IB Chemistry SLTopic 3 — Classifying the ElementsPaper 1 & 2Trends~14 min read
Trends Across the Periodic Table
There are five trends on this page and only one explanation. Every single one comes down to a tug of war between the nucleus and the outer electrons, decided by three factors. Learn the three factors properly and you never have to memorise a trend again — you derive it on the spot.
📘 What you need to know
Three factors control how tightly the outer electrons are held: nuclear charge, shielding by inner electrons, and distance from the nucleus.
Across a period: nuclear charge rises, shielding stays roughly the same, so the pull gets stronger.
Down a group: an extra shell is added each time, so distance and shielding both rise and the pull gets weaker.
Atomic radius decreases across a period and increases down a group.
Cations are smaller than their parent atoms; anions are larger.
Ionisation energy, electron affinity (more negative) and electronegativity all increase across a period and decrease down a group.
Electronegativity applies to atoms in a covalent bond; electron affinity is a measurable energy change for gaseous atoms. They are not the same thing.
The three-factor engine
Picture the outer electron being pulled inwards by the protons and pushed outwards by nothing in particular — but partly screened from the nucleus by the electrons in between. How firmly it is held depends on three things at once.
Across a period, shielding and distance barely change, so rising nuclear charge wins unopposed. Down a group, distance and shielding both rise and together they beat the extra protons.
Atomic radius
The atomic radius is the distance from the nucleus to the outermost electrons. It is awkward to define exactly — an atom has no hard edge — so it is usually measured as half the distance between the nuclei of two touching atoms of the same element.
Across a period it decreases. Each step right adds a proton but puts the new electron in the same shell. Shielding is unchanged, so the stronger nucleus pulls the whole shell in tighter.
Down a group it increases. Each step down adds a whole new shell, so the outer electrons start further out and are shielded by more inner electrons.
The across-a-period trend surprises people: you are adding electrons, so surely the atom grows? No — you are adding them to a shell that is already there, while the nucleus gets stronger. The atom shrinks.
Ionic radius
Turning an atom into an ion changes its size, sometimes dramatically.
Cations are smaller than their atoms. Losing electrons often empties the outer shell completely, and the electrons that remain feel the same nuclear charge shared between fewer of them, so they are pulled in harder.
Anions are larger than their atoms. The extra electrons add repulsion within the same shell, so the electron cloud spreads out, while the nuclear charge has not changed at all.
The neatest test of all this is an isoelectronic series — ions with identical electron configurations but different nuclei. Since shielding and the number of electrons are fixed, only nuclear charge varies, and the size falls steadily as the charge climbs.
This is the cleanest possible demonstration of the argument. With electrons and shielding held constant, size depends on nuclear charge alone — and it does exactly what the theory predicts.
Ionisation energy
First ionisation energy
X(g) → X+(g) + e−the energy to remove one mole of electrons from one mole of gaseous atoms
Notice the state symbols — everything must be gaseous, or you would also be paying to separate the particles from each other. Values are in kJ mol−1 and are always positive, because you always have to pull against an attraction.
Across a period ionisation energy rises: stronger nucleus, same shielding, so the electron is harder to remove. Down a group it falls: the outer electron is further out and better shielded, so it leaves more easily. There is a sharp drop between the end of one period and the start of the next, because a brand new shell has just opened up.
Electron affinity
Electron affinity is the mirror image: the energy change when a gaseous atom gains an electron.
First electron affinity
X(g) + e− → X−(g)
For most non-metals this releases energy, so the value is negative. Across a period it becomes more negative, because a smaller atom with a stronger nucleus attracts the incoming electron more strongly. Group 2 and group 15 break the pattern, because their subshells are full or exactly half full and an extra electron would have to pair up.
Second electron affinities are always endothermic. You are pushing a negative electron towards an ion that is already negative, and that costs energy no matter what.
