Periodic Trends: Atomic Radius, Ionization Energy, Electronegativity and Electron Affinity

The great power of the periodic table is that it lets you infer an element's properties from its neighbors. The basis for that inference is the set of periodic trends. Atomic size, the energy needed to remove an electron and the pull an atom exerts on shared electrons all change in consistent directions across rows and down columns. This guide covers the four main trends, the reasons behind them, and the exceptions that break the pattern.

The root of every trend: effective nuclear charge and shielding

An outer electron is attracted by the positive nucleus, but the inner electrons between them cancel part of that attraction. The positive charge the outer electron actually feels is therefore less than the full nuclear charge. This is the effective nuclear charge, Z_eff, and it can be approximated as:

Z_eff ≈ Z − S

Here Z is the atomic number and S is a shielding constant. Electrons in the same shell shield each other poorly, while electrons in inner shells shield outer electrons effectively.

Using Slater's rules, a simple set of estimates, the 2s electron of lithium feels a Z_eff of about 1.3, while a 2p electron of fluorine, in the same period, feels about 5.2. Moving right along a period adds one proton at a time, but the added electron goes into the same shell and adds little shielding. So effective nuclear charge increases from left to right across a period. Moving down a group, a new shell is added each time, and the outer electrons sit farther from the nucleus.

Those two facts, a stronger pull toward the right and more distant outer electrons toward the bottom, explain nearly all of the trends below.

1. Atomic radius

• Across a period (left to right): decreases, because the rising Z_eff pulls the electron cloud inward.
• Down a group: increases, because electron shells are added.

The largest atoms sit at the lower left (caesium, francium) and the smallest at the upper right (helium, fluorine). Atoms do not have sharp edges, so "radius" depends on how it is measured: covalent, metallic and van der Waals radii all differ. When comparing numbers from different sources, make sure they refer to the same kind of radius.

Ions differ in size from their parent atoms. Cations are smaller, having lost electrons; anions are larger, because the added electrons repel each other and the cloud expands. Among isoelectronic ions, which have the same number of electrons, more protons means a smaller ion. Of the ten-electron ions O²⁻, F⁻, Na⁺ and Mg²⁺, O²⁻ is the largest and Mg²⁺ the smallest.

Across the d-block the radius changes only gradually, and poor shielding by d electrons produces anomalies such as gallium being almost the same size as aluminium. In period 6 the lanthanide contraction makes hafnium nearly the same size as zirconium.

2. Ionization energy

Ionization energy is the minimum energy needed to remove one electron from a gaseous atom. The more tightly an electron is held, the larger it is.

• Across a period: generally increases.
• Down a group: decreases.

First ionization energies for period 2, in electronvolts, show both the trend and its exceptions.

The overall rise is interrupted twice.

• Beryllium > boron: the electron removed from boron is a 2p electron, which is higher in energy than 2s and partly shielded by the 2s pair, so it comes off more easily.
• Nitrogen > oxygen: nitrogen has one electron in each of its three 2p orbitals. In oxygen the fourth 2p electron has to pair up in an occupied orbital, and the extra repulsion makes it easier to remove.

The same pattern repeats in period 3, between magnesium and aluminium and between phosphorus and sulfur.

Removing electrons one after another gives the second, third and higher ionization energies. For magnesium the first is about 7.6 eV and the second about 15.0 eV, but the third jumps to over 80 eV, because the third electron must come from the stable neon-like inner shell. That sudden leap is the numerical signature of magnesium's preference for forming Mg²⁺.

3. Electronegativity

Electronegativity describes how strongly an atom in a chemical bond pulls the shared electron pair toward itself. Unlike ionization energy it is not measured directly; it is a relative scale derived from other data. The most widely used is the Pauling scale, proposed by Linus Pauling.

• Across a period: increases.
• Down a group: decreases.

Fluorine has the highest value, 3.98, followed by oxygen (3.44), chlorine (3.16) and nitrogen (3.04). At the other end are caesium (0.79) and francium (0.7). Most noble gases form few or no bonds and are often left without a value.

The difference in electronegativity between two atoms tells you about the bond. A small difference gives a nonpolar covalent bond with evenly shared electrons, a larger difference a polar covalent bond, and a very large difference a bond that is essentially ionic. Water is polar, and dissolves salt so well, because of the gap between oxygen (3.44) and hydrogen (2.20).

4. Electron affinity

Electron affinity is the energy released when a gaseous atom gains one electron. A larger value means the atom accepts an electron more readily.

• Across a period: generally increases, peaking in group 17.
• Down a group: generally decreases, but with many exceptions.

The most famous exception is that chlorine has a higher electron affinity than fluorine (about 3.6 eV for chlorine versus about 3.4 eV for fluorine). The fluorine atom is so small that an incoming electron is crowded into the compact 2p subshell and suffers strong repulsion. Fluorine still has the higher electronegativity and reactivity because other factors, such as bond energies and atomic size, also play a part.

Group 2 elements (filled s subshell), nitrogen in group 15 (half-filled p subshell) and the noble gases (filled shell) release little or no energy on gaining an electron, and may even require energy. Their electron affinities are near zero or negative, and some data sources leave them blank.

Linking the trends to metallic character

Put the four trends together and you get the trend in metallic character. Elements with low ionization energy and low electronegativity lose electrons easily and are strongly metallic. So metallic character increases toward the lower left, and nonmetallic character toward the upper right. That diagonal flow is what produces the staircase of metalloids across the p-block.

Summary

Rather than memorizing the table, practise deriving each trend from two sentences: effective nuclear charge grows toward the right, and shells are added toward the bottom. The exceptions then make sense too. Compare the ionization energies and electronegativities of elements along one period or down one group to see the pattern for yourself.