Electron Affinity

Electron affinity (E_ea or EA) is the quantitative measure of an isolated gaseous atom's or molecule's thermodynamic tendency to capture an additional electron to form a negatively charged ion (anion). By strict IUPAC convention, electron affinity is defined as the energy required to detach an electron from the singly charged gaseous anion A⁻(g) at T = 0 K—making it the zeroth ionization energy of the species:

A⁻(g) → A(g) + e⁻(g) (E_ea = Δ*U*°_{0 K})

Equivalently, in terms of the forward electron-attachment reaction:

A(g) + e⁻(g) → A⁻(g) (Δ*U*°_{attach, 0 K} = −*E*_ea; Δ*H*°_{attach, 298 K} ≈ −*E*_ea − (5/2)*RT*)

Thus, a positive electron affinity (E_ea > 0) signifies that electron attachment is exothermic (Δ*H*_attach < 0) and that the resulting anion A⁻(g) lies lower in energy than the neutral atom A(g) plus a free electron at rest. Conversely, a negative or zero electron affinity (E_ea ≤ 0) indicates that adding an electron is endothermic and that the isolated gas-phase anion is unstable with respect to spontaneous autodetachment of the electron.

Across the periodic table, positive first electron affinities range from near zero in calcium (Ca, E_ea = +2.37 kJ mol⁻¹ = 0.0245 eV) to the periodic maximum in chlorine (Cl, E_ea = +348.575 kJ mol⁻¹ = 3.6127 eV), playing a primary role in ionic bonding, Mulliken electronegativity, and biological electron-transfer chains.

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Laser Photodetachment Spectroscopy of Anions

Unlike neutral atoms, which possess an infinite series of bound Rydberg excited states converging on the ionization limit (because the ejected electron moves in a long-range −*e*² / 4π*ε*₀r Coulomb potential due to the residual M⁺ cation), an electron departing from a neutral atom A(g) in an anion A⁻(g) experiences only a short-range induced-dipole polarization potential (V(r) = −*αe*² / [2(4π*ε*₀)²r⁴]). Consequently, most atomic anions possess no bound electronically excited states below the detachment threshold, preventing measurement of E_ea by conventional optical emission spectroscopy.

Modern high-precision electron affinities (accurate to ±0.001 kJ mol⁻¹) are measured via laser photodetachment threshold spectroscopy and laser photoelectron spectroscopy (LPES), pioneered by Jan L. Hall, Lineberger, and Hotop in the 1970s. A mass-selected beam of negative ions A⁻(g) is crossed in vacuum with a tunable monochromatic laser beam (hν). When hν exceeds the detachment threshold E_ea, photoelectrons are ejected with translational kinetic energy E_k = hν − E_ea, allowing direct determination of the energy gap between the ground states of A⁻(g) and A(g).

Endothermic Second Electron Affinities (E_ea,2 < 0)

While attaching a first electron to a neutral oxygen or sulfur atom is exothermic because the incoming electron is attracted by the unshielded effective nuclear charge (Z_eff > 0) of the incomplete valence shell, attaching a second electron to an already negatively charged monoanion (A⁻ + e⁻ → A²⁻) must overcome long-range Coulombic electrostatic repulsion between two like negative charges:

• First attachment to oxygen: O(g) + e⁻(g) → O⁻(g)

E_ea,1 = +141.0 kJ mol⁻¹ (+1.461 eV); Δ*H*°₁ = −141.0 kJ mol⁻¹ (exothermic)

• Second attachment to oxide: O⁻(g) + e⁻(g) → O²⁻(g)

E_ea,2 ≈ −744 to −844 kJ mol⁻¹; Δ*H*°₂ = +744 to +844 kJ mol⁻¹ (strongly endothermic)

• Net formation of gaseous O²⁻(g) from O(g): O(g) + 2 e⁻(g) → O²⁻(g)

Δ*H*°_net ≈ +603 to +703 kJ mol⁻¹

Isolated gas-phase dianions such as O²⁻(g), S²⁻(g), and N³⁻(g) are therefore unstable in vacuum and spontaneously eject an electron; they exist in nature only within condensed phases (such as crystalline MgO or CaO), where the surrounding shell of positively charged cations provides a stabilizing Madelung electrostatic potential well (U_L ≈ −3,400 to −3,800 kJ mol⁻¹).

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Why Group 17 Halogens Peak and Chlorine Exceeds Fluorine

Across any period, E_ea generally increases with rising effective nuclear charge Z_eff and reaches its maximum in Group 17 (halogen elements, ns²np⁵), where adding one electron completes the stable noble gas octet (ns²np⁶, ¹S₀). However, a universal anomaly occurs between Period 2 and Period 3 across Groups 14, 15, 16, and 17: the electron affinity of every second-row p-block element is substantially smaller than that of its third-row congener directly below it:

• Group 17: F (+328.2 kJ mol⁻¹) < Cl (+348.6 kJ mol⁻¹) > Br (+324.6 kJ mol⁻¹) > I (+295.2 kJ mol⁻¹)
• Group 16: O (+141.0 kJ mol⁻¹) < S (+200.4 kJ mol⁻¹) > Se (+195.0 kJ mol⁻¹) > Te (+190.2 kJ mol⁻¹)

This Period 2 < Period 3 inversion arises from the exceptionally small atomic radius of the n = 2 valence shell (covalent radius of F = 57 pm vs. Cl = 102 pm; O = 66 pm vs. S = 105 pm). Because the 2*p* subshell lacks an inner p core and has zero radial nodes (n − l − 1 = 0), its seven (in F) or six (in O) valence electrons are crowded into a tiny spatial volume of high charge density. Introducing an extra electron into the compact 2*p* subshell incurs strong inter-electronic Coulombic repulsion (J_2*p*,2*p*) that offsets approximately 70–80 kJ mol⁻¹ of nuclear attraction, whereas in the more spatially diffuse 3*p* subshell of chlorine and sulfur, electron–electron repulsion is substantially reduced.

