Atomic Radius
Atomic radius is the first periodic trend you should master, because once you know it, you essentially know three others for free. Ionization energy, electron affinity, and electronegativity all trend in the opposite direction. Nail atomic radius, and the rest falls into place.
Defining Atomic Radius
Atomic radius is defined as half the distance between the nuclei of two identical bonded atoms. Because electron clouds do not have sharp boundaries, we cannot simply measure where an atom “ends.” Instead, we measure the internuclear distance in a diatomic molecule (or a metallic crystal) and divide by two.
For example, the distance between the two chlorine nuclei in Cl₂ is about 198 pm. The atomic radius of chlorine is therefore approximately 99 pm.
Trend Across a Period: Decreases Left to Right
As you move from left to right across a period, atomic radius decreases. This happens because each step to the right adds one proton to the nucleus and one electron to the same principal energy level. The extra proton increases the effective nuclear charge (Zeff), pulling all the electrons in that shell closer to the nucleus. The added electron does not shield effectively because it sits in the same shell as the others.
The result is a stronger net pull on the electron cloud with no new shell to offset it. The atom gets smaller.
Trend Down a Group: Increases Top to Bottom
As you move down a group, atomic radius increases. Each row adds an entirely new principal energy level (a new electron shell). Even though the nuclear charge also increases, the inner-shell electrons shield the outer electrons from the full nuclear pull. The new valence electrons occupy orbitals that are farther from the nucleus, so the atom is physically larger.
Periodic trends, and the single cause behind them
Scroll sideways to see the whole map.
The one cause underneathAcross a period, protons are added but the new electrons go into the same shell, so they shield each other poorly and effective nuclear charge climbs. Down a group, a whole new shell is added and the outer electrons are both further out and screened by everything beneath. Every trend on this page is one of those two sentences.
The two ionisation-energy dipsBe to B falls because boron's outermost electron is in a higher-energy 2p rather than 2s. N to O falls because oxygen is the first period-2 element forced to put two electrons in one 2p orbital, and their mutual repulsion makes one easier to remove. Both dips are on the plot, not smoothed away.
Ionic radiusCations are always smaller than their parent atom, often dramatically so when a whole shell is lost: Na is 186 pm, Na⁺ is 102 pm. Anions are always larger, because added electrons increase repulsion without adding any protons.
This effect dominates over the increase in Z. More shells always means a bigger atom.
Atomic Radius Across Period 3
The following table shows how atomic radius shrinks as you cross Period 3 from sodium to chlorine.
| Element | Symbol | Atomic Number | Atomic Radius (pm) |
|---------|--------|---------------|---------------------|
| Sodium | Na | 11 | 186 |
| Magnesium | Mg | 12 | 160 |
| Aluminum | Al | 13 | 143 |
| Silicon | Si | 14 | 117 |
| Phosphorus | P | 15 | 110 |
| Sulfur | S | 16 | 104 |
| Chlorine | Cl | 17 | 99 |
Every step to the right adds a proton and an electron to the same shell. Zeff climbs, and the radius shrinks.
Atomic Radius Down Group 1
| Element | Symbol | Period | Atomic Radius (pm) |
|---------|--------|--------|---------------------|
| Lithium | Li | 2 | 152 |
| Sodium | Na | 3 | 186 |
| Potassium | K | 4 | 227 |
| Rubidium | Rb | 5 | 248 |
| Cesium | Cs | 6 | 265 |
Each step down adds a new electron shell, and the atom balloons in size. Cesium, at the bottom of Group 1, is the largest non-radioactive atom at approximately 265 pm. At the opposite extreme, helium is the smallest atom at roughly 25 pm.
The Opposite Trend Shortcut
Here is the payoff for learning atomic radius first. Ionization energy, electron affinity, and electronegativity all increase in the opposite direction - they increase going up and to the right on the periodic table.
This makes intuitive sense. Smaller atoms hold their electrons more tightly (higher ionization energy), attract new electrons more strongly (higher electron affinity), and pull on shared bonding electrons more fiercely (higher electronegativity).
Why This Matters on the MCAT
The MCAT rarely asks you to recall a specific atomic radius in picometers. Instead, you will be asked to compare atoms: “Which is larger, Na or Cl?” or “Rank these elements by increasing atomic radius.” The two rules are all you need:
- Across a period: radius decreases (higher Zeff, same shell)
- Down a group: radius increases (new shell added)
If two atoms are not in the same period or group, use both rules together. For example, comparing Li (Period 2, Group 1) to K (Period 4, Group 1) is straightforward - K is larger because it is lower in the same group. Comparing Na (Period 3, Group 1) to Cl (Period 3, Group 17) is also straightforward - Na is larger because it is farther left in the same period.
Atomic radius increases going down a group (new electron shells are added, placing valence electrons farther from the nucleus) and going left across a period (lower effective nuclear charge means a weaker pull on the outermost electrons). These two rules let you compare the size of any two atoms on the table.
Mg < Na < K. Mg and Na are in Period 3, but Mg is farther right (higher Zeff), so Mg is smaller than Na. K is in Period 4, one shell below Na, so K is the largest. The order from smallest to largest is Mg, Na, K.