Periodic Trends — Atomic Radius, Ionisation Energy and Electronegativity
Most students memorise periodic trends as arrows: "size decreases across, increases down." That gets you through a one-mark question and fails you the moment an exam asks why boron has a lower ionisation energy than beryllium. Every trend on the periodic table comes from just two competing effects. Learn those two and you can rebuild every arrow from scratch — and explain the exceptions instead of memorising them separately.
The two effects that control everything
1. Effective nuclear charge (Zeff). An outer electron does not feel the full nuclear charge. Inner electrons stand between it and the nucleus and cancel part of the pull. What is left over is the effective nuclear charge:
2. Distance (the principal quantum number n). Electrostatic attraction falls off with distance. An electron in the fourth shell is far from the nucleus and is held weakly no matter how large Z is.
Across a period, n stays the same while Z rises by one at every step. The electrons being added go into the same shell, and electrons in the same shell shield each other poorly. So Zeff climbs steadily and the outer electrons are pulled in harder. Down a group, a whole new shell opens and each filled inner shell shields almost completely, so Zeff barely changes while distance jumps.
Zeff by Slater's rules — sodium vs chlorine.
Na (Z = 11), outer 3s electron: inner shells shield with 8 × 0.85 + 2 × 1.00 = 8.80, so
S = 8.80 and Zeff = 11 − 8.80 = 2.20.
Cl (Z = 17), outer 3p electron: 6 same-shell electrons × 0.35 = 2.10, plus 8 × 0.85 = 6.80,
plus 2 × 1.00 = 2.00, so S = 10.90 and Zeff = 17 − 10.90 = 6.10.
Both outer electrons sit in the third shell, but chlorine's feels nearly three times the pull.
That single number explains why Cl is smaller, harder to ionise and far more electronegative
than Na.
Atomic radius
Radius falls left to right because Zeff rises and squeezes the same shell inward. It rises down a group because each new period adds a shell with a larger n.
| Across period 2 — covalent radius (pm) | Down group 17 — covalent radius (pm) |
|---|---|
| Li 152 · Be 112 · B 85 · C 77 · N 75 · O 73 · F 72 | F 72 · Cl 99 · Br 114 · I 133 |
Notice the contraction across period 2 is large at the start (Li → Be drops 40 pm) and small at the end (N → O → F drops 1–2 pm each). Early in a period the added electron enters a half-empty shell; later the shell is crowded and electron–electron repulsion partly cancels the extra nuclear pull.
Ions break the pattern deliberately. A cation has lost its outermost shell entirely, so it is much smaller than the parent atom (Na 186 pm metallic → Na⁺ 102 pm). An anion has added electrons to the same shell, increasing repulsion with no extra protons, so it swells (Cl 99 pm covalent → Cl⁻ 181 pm). For isoelectronic species such as N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺, every ion has 10 electrons and radius simply falls as Z rises.
Ionisation energy
The first ionisation energy is the minimum energy needed to remove the loosest electron from a gaseous atom:
Higher Zeff and smaller radius both make removal harder, so IE rises across a period and falls down a group — the mirror image of atomic radius.
| Element | Li | Be | B | C | N | O | F | Ne |
|---|---|---|---|---|---|---|---|---|
| IE₁ (kJ/mol) | 520 | 899 | 801 | 1086 | 1402 | 1314 | 1681 | 2081 |
The overall climb from 520 to 2081 is obvious. So are the two dips — and these are exam favourites.
Exception 1 — B (801) is lower than Be (899). Beryllium is 1s²2s²; boron is 1s²2s²2p¹. The electron removed from boron is a 2p electron, which is higher in energy and is shielded by the filled 2s pair, so it comes off more easily than a 2s electron. The same dip repeats in period 3: Al (578) is below Mg (738).
Exception 2 — O (1314) is lower than N (1402). Nitrogen is 2p³ with one electron in each of the three p orbitals — a stable half-filled set. Oxygen is 2p⁴, so one orbital holds a pair, and the repulsion between those two paired electrons helps push one out. Repeats in period 3: S (1000) is below P (1012).
