Electron gain enthalpy (also known as electron affinity) is the enthalpy change when a neutral gaseous atom in its ground state gains an electron to form a negatively charged ion (anion). It measures an atom's tendency to accept an electron. The reaction is: X(g) + e⁻ → X⁻(g) + ΔH. For most non-metals, the process releases energy (exothermic) — the enthalpy change is negative.
Electron gain enthalpy = energy change when X(g) + e⁻ → X⁻(g).
Negative ΔegH: exothermic (energy released), atom readily gains electron.
Positive ΔegH: endothermic (energy required), atom resists gaining electron.
Most negative electron gain enthalpy: Cl (−349 kJ/mol).
F has less negative ΔegH (−328) than Cl due to small size and inter-electronic repulsion.
Down a group: ΔegH becomes less negative (increasing size reduces attraction).
Across a period: ΔegH generally becomes more negative (left to right).
Noble gases and N, Be, Mg have positive or near-zero ΔegH (stable configurations).
Definition: Electron gain enthalpy (ΔegH) is the energy change when a neutral gaseous atom in its ground state accepts an electron to form a gaseous anion.
Reaction: X(g) + e⁻ → X⁻(g) ΔegH
Sign convention: • Negative ΔegH: energy is RELEASED → exothermic → atom has high tendency to accept electron • Positive ΔegH: energy is ABSORBED → endothermic → atom does not readily accept electron
Note on terminology: • 'Electron gain enthalpy' = term used in modern NCERT Chemistry • 'Electron affinity' = older term (still used widely) • When ΔegH is negative, electron affinity is said to be positive (and vice versa) — signs are opposite in older convention
First vs second electron gain enthalpy: • First: X(g) + e⁻ → X⁻(g) — usually exothermic • Second: X⁻(g) + e⁻ → X²⁻(g) — always endothermic (adding e⁻ to already negative ion requires energy) • Example: O⁻ → O²⁻ requires energy input (ΔegH₂ is positive)
Across a period (left to right): • Electron gain enthalpy generally becomes more negative (more energy released) • Nuclear charge increases → more attraction for incoming electron • Trend: more negative from left to right (non-metals have more negative values) • Exception: Noble gases (He, Ne, Ar) have highly positive ΔegH — they resist gaining electrons • Exception: N, Be, Mg have less negative values than expected (half-filled/fully-filled stable configurations)
Down a group: • Electron gain enthalpy generally becomes less negative (less energy released) • As size increases, incoming electron is farther from nucleus → less attracted • Exception: F (fluorine) has LESS negative ΔegH than Cl (chlorine) despite F being higher in group
Halogens (Group 17) electron gain enthalpy: Element | ΔegH (kJ/mol) F | -328 Cl | -349 ← most negative among halogens (highest electron affinity) Br | -325 I | -295 At | -270
Why Cl > F: • F has a very small atomic radius → high electron density → significant interelectronic repulsion when a new electron is added • Cl has slightly larger size → less repulsion → more easily accommodates the extra electron • Therefore Cl has MORE negative ΔegH despite F being above it
Highest (most negative) electron gain enthalpy among all elements: Cl (−349 kJ/mol)
Factors that influence electron gain enthalpy:
Atomic size (radius): • Smaller atom → electron added closer to nucleus → more attracted → more negative ΔegH • Exception: F (too small, inter-electronic repulsion reduces ΔegH)
Nuclear charge (Z): • Higher Z → stronger attraction for incoming electron → more negative ΔegH
Electronic configuration: • Atoms with stable configurations (noble gas, half-filled sub-shells) resist gaining electrons • N (2s²2p³): half-filled 2p → ΔegH is nearly zero or slightly positive • Be (2s²): fully filled 2s → positive ΔegH • Ne, Ar: filled shells → very positive ΔegH
Elements with nearly zero or positive ΔegH: • Noble gases (He, Ne, Ar, Kr, Xe) • N (nitrogen) — half-filled 2p³ • Be, Mg — filled s²
Practical significance: • High (very negative) ΔegH → more reactive non-metal, stronger oxidising agent • Cl has higher ΔegH than F → practical significance in displacement reactions • O has two-step electron gain (O → O⁻ is exothermic; O⁻ → O²⁻ is endothermic)
Electron gain enthalpy (ΔegH) is the enthalpy change when a neutral gaseous atom in its ground state gains an electron to form a gaseous anion: X(g) + e⁻ → X⁻(g). When energy is released (exothermic), ΔegH is negative — the atom readily accepts an electron. When energy is absorbed, ΔegH is positive — the atom resists gaining an electron.
Chlorine (Cl) has the most negative electron gain enthalpy among all elements: −349 kJ/mol. Although fluorine is above chlorine in Group 17 and is smaller, its very small size causes significant inter-electronic repulsion when an extra electron is added, making its ΔegH (−328 kJ/mol) less negative than that of chlorine.
Fluorine has a very small atomic radius. When an extra electron is added to F, the electron cloud is very compact and the incoming electron experiences significant repulsion from the already dense electron cloud. This reduces the energy released. Chlorine has a slightly larger size, accommodates the extra electron more comfortably, and releases more energy — hence Cl (−349 kJ/mol) has a more negative ΔegH than F (−328 kJ/mol).
Across a period (left to right), electron gain enthalpy generally becomes more negative (more energy is released). This is because nuclear charge increases while atomic size decreases, leading to stronger attraction for the incoming electron. Exceptions: N (half-filled 2p³), Be (filled 2s²), and noble gases (filled shells) have less negative or positive ΔegH.
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