Electronegativity
Electronegativity is the ability of an atom in a covalent bond to attract the shared pair of electrons towards itself. It is a comparative scale with no units, and the values are in Section 9 of the data booklet.
It increases across a period and decreases down a group, for exactly the same reasons as everything else on this page. Fluorine, at the top right of the reactive elements, is the most electronegative element of all.
Do not confuse the three. Ionisation energy is about losing an electron. Electron affinity is about gaining one. Electronegativity is about sharing one inside a bond. Only electronegativity requires a bond to exist.
Property
Across a period (left to right)
Down a group
Main reason
Atomic radius
Decreases
Increases
Nuclear charge across; extra shells down
Ionic radius
Decreases within cations, jumps up at the first anion
Increases
Charge on the ion, then the same factors as radius
Ionisation energy
Increases
Decreases
Stronger pull on the outer electron across
Electron affinity
More negative
Less negative
Smaller atom attracts the extra electron better
Electronegativity
Increases
Decreases
Bonding pair pulled closer to a small, high-charge nucleus
Worked examples
WORKED EXAMPLE
Arrange these atoms in order of decreasing atomic radius: N, Be, Mg, C.
Step 1: separate them by period
Be, C and N are all in period 2. Mg is in period 3, so it has an extra shell.
Step 2: place Mg
An extra shell beats everything else, so Mg is the largest.
Step 3: order the period 2 atoms
Radius decreases across, and their groups are 2, 14 and 15.
Be > C > NMg > Be > C > Nalways sort by period first, then by position across the period
WORKED EXAMPLE
O2−, F−, Na+ and Mg2+ all have 10 electrons. Explain why they are not the same size.
Step 1: state what is constant
All four have the configuration 1s² 2s² 2p⁶, so the number of electrons and the shielding are identical.
Step 2: state what variesO: 8 protons, F: 9, Na: 11, Mg: 12Step 3: apply the argument
More protons pulling on the same ten electrons means a stronger attraction and a tighter electron cloud.
Radius decreases: O2− > F− > Na+ > Mg2+say “same shielding, increasing nuclear charge” explicitly — that is the mark
WORKED EXAMPLE
The first ionisation energy of caesium is far lower than that of lithium. Explain why, in terms of the three factors. [3]
Mark 1: distance
Caesium’s outer electron is in the 6th shell, lithium’s is in the 2nd, so it is much further from the nucleus.
Mark 2: shielding
Caesium has far more inner shells, so the outer electron is much better shielded from the nuclear charge.
Mark 3: the balance
Caesium does have more protons, but distance and shielding together outweigh that.
The outer electron is held far more weakly, so much less energy is needed to remove itnotice that all three factors are named. That is what a three-mark trend answer looks like.
💡 Exam tip
Trend questions are marked for reasons, not directions. “It increases” is worth nothing on its own.
Use the three factors as a checklist every time: nuclear charge, shielding, distance. Say which one wins.
Include state symbols and the (g) in ionisation energy and electron affinity equations.
Electron affinity values are usually negative. Say “more negative” rather than “bigger” to avoid ambiguity.
Electronegativity has no units and no equation — it is a comparative scale, so only ever compare with it.
For isoelectronic ions, sort by number of protons. It is the only thing that differs.
⚠ Common mix-up
Saying atoms get bigger across a period because electrons are added. They get smaller — same shell, stronger nucleus.
Claiming shielding increases across a period. It is roughly constant; no new inner shell is added.
Confusing electronegativity with electron affinity. One describes sharing inside a bond, the other is a measurable energy change for a gaseous atom.
Thinking cations and anions both shrink. Cations shrink, anions grow.
Forgetting the drop between periods. Ionisation energy falls sharply at the start of each new period because a new shell opens.
Writing “the nucleus gets bigger”. Say the nuclear charge increases — the nucleus itself barely changes in size.
Up next: Group 1 Metals Reacting with Water — the first place these trends stop being abstract, because a lump of potassium skating across a water bath is the ionisation energy trend made visible.
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