Successive ionisation energies always increase (IE₁ < IE₂ < IE₃ …) because you are pulling an electron away from an ion that is already positive. A huge jump reveals the group: for magnesium IE₁ = 738 and IE₂ = 1451, but IE₃ = 7733 because the third electron must come out of the stable neon core.
Electronegativity
Electronegativity is the tendency of an atom in a bond to pull the shared pair towards itself. It is a relative scale with no unit — Pauling fixed fluorine at 3.98 and everything else follows.
| Element | F | O | N | Cl | Br | C | I | Na | Cs |
|---|---|---|---|---|---|---|---|---|---|
| Pauling value | 3.98 | 3.44 | 3.04 | 3.16 | 2.96 | 2.55 | 2.66 | 0.93 | 0.79 |
Same reasoning: high Zeff plus a small radius means the bonding pair sits close to the nucleus and is pulled hard. Fluorine wins on both counts and is the most electronegative element. Caesium, with a huge radius and a Zeff near 2, is the least.
Two practical uses. The difference in electronegativity predicts bond character — roughly, a gap above 1.7 gives a largely ionic bond, 0.4 to 1.7 a polar covalent bond and below 0.4 an essentially non-polar one. And the more electronegative atom carries the δ− end of the dipole, which is what sets up hydrogen bonding and molecular polarity.
The exceptions worth knowing by name
- d-block contraction (Al vs Ga). Gallium comes after the first transition series, and 3d electrons shield very poorly. Ga's atomic radius (135 pm) is actually slightly smaller than Al's (143 pm), and its electronegativity is higher (1.81 vs 1.61) — a group trend running backwards.
- Lanthanide contraction (Zr vs Hf). The 4f electrons shield badly too, so hafnium (159 pm) is essentially the same size as zirconium (160 pm) despite being a full period lower. This is why Zr and Hf are notoriously hard to separate chemically.
- Electron affinity of F vs Cl. Chlorine releases more energy on gaining an electron (−349 kJ/mol) than fluorine does (−328 kJ/mol), because fluorine's 2p shell is so small that the incoming electron faces severe repulsion. Electronegativity still puts F first — the two properties are not the same thing.
- Noble gas radii are quoted as van der Waals radii, not covalent radii, so they look anomalously large. Do not compare them directly with the halogens next door.
- Saying "shielding increases across a period". It barely changes — electrons in the same shell shield each other only about 35%. It is the rise in Z that drives the trend.
- Treating ionisation energy and electron affinity as opposite trends. Both generally rise across a period; they measure removing and adding an electron, not one forward and one backward version of the same thing.
- Comparing radii measured different ways. Covalent, metallic, ionic and van der Waals radii are different quantities. Use one consistent set within a comparison.
- Forgetting that ionisation energy is defined for the gaseous atom. Drop the "(g)" and the definition is wrong, and boards do deduct for it.
Where periodic trends appear in exams
| Exam | Typical use |
|---|---|
| CBSE/ICSE Class 11 | Classification of elements, order-the-elements questions, reason-based answers |
| JEE / NEET | Ionisation-energy anomalies, isoelectronic radius ordering, bond polarity |
| IIT-JAM / CUET-PG | Slater's rules, Zeff calculations, d-block and lanthanide contraction |
| GATE / CSIR-NET | Inorganic reactivity trends, hard–soft acid–base reasoning, ionic radii in solid state |
See every trend on one screen. The Interactive Periodic Table lets you click any element and read its atomic number, atomic mass, electron configuration and group position, so you can compare neighbours instead of trusting memory.
Open the Interactive Periodic Table →Working through Class 11–12 inorganic chemistry, or preparing for IIT-JAM, GATE or CSIR-NET? ABC Chemistry runs Class 11–12 coaching at its Gurugram centre and online across India, along with dedicated competitive-exam batches — details at abcchemistry.in.