Classifying the Elements

The periodic table is organised so that elements with similar chemical behaviour line up in vertical groups.

It is structured by atomic number — the number of protons — and by electron configuration.

Elements are classified into blocks according to the sub-shell that the highest-energy electron occupies.

The s-block contains Groups 1 and 2: elements whose outermost electron is in an s sub-shell.

The p-block contains Groups 3 through 0 (or 13 through 18 in modern numbering), with the outermost electron in a p sub-shell.

The d-block sits between the s and p blocks, with the outermost electron in a d sub-shell — these are the transition metals and Sc/Zn.

The f-block (lanthanides and actinides) sits below the main table, with the outermost electron in an f sub-shell.

Periodic trends arise because elements in the same period have their outer electrons in the same shell, and elements in the same group have similar outer-shell configurations.

First Ionisation Energy

First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly-positively-charged gaseous ions.

Four conditions to include in the definition: one mole, gaseous atoms, gaseous ions, +1 charge.

The equation is always of the form:

X(g) → X⁺(g) + e⁻

State symbols are essential — AQA penalises any answer that omits them.

First ionisation energy is measured in kJ mol⁻¹ and is always positive, because energy must be supplied to overcome the attraction between the electron and the nucleus.

Factors Affecting Ionisation Energy

Three factors determine how much energy is required to remove an electron from an atom.

Nuclear charge.

More protons in the nucleus means a stronger electrostatic pull on every electron.

All else being equal, higher nuclear charge means higher ionisation energy.

Distance of the outer electron from the nucleus.

The further out the electron sits, the weaker the attraction (electrostatic force falls off with distance squared).

Greater distance means lower ionisation energy.

Shielding by inner electron shells.

Inner-shell electrons repel the outer electrons and partially cancel the attraction of the nucleus.

More inner shells means more shielding, and lower ionisation energy.

These three factors compete in opposite directions, and the actual trend depends on which dominates.

Trend Down a Group

Going down a group, first ionisation energy decreases.

Each step down adds another full shell of electrons, so the outer electron is in a higher principal quantum shell, further from the nucleus.

There are also more inner shells to provide shielding.

Nuclear charge does increase down the group, but the increase in distance and shielding more than compensates.

The outer electron is therefore easier to remove, and IE drops.

This is the trend behind the increasing reactivity of Groups 1 and 2 down the group — alkali metals like caesium lose their outer electron far more readily than lithium.

Trend Across a Period

Going across a period, first ionisation energy generally increases.

Each step across adds one proton (so nuclear charge rises) and one electron to the same shell, so shielding and distance from the nucleus stay roughly the same.

With nuclear charge dominating, the outer electron is held more tightly and harder to remove.

But the trend is not perfectly smooth. There are two well-known dips in Period 3.

The Two Famous Dips in Period 3

The first dip is between Group 2 and Group 3 — magnesium to aluminium.

Magnesium’s outer electron is in a 3s sub-shell.

Aluminium’s outer electron is in a 3p sub-shell.

The 3p sub-shell is at a slightly higher energy than 3s, so the 3p electron in aluminium is easier to remove than the 3s electron in magnesium — even though aluminium has one more proton.

Mg has a higher first IE than Al.

The second dip is between Group 5 and Group 6 — phosphorus to sulfur.

Phosphorus has three 3p electrons, one in each of the three 3p orbitals. They sit singly.

Sulfur has four 3p electrons, so one of the 3p orbitals now contains a paired set of electrons.

Paired electrons in the same orbital repel each other.

That extra repulsion makes one of sulfur’s 3p electrons easier to remove than would otherwise be expected, so sulfur has a lower first IE than phosphorus.

Both dips need to be explained in terms of sub-shell or orbital occupation — vague references to ‘energy levels’ won’t score full marks.

Successive Ionisation Energies

After the first electron has been removed, the second, third, fourth and so on can also be removed in turn.

Each successive ionisation energy is greater than the last, because each electron is being removed from an increasingly positive ion.

Plotting successive ionisation energies against the number of electrons removed reveals jumps — sudden, much larger increases.

Each jump corresponds to the next electron coming from a shell closer to the nucleus.

An electron in a more-inner shell experiences less shielding and is much closer to the nucleus, so it requires far more energy to remove.

The pattern of jumps tells you which group an element belongs to.

If there is a big jump between the first and second ionisation energies, the element is in Group 1.

If the jump is between the second and third, it is Group 2.

Between third and fourth, Group 3, and so on.

Successive ionisation energies are therefore strong experimental evidence for the existence of electron shells.

Atomic Radius Across Period 3

Atomic radius decreases steadily across Period 3 from sodium to argon (though noble gas radii are estimated).

Going across the period, the number of protons increases by one each time, but the outer electrons are being added to the same third shell.

Higher nuclear charge pulls all the electrons in more tightly, shrinking the atom.

Down a group, atomic radius increases as electrons occupy higher shells further from the nucleus.

Atomic radius is an important factor underlying ionic radii, lattice energies, and reactivity trends.

Melting Points Across Period 3

The trend in melting points across Period 3 is dramatic and depends on the structure and bonding of each element.

Sodium, magnesium and aluminium are metallic.

Their melting points increase from Na to Al because each metal donates more delocalised electrons (one, two, three respectively) and has smaller ionic radii — both of which strengthen the metallic bond.

Silicon is the giant covalent element — a macromolecular structure of Si atoms each covalently bonded to four neighbours.

Breaking such a structure requires breaking strong covalent bonds throughout the lattice.

Silicon has by far the highest melting point in Period 3.

Phosphorus, sulfur, chlorine and argon are all molecular.

Phosphorus exists as P₄, sulfur as S₈, chlorine as Cl₂, argon as single atoms.

Their melting points depend on the strength of the van der Waals forces between molecules.

Larger molecules with more electrons have stronger van der Waals forces — so S₈ (the largest) has the highest molecular melting point in the run, P₄ next, then Cl₂, then Ar with the lowest.

Crucially, melting molecular substances does not break covalent bonds. It only overcomes the weaker intermolecular forces.

This is why the molecular elements all have much lower melting points than silicon, despite some of them having quite large molecules.