Electron Affinity Calculator: Energy Released When Atoms Accept Electrons

In 1905, the American chemist Gilbert N. Lewis introduced the concept of electron affinity while studying the stability of ions in the gas phase. Electron affinity is the energy change that occurs when a neutral gaseous atom gains an electron to form a negatively charged ion. Unlike ionization energy, which measures the energy required to remove an electron, electron affinity measures the energy change associated with adding one. This property is fundamental to understanding chemical bonding, reactivity, and the formation of ionic compounds. For chemists, materials scientists, and physicists, electron affinity provides critical insight into why some elements form anions readily while others resist doing so.

Table of Contents

1. What Electron Affinity Measures
2. Periodic Trends in Electron Affinity
3. First vs. Second Electron Affinity
4. Electron Affinity vs. Ionization Energy vs. Electronegativity
5. Worked Examples
6. Applications in Chemistry and Materials Science
7. Frequently Asked Questions

What Electron Affinity Measures

Electron affinity (EA or Eea) is defined as the energy change when an electron is added to a neutral gaseous atom to form a negative ion:

A(g) + e⁻ → A⁻(g)

By convention:

• If energy is released (exothermic), the electron affinity is positive.
• If energy is absorbed (endothermic), the electron affinity is negative.

This sign convention can be confusing because some older textbooks and physics contexts define electron affinity as the energy required to remove an electron from the anion, which reverses the sign. In modern chemistry, the convention above is standard: positive EA means the anion is more stable than the neutral atom plus a free electron.

Key points:

• Electron affinity is measured for atoms in the gas phase to avoid complications from intermolecular forces and lattice energies.
• The process is the reverse of ionization energy for the anion: IE(A⁻) = EA(A) when using consistent sign conventions.
• Electron affinity depends on the effective nuclear charge, atomic radius, and electron-electron repulsion in the resulting anion.

Periodic Trends in Electron Affinity

Electron affinity follows general periodic trends, but with more exceptions than ionization energy.

Across a period (left to right): Electron affinity generally becomes more exothermic. The nuclear charge increases, pulling the added electron more tightly, while atomic radius decreases. Halogens (Group 17) have the most exothermic electron affinities because adding one electron completes a stable noble-gas p⁶ configuration.

Down a group: Electron affinity generally becomes less exothermic. The added electron enters a larger orbital farther from the nucleus, and shielding by inner electrons reduces the effective nuclear charge felt by the new electron.

Important exceptions:

• Group 2 elements (alkaline earth metals) have endothermic electron affinities because the added electron would enter a higher-energy p subshell after the filled s² configuration.
• Group 15 elements (N, P) have less exothermic electron affinities than expected due to the stability of the half-filled p³ configuration and electron-electron repulsion when adding a fourth p electron.
• Group 18 elements (noble gases) have endothermic electron affinities because the added electron must enter a new, higher-energy shell.
• Fluorine has a less exothermic electron affinity than chlorine due to greater electron-electron repulsion in the small F⁻ ion.

First vs. Second Electron Affinity

Just as there are successive ionization energies, there are successive electron affinities. The first electron affinity adds an electron to a neutral atom. The second electron affinity adds a second electron to an already negatively charged ion and is almost always endothermic because of electrostatic repulsion.

For oxygen:

• First EA: O(g) + e⁻ → O⁻(g) ΔH = −141 kJ/mol (exothermic, EA positive)
• Second EA: O⁻(g) + e⁻ → O²⁻(g) ΔH = +744 kJ/mol (endothermic, EA negative)

The large positive second electron affinity is why O²⁻ ions are not observed in the gas phase. In ionic compounds, the formation of O²⁻ is driven by the large lattice energy released when cations and anions come together in a crystal lattice.

Electron Affinity vs. Ionization Energy vs. Electronegativity

These three properties are related but distinct:

Ionization energy measures the energy required to remove an electron from a gaseous atom. It is always positive for the first ionization (energy must be supplied).

Electron affinity measures the energy change when an electron is added to a gaseous atom. It can be positive or negative.

Electronegativity is a dimensionless measure of an atom's ability to attract shared electrons in a chemical bond. It combines ionization energy and electron affinity effects in a bonded environment.

Worked Examples

Example 1: Comparing halogen electron affinities

Chlorine has an electron affinity of 349 kJ/mol, while fluorine has 328 kJ/mol. Why is chlorine's electron affinity more exothermic despite fluorine being more electronegative?

Fluorine is smaller, so the incoming electron experiences a stronger effective nuclear charge. However, the added electron also experiences stronger repulsion from the seven existing electrons in fluorine's compact 2p subshell. In chlorine, the 3p subshell is larger, so repulsion is lower, and the net energy release is greater.

Example 2: Predicting anion stability

Sodium (Na) has a first ionization energy of 496 kJ/mol and chlorine (Cl) has an electron affinity of 349 kJ/mol. When Na donates an electron to Cl, the energy cost is 496 − 349 = 147 kJ/mol. This positive value means the separated gas-phase ions are higher in energy than the separated atoms. In solid NaCl, however, the lattice energy of approximately 787 kJ/mol more than compensates, making NaCl stable.

Example 3: Energy cycle

For the reaction Na(g) + Cl(g) → NaCl(g), the energy change can be estimated from ionization energy and electron affinity. The actual formation of ionic NaCl in the gas phase is endothermic, but in the solid state the lattice energy drives stability. This illustrates why electron affinity alone cannot predict compound stability — lattice energy and solvation energy must also be considered.

Applications in Chemistry and Materials Science

Predicting ionic bond formation: Elements with highly exothermic electron affinities (halogens, oxygen) readily form anions. Elements with low ionization energies (alkali metals, alkaline earth metals) readily form cations. The combination forms stable ionic compounds.

Semiconductor design: Electron affinity determines the work function of materials and the band alignment at interfaces. Materials scientists tune electron affinity to optimize photovoltaic cells, OLEDs, and transistors.

Catalysis and surface science: Adsorption of electrons at metal surfaces depends on the surface electron affinity. Heterogeneous catalysts often exploit differences in electron affinity between substrates and catalyst surfaces.

Photochemistry: Electron transfer reactions depend on the electron affinity of acceptor species. Photosynthesis, solar cells, and battery electrodes all involve processes where electron affinity plays a central role.