Electron Gain Enthalpy ($\Delta_{eg}H$)

Electron gain enthalpy is defined for the process: $\ce{X(g) + e^- -> X^-(g)}$. When energy is released (exothermic), $\Delta_{eg}H$ is negative. When energy is absorbed (endothermic), it is positive. For most nonmetals, the first electron gain enthalpy is negative (energy-releasing), while the second electron gain enthalpy (adding a second electron to an already negatively charged ion) is always positive because of electron–electron repulsion.

Moving from left to right across Period 2 or Period 3, nuclear charge (Z) increases from, for example, Na (Z=11) to Cl (Z=17). With the same principal quantum number shielding, the effective nuclear charge ($Z_{eff}$) rises. The incoming electron experiences a stronger pull, so more energy is released. For example: $\Delta_{eg}H$(Na) ≈ −53 kJ mol⁻¹, while $\Delta_{eg}H$(Cl) ≈ −349 kJ mol⁻¹. Notable interruptions occur at N, O, and the noble gases due to half-filled stability and fully filled shells, respectively.

As you descend Group 17 from F to Cl to Br to I, the atomic radius increases and more electron shells are present. The added electron enters an orbital that is farther from the nucleus and more shielded from its attraction. As a result, $\Delta_{eg}H$ generally becomes less negative: $\Delta_{eg}H$(F) ≈ −328 kJ mol⁻¹, $\Delta_{eg}H$(Cl) ≈ −349 kJ mol⁻¹, $\Delta_{eg}H$(Br) ≈ −325 kJ mol⁻¹, $\Delta_{eg}H$(I) ≈ −295 kJ mol⁻¹. The anomaly of Cl > F is central to this trend.

Fluorine has the highest electronegativity of all elements, yet its electron gain enthalpy (−328 kJ mol⁻¹) is less negative than that of chlorine (−349 kJ mol⁻¹). The reason lies in the unusually small size of fluorine's 2p orbitals. When a 9th electron is added to the $2p$ subshell (which already contains 5 electrons across three compact orbitals), there is severe interelectronic repulsion. Chlorine's $3p$ orbitals are significantly larger, so the same number of electrons occupies a greater volume, and repulsion is reduced. The net energy released is therefore greater for chlorine.

Noble gases (He, Ne, Ar, Kr, Xe) possess completely filled valence shells ($ns^2np^6$). An incoming electron would have to enter the next higher-energy principal quantum shell, which is energetically very unfavorable, giving positive $\Delta_{eg}H$ values. Similarly, nitrogen ($[He]\,2s^2\,2p^3$) has a half-filled 2p subshell, which is extra-stable due to exchange energy and minimal repulsion (one electron per 2p orbital). The added electron would be forced to pair up in an already occupied 2p orbital, causing repulsion and leading to $\Delta_{eg}H$(N) ≈ +7 kJ mol⁻¹ (slightly positive), while oxygen immediately to its right has $\Delta_{eg}H$ ≈ −141 kJ mol⁻¹.

• X(g) — Gaseous atom (mol): One mole of isolated gaseous atoms of element X in their ground state electronic configuration.
• e^- — Electron (mol): One mole of electrons added from outside, representing the electron being gained.
• X^-(g) — Gaseous anion (mol): One mole of singly negatively charged gaseous ions formed after electron addition.
• \Delta_{eg}H — Electron gain enthalpy (kJ mol⁻¹): The enthalpy change for the process. Negative means energy is released (exothermic); positive means energy is absorbed (endothermic).

Problem. The electron gain enthalpy of fluorine is −328 kJ mol⁻¹ and that of chlorine is −349 kJ mol⁻¹. (a) Write the equation for the electron gain of chlorine. (b) Explain why chlorine has a more negative electron gain enthalpy than fluorine, even though fluorine is the more electronegative element. (c) Predict whether the electron gain enthalpy of bromine will be more or less negative than chlorine, and explain your reasoning.

Answer. Chlorine has a more negative electron gain enthalpy (−349 kJ mol⁻¹) than fluorine (−328 kJ mol⁻¹) because fluorine's 2p orbitals are very compact, causing greater interelectronic repulsion when the incoming electron is added. Bromine (−325 kJ mol⁻¹) has a less negative value than chlorine because the added electron enters the larger 4p orbital, farther from the nucleus with increased shielding.

• Ionization Enthalpy — Ionization enthalpy is the complementary concept to electron gain enthalpy — it measures the energy required to remove an electron from a gaseous atom, and both follow similar periodic trends driven by effective nuclear charge and atomic radius.
• Electronegativity — Electronegativity (especially the Pauling scale) is directly related to electron gain enthalpy; elements with highly negative ΔegH tend to be highly electronegative, making this concept essential for understanding bond polarity and chemical reactivity.
• Atomic Radius — Atomic radius governs how close an incoming electron can get to the nucleus, directly affecting the magnitude of electron gain enthalpy; the F–Cl anomaly is entirely explained by fluorine's unusually small atomic radius.
• Electron Configuration — The exceptions in electron gain enthalpy (noble gases, nitrogen, phosphorus) are explained by electron configuration concepts such as fully filled shells, half-filled subshells, and Hund's rule.
• Effective Nuclear Charge (Z_eff) — Effective nuclear charge is the net positive charge experienced by an electron after accounting for shielding; it is the primary driver of trends in electron gain enthalpy across periods and down groups.