The building blocks

• An atom is the smallest part of an element that can exist. Each element has a chemical symbol — e.g. O for oxygen, Na for sodium.
• There are about 100 different elements, shown in the periodic table.
• A compound contains two or more elements chemically combined in fixed proportions (e.g. H₂O). Compounds form from elements by chemical reactions, which make one or more new substances and often involve a detectable energy change.
• A compound can only be split back into its elements by a chemical reaction — not by physical means.

Element, compound or mixture?

Word & balanced symbol equations

Reactions are written as word equations or, more precisely, as balanced symbol equations. Atoms are never created or destroyed, so the same number of each type of atom must appear on both sides.

The big numbers in front (2H₂) are balancing numbers — you may change these but never the small formula subscripts. Adding state symbols shows the physical state: 2H₂(g) + O₂(g) → 2H₂O(l).

HT only: you also write ionic equations and balanced half equations where appropriate.HT

Your turn — balance the equation

Mixtures and how to separate them

A mixture is two or more elements or compounds not chemically combined — each keeps its own properties. Because nothing is bonded, mixtures are separated by physical processes (no new substances are made):

• Filtration — separates an insoluble solid from a liquid.
• Crystallisation — evaporates a solvent to leave soluble solid crystals.
• Simple distillation — separates a solvent from a dissolved solid by boiling and condensing.
• Fractional distillation — separates liquids with different boiling points (e.g. crude oil, air).
• Chromatography — separates dissolved substances by how they travel up a paper.

Pick the right technique

Tap a job on the left, then its separation technique on the right.

How our model of the atom changed

New experimental evidence keeps reshaping the model:

• Dalton — atoms are tiny solid spheres that cannot be divided.
• Plum pudding — after the electron was discovered, the atom was pictured as a ball of positive charge with negative electrons embedded in it.
• Rutherford (nuclear model) — the alpha-particle scattering experiment showed mass and positive charge are concentrated in a tiny central nucleus.
• Bohr — electrons orbit the nucleus at fixed distances (energy levels); his calculations matched observations.
• Protons & the neutron — the nuclear charge was found to be made of whole-number protons; later, James Chadwick provided evidence for the neutron.

Reading the scattering result

Protons, neutrons & electrons

Atoms contain three subatomic particles. The exam expects their relative charges and relative masses:

• Atomic number = number of protons (defines the element).
• Mass number = protons + neutrons.
• Atoms are tiny: radius about 0.1 nm (1 × 10⁻¹⁰ m). The nucleus radius is less than 1/10 000 of the atom (about 1 × 10⁻¹⁴ m), yet holds almost all the mass.

Watch out: the atom is mostly empty space — if the nucleus were a marble, the atom would be the size of a stadium.

Your turn — count the neutrons

Isotopes

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons — so they share the atomic number but have different mass numbers.

Misconception: isotopes differ in neutrons, not protons. Change the protons and you change the element entirely.

Relative atomic mass (Aᵣ)

Because an element is a mix of isotopes, its relative atomic mass (Aᵣ) is an average that takes account of the abundance of each isotope — a weighted mean, not a simple average.

Misconception: Aᵣ is a weighted mean, so it lands closer to the most abundant isotope — that's why chlorine's Aᵣ (35.5) is nearer 35 than 37.

Your turn — relative atomic mass

Electron shells (energy levels)

Electrons fill the lowest available energy levels (innermost shells) first. The shells hold up to 2, then 8, then 8 electrons. We write the configuration as numbers, e.g. sodium = 2,8,1.

Link: the number of outer electrons equals the group number (so sodium, with 1 outer electron, sits in Group 1).

Build a configuration

How the table is arranged

Elements are arranged in order of atomic (proton) number, so that elements with similar properties line up in columns:

• A group is a column. Elements in a group have the same number of outer electrons, giving similar chemical properties.
• A period is a row. Going across, an outer shell is being filled.

From Newlands to the modern table

• Early tables ordered elements by atomic weight. Newlands' law of octaves grouped every eighth element, but it forced unlike elements together and was rejected.
• Mendeleev also ordered by atomic weight but left gaps for undiscovered elements and swapped some orders so properties matched. His predictions for the gaps were later confirmed.
• Knowledge of isotopes later explained why ordering by atomic weight sometimes failed — the modern table is ordered by atomic number.

Why Mendeleev is celebrated: leaving gaps and predicting properties of missing elements meant his table could be tested — and it passed.

What made Mendeleev's table work?

Metals vs non-metals

• Elements that react to form positive ions are metals.
• Elements that do not form positive ions are non-metals.
• Most elements are metals — found to the left and towards the bottom. Non-metals sit towards the right and top.

This links directly to electronic structure: an atom's position in the table reflects its electron arrangement, which controls how it reacts.

Group 0 — the noble gases

• The noble gases are unreactive because their atoms have stable, full outer shells — 8 outer electrons (helium has only 2).
• They don't easily form molecules.
• Boiling points increase going down the group (with increasing relative atomic mass).

Misconception: noble gases are unreactive because their outer shell is full — not because they're "heavy" or "rare".

Group 1 — the alkali metals

The alkali metals (Li, Na, K…) each have a single outer electron, giving characteristic reactions. They react with:

• Oxygen → metal oxides (e.g. 4Na + O₂ → 2Na₂O).
• Chlorine → metal chlorides (e.g. 2Na + Cl₂ → 2NaCl).
• Water → metal hydroxide + hydrogen (e.g. 2Na + 2H₂O → 2NaOH + H₂), fizzing on the surface.

Going down Group 1, reactivity increases — the outer electron is further from the nucleus and more easily lost.

Group 1 reactivity

Group 7 — the halogens

The halogens are non-metals with 7 outer electrons, existing as diatomic molecules (Cl₂, Br₂, I₂). Going down the group:

• Relative molecular mass, melting point and boiling point all increase.
• Reactivity decreases (the opposite trend to Group 1).

A more reactive halogen displaces a less reactive one from a solution of its salt — a displacement reaction:

Watch out: Group 7 reactivity goes down the group, while Group 1 goes up — don't mix the trends.

Which group?

Tap a property, then the group it describes.

Transition metals vs Group 1

The transition elements (Cr, Mn, Fe, Co, Ni, Cu…) are metals, but very different from the Group 1 alkali metals:

Typical properties (HT): many transition elements form ions with different charges, make coloured compounds, and act as catalysts — exemplified by Cr, Mn, Fe, Co, Ni, Cu.HT

Spot the transition metal

The facts to know

You've covered all of AQA 4.1 — atoms & equations, mixtures & separation, the atomic model, particles, isotopes, Aᵣ, electronic structure and the periodic table. Press Finish to see your score.