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

Condensed for the final weeks. For the full explanation, use the Atomic Structure and the Periodic Table study guide.

Development of the model

Dalton (indivisible spheres) → Thomson (plum pudding) → Rutherford (nuclear model) → Bohr (shells) → Chadwick (neutron).

Alpha scattering — each observation with its conclusion:

1. Most particles passed straight through → the atom is mostly empty space.
2. Some were deflected → the nucleus is positively charged.
3. A very few rebounded → the nucleus is very small and holds most of the mass.

Listing the observations without the conclusions earns nothing. This is also the standard example of a model changing when new evidence cannot be explained by the existing one.

Atoms and isotopes

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

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

Isotopes have the same protons, different neutrons, and identical chemical properties — chemistry depends on electrons.

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

That weighted average is why relative atomic masses are rarely whole numbers.

Worked example. Chlorine has two isotopes, ³⁵Cl (75%) and ³⁷Cl (25%). Relative atomic mass = (35 × 75 + 37 × 25) ÷ 100 = (2625 + 925) ÷ 100 = 35.5.

Electron configuration

Shells fill 2, 8, 8.

• Group number = outer-shell electrons.
• Period number = number of occupied shells.

Elements in the same group react similarly because they have the same number of outer electrons, and outer electrons determine chemical behaviour. That sentence answers a large share of periodic-table questions.

The periodic table

Arranged by increasing atomic number. Mendeleev left gaps for undiscovered elements and predicted their properties — the discovery of gallium and germanium confirmed his arrangement, which is why his table succeeded where earlier ones failed.

Metals lose electrons to form positive ions (left); non-metals gain or share (right).

Group 1 — alkali metals

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

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

1. The outer electron is further from the nucleus.
2. There is more shielding from inner shells.
3. So the attraction is weaker and the electron is lost more easily.

Group 7 — halogens

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

The reason is the same physics applied to the opposite process: a halogen must gain 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.

Displacement: a more reactive halogen displaces a less reactive one from solution — chlorine displaces bromine, bromine displaces iodine. Worked observation: adding chlorine water to potassium iodide solution, the more reactive chlorine displaces iodine, and the colourless solution turns brown as iodine forms.

Group 0 — noble gases

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

Separation techniques

Technique	Separates
Filtration	Insoluble solid from liquid
Crystallisation	Soluble solid from solution
Simple distillation	Solvent from solution
Fractional distillation	Liquids with different boiling points
Chromatography	Substances by different solubility in the solvent

Choose by the property being exploited — solubility, particle size, or boiling point. Naming the technique without the property is a partial answer.

Worked example — sand and salt solution. To get dry sand, use filtration: the insoluble sand is trapped by the filter paper, exploiting the difference in particle size between the solid and the solution. To get pure, dry salt crystals from the filtered solution, use crystallisation: heat to evaporate some of the water until the solution becomes saturated, then leave the rest to evaporate (or cool) so the solute the solution can no longer hold comes out of solution as crystals.

Exam traps

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

Self-test

1. Give the three alpha-scattering observations and their conclusions.
2. Why is relative atomic mass rarely a whole number?
3. Explain fully why Group 1 reactivity increases down the group.
4. Why does Group 7 reactivity decrease down the group?
5. Which technique separates two liquids, and on what property?

Answers: 1. Most passed through — the atom is mostly empty space; some deflected — the nucleus is positive; a few rebounded — the nucleus is tiny and holds most of the mass. 2. It is a weighted average of the masses of the element’s isotopes, allowing for their relative abundances. 3. The outer electron is further from the nucleus, there is more shielding by inner shells, so the attraction between nucleus and outer electron is weaker and it is lost more easily. 4. A halogen must gain an electron, and further down the group the outer shell is further from the nucleus with more shielding, so the incoming electron is attracted less strongly. 5. Fractional distillation, exploiting their different boiling points.