Closed-Shell and Half-Filled Subshell Minima (Groups 2, 15, 18)

Sharp minima (E_ea ≤ 0 or small positive values) occur at three electron configuration milestones:

1. Group 18 Noble Gases (ns²np⁶): All noble gases (helium, neon, argon, krypton, xenon) have negative electron affinities (E_ea < 0) because the added electron must enter the next higher principal shell (n + 1)*s*, where it is almost fully screened by the closed core.
2. Group 2 Alkaline Earths (ns²) and Group 12 (nd¹⁰ns²): Beryllium, magnesium, zinc, cadmium, and mercury have closed ns² subshells; an added electron must occupy the higher-energy, poorly penetrating np subshell, giving E_ea ≤ 0 (whereas Group 1 alkali metals ns¹ have positive affinities of +47 to +73 kJ mol⁻¹ because the added electron pairs into the penetrating ns orbital to form ns² alkalide ions such as Na⁻).
3. Group 15 Pnictogens (np³): Nitrogen (E_ea = −6.8 kJ mol⁻¹) has a half-filled 2*p*³ subshell obeying Hund's rules; adding a fourth 2*p* electron forces spin pairing against intra-orbital Coulombic repulsion, whereas carbon (2*p*² → C⁻ 2*p*³, E_ea = +121.8 kJ mol⁻¹) gains half-filled exchange stability upon electron capture.

Mulliken Electronegativity and Chemical Hardness

In 1934, Robert S. Mulliken combined first ionization energy (IE₁, resistance to losing an electron) and first electron affinity (E_ea, drive to gain an electron) into an absolute physical scale of electronegativity:

χ_M = ½ (IE₁ + E_ea)

In modern conceptual density functional theory (DFT), Ralph G. Pearson and Robert G. Parr identified χ_M = −*μ* as the negative of the electronic chemical potential (∂*E*/∂*N*)_v and defined the half-difference η = ½ (IE₁ − E_ea) as absolute chemical hardness (half the HOMO–LUMO gap), quantifying resistance to polarization in Lewis acid and base interactions.

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If fluorine is the most electronegative element, why is its electron affinity lower than chlorine's?

Pauling electronegativity (χ_P = 3.98 for F vs. 3.16 for Cl) and Mulliken electronegativity (χ_M = ½[IE₁ + E_ea] = 10.41 eV for F vs. 8.30 eV for Cl) depend on both ionization energy and electron affinity (or covalent bond dissociation energies). Although fluorine has a slightly lower electron affinity than chlorine (3.40 eV vs. 3.61 eV) due to electron–electron crowding in the tiny 2*p* subshell, fluorine has a much higher first ionization energy (17.42 eV vs. 12.97 eV for Cl) and forms much shorter, stronger covalent and ionic bonds.

How can gold form a stable cesium auride (Cs⁺Au⁻) ionic crystal?

Gold ([Xe] 4*f*¹⁴ 5*d*¹⁰ 6*s*¹) undergoes intense Dirac relativistic contraction of its penetrating 6*s* orbital, lowering the 6*s* energy level and giving gold an electron affinity of +222.8 kJ mol⁻¹ (2.309 eV)—higher than that of oxygen (+141 kJ mol⁻¹) or sulfur (+200 kJ mol⁻¹) and comparable to iodine (+295 kJ mol⁻¹). When equimolar cesium (IE₁ = 375.7 kJ mol⁻¹) and gold are melted together, an electron transfers from Cs to Au to form yellow, transparent semiconducting cesium auride (Cs⁺Au⁻) crystallizing in the ionic CsCl lattice with closed-shell Au⁻ (5*d*¹⁰ 6*s*²) anions.

Why do textbooks sometimes list electron affinities with opposite algebraic signs?

Physical chemistry and atomic physics define electron affinity E_ea as the positive energy liberated when an electron attaches (or equivalently, the positive Δ*U*° required to photodetach the electron from the anion A⁻ → A + e⁻, so Cl has E_ea = +348.6 kJ mol⁻¹). Introductory general chemistry textbooks often tabulate the standard enthalpy of electron attachment (Δ*H*°_EA for A + e⁻ → A⁻), which uses standard thermodynamic sign convention where exothermic heat release is negative (so Cl has Δ*H*°_EA = −348.6 kJ mol⁻¹). Both refer to the identical physical quantity with opposite signs: E_ea = −Δ*U*°_attach.

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1. Rienstra-Kiracofe, J. C., Tschumper, G. S., Schaefer, H. F., Nandi, S., & Ellison, G. B. (2002). "Atomic and molecular electron affinities: Photoelectron experiments and theoretical computations." Chemical Reviews, 102(1), 231–282. https://doi.org/10.1021/cr990044u
2. Andersen, T., Haugen, H. K., & Hotop, H. (1999). "Binding energies in atomic negative ions: III." Journal of Physical and Chemical Reference Data, 28(6), 1511–1533. https://doi.org/10.1063/1.556047
3. Mulliken, R. S. (1934). "A new electroaffinity scale; together with data on valence states and on valence ionization potentials and electron affinities." The Journal of Chemical Physics, 2(11), 782–793. https://doi.org/10.1063/1.1749394
4. Parr, R. G., & Pearson, R. G. (1983). "Absolute hardness: companion parameter to absolute electronegativity." Journal of the American Chemical Society, 105(26), 7512–7516. https://doi.org/10.1021/ja00364a005