Periodic trends

The key idea: nuclear charge and shielding

An outer electron is attracted by the positive nucleus but repelled by all the electrons between it and the nucleus. Inner (core) electrons shield the outer electrons from part of the nuclear charge. The net pull felt by an outer electron is called the effective nuclear charge, roughly the nuclear charge minus the number of core electrons.

• Across a period (left to right): protons are added to the nucleus, but the extra electrons go into the same outer shell, which shields very little. Effective nuclear charge rises, so outer electrons are pulled closer and held more tightly.
• Down a group: each element has one more shell. The outer electrons are further from the nucleus and shielded by more inner shells, so they are held less tightly, even though the nuclear charge is larger.

Every explanation of a trend should use these two ideas: the size of the attraction (nuclear charge versus shielding) and the distance of the outer electrons from the nucleus.

Atomic radius

Atoms have no sharp edge, so the atomic radius is usually taken as half the distance between the nuclei of two identical atoms bonded together (the covalent radius).

• Decreases across a period: increasing effective nuclear charge pulls the same shell closer (Na 166 pm to Cl 102 pm).
• Increases down a group: each extra shell puts the outer electrons further out (Li 128 pm, Na 166 pm, K 203 pm).

Ions. A cation is smaller than its atom: losing the outer electrons often removes a whole shell, and the remaining electrons are pulled in more tightly. An anion is larger: added electrons repel each other and the same nuclear charge holds more electrons. In an isoelectronic series, more protons means a smaller ion: O²⁻ (140 pm) > F⁻ (133 pm) > Na⁺ (102 pm) > Mg²⁺ (72 pm), all with 10 electrons.

Ionisation energy

The first ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms:

$$\text{Na}(g) \rightarrow \text{Na}^+(g) + e^- \qquad \Delta H = +496\text{ kJ mol}^{-1}$$

• Increases across a period: higher effective nuclear charge and smaller radius hold the outer electron more tightly.
• Decreases down a group: the outer electron is further away and better shielded.
• Noble gases sit at the peaks; group 1 metals at the troughs.

Two small dips reveal subshells. Boron (801) is lower than beryllium (900) because its outer electron is in 2p, slightly higher in energy and better shielded than 2s. Oxygen (1314) is lower than nitrogen (1402) because oxygen's fourth 2p electron has to pair up in an orbital that already holds one electron, and the repulsion between the pair makes it easier to remove. The same dips appear at aluminium and sulfur in period 3.

Successive ionisation energies

Electrons can be removed one at a time: the first, second, third ionisation energies and so on. Each is larger than the one before, because each electron is removed from an increasingly positive ion. But the increase is not steady. When all the valence electrons have gone, the next electron must come from a full inner shell that is much closer to the nucleus and far less shielded, and the ionisation energy jumps. Counting the electrons removed before the big jump gives the number of valence electrons, and so the group.

Electronegativity

Electronegativity is the ability of an atom in a covalent bond to attract the shared electrons towards itself. The most common scale is Linus Pauling's, on which fluorine has the highest value (3.98) and caesium and francium the lowest (about 0.8). Electronegativity increases across a period (greater effective nuclear charge, smaller atom, so shared electrons are closer to a stronger pull) and decreases down a group (bonding electrons further from the nucleus and better shielded). Noble gases are usually given no value because helium, neon and argon form no stable compounds. Electronegativity is central to bonding in lessons 1G to 1I.

Metallic character, reactivity and state

Metals tend to lose electrons; non-metals tend to gain or share them. Metallic character decreases across a period and increases down a group, following ionisation energy in reverse: the easier it is to lose an outer electron, the more metallic the element. A diagonal band of metalloids (B, Si, Ge, As, Sb, Te) separates metals from non-metals.

• Group 1 and 2 metals with water: reactivity increases down the group because the outer electron is lost more easily. 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g). Group 2 metals are less reactive than group 1 in the same period because they must lose two electrons, held by a larger nuclear charge: magnesium barely reacts with cold water but burns in steam.
• Halogens: reactivity decreases down group 17, because a halogen reacts by gaining an electron, and the incoming electron is attracted less strongly by larger atoms. Fluorine is the most reactive non-metal. A more reactive halogen displaces a less reactive one from its salts: Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq).

State and melting point across period 3 depend on structure, not just on the atoms: Na, Mg and Al are metallic solids with rising melting points (more delocalised electrons and smaller cations); silicon is a covalent network with the highest melting point (1414 °C); P₄, S₈, Cl₂ and Ar are molecular, held only by dispersion forces, so their melting points are low (S₈, the largest molecule, is highest of these).