AQA GCSE Chemistry 8462: The periodic table – Study Guide

This guide teaches section 4.1.2 The periodic table of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 and exams from June 2018 onwards. It covers every statement from 4.1.2.1 to 4.1.2.6: how the table is arranged, how it was developed, metals and non-metals, and Groups 0, 1 and 7. The topic is assessed on Paper 1 (topics 1 to 5), set at Foundation and Higher Tier. Section 4.1.2 itself has no (HT only) statements; the one skill below that is Higher tier only is writing ionic equations, which the specification marks (HT only) in 4.1.1.1.

For quick recall, use the periodic table revision notes. To test yourself, use the periodic table practice questions. For the whole of topic 4.1 in one place, see the Atomic structure and the periodic table overview and its revision notes. The course hub is AQA GCSE Chemistry, and the printable checklist lists every specification point.

What this unit covers

Spec point	What you must be able to do	Tier
4.1.2.1 The periodic table	Explain how an element’s position links to its electron arrangement and atomic number; predict reactions and reactivity from position	Both
4.1.2.2 Development of the periodic table	Describe the steps from atomic-weight tables to Mendeleev’s table and explain the role of isotopes	Both
4.1.2.3 Metals and non-metals	Explain differences in physical and chemical properties; link atomic structure and reactions to position	Both
4.1.2.4 Group 0	Explain properties from the outer shell; predict properties from trends	Both
4.1.2.5 Group 1	Describe reactions of Li, Na and K with oxygen, chlorine and water; explain and predict trends	Both
4.1.2.6 Group 7	Describe compounds of Cl, Br and I with metals and non-metals; explain trends and displacement	Both

There is no required practical in 4.1.2. The specification flags halogen displacement as an opportunity for AT 6 (safe handling and careful mixing of reagents).

4.1.2.1 How the table is arranged

The elements are arranged in order of atomic (proton) number. They are placed so that elements with similar properties fall in the same column, called a group. Rows are called periods. The table is “periodic” because similar properties recur at regular intervals as atomic number increases.

The link to electrons is the key idea. For the first 20 elements:

• the period number equals the number of occupied shells
• the group number (Groups 1 to 7) equals the number of electrons in the outer shell
• Group 0 elements have a full outer shell (8 electrons, or 2 for helium).

Elements in a group have the same number of outer electrons, which decide how an atom reacts, so they have similar chemical properties.

Worked example 1. An atom has electronic structure 2,8,6. Identify the element and its position.

Total electrons = 2 + 8 + 6 = 16
Neutral atom, so protons = 16, atomic number = 16 -> sulfur
Three occupied shells -> period 3
Six outer electrons -> Group 6

Position also predicts reactivity: lower in Group 1 means more reactive, lower in Group 7 means less reactive.

4.1.2.2 Development of the periodic table

Before protons, neutrons and electrons were discovered, scientists could only sort elements by their atomic weights (what we now call relative atomic masses). The steps you need to describe are these.

1. Early tables used strict atomic-weight order. They were incomplete, and strict order forced some elements into groups they did not belong in.
2. Mendeleev (1869) left gaps for elements he thought had not yet been discovered, so that each column kept elements with similar properties.
3. He changed the order in some places. Tellurium has a higher atomic weight than iodine, but he placed tellurium first so that iodine sat with chlorine and bromine, which it resembles.
4. His predictions were tested. He used the gaps to predict the properties of missing elements. Gallium (1875), scandium (1879) and germanium (1886) were discovered and fitted his gaps.
5. Isotopes explained the swaps. The table is really in order of atomic number. Atomic weight is an average over isotopes, so an element with heavy isotopes can outweigh the next element.

Worked example 2. Decide whether this data supports Mendeleev’s table. He predicted an element he called eka-silicon, below silicon. Germanium was later found to fill that gap.

Property	Mendeleev’s prediction	Germanium (measured)
Relative atomic mass	72	72.6
Density in g/cm³	5.5	5.32
Oxide	a dioxide	GeO₂, a dioxide

Mass:    (72.6 - 72) / 72.6 x 100 = 0.8% difference
Density: (5.5 - 5.32) / 5.32 x 100 = 3.4% difference

The predictions are close and the oxide formula matches. A prediction made before a discovery that turns out correct supports the idea behind it; a badly wrong one would refute it. This is how testing a prediction works.

Worked example 3 (isotopes). Tellurium (atomic number 52) has relative atomic mass 127.6. Iodine (atomic number 53) has 126.9. Explain why tellurium comes first.

Elements are ordered by atomic number, and 52 is less than 53. Iodine has only one natural isotope, iodine-127. Tellurium’s most common isotopes are tellurium-128 and tellurium-130, both heavier than 127, so tellurium’s average mass is higher even though it has fewer protons. The same reasoning places argon (39.95) before potassium (39.10): almost all argon is argon-40, while most potassium is potassium-39.

4.1.2.3 Metals and non-metals

The specification defines the two groups by ions:

• metals are elements that react to form positive ions
• non-metals are elements that do not form positive ions.

Most elements are metals. Metals are on the left and towards the bottom of the table; non-metals are on the right and towards the top.

Property	Typical metal	Typical non-metal
Electrical and thermal conduction	Good	Poor (graphite is an exception)
Appearance	Shiny	Dull
Melting and boiling points	Mostly high	Mostly low
Solid shape change	Malleable and ductile	Brittle
Reaction in compounds	Loses electrons, forms positive ions	Gains electrons (negative ions) or shares electrons
Oxide	Reacts with acids (a base)	Often acidic

Why position matches atomic structure. Metal atoms have only a few outer electrons (usually 1, 2 or 3). They lose these to reach a stable arrangement, forming positive ions, so metals sit on the left. Non-metal atoms have nearly full outer shells (4 to 7 outer electrons). They gain or share electrons instead, so they sit on the right. Going down a group, the outer electrons are further from the nucleus and are lost more easily, so elements become more metallic. In Group 4, carbon at the top is a non-metal, while tin and lead lower down are metals.

