GCSE Chemistry Atomic structure revision

Every atom has a tiny, dense, positively charged nucleus made of protons and neutrons, around which electrons move in shells. The numbers are extreme: an atom’s radius is about 1 × 10⁻¹⁰ m, while the nucleus is roughly ten thousand times smaller still — and yet the nucleus holds almost all of the mass. If an atom were a stadium, the nucleus would be a pea on the centre spot.

Three numbers label every atom. The atomic number counts protons and fixes the element’s identity and place in the periodic table. The mass number counts protons plus neutrons. In a neutral atom the electron count equals the proton count; in an ion, electrons are gained or lost but the nucleus never changes.

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons — carbon-12 with 6 neutrons and carbon-14 with 8, for example. Because chemistry depends on electrons, not neutrons, isotopes of an element react identically; only their masses (and, for some, their radioactivity) differ.

1. 1Dalton (early 1800s): solid spheres

   Matter was pictured as indivisible spheres, with each element made of its own identical balls. No subatomic particles were known.

2. 2Thomson (1897): the plum pudding

   Discovering the electron showed atoms could be split. Thomson pictured a ball of positive charge with electrons dotted through it like fruit in a pudding.

3. 3Rutherford, Geiger and Marsden (1909–1911): the nuclear model

   Firing alpha particles at thin gold foil, most sailed straight through, but a small number bounced back or deflected sharply. A spread-out positive “pudding” could not do that — so the positive charge and mass must sit in a tiny nucleus, with the rest mostly empty space.

4. 4Bohr (1913): shells

   Bohr proposed that electrons orbit the nucleus only at fixed distances — shells — which matched the light spectra and explained why atoms absorb and emit energy in fixed packets.

5. 5Chadwick (1932): the neutron

   The mass of atoms did not add up from protons alone. Chadwick proved the existence of uncharged neutrons, completing the modern nuclear model.

Electrons occupy shells (energy levels) that fill in the order 2, 8, 8. Sodium, with 11 electrons, is 2,8,1. Potassium, with 19, is 2,8,8,1. The electronic structure controls everything chemical: the number of outer-shell electrons equals the group number (for Groups 1–7 and 0), and the number of shells equals the period number.

Elements in the same group share their outer-shell count, which is why they react in similar ways — all Group 1 metals lose one electron, all Group 7 halogens gain one. Reactivity in each group is set by how easily that outer electron is lost or gained, which is exactly what the trends below explain.

Nineteenth-century chemists spotted repeating patterns of properties and tried to tabulate the elements by increasing atomic weight. Mendeleev (1869) went further: he left gaps for elements not yet found and swapped the order of a few pairs so that columns held elements of matching chemistry. When gallium and germanium turned up filling his gaps with the properties he predicted, the table became accepted as a scientific tool, not just a list.

The anomalies Mendeleev had to dodge — tellurium weighing more than iodine but sitting before it, for instance — were resolved only when protons were discovered and the table was reordered by atomic number. In the modern table, every element sits in order of proton count. Isotopes explain the old weight paradox too: tellurium’s average mass is raised by its heavy, abundant isotopes, but its identity — and therefore its place — is set by its 52 protons, one fewer than iodine.

The bold zig-zag line on the periodic table splits metals (left and centre) from non-metals (right, plus hydrogen). Atoms of metal elements have few outer electrons, which they lose to form positive ions; atoms of non-metals have nearly full outer shells, which they gain or share electrons to fill. That one difference explains the whole property list.

The Group 1 metals — lithium, sodium, potassium, rubidium, caesium, francium — are shiny when freshly cut but tarnish fast in air, and unlike most metals they are soft and have low densities (lithium, sodium and potassium float on water). Their single outer electron is easily lost, so they react vigorously with:

• Water → a metal hydroxide plus hydrogen (lithium fizzes steadily; potassium ignites its own hydrogen with a lilac flame).
• Oxygen → a metal oxide.
• Chlorine → a white metal chloride salt.

Reactivity increases down the group. The outer electron sits in a shell further from the nucleus, so the attraction pulling it in is weaker — the electron is lost more easily and the metal reacts harder. Higher-tier answers should say that in full: greater distance, weaker electrostatic attraction, easier electron loss.

The halogens — fluorine, chlorine, bromine, iodine — exist as diatomic molecules (F₂, Cl₂, Br₂, I₂) and need one electron to complete their outer shell. At room temperature chlorine is a pale green gas, bromine a red-brown liquid, iodine a grey solid: state and colour shift down the group as melting and boiling points rise.

Reactivity decreases down the group: the incoming electron must fall into a shell further from the nucleus, where attraction is weaker, so the pull on that electron is less. Consequences: fluorine is ferocious, iodine is tame; and a more reactive halogen displaces a less reactive one from its salts — chlorine water turns colourless potassium bromide solution orange as orange bromine is displaced.

The noble gases — helium, neon, argon, krypton, xenon, radon — are colourless, monatomic, unreactive gases. Their outer shells are full (helium 2, the rest 8), so they have no drive to gain, lose or share electrons: their lack of reactivity is exactly why helium fills balloons and argon shields welds and fills light bulbs.

Their boiling points increase down the group as the atoms grow heavier: bigger atoms have more electrons, so the intermolecular forces between them are stronger and more energy is needed to boil them. It is a neat inverse of Group 7: reactivity barely features here — mass does.

• Compared with Group 1

  Hard and strong

  Transition metals are mechanically tough — think girders and tools — while Group 1 metals are soft enough to cut with a knife.

• Compared with Group 1

  High melting points

  Iron melts at 1538 °C; mercury is the lone low-melting exception. Group 1 metals melt at low temperatures.

• Compared with Group 1

  Dense

  Transition metals pack heavy atoms tightly; lithium, sodium and potassium float.

• Chemistry

  Slow and steady

  Transition metals react slowly, if at all, with water and oxygen at room temperature — a world away from potassium on a pond.

Three more signatures complete the set. Transition metals form ions with several different charges (iron(II) Fe²⁺ and iron(III) Fe³⁺, copper(I) and copper(II)), their compounds are often coloured (blue copper(II) sulfate, green iron(II) sulfate, purple potassium manganate(VII)), and they or their compounds work as catalysts — iron in the Haber process, nickel in hydrogenation, manganese(IV) oxide speeding up hydrogen peroxide decomposition.