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Chemistry courses usually present periodic trends as a list to memorize: radius goes down across a period and up down a group; ionization energy does the opposite; electronegativity does the same as ionization energy… It’s much easier to understand them as consequences of two ideas. Get those two ideas and you can work out every trend yourself.
Idea 1: effective nuclear charge
An outer electron is pulled towards the nucleus by all its protons, but it’s also pushed away by the other electrons, especially the ones in inner shells between it and the nucleus. The net pull it actually feels is called the effective nuclear charge (Z_eff).
A rough estimate: Z_eff ≈ (number of protons) − (number of inner-shell electrons).
- Sodium: 11 protons − 10 inner electrons ≈ +1
- Chlorine: 17 protons − 10 inner electrons ≈ +7
Both are in period 3, with outer electrons in the same shell, but chlorine’s outer electrons feel a much stronger pull.
Across a period, Z_eff increases, because each step adds a proton but the new electron goes into the same shell and hardly shields its neighbours.
Idea 2: distance (the number of shells)
The further an electron is from the nucleus, the weaker the attraction — and the more inner shells there are to shield it.
Down a group, the outer electrons are in a higher shell each time, so they’re further away and more shielded. Z_eff stays roughly similar, but distance wins.
Now derive the trends
Atomic radius
- Across a period → smaller. More Z_eff pulls the same shell in tighter.
- Down a group → larger. Each new shell is further out.
Largest atoms: bottom left (cesium, francium). Smallest: top right (helium, neon, fluorine). Details in the atomic radius trend.
Ionization energy (energy to remove an electron)
- Across a period → higher. The electron is held more tightly (more Z_eff, smaller atom).
- Down a group → lower. The electron is further out and easier to remove.
Highest: helium. Lowest: cesium and francium. See the ionization energy trend.
Electronegativity (pull on shared electrons in a bond)
Same reasoning, same direction as ionization energy:
- Across a period → higher.
- Down a group → lower.
Highest: fluorine. Lowest: francium and cesium. See the electronegativity trend.
Electron affinity (energy released when an atom gains an electron)
Generally more energy released across a period and less down a group, for the same reasons — but with more exceptions than the others. See electron affinity.
Metallic character
Metals are elements that lose electrons easily. So metallic character is the opposite of ionization energy:
- Across a period → less metallic.
- Down a group → more metallic.
That’s why metals fill the left and bottom of the table, nonmetals the top right, and the metalloids form a diagonal staircase between them. See metals, nonmetals and metalloids.
The summary table
| Property | Across a period (→) | Down a group (↓) | Maximum |
|---|---|---|---|
| Atomic radius | decreases | increases | bottom left |
| Ionization energy | increases | decreases | top right |
| Electronegativity | increases | decreases | top right (F) |
| Electron affinity | generally increases | generally decreases | top right (Cl) |
| Metallic character | decreases | increases | bottom left |
A handy shortcut: radius and metallic character point to the bottom left; everything else points to the top right.
Why the trends have exceptions
Real data has bumps, because the simple picture ignores the difference between s, p, d and f subshells:
- Ionization energy dips from beryllium to boron and from nitrogen to oxygen, because of subshell and electron-pairing effects.
- The lanthanide contraction makes period 6 transition metals nearly the same size as period 5.
- Electron affinities of noble gases, and of elements like nitrogen and beryllium, are unusually low or negative.
- Relativistic effects alter the heaviest elements — part of why gold is yellow and mercury is liquid.
The exceptions don’t break the rules; they reveal the finer structure of electron shells.
See the trends for yourself
The site has interactive trend maps that colour the whole table by property: atomic radius, ionization energy, electronegativity, electron affinity, melting point and more. On the periodic table, switch on heatmap mode to see any property at a glance.
Quick answers
Which trends increase across a period? Ionization energy, electronegativity and (generally) electron affinity.
Why do atoms get smaller across a period if they gain electrons? Because the added electrons go into the same shell while the nucleus gains protons, so the whole shell is pulled in.
What causes periodic trends? Effective nuclear charge and the distance of the outer electrons from the nucleus.
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