Q.How do atomic radius vary in a period and in a group? How do you explain the variation?
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Start your 14-day free trial to unlock the full solution →Atomic radius decreases across a period due to increasing nuclear charge pulling electrons closer, and increases down a group due to the addition of new electron shells.
The Core Idea: What "Atomic Radius" Really Means
Atomic radius isn't a fixed, hard boundary — it's the average distance from the nucleus to the outermost electron. Think of it as the "size" of an atom, measured in picometers (pm) or angstroms (Å). Two major trends govern how this size changes across the periodic table: one across a period (left to right), and one down a group (top to bottom).
The key to understanding both trends lies in two competing forces: the nuclear charge (how many protons pull inward) and the shielding effect (how inner electrons block that pull). Let's see how these play out.
1. Variation Across a Period (Left to Right)
Trend: Atomic radius decreases as you move from left to right across a period.
Why? Consider sodium (Na) at the start of period 3 and chlorine (Cl) near the end.
- As you go from Na to Cl, the number of protons in the nucleus increases (11 → 17). This means the effective nuclear charge — the net positive pull felt by the outermost electrons — increases.
- Crucially, all these elements have electrons in the same principal energy level (n=3 for period 3). The inner shells (n=1 and n=2) remain the same, so the shielding effect from inner electrons is nearly constant.
- With more protons pulling and the same shielding, the outermost electrons are pulled closer to the nucleus. The atom shrinks.
A common mistake is to think that adding more electrons makes the atom bigger. But across a period, you're adding electrons to the same shell — they don't create a new layer. The increasing nuclear charge wins, so the atom contracts.
Example: In period 3, atomic radii (in pm) are:
| Element | Na | Mg | Al | Si | P | S | Cl |
|---|---|---|---|---|---|---|---|
| Radius | 186 | 160 | 143 | 117 | 110 | 104 | 99 |
Notice the steady drop.
The decrease is not perfectly smooth — there are small bumps (e.g., Al is slightly larger than Mg) due to changes in electron-electron repulsion when subshells fill. But the overall trend is clear: size decreases left to right.
2. Variation Down a Group (Top to Bottom)
Trend: Atomic radius increases as you move down a group.
Why? Compare lithium (Li) at the top of group 1 with cesium (Cs) at the bottom.
- As you go down, you add entirely new electron shells. Li has electrons in n=2, Na in n=3, K in n=4, and so on. Each new shell is farther from the nucleus.
- The nuclear charge also increases (3 → 11 → 19 → ...), but the shielding effect from the many inner electrons almost completely cancels this out. The outermost electron feels only a small net pull.
- The dominant effect is the addition of a new shell, which pushes the outermost electron much farther away.
The effective nuclear charge () felt by an outer electron is roughly:
where is the atomic number and is the shielding constant. Down a group, increases but increases almost as much, so changes little — the shell number is what matters.
Example: In group 1, atomic radii (in pm) are:
| Element | Li | Na | K | Rb | Cs |
|---|---|---|---|---|---|
| Radius | 152 | 186 | 227 | 248 | 265 |
The jump from Li to Na is large (new shell), and the increase continues down. …
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