Q.Discuss the factors affecting electron gain enthalpy and the trend in its variation in the periodic table.

Concept: Standard State Conditions — Electron gain enthalpy is the energy change when an isolated gaseous atom gains an electron to form a negative ion. It is measured under standard conditions (1 atm, 298 K) for the gaseous state.

Reasoning:

1. Atomic size: As size increases down a group, the added electron is farther from the nucleus, experiencing less attraction — electron gain enthalpy becomes less negative (more positive).
2. Nuclear charge: Across a period, increasing nuclear charge pulls the added electron more strongly, making electron gain enthalpy more negative.
3. Electron-electron repulsion: Adding an electron to a half-filled or fully-filled subshell (e.g., N, Be, noble gases) faces extra repulsion, making the process less favourable (positive or less negative values).

Trends: …

The Concept: Why Atoms Gain Electrons

When a neutral gaseous atom captures an electron, it releases energy — this is called electron gain enthalpy ($\Delta_{eg}H$). A more negative value means the atom has a stronger tendency to accept an electron. Think of it as the "electron hunger" of an atom. The process is:

The sign convention is important: most first electron gain enthalpies are exothermic (negative), because the added electron experiences attraction from the nucleus. But adding a second electron to an already negative ion is always endothermic (positive), because you must overcome electrostatic repulsion.

Factors That Affect Electron Gain Enthalpy

Four main factors determine how much energy is released when an atom gains an electron:

1. Atomic size — Larger atoms have their outermost electrons farther from the nucleus. The incoming electron feels weaker attraction, so less energy is released. Smaller atoms have a tighter hold, giving more negative $\Delta_{eg}H$.

2. Nuclear charge — A higher positive charge in the nucleus pulls the incoming electron more strongly, releasing more energy. This makes $\Delta_{eg}H$ more negative.

3. Electron configuration — This is the most subtle factor. Atoms with a half-filled or fully-filled subshell (like $p^3$, $p^6$, $d^5$, $d^{10}$, $f^7$, $f^{14}$) are unusually stable. Adding an electron disrupts this stability, so the energy released is much smaller — sometimes the process even becomes endothermic. Conversely, atoms one electron short of such a stable configuration release a lot of energy when they gain that electron.

4. Shielding effect — Inner electrons partially block the nuclear charge felt by outer electrons. More inner shells mean weaker effective nuclear charge ($Z_{eff}$), making $\Delta_{eg}H$ less negative.

Periodic Trends

Across a Period (Left to Right)

As you move from left to right across a period, atomic size decreases and nuclear charge increases. Both factors make the atom hold electrons more tightly. The incoming electron feels stronger attraction, so $\Delta_{eg}H$ becomes more negative (more exothermic).

Notice the exceptions: Be and N have positive values because they have stable configurations ($2s^2$ and $2p^3$ respectively). Ne also has a positive value because its octet is already complete.

Down a Group (Top to Bottom)

As you go down a group, atomic size increases significantly, and shielding increases because more inner shells are present. Even though nuclear charge also increases, the effective nuclear charge felt by the incoming electron decreases. The electron is added farther from the nucleus, so less energy is released. …