GCSE Combined Science Atoms & periodic table revision

An element is a substance made of only one kind of atom — about 100 exist. A compound forms when atoms of different elements bond chemically in fixed ratios (water is always H₂O); compounds can only be split by chemical reactions. A mixture is elements or compounds jumbled together without bonding — so the substances keep their own properties and can be retrieved by physical separation.

All separation methods are physical — no new substances form. That is the definition-level difference between separating a mixture and decomposing a compound.

The alpha-scattering logic is worth remembering: if the positive charge were spread out, alpha particles would sail through undeflected. The rare large deflections proved a concentrated positive nucleus. Most of the atom is empty space.

An atom’s atomic number counts protons (and, in a neutral atom, electrons); its mass number counts protons plus neutrons. Isotopes are atoms of the same element with the same proton number but different neutron numbers — chemically identical, slightly different in mass.

Electrons occupy shells: the first holds 2, the second holds 8, the third holds 8 (up to element 20). The configuration tells you an element’s story — the number of shells is its period (row), and the electrons in the outer shell are its group (column). Potassium, atomic number 19, fills 2, 8, 8, 1: four shells (period 4), one outer electron (Group 1).

Early tables lined elements up by atomic mass and spotted repeating patterns of properties. Mendeleev (1869) made three decisive moves: he ordered elements so that similar properties fell into columns, he left gaps for elements not yet discovered — predicting their properties (germanium, gallium) — and he swapped a few orders where properties demanded it. When those elements were found with the predicted properties, his table was accepted. The modern table orders elements by atomic (proton) number — a fix that explained the few mass-order anomalies, since protons, not mass, define an element.

Group 1 metals — lithium, sodium, potassium and below — have one outer electron, which they lose easily in reactions. They are soft (cuttable with a knife), low density, and react vigorously with water: they float, fizz (hydrogen gas) and form an alkaline metal hydroxide. Lithium fizzes steadily, sodium races about with a orange flame, potassium ignites instantly — reactivity increases down the group because the outer electron sits further from the nucleus and is lost more easily.

Halogens have seven outer electrons — one short of a full shell — so they react by gaining an electron. Going down the group their boiling points rise, and their reactivity falls: the outer shell is further from the nucleus, so the pull on an extra electron weakens. The displacement rule follows: a more reactive halogen displaces a less reactive one from its salt solution. Chlorine water turns colourless potassium bromide solution orange (chlorine takes bromide’s place); adding iodine solution to chlorine bromide does nothing — iodine is the weakest of the three.

The noble gases — helium, neon, argon, krypton, xenon — have full outer shells (helium 2; the rest 8), so they need neither to give nor take electrons: they are monatomic and strikingly unreactive. Their boiling points increase down the group as the atoms grow heavier and their intermolecular attractions strengthen — argon condenses far colder than helium, and only the heaviest form solids near absolute zero. Uses follow properties: helium’s low density lifts balloons; argon’s inertness shields filament bulbs and welding from oxygen; neon glows in signs.