly and gained less easily.

Halogen	Mr	State at room temperature	Colour
Chlorine, Cl₂	71	Gas	Pale green
Bromine, Br₂	160	Liquid	Red-brown
Iodine, I₂	254	Solid	Grey-black

Compounds:

Reacts with	Type of compound	Example
Metal	Ionic; halide ion with 1− charge	Lithium + chlorine → lithium chloride, LiCl
Non-metal	Covalent molecule	Hydrogen + bromine → hydrogen bromide, HBr

Displacement rule: more reactive halogen + salt of a less reactive halogen → swap. Otherwise no reaction.

Reminder (not the same pair as the study guide): bromine water added to sodium iodide solution:

Br₂ + 2NaI → 2NaBr + I₂     solution turns from orange to brown

Higher tier only: ionic equation Br₂ + 2I⁻ → 2Br⁻ + I₂ (sodium ions are spectators).

Must-know distinctions

• Group 1 vs Group 7 reactivity: Group 1 increases down (losing an electron gets easier); Group 7 decreases down (gaining an electron gets harder).
• Melting point vs reactivity in Group 7: melting point depends on forces between molecules; reactivity depends on attraction for an extra electron. Different reasons.
• Chlorine vs chloride: Cl₂ is the element; Cl⁻ is the ion in a compound.
• Atomic number vs atomic weight: the modern table uses atomic number; Mendeleev only had atomic weights.
• Predicted vs discovered: Mendeleev predicted elements; other chemists discovered them.

Quick self-test

1. What is the order of elements in the modern periodic table based on?
2. An element is in period 2, Group 6. Give its electronic structure and name.
3. Give two things Mendeleev did that earlier tables did not.
4. Why is argon placed before potassium although its relative atomic mass is higher?
5. Define a metal in terms of ions.
6. Why do noble gases not easily form molecules?
7. Predict whether radon’s boiling point is above or below that of xenon.
8. Name the products when potassium reacts with water. What colour is universal indicator in the solution?
9. Write a balanced equation for lithium reacting with chlorine.
10. What type of compound forms when iodine reacts with sodium? Name it.
11. Will iodine solution react with potassium bromide solution? Explain.
12. Calculate Mr of bromine, Br₂ (Ar Br = 80).

Answers

1. Atomic (proton) number.
2. Two shells, six outer electrons: 2,6, which is oxygen.
3. Left gaps for undiscovered elements; changed the atomic-weight order where properties required it (also: predicted properties of missing elements).
4. Atomic number of argon (18) is less than potassium (19). Nearly all argon is argon-40, while most potassium is potassium-39, so argon’s average mass is higher.
5. An element that reacts to form positive ions.
6. Their atoms already have a stable arrangement of electrons (full outer shell), so they do not need to share, gain or lose electrons.
7. Above (−62 °C compared with −107 °C): boiling point increases down the group.
8. Potassium hydroxide and hydrogen; purple.
9. 2Li + Cl₂ → 2LiCl.
10. Ionic; sodium iodide, NaI.
11. No. Iodine is less reactive than bromine, so it cannot displace bromine.
12. 2 × 80 = 160.

Where marks are usually lost

• Saying “more shells” for the Group 1 trend and stopping. Link it to distance, weaker attraction and an electron lost more easily.
• Using the Group 1 explanation word for word for Group 7. Halogens gain an electron, so the argument is about attracting an incoming electron.
• Explaining halogen melting points with shielding. Melting point is about forces between molecules.
• Writing Na₂Cl or NaCl₂. Group 1 ions are 1+ and halide ions are 1−, so the formula is NaCl.
• Forgetting to balance the water reaction: it is 2M + 2H₂O → 2MOH + H₂, not M + H₂O → MOH + H₂.
• Describing a colour change for a halogen added to a salt of a more reactive halogen. There is no reaction.
• Stating the noble gases have “no outer electrons” or “an empty shell”. The outer shell is full.
• Giving a predicted value outside the range of its neighbours on a trend question.
• Saying isotopes have different chemical properties when explaining Mendeleev’s swaps. Only their masses differ.

Official syllabus

AQA GCSE Chemistry (8462) specification, Version 1.1 (October 2019), for teaching from September 2016 and exams from June 2018 onwards (AQA), section 4.1.2 The periodic table (4.1.2.1 to 4.1.2.6).