Five trends, one explanation. Atomic radius, ionic radius, ionisation energy, electron affinity and electronegativity all come from the same tug of war between the nucleus pulling electrons in and distance plus shielding letting them escape. Learn the tug of war once and you can reason your way to all five.
Everything on this page reduces to a competition between two things.
Now watch what happens in each direction. Across a period you add protons but you keep adding electrons to the same shell, so shielding barely changes. The nucleus wins. Down a group you add protons too, but you also add a whole new shell, so distance and shielding jump. Shielding wins.
The atomic radius is the distance from the nucleus to the outermost electron shell. Atoms have no hard edge, so in practice it is measured as half the distance between the centres of two touching atoms.
Across a period it decreases. Protons are added, electrons go into the same shell, shielding stays roughly constant, so the growing nuclear charge drags the outer shell in tighter.
Down a group it increases. Each element has one more occupied shell, so the outer electrons are further out and better shielded. Nuclear charge rises too, but distance and shielding outweigh it.
There is also a big jump between the noble gas at the end of one period and the alkali metal at the start of the next — that is a brand-new shell beginning.
Making an ion changes the size, sometimes dramatically.
| Ion type | Size vs the atom | Why |
|---|---|---|
| Cation (e.g. Na+) | Smaller | Electrons are lost, often losing a whole outer shell; the remaining electrons feel less repulsion and more pull each |
| Anion (e.g. Cl–) | Larger | Electrons are added with no extra protons, so electron–electron repulsion spreads the cloud out |
A neat consequence: in an isoelectronic series (ions with the same number of electrons, such as N3–, O2–, F–, Na+, Mg2+), the one with the most protons is the smallest, because the same electron cloud is being pulled by a bigger charge.
The first ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms, forming one mole of gaseous 1+ ions. It is always endothermic — you are pulling a negative electron away from a positive nucleus.
Those (g) state symbols are not decoration — the definition specifies gaseous atoms, and examiners look for them.
The graph shows the pattern beautifully. Values climb across each period as the nuclear charge grows, spike at the noble gases where the outer shell is full and tightly held, then crash at the next alkali metal, where a single electron sits alone in a brand-new, well-shielded shell.
Electron affinity is the energy change when one mole of gaseous atoms each gains an electron, forming 1– ions. For most non-metals it is exothermic (negative), because the incoming electron is attracted to the nucleus.
Across a period it becomes more negative: smaller atoms with a bigger nuclear charge hold the extra electron more strongly. Down a group it becomes less negative: the incoming electron ends up further out and more shielded.
A few elements buck the trend and have positive (endothermic) first electron affinities — Group 2 and Group 15 in particular — because the incoming electron has to squeeze into an already stable full or half-filled sub-shell.
Electronegativity is the ability of an atom to attract a shared pair of electrons in a covalent bond. It is not measured in energy units — it is a relative scale, with fluorine the highest at 4.0.
So the most electronegative elements sit in the top right (ignoring the noble gases) and the least electronegative in the bottom left. This is the trend you use later to decide whether a bond is ionic, polar covalent or non-polar.
| Property | Across a period → | Down a group ↓ | Deciding factor |
|---|---|---|---|
| Atomic radius | Decreases | Increases | Nuclear charge vs extra shells |
| Ionisation energy | Increases | Decreases | How tightly the outer electron is held |
| Electron affinity | More negative | Less negative | How strongly an incoming electron is attracted |
| Electronegativity | Increases | Decreases | Pull on a shared pair |
| Metallic character | Decreases | Increases | How easily electrons are lost |
Place these atoms in order of increasing atomic radius: S, Mg, Ca, Cl.
Explain why the first ionisation energy of potassium is lower than that of sodium.
Which is smaller, K+ or Cl–? Both have 18 electrons.
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