The periodic table arranged by atomic number into periods and groups, the s, p and d blocks, and the periodic trends in atomic radius, first ionisation energy and melting point across Periods 2 and 3.

What this topic is asking

OCR specification point 3.1.1 wants you to explain how the modern periodic table is arranged by atomic number into periods and groups, to assign elements to the s, p and d blocks from their electron configuration, and to describe and explain the periodic trends in atomic radius, first ionisation energy and melting point across Periods 2 and 3. Periodicity is the repeating pattern of properties that follows directly from electron structure.

Structure of the periodic table

Elements are placed in blocks named after the sub-shell that holds the highest-energy electron: the s-block (Groups 1 and 2), the p-block (Groups 13 to 18), and the d-block (the transition metals). Knowing the block lets you write the electron configuration directly from an element's position.

Trend in first ionisation energy

The first ionisation energy is the energy to remove one electron from each atom in one mole of gaseous atoms:

Across a period the nuclear charge increases while the electrons enter the same shell with similar shielding, so the outer electrons are held more tightly. Down a group the outer electron is in a higher shell, further from the nucleus and better shielded, so it is easier to remove.

Successive ionisation energies

Removing electrons one after another gives successive ionisation energies that always rise, because each electron is pulled from an increasingly positive ion. Large jumps mark the move to a new, inner shell, which is evidence for shell structure and lets you identify the group of an element.

Trend in atomic radius

Atomic radius decreases across a period because the rising nuclear charge pulls the same outer shell inward. It increases down a group because each new period adds a shell, and the extra shielding outweighs the higher nuclear charge.

Trend in melting point

Examples in context

Example 1. Predicting reactivity. Because first ionisation energy falls down Group 1, caesium loses its outer electron far more readily than lithium, which is why the Group 1 metals get more reactive down the group.

Example 2. Identifying an element from data. A jump between the second and third successive ionisation energies tells you the element is in Group 2, because the third electron must come from a stable inner shell.

Try this

Q1. State the block of the periodic table that contains iron and explain your answer. [2 marks]

• Cue. The d-block, because iron's highest-energy electrons fill the $3\text{d}$ sub-shell ($[\text{Ar}]3\text{d}^6 4\text{s}^2$).

Q2. Explain why the first ionisation energy of sulfur is lower than that of phosphorus. [2 marks]

• Cue. Sulfur's $3\text{p}^4$ has a paired electron in one $3\text{p}$ orbital; the pair repulsion makes one electron easier to remove than from phosphorus's half-filled $3\text{p}^3$.