Electron Affinity
Electron affinity (EA) is the energy change that occurs when a neutral gaseous atom gains one electron to form a negative ion (anion). A negative EA value indicates an exothermic process — the atom releases energy and the anion is more stable than the separated atom and electron — which is the case for most halogens. Electron affinity generally increases (becomes more negative) across a period and decreases down a group, though there are notable exceptions such as the anomalously low EA of fluorine compared to chlorine due to electron–electron repulsion in fluorine's compact 2p orbitals.
Key Formula
X(g) + e⁻ → X⁻(g) ΔH = EA
LaTeX: X(g) + e^- \rightarrow X^-(g) \quad \Delta H = EA
| Symbol | Meaning | Unit |
|---|---|---|
| X(g) | Neutral gaseous atom | — |
| e⁻ | Added electron | — |
| X⁻(g) | Gaseous anion formed | — |
| EA | Electron affinity (negative = exothermic) | kJ mol⁻¹ or eV |
Electron Affinities of Selected Non-Metal Elements (kJ mol⁻¹)
| Element | Symbol | Z | EA (kJ mol⁻¹) | Anion Formed |
|---|---|---|---|---|
| Fluorine | F | 9 | −328 | F⁻ |
| Chlorine | Cl | 17 | −349 | Cl⁻ |
| Bromine | Br | 35 | −325 | Br⁻ |
| Oxygen | O | 8 | −141 | O⁻ |
| Sulfur | S | 16 | −200 | S⁻ |
| Nitrogen | N | 7 | +7 | N⁻ (endothermic) |
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Related Terms
Ionization Energy
Ionization energy (IE) is the minimum energy required to remove the most loosely bound electron from a gaseous atom or ion in its ground state, producing a positive ion. The first ionization energy (IE₁) removes the first electron; successive ionization energies increase because each removal leaves behind a more positively charged species that holds remaining electrons more tightly. Ionization energy increases across a period (due to greater effective nuclear charge) and decreases down a group (due to greater atomic radius and electron shielding), making it a key periodic trend.
Hund's Rule
Hund's rule of maximum multiplicity states that when electrons occupy degenerate (equal-energy) orbitals, one electron fills each orbital with parallel spins before any orbital receives a second electron. This arrangement minimises electron–electron repulsion because electrons in separate orbitals are farther apart, resulting in the lowest-energy (most stable) ground state. The rule also means that atoms with partially filled degenerate orbitals possess unpaired electrons, making them paramagnetic.
Effective Nuclear Charge
Effective nuclear charge (Z_eff) is the net positive charge experienced by a valence electron after accounting for the shielding (screening) effect of inner electrons, which partially cancel the attraction from the nucleus. It is calculated as Z_eff = Z − S, where Z is the actual atomic number and S is the shielding constant. Effective nuclear charge increases across a period because additional protons are added while shielding remains approximately constant, explaining trends in atomic radius, ionization energy, and electron affinity.
From Latin "electron" (amber, referring to the material that exhibits static electricity) and "affinis" (related to, bordering on). The term was coined in the early 20th century as chemists quantified the tendency of atoms to attract electrons.