Q.Define IE1 and IE2. Why is IE2 > IE1 for a given atom? Discuss the factors that effect IE of an element.

IE1 and IE2 are the energies to remove successive electrons from a gaseous atom/ion; IE2 exceeds IE1 because removing an electron shrinks the species and increases the effective pull of the nucleus on the remaining electrons.

Definitions

• First ionization enthalpy (IE1): the minimum energy required to remove the most loosely bound (outermost) electron from an isolated, neutral atom in the gaseous state, forming a unipositive gaseous ion: $M(g) \rightarrow M^{+}(g) + e^{-}, \quad \Delta H = IE_1$
• Second ionization enthalpy (IE2): the minimum energy required to remove an electron from the already-formed unipositive gaseous ion, forming a dipositive gaseous ion: $M^{+}(g) \rightarrow M^{2+}(g) + e^{-}, \quad \Delta H = IE_2$

Why is IE2 always greater than IE1?

After the first electron is removed, the resulting cation M+ has one fewer electron but the same number of protons as the neutral atom. This means:

1. The effective nuclear charge (Zeff) experienced by each remaining electron INCREASES, because the same nuclear charge now attracts fewer electrons.
2. Electron-electron repulsion DECREASES (fewer electrons in the same shells), so the electron cloud contracts — the ion is smaller than the neutral atom. Both effects mean the remaining electrons are held more tightly by the nucleus, so more energy is needed to remove the next electron. Hence IE2 > IE1 always, and generally IE3 > IE2 > IE1, and so on for successive ionizations.

Factors affecting ionization enthalpy

1. Atomic/ionic size: Larger atoms have their outermost electrons farther from the nucleus, so they are held less strongly → lower ionization enthalpy. IE generally decreases down a group as atomic size increases.
2. Nuclear charge: A greater nuclear charge (more protons) attracts electrons more strongly → higher ionization enthalpy. IE generally increases across a period as nuclear charge increases (at roughly constant shielding). …