AQA GCSE Chemistry: Atomic Structure and the Periodic Table (8462)

This guide covers Topic 4.1 Atomic Structure and the Periodic Table, the first of eleven topics in AQA GCSE Chemistry (8462), for teaching from September 2016. The qualification is tiered (Foundation/Higher) and co-teachable with AQA GCSE Combined Science – content marked “chemistry only” in the specification is unique to this separate-science qualification and not shared with Combined Science.

Where this fits in 8462

Topic 4.1 opens the syllabus by explaining how atomic structure determines an element’s position and behaviour in the periodic table, setting up the foundation for Topic 4.2 (Bonding, structure and the properties of matter) and every later topic that references specific elements or groups.

Syllabus coverage

AQA GCSE CHEMISTRY (8462) — TOPIC 4.1 ATOMIC STRUCTURE AND THE PERIODIC TABLE

• 4.1.1 Atomic structure — the structure of atoms, relative electrical charges and masses of subatomic particles, the separation of mixtures (4.1.1.2), and the development of the model of the atom (shared content with physics)
• 4.1.2 The periodic table — how elements are arranged in the periodic table, its historical development, the distinction between metals and non-metals, the properties of Group 0, Group 1 and Group 7 elements (4.1.2.4 to 4.1.2.6)
• 4.1.3 Properties of transition metals (chemistry only) — a comparison of transition metals with Group 1 elements on melting point, density and hardness, and their typical properties of variable ionic charge, coloured compounds and use as catalysts (4.1.3.2), unique to the separate Chemistry qualification

How to approach it

Because atomic structure underpins how the periodic table is organised, practise linking the two directly – explaining why elements in the same group share similar properties in terms of electron configuration – rather than learning atomic structure and the periodic table as separate facts. The historical development of the periodic table (4.1.2) is a common source of extended-response questions, so practise explaining how scientific models change as new evidence emerges, using this topic as a specific example. Since 4.1.3 (transition metals) is chemistry-only content not shared with Combined Science, make sure you know which parts of this topic apply to your specific qualification if you are studying alongside Combined Science students.

Official syllabus

AQA GCSE Chemistry (8462) specification, for teaching from September 2016 — aqa.org.uk.

The structure of the atom

An atom has a small central nucleus containing protons and neutrons, surrounded by electrons in shells. The nucleus is about 1/10,000 the diameter of the whole atom, yet holds virtually all of its mass — the radius of an atom is roughly 0.1 nm (1 x 10^-10 m), while the nucleus is around 1 x 10^-14 m.

Particle	Relative mass	Relative charge	Location
Proton	1	+1	Nucleus
Neutron	1	0	Nucleus
Electron	Very small (1/1836)	-1	Shells

The atomic number is the number of protons, and it defines the element. The mass number is the total number of protons and neutrons. Atoms are neutral overall, so the number of electrons equals the number of protons.

Isotopes and relative atomic mass

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. Because chemical behaviour is determined by electrons, isotopes of an element react identically — only their masses differ.

Relative atomic mass is the weighted mean mass of the isotopes present:

Ar = [(mass 1 x abundance 1) + (mass 2 x abundance 2)] / 100

Electronic structure

Electrons occupy shells from the innermost outwards, holding a maximum of 2, then 8, then 8. Chlorine (17 electrons) is 2,8,7. The number of outer-shell electrons determines the group, and the number of occupied shells determines the period.

Development of the model

Ideas about the atom changed as experimental evidence accumulated: Dalton’s solid spheres, then Thomson’s plum pudding model, then Rutherford’s alpha-scattering experiment which showed most of an atom is empty space with a dense positive nucleus, then Bohr’s shells, and finally the discovery of the neutron. This sequence is a favourite context for questions about how scientific models respond to new evidence.

The periodic table

Elements are arranged in the periodic table by increasing atomic number. Mendeleev built an early version by ordering elements by atomic mass and grouping them by properties, but crucially he left gaps for elements not yet discovered and used those gaps to predict their properties – the later discovery of gallium and germanium, matching his predictions closely, is what confirmed his arrangement and is why his table succeeded where earlier attempts failed.

