The big idea: As you move around the periodic table, five atomic properties change in predictable ways. Almost every trend comes down to a tug-of-war between two factors:
- Nuclear charge — the pull of the protons in the nucleus on the outer electrons. More protons → stronger pull. - Shielding (and the number of shells) — inner electrons block the nuclear pull, and extra shells put the outer electrons further from the nucleus.
Get these two ideas right and you can explain every trend in this micro.
Define each property: - Atomic radius — half the distance between the nuclei of two bonded atoms (a measure of an atom's size). - Ionic radius — the size of the atom's ion. - First ionisation energy — the energy needed to remove one mole of electrons from one mole of gaseous atoms: X(g) → X⁺(g) + e⁻. - Electron affinity — the energy change when one mole of gaseous atoms gains an electron: X(g) + e⁻ → X⁻(g). - Electronegativity — how strongly an atom attracts a bonding pair of electrons (Pauling scale).
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Going left to right across a period, each element has one more proton (greater nuclear charge), but the new electron goes into the same outer shell — so the shielding stays roughly the same. The result is a stronger net pull on the outer electrons.
Across a period (e.g. period 3, highlighted): atomic radius decreases, while first ionisation energy and electronegativity increase.
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So across a period: - Atomic radius decreases — the stronger pull draws the outer shell in closer. - First ionisation energy increases — the outer electron is held more tightly, so it takes more energy to remove. - Electronegativity increases — the atom pulls a bonding pair more strongly.
All three follow from the same cause: increasing nuclear charge with near-constant shielding.
Small dips are normal: The ionisation-energy trend is a general rise, not a perfectly smooth line — there are small dips (e.g. between groups 2→13 and 15→16) caused by sub-shell structure. At SL you describe the overall increase across the period.
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Going down a group, each element has an extra electron shell. The outer electrons are further from the nucleus and more shielded by the inner shells, so the net pull on them is weaker — even though the nuclear charge is larger.
Down a group (e.g. group 1, highlighted): atomic radius increases, while first ionisation energy and electronegativity decrease.
Interactive diagram
Explore the labelled diagram, charts and maps for this topic in full study mode.
So down a group: - Atomic radius increases — more shells mean a bigger atom. - First ionisation energy decreases — the outer electron is further out and shielded, so it is easier to remove. - Electronegativity decreases — the nucleus pulls a bonding pair less strongly.
The extra shells + shielding outweigh the rise in nuclear charge.
Ionic radius: When an atom forms an ion, its radius changes:
- A cation (lost electron[s]) is smaller than its atom — often a whole shell is lost. - An anion (gained electron[s]) is larger than its atom — extra electron–electron repulsion spreads the shell out.
Down a group, ionic radius still increases (more shells).
| Property | Across a period → | Down a group ↓ | Why |
|---|---|---|---|
| Atomic radius | decreases | increases | more protons pull shells in / more shells added |
| Ionic radius | cations smaller than atom; anions larger | increases | lost shell (cation) or extra repulsion (anion) / more shells |
| 1st ionisation energy | increases | decreases | stronger nuclear pull / outer electron further out and shielded |
| Electron affinity | more exothermic (to chlorine) | less exothermic | stronger pull on the added electron / further out |
| Electronegativity | increases | decreases | stronger pull on a bonding pair / outer shell further from nucleus |
How this is tested: This micro is examined as an explain or an order/deduce task:
- Paper 1A (MCQ): order species or elements by radius or by first ionisation energy. - Paper 2: explain a trend (e.g. why first ionisation energy rises across a period, or why one element has a lower first ionisation energy than another) in terms of nuclear charge and shielding.
The markers want both factors named — nuclear charge and shielding/distance — not just 'it gets bigger'.
The marking phrase: An explanation scores when it links cause → effect: compare the nuclear charge, compare the shielding/number of shells, then state the net effect on the electron (held more or less tightly).
IB-style question — ionisation energy across a period (a)
(a) Explain why the first ionisation energy of sulfur is higher than that of magnesium, both in period 3. [2]
How to score the marks
- Mark 1 — nuclear charge. Sulfur has more protons (greater nuclear charge) than magnesium, but the outer electrons are in the same main energy level, so the shielding is similar.
- Mark 2 — the net effect. The greater nuclear charge pulls the outer electrons more strongly (and the atom is smaller), so more energy is needed to remove an outer electron from sulfur.
Final answer
Sulfur has a greater nuclear charge with similar shielding, so its outer electrons are held more tightly and more energy is needed to remove one.
IB-style question — comparing two elements (b)
(b) Explain, in terms of nuclear charge and shielding, why potassium (group 1, period 4) has a lower first ionisation energy than sodium (group 1, period 3). [2]
How to score the marks
- Mark 1 — shells and shielding. Potassium's outer electron is in a higher main energy level (further from the nucleus) and is more shielded by the extra inner shell.
- Mark 2 — the net effect. Although potassium has more protons, the extra distance and shielding outweigh the larger nuclear charge, so its outer electron is held less tightly and is easier to remove.
Final answer
Potassium's outer electron is further out and more shielded, so the net pull is weaker and less energy is needed to remove it.