Reactions follow electrons. Sodium (2,8,1) loses one electron to form Na⁺; chlorine (2,8,7) gains one to form Cl⁻, giving ionic sodium chloride. For bonding, see the ionic bonding revision notes.

4.1.2.4 Group 0: the noble gases

The noble gases are unreactive and do not easily form molecules. They exist as single atoms, because their atoms already have a stable arrangement of electrons: eight outer electrons, except helium, which has two (a full first shell). There is nothing to gain by losing, gaining or sharing electrons.

Their boiling points increase with relative atomic mass, going down the group.

Element	Boiling point (°C)
Helium	−269
Neon	−246
Argon	−186
Krypton	?
Xenon	−107

Worked example 4. Predict the boiling point of krypton.

Midpoint of argon and xenon = (-186 + -107) / 2 = -146.5 °C

A sensible prediction is about −150 °C (the measured value is −152 °C). Any value between the two neighbours fits the trend; a value outside them does not.

4.1.2.5 Group 1: the alkali metals

All Group 1 elements have one electron in the outer shell, which gives them their characteristic properties. They lose that electron to form ions with a 1+ charge. Their reactivity increases down the group.

Explaining the trend. Going down the group there are more shells, so the outer electron is further from the nucleus and more shielded. It is less strongly attracted, so it is lost more easily and the reaction is more vigorous.

You must describe the reactions of the first three alkali metals with three substances.

With oxygen. Freshly cut lithium, sodium and potassium are shiny but tarnish quickly in air as an oxide forms, so they are stored under oil. Heated in oxygen they burn (lithium red flame, sodium yellow-orange, potassium lilac). At GCSE the product is written as the oxide.

4Li + O₂ → 2Li₂O    lithium oxide
4Na + O₂ → 2Na₂O    sodium oxide
4K + O₂ → 2K₂O      potassium oxide

With chlorine. Each metal reacts vigorously when heated in chlorine to form a white solid chloride, more vigorously going down the group.

2Na + Cl₂ → 2NaCl    sodium chloride

With water. Each metal forms a metal hydroxide and hydrogen. The hydroxide dissolves, so the solution is alkaline and turns universal indicator purple.

Metal	Observations
Lithium	Floats; fizzes steadily; slowly gets smaller
Sodium	Floats; melts into a ball; fizzes rapidly; moves about on the surface
Potassium	Floats; melts; fizzes violently; hydrogen ignites with a lilac flame

2K + 2H₂O → 2KOH + H₂

Worked example 5 (predicting). Rubidium is below potassium. Predict its reaction with water.

It forms rubidium hydroxide and hydrogen, 2Rb + 2H₂O → 2RbOH + H₂, and reacts more vigorously than potassium, because its outer electron is further from the nucleus and lost more easily.

4.1.2.6 Group 7: the halogens

The halogens are non-metals made of diatomic molecules (F₂, Cl₂, Br₂, I₂). They react in similar ways because each atom has seven outer electrons and needs one more.

Trends down the group:

• relative molecular mass, melting point and boiling point all increase
• reactivity decreases.

At room temperature chlorine is a pale green gas, bromine a red-brown liquid and iodine a grey-black solid. Larger molecules mean stronger forces between molecules, so more energy is needed to separate them.

Explaining the reactivity trend. A halogen reacts by gaining one electron. Further down, the outer shell is further from the nucleus, so an incoming electron is attracted less strongly and is gained less easily.

Compounds with metals and non-metals. This is a separate learning outcome, so learn both cases.

• With metals, chlorine, bromine and iodine form ionic compounds (salts) containing halide ions with a 1− charge: chloride Cl⁻, bromide Br⁻, iodide I⁻. Example: 2K + Br₂ → 2KBr, potassium bromide.
• With non-metals, they form covalent molecules, sharing electrons. Example: H₂ + Cl₂ → 2HCl, hydrogen chloride, a gas made of small molecules. It dissolves in water to form an acidic solution.

Displacement. A more reactive halogen displaces a less reactive halogen from an aqueous solution of its salt. If the halogen added is less reactive than the one in the salt, nothing happens.

Halogen added	Potassium chloride (aq)	Potassium bromide (aq)	Potassium iodide (aq)
Chlorine water	–	Turns orange (bromine formed)	Turns brown (iodine formed)
Bromine water	No reaction	–	Turns brown (iodine formed)
Iodine solution	No reaction	No reaction	–

Worked example 6. Write the equation for chlorine reacting with potassium bromide solution.

Word:     chlorine + potassium bromide → potassium chloride + bromine
Symbol:   Cl₂ + 2KBr → 2KCl + Br₂

Higher tier only. Potassium ions do not change, so the ionic equation leaves them out: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂. Check charges: 2− on each side.

Predicting. Astatine, below iodine, should be a dark solid with a higher melting point, less reactive than iodine, and unable to displace iodine from potassium iodide.

Common errors

• Saying Mendeleev “discovered” gallium and germanium. He predicted them; others found them.
• Explaining tellurium and iodine without isotopes. The mark is for “tellurium has heavier isotopes, so a higher average mass”.
• Writing that noble gases are unreactive “because they are gases”. The reason is a stable electron arrangement.
• Giving bromine water and potassium chloride a colour change. There is no reaction.

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).