Metals are found on the left and centre of the table and lose electrons to form positive ions; non-metals are found on the right and gain or share electrons. Elements in the same group have the same number of outer-shell electrons and therefore react in similar ways.

Group 1 – the alkali metals

Group 1 elements react with water to form a metal hydroxide and hydrogen: metal + water → metal hydroxide + hydrogen.

Reactivity increases going down the group, and a full explanation needs all three linked steps: the outer electron is further from the nucleus as more shells are added; there is therefore more shielding from the inner shells; so the attraction between the nucleus and the outer electron is weaker, and that electron is lost more easily.

Group 7 – the halogens

Melting and boiling points increase going down Group 7, because the molecules get larger and the intermolecular forces between them get stronger – but reactivity decreases, the opposite trend to Group 1. The reasoning is the same physics applied to the opposite process: a halogen reacts by gaining an electron, and further down the group the outer shell is further from the nucleus with more shielding, so an incoming electron is attracted less strongly and gained less readily.

A more reactive halogen will displace a less reactive one from a solution of its salt – chlorine displaces bromine, and bromine displaces iodine. Adding chlorine water to potassium iodide solution, for example, the more reactive chlorine displaces iodine from solution, and the colourless mixture turns brown as iodine is formed.

Group 0 – the noble gases

Group 0 elements are unreactive because their outer shell is already full, and this is why they exist as single, monatomic atoms rather than forming molecules. Boiling point increases going down the group as the atoms get larger and the intermolecular forces between them get stronger. The number of electrons in the outer shell equals the group number pattern used across the table (8 for Group 0, except helium with 2), which is why the noble gases sit in their own group rather than being folded into Group 7.

Topic 4.1.1.2 – separating mixtures

A mixture is two or more elements or compounds not chemically combined together, so the substances keep their own properties and can be separated by physical methods rather than chemical reactions: filtration, crystallisation, simple distillation, fractional distillation and chromatography, chosen according to the mixture’s components and the property being exploited (particle size, solubility, or boiling point).

Properties of transition metals (chemistry only)

Transition metals sit in the central block of the periodic table, between Group 2 and Group 3. Compared with the Group 1 alkali metals, transition metals are typical metals: they have much higher melting points (with the exception of mercury, which is liquid at room temperature), are considerably denser, and are harder and stronger, whereas Group 1 metals are soft enough to cut with a knife, have low melting points and low density. Transition metals are also much less reactive than Group 1 metals – they do not react vigorously with water or oxygen the way sodium or potassium do, which is why they are useful as structural and everyday materials.

Beyond the metallic comparison, transition metals share three typical properties: they can form ions with variable (more than one) ionic charge, for example iron forming Fe2+ and Fe3+; their compounds are typically coloured (copper compounds are often blue, iron(II) compounds green and iron(III) compounds orange-brown); and the elements and their compounds are often useful as catalysts, for example iron in the Haber process and manganese(IV) oxide as a catalyst for the decomposition of hydrogen peroxide.

Worked example

Chlorine exists as 75% chlorine-35 and 25% chlorine-37. Calculate its relative atomic mass.

Ar = [(35 x 75) + (37 x 25)] / 100
   = (2625 + 925) / 100
   = 3550 / 100
   = 35.5

The answer is not a whole number, which is expected for a weighted mean.

Common mistakes

Confusing atomic number with mass number when writing electronic structures. Stating that isotopes have different chemical properties — they do not, because they have identical electron arrangements. Forgetting that the number of electrons changes in an ion but the number of protons never does. Describing the alpha-scattering result as “most particles bounced back” when in fact the great majority passed straight through, which is the whole point of the evidence.

Quick revision checklist

• State the relative mass and charge of a proton, neutron and electron.
• Work out protons, neutrons and electrons from atomic and mass number.
• Define an isotope and explain why isotopes behave identically in reactions.
• Calculate relative atomic mass from isotopic abundances.
• Write the electronic structure of the first 20 elements and link it to position in the Periodic Table.
• Describe how the model of the atom changed, and what evidence drove each change.