Four properties change in predictable patterns as you move across or down the periodic table. Memorizing the directions is easy; the goal here is to understand why they happen, because every trend comes down to a tug-of-war between two things: how much positive charge the nucleus has, and how well the inner electrons shield the outer ones from it.
The two forces behind every trend
- Effective nuclear charge (): the net positive pull a valence electron actually feels, after accounting for the nucleus's charge and the shielding from inner electrons.
- Shielding: inner-shell (core) electrons repel and block outer valence electrons from feeling the full pull of the nucleus.
- Distance: electrical attraction weakens quickly with distance, so an electron in a higher shell (farther out) is held more weakly than one in a lower shell.
This rough estimate turns the trends into arithmetic. Sodium and chlorine both have 10 core electrons (shells 1 and 2), but sodium's valence electron feels a pull of about +1 while chlorine's seven valence electrons each feel about +7. Chlorine therefore holds its electrons in much closer: its radius is about 99 pm compared with 186 pm for sodium, and removing one of its electrons takes 1251 kJ/mol compared with only 496 kJ/mol for sodium. Down a group the estimate barely changes — lithium () and sodium () feel about the same net charge — but sodium's valence electron sits one shell farther out, so distance wins and it is held more weakly. That is why ionization energy falls from 520 kJ/mol for lithium to 376 kJ/mol for cesium.
The four trends
| Property | Definition | Across a period → | Down a group ↓ |
|---|---|---|---|
| Atomic radius | Size of the atom | Decreases ( ↑, pulls electrons in) | Increases (new shells added, more shielding) |
| Ionization energy | Energy to remove the outermost electron | Increases (electron held tighter) | Decreases (electron farther, more shielded) |
| Electron affinity | Energy change when an atom gains an electron | Generally more negative / exothermic (noble gases excepted) | Generally less negative (with exceptions, e.g. Cl releases more than F) |
| Electronegativity | Ability to attract shared electrons in a bond | Increases | Decreases |
Electronegativity ranges roughly from 0.7 (francium/caesium) to 3.98 (fluorine) on the Pauling scale — fluorine is the most electronegative element that exists, sitting in the top-right corner where is high and the atom is small enough to pull shared electrons in strongly.
Ion size, reactivity, and answering trend questions
- Na (186 pm) → Na⁺ (102 pm)
- Losing all the valence electrons empties the outer shell entirely
- The same number of protons now pulls on fewer electrons
- Metals shrink when they ionize
- Cl (99 pm) → Cl⁻ (181 pm)
- The added electron goes into the same outer shell
- More electron–electron repulsion spreads the cloud out
- The same number of protons is shared among more electrons
Reactivity follows from the same ideas, but in opposite directions for metals and nonmetals. Metals react by losing electrons, so the lower the ionization energy, the more reactive the metal: reactivity increases down a group, which is why cesium reacts explosively with water while lithium only fizzes. Nonmetals react by gaining electrons, so the stronger the pull on incoming electrons, the more reactive the nonmetal: reactivity increases up a group and to the right, which makes fluorine the most reactive nonmetal of all. Put together, metallic character increases toward the bottom left of the table and nonmetallic character toward the top right.