GCSE Chemistry: Atomic Structure and the Periodic Table — Revision Notes

Condensed for the final weeks. For the full explanation of the atomic-structure half, use the Atomic Structure study guide.

This file spans two separate specification topics. The atomic models and isotopes sections below cover C1.2 Atomic structure, assessed on Papers 1 and 3. The periodic table, Group 1, Group 7 and Group 0 sections cover C4.1 Predicting chemical reactions, assessed on Papers 2 and 4 – revise that half alongside your Paper 2/4 preparation, not your Paper 1/3 preparation.

Development of the atomic model

Model	Contribution
Dalton	Atoms are tiny indivisible spheres
Thomson	Plum pudding — electrons embedded in positive charge
Rutherford	Nuclear model — small dense positive nucleus, mostly empty space
Bohr	Electrons in fixed shells at set distances
Chadwick	Discovery of the neutron

The alpha-scattering experiment is the standard question, and the reasoning must be given in three parts:

1. Most alpha particles passed straight through → the atom is mostly empty space.
2. Some were deflected → the nucleus is positively charged, repelling the positive alpha particles.
3. A very few bounced back → the nucleus is very small but contains most of the mass.

Each observation maps to one conclusion; giving the observations alone earns nothing.

This is also the standard example of how scientific models change when new experimental evidence cannot be explained by the existing model.

Atoms and isotopes

Particle	Relative mass	Charge	Location
Proton	1	+1	Nucleus
Neutron	1	0	Nucleus
Electron	Very small	−1	Shells

Atomic number = protons; mass number = protons + neutrons; neutrons = mass number − atomic number.

Isotopes have the same number of protons but different numbers of neutrons. They react identically because chemical behaviour depends on electrons.

relative atomic mass = sum of (isotope mass x abundance) / 100

That weighted average is why relative atomic masses are not whole numbers. Chlorine is the standard example: it exists as two isotopes, chlorine-35 and chlorine-37, and its relative atomic mass of 35.5 is the weighted average of the two, allowing for how abundant each isotope actually is — not a simple average of 35 and 37.

Atoms and small molecules have a typical radius on the order of 10⁻¹⁰ metres, a scale worth recalling directly since questions may ask you to compare it with other sizes, such as a human hair or a bacterium.

Electron configuration

Shells fill 2, 8, 8. Write configurations as 2,8,1 for sodium.

• Group number = number of outer-shell electrons (for the main groups).
• Period number = number of occupied shells.

Elements in the same group react similarly because they have the same number of outer electrons, and it is the outer electrons that take part in reactions. That sentence answers a large share of periodic-table questions.

The periodic table (C4.1 – Papers 2 and 4)

Arranged by increasing atomic number. Mendeleev left gaps for undiscovered elements and predicted their properties — which is why his table is remembered while earlier attempts are not, and why the discovery of gallium and germanium confirmed it.

• Metals lose electrons to form positive ions; found on the left.
• Non-metals gain or share electrons; found on the right.

Group 1 — alkali metals

React vigorously with water: metal + water → metal hydroxide + hydrogen.

Reactivity increases down the group. The explanation must have all three parts:

1. The outer electron is in a shell further from the nucleus.
2. There is more shielding by inner shells.
3. So the attraction between nucleus and outer electron is weaker, and the electron is lost more easily.

Group 7 — halogens

Diatomic non-metals. Melting and boiling points increase down the group (larger molecules, stronger intermolecular forces), but reactivity decreases — the opposite direction to Group 1.

Why: a halogen atom must gain an electron. Further down the group, the outer shell is further from the nucleus with more shielding, so the incoming electron is attracted less strongly and is gained less readily.

Displacement reactions: a more reactive halogen displaces a less reactive one from its salt solution — chlorine displaces bromine, bromine displaces iodine.

Group 0 — noble gases

Unreactive because they have full outer shells, so they have no tendency to lose, gain or share electrons. Boiling points increase down the group as the atoms get larger and intermolecular forces strengthen.

Exam traps

• Giving the alpha-scattering observations without the conclusions.
• Saying reactivity increases down both Group 1 and Group 7.
• Giving only one reason for the Group 1 trend — all three parts are needed.
• Saying isotopes react differently.
• Saying noble gases are unreactive because they are gases.
• Confusing group and period numbers.

Self-test

1. Give the three alpha-scattering observations and what each shows.
2. Why do isotopes have identical chemical properties?
3. Explain fully why Group 1 reactivity increases down the group.
4. Why does Group 7 reactivity decrease down the group?
5. Why are elements in the same group chemically similar?

Answers: 1. Most particles passed straight through — the atom is mostly empty space; some were deflected — the nucleus is positively charged; a very few bounced back — the nucleus is very small and holds most of the mass. 2. Chemical properties depend on electron arrangement, which is identical; only neutron number differs. 3. The outer electron is further from the nucleus, there is more shielding from inner shells, so the attraction between the nucleus and the outer electron is weaker and the electron is lost more easily. 4. A halogen must gain an electron; further down the group the outer shell is further from the nucleus with more shielding, so an incoming electron is attracted less strongly. 5. They have the same number of outer-shell electrons, and it is the outer electrons that determine chemical behaviour.