The periodic table is not just a list of elements; its layout predicts how atoms behave. Three trends show up on nearly every chemistry test: atomic radius, ionization energy and electronegativity. Instead of memorizing arrows, learn the two ideas behind all of them and you can work out any comparison.
The two ideas that explain every trend
- Effective nuclear charge is the pull the outer electrons actually feel from the nucleus. Moving left to right across a period, protons are added to the nucleus while new electrons go into the same energy level, which barely shields each other. The pull on the outer electrons gets stronger.
- Energy levels (shells). Moving down a group, each element adds a whole new shell. Outer electrons are farther from the nucleus and shielded by more inner electrons, so the pull on them is weaker.
Trend 1: atomic radius
Atomic radius is roughly half the distance between the nuclei of two bonded identical atoms.
- Across a period (left to right): decreases. The stronger effective nuclear charge pulls the electron cloud in tighter.
- Down a group: increases. Each new shell puts the outer electrons farther out.
So the largest atoms are at the bottom left (cesium and francium) and the smallest at the top right (helium, then fluorine and neon).
Trend 2: ionization energy
First ionization energy is the energy needed to remove one electron from a gaseous atom.
- Across a period: increases. Electrons are held more tightly, so it takes more energy to pull one away.
- Down a group: decreases. The outer electron is farther away and better shielded, so it comes off more easily.
Two exceptions are worth knowing for advanced classes. Boron's ionization energy is slightly lower than beryllium's, because boron's outer electron sits in a higher-energy p orbital. Oxygen's is slightly lower than nitrogen's, because oxygen's fourth p electron is paired, and electron-electron repulsion makes it easier to remove.
Trend 3: electronegativity
Electronegativity is how strongly an atom attracts shared electrons in a bond. It is measured on the Pauling scale, where fluorine is the highest at about 4.0.
- Across a period: increases.
- Down a group: decreases.
Noble gases are usually left out because they rarely form bonds. The difference in electronegativity between two atoms predicts bond type: a large difference (above about 1.7) suggests an ionic bond, a moderate one a polar covalent bond, and a very small one a nonpolar covalent bond.
Bonus: ion size
Cations (positive ions) are smaller than their neutral atoms, because losing electrons often removes a whole shell and the remaining electrons are pulled in more tightly. Anions (negative ions) are larger, because extra electrons repel each other and spread out. For example, Na⁺ is much smaller than Na, and Cl⁻ is larger than Cl.
Worked example: rank three elements
Rank Mg, S and Ba by atomic radius from smallest to largest.
- Mg and S are in the same period (period 3). S is farther right, so S is smaller than Mg.
- Ba is in the same group as Mg but much lower (period 6), so Ba is larger than Mg.
- Order: S < Mg < Ba.
For ionization energy, the order flips: Ba < Mg < S, because smaller atoms hold their electrons more tightly.
Common mistakes
- Thinking more protons always means a bigger atom. Across a period, more protons make the atom smaller.
- Mixing up electronegativity and electron affinity. Electronegativity is about shared electrons in a bond; electron affinity is the energy change when a free atom gains an electron.
- Including noble gases in electronegativity rankings.
Practice questions
- Which has the larger radius, K or Br?
- Which has the higher first ionization energy, Na or Cl?
- Rank N, O and F by electronegativity.
- Which is larger, O or O²⁻?
Answers: (1) K, because it is farther left in the same period. (2) Cl, because it is farther right. (3) N < O < F. (4) O²⁻, because the two extra electrons repel each other and spread out.
Last reviewed 2026-10-23 by the BrainBoost team.