3.1.1 Periodicity

Read the modern periodic table from left to right in order of increasing proton number. When you reach the end of a period, the pattern of outer-electron arrangements starts again, which is why chemical properties repeat periodically.

Elements in the same group have similar outer-electron arrangements, so you should expect similar chemistry. Use electron structure to explain the similarity; the group number on its own is only a label.

Across Period 2, you are filling the 2s and then the 2p sub-shell. Across Period 3, the same pattern repeats with 3s then 3p. Each step to the right adds one proton to the nucleus and one electron to the atom.

For example, Na is [Ne]3s¹, Mg is [Ne]3s², then Al starts the 3p sub-shell: [Ne]3s²3p¹. That change from s to p becomes useful when you explain the ionisation-energy dip from Mg to Al.

To decide the block, look at the sub-shell receiving the highest-energy electron in the ground-state atom. If electrons are being added to an s sub-shell, you are in the s block; p gives the p block; d gives the d block.

So Mg, ending 3s², is s-block. Chlorine, ending 3p⁵, is p-block. Scandium starts filling 3d, so it sits in the d block.

For first ionisation energy, I’d learn the definition accurately: it is the energy required to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1+ ions.

The words “gaseous atoms” and “one mole” matter. This is a molar energy change, so the usual unit is kJ mol⁻¹.

When you explain an ionisation-energy trend, compare the attraction between the nucleus and the electron being removed. Across a period, nuclear charge increases while shielding is broadly similar and the outer electron is in the same main shell, so the attraction generally gets stronger and first ionisation energy rises.

Down a group, the outer electron is in a higher shell, further from the nucleus and more shielded. Those effects outweigh the larger nuclear charge, so the attraction is weaker and first ionisation energy falls.

The overall rise across a period has two familiar dips. From Mg to Al, the electron removed from Al is in a 3p orbital, which is higher in energy than Mg’s 3s electron, so it is easier to remove despite the increased nuclear charge.

From P to S, sulfur has a pair of electrons in one 3p orbital. Repulsion within that pair makes one electron easier to remove than you might expect. Name the actual sub-shell or paired orbital, not just “more repulsion”.

Look for the big jump in successive ionisation energies. Before the jump, you are removing electrons from the same outer shell; after it, the next electron comes from an inner shell, where it is much more strongly attracted to the nucleus.

If the large jump comes between the second and third ionisation energies, the atom had two outer-shell electrons, so it is consistent with Group 2. Count how many electrons were removed before the jump.

In a metal, picture a giant lattice of positive metal ions surrounded by delocalised electrons. Metallic bonding is the electrostatic attraction between those positive ions and the delocalised electrons.

The electrons are free to move through the giant structure. That mobility is central when you explain electrical conductivity, while the strong attraction throughout the lattice helps explain high melting temperatures for many metals.

When you compare these structures, keep asking how many strong covalent bonds each atom makes and whether any electrons can move. Diamond has a three-dimensional carbon network with each carbon bonded to four others. Graphite has layers: each carbon bonds to three others, leaving one electron per carbon delocalised within the layer.

Graphene is essentially one carbon layer, one atom thick, with the same three-bonds-plus-delocalised-electron pattern. Silicon forms a giant covalent network similar in broad shape to diamond.

• diamond: 4 bonds per C, 3D network
• graphite: 3 bonds per C, layers, delocalised electrons
• graphene: single layer, 3 bonds per C, delocalised electrons
• silicon: giant covalent network

Build each property explanation from the structure. If a giant lattice has many strong bonds or attractions that must be overcome, you expect a high melting temperature. If charged particles or delocalised electrons can move, you can explain conductivity.

For graphite, the weak attractions between layers let the layers slide, while delocalised electrons conduct along the layers. Diamond has no mobile charged particles, so it does not conduct electricity. Metals conduct because their delocalised electrons can move through the lattice.

Across a period, you should not try to force one ionisation-energy-style trend onto the melting points. First identify the structure of each element, then compare the attractions or bonds that melting has to overcome.

In Period 3, Na, Mg and Al are giant metallic; stronger metallic bonding generally raises the melting point across those three. Silicon has a giant covalent lattice, so many strong covalent bonds must be broken. P₄, S₈, Cl₂ and Ar are simple molecular or atomic, so their much lower melting points depend mainly on London forces; larger electron clouds usually give stronger London forces.