Electron affinity generally decreases (becomes less negative) as you move down a group in the periodic table. In other words, atoms lower in a group release less energy when they gain an electron. The main reason is that larger atoms hold an incoming electron farther from the nucleus, where the attraction is weaker. There is one notable exception: the very first element in many groups actually has a lower electron affinity than the element directly below it.
What Electron Affinity Measures
Electron affinity is the energy change when a neutral atom in the gas phase gains one electron to form a negative ion. When the atom readily accepts the electron, energy is released and the value is negative. Chlorine, for example, has an electron affinity of about −349 kJ/mol, meaning each chlorine atom gives off that much energy when it picks up an extra electron. The more negative the value, the stronger the atom’s “desire” for that additional electron.
Some atoms resist gaining an electron entirely. Noble gases and alkaline earth metals have full or half-full electron configurations that are already stable, so forcing an extra electron onto them would require energy rather than release it. Their electron affinities are effectively zero or positive.
Why Electron Affinity Decreases Down a Group
Two factors work together to weaken the pull on an incoming electron as you move to larger atoms.
Larger atomic radius. Each new period adds a principal energy shell, pushing the outermost electrons farther from the nucleus. An incoming electron slots into that outer shell, where the distance alone reduces the electrostatic attraction to the positive protons in the nucleus.
Greater electron shielding. As you go down a group, more inner-shell (core) electrons sit between the nucleus and the valence shell. These core electrons repel the outer electrons and cancel out a portion of the nuclear charge. The result is a lower effective nuclear charge felt by the incoming electron. Effective nuclear charge can be approximated as the total number of protons minus the number of core electrons shielding them. Even though the total nuclear charge increases down a group, the shielding increases at roughly the same pace, so the effective pull on a new valence electron stays flat or slightly drops.
Together, these effects mean the atom simply has less grip on an additional electron the farther down the group you go. Less energy is released, and the electron affinity becomes less negative.
The Period 2 Exception
If you look at actual data, the trend down a group is not perfectly smooth. The first element in many groups (the Period 2 element) has a surprisingly low electron affinity compared to the element just below it in Period 3. The second element in the group most often has the greatest electron affinity of the entire column.
The halogens are the classic example. Fluorine’s electron affinity is 3.339 eV, while chlorine’s is 3.617 eV. Chlorine releases more energy per atom when it gains an electron, even though fluorine is smaller and has a higher effective nuclear charge. The explanation comes down to size: fluorine’s 2p orbitals are extremely compact, and its seven existing electrons are crammed into a very small space. When an eighth electron arrives, it faces intense electron-electron repulsion that partially offsets the strong nuclear attraction. Chlorine’s 3p orbitals are larger, giving the electrons more room and reducing that repulsion.
The same pattern appears in Group 16. Oxygen’s electron affinity (about 1.46 eV) is lower than sulfur’s, for the same reason: oxygen’s small shell creates crowding that counteracts its nuclear pull.
Numerical Example: Group 1
The alkali metals illustrate the general downward trend clearly. Lithium, sodium, potassium, rubidium, and cesium all have modest electron affinities because their incoming electron enters a new, relatively large s orbital. But the values do shrink as the atoms get bigger. Hydrogen, sitting at the top, is a special case with its own unique electron configuration, but the metals from lithium onward show a gradual decline in how much energy is released when they accept an electron.
The trend is less dramatic in Group 1 than in the halogens because alkali metals already have low electron affinities to begin with. They far prefer to lose an electron than to gain one, which is why they form positive ions so readily.
How This Compares to Other Periodic Trends
Electron affinity’s group trend runs in the same direction as electronegativity and ionization energy: all three generally decrease going down a group. This makes intuitive sense because all three depend on how tightly an atom holds onto electrons. A larger atom with more shielding holds electrons less tightly, so it has lower ionization energy, lower electronegativity, and a less negative electron affinity.
The key difference is precision. Electronegativity is a relative scale describing how strongly an atom pulls on shared electrons within a bond. Electron affinity is a measurable quantity, reported in kJ/mol or eV, for an isolated atom gaining an electron. They track together in most cases, but electron affinity is more sensitive to the small shell-size effects that cause anomalies like the fluorine-chlorine inversion.
Quick Summary of the Trend
- General direction: Electron affinity becomes less negative (weaker) going down a group.
- Main causes: Increasing atomic radius and greater electron shielding reduce the effective nuclear charge on the incoming electron.
- Notable exception: Period 2 elements (fluorine, oxygen, nitrogen) often have lower electron affinities than the Period 3 element below them, because their small orbitals create strong electron-electron repulsion.
- Groups with near-zero values: Noble gases and alkaline earth metals do not form stable negative ions, so their electron affinities are essentially zero regardless of position in the group.