Revision notes: The Periodic Table of Elements

These notes work through The Periodic Table of Elements one content standard at a time. The whole chapter rewards a single skill: reading an element’s position and using it to predict structure, properties and reactivity. Keep these notes beside past-paper questions and use them to check your reasoning, not just your answers.

4.1 Development of the periodic table

The exam expects you to sequence the scientists who built the table and to say what each contributed.

• Antoine Lavoisier grouped the known substances into metals and non-metals, a first, rough sort.
• Johann Döbereiner arranged elements with similar properties into triads of three, where the middle element’s relative atomic mass is about the average of the other two.
• John Newlands set out his law of octaves: arranged by increasing atomic mass, every eighth element had similar properties, like notes on a scale.
• Dmitri Mendeleev arranged elements by increasing relative atomic mass, grouped those with similar chemical properties in columns, and, crucially, left gaps for elements not yet discovered and predicted their properties. His predictions were later confirmed.
• Henry Moseley re-ordered the table by increasing proton number (atomic number), which is the basis of the modern periodic table and fixed the few places where mass order gave the wrong sequence.

4.2 Groups and periods

In the modern table, elements are arranged by increasing proton number. Vertical columns are groups and horizontal rows are periods.

• For the main groups, the group number tells you the number of valence electrons; the period number tells you the number of occupied shells. Sodium (2.8.1) is in Group 1, Period 3.
• Elements in the same group have the same number of valence electrons, so they have similar chemical properties.

Trends across a period (left to right): proton number rises by one each step; a valence electron is added to the same shell; the number of shells stays the same; the atomic radius decreases because the growing nuclear charge pulls the shells inward; and the elements change from metallic to non-metallic.

Trends down a group (top to bottom): the number of shells increases by one each step; the atomic size increases; but the number of valence electrons stays the same, which is why the group keeps its family character.

4.3 Group 1: alkali metals

Lithium, sodium, potassium and the rest each have one valence electron. They are soft, silvery, low-density metals that conduct electricity, and they are very reactive, so they are stored under paraffin oil.

• With water they form an alkaline metal hydroxide and hydrogen gas: for example, sodium gives sodium hydroxide and hydrogen and skates across the surface.
• With oxygen they form a white metal oxide that dissolves in water to give an alkali.
• Down the group, reactivity increases. The atom gets larger, the single valence electron is further from the nucleus and less strongly held, so it is lost more easily, and losing that electron is what the reaction requires.

4.4 Group 17: halogens

Fluorine, chlorine, bromine and iodine each have seven valence electrons and exist as diatomic molecules.

• Their colour deepens and physical state changes down the group: chlorine is a greenish-yellow gas, bromine a reddish-brown liquid, iodine a purplish-black solid, so melting and boiling points rise down the group.
• They react with water, with iron to form a halide, and with alkali.
• Down the group, reactivity decreases. The atom gets larger, so the nucleus attracts an incoming electron less strongly, and gaining an electron is what a halogen reaction requires.
• In a displacement reaction, a more reactive halogen displaces a less reactive one from its salt solution, chlorine displaces bromine from potassium bromide, and the solution turns brown.

4.5 Group 18: noble gases

Helium, neon, argon and the others have a stable, full valence shell (helium has 2 electrons, the rest have 8). Because their arrangement is already stable, they are chemically inert, exist as single atoms (monatomic) and are colourless gases.

• Uses follow directly from being unreactive and, for helium, low-density: helium in balloons and airships, neon in display lighting, and argon as an inert filling in light bulbs and in welding.
• Down the group atomic size, boiling point and density all increase.

4.6 Period 3 and transition elements

Across Period 3, sodium, magnesium, aluminium, silicon, phosphorus, sulfur, chlorine, argon, the proton number rises, the atomic size decreases, and the elements shift from metal to metalloid to non-metal. The oxides change from basic (sodium, magnesium) through amphoteric (aluminium) to acidic (phosphorus, sulfur).

Transition elements are the block of metals such as iron, copper, zinc, chromium and manganese. Their special properties are worth memorising:

• they form coloured ions and compounds;
• they show variable oxidation numbers (iron as Fe²⁺ and Fe³⁺, copper as Cu⁺ and Cu²⁺);
• they act as catalysts (iron in the Haber process, vanadium(V) oxide in the Contact process, nickel in hydrogenation);
• they form complex ions, and are hard, strong metals with high melting points and densities.

How to use these notes: predict from position

The single most valuable habit in this chapter is turning a position into a prediction. Given an element’s group and period, state its number of valence electrons and shells, decide whether it is a metal or non-metal, and predict how its size and reactivity compare with a neighbour. For a Group 1 or Group 17 element, always justify a reactivity trend with the same three ideas, atomic size, distance of the valence shell from the nucleus, and how easily an electron is lost or gained.

A one-to-one teacher can check that every trend you state is backed by an electron-arrangement reason rather than left as a bare fact, that is exactly where marks are won and lost across the SPM Chemistry written papers. Because groups and periods reappear in Chemical Bond and in every reactivity comparison later, securing this reasoning now pays back through Form 4 and Form 5.

Worked reasoning: comparing two metals

Suppose you must compare the reactivity of sodium and potassium with water. Both are Group 1 with one valence electron, so both react to give a hydroxide and hydrogen. Potassium is lower in the group, so it has more shells and a larger atom; its valence electron is further from the nucleus and held less tightly, so it is lost more easily and potassium is more reactive. Write the answer in exactly that order, position, structure, then prediction, and the marks follow.