Q.Define ionisation enthalpy. Discuss the factors affecting ionisation enthalpy of the elements and its trends in the periodic table.

Ionisation Enthalpy

Ionisation enthalpy is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state to form a gaseous cation:

It is always endothermic and measured in kJ mol⁻¹.

Factors Affecting Ionisation Enthalpy

1. Nuclear charge: Higher $Z_{\text{eff}}$ pulls electrons more strongly, increasing ionisation enthalpy.
2. Atomic size: Larger radius weakens the nucleus–electron attraction, lowering ionisation enthalpy.
3. Shielding/screening: Inner electrons reduce the effective nuclear charge felt by valence electrons; more shielding lowers ionisation enthalpy.
4. Electronic configuration: Half-filled ($p^3, d^5$) and fully filled ($p^6, d^{10}$) subshells are extra stable, requiring more energy to remove an electron.

Periodic Trends

What is Ionisation Enthalpy?

Ionisation enthalpy (or ionisation energy) is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state, converting it into a gaseous cation:

The process is always endothermic because energy must be supplied to overcome the electrostatic attraction between the nucleus and the electron. The value is typically expressed in $\text{kJ mol}^{-1}$ or $\text{eV atom}^{-1}$.

Successive ionisation enthalpies ($\Delta_i H_1, \Delta_i H_2, \Delta_i H_3, \ldots$) refer to the removal of the first, second, third electrons, and so on. Each successive ionisation requires more energy because you are removing an electron from an increasingly positive ion.

Factors Affecting Ionisation Enthalpy

Understanding why some atoms hold their electrons more tightly than others requires examining four interrelated factors:

1. Nuclear Charge ($Z$)

As the number of protons in the nucleus increases, the positive charge pulling on the electrons grows stronger. Higher nuclear charge means a tighter grip on the outermost electron, so ionisation enthalpy increases.

For example, across the second period from lithium to neon, the nuclear charge rises from $+3$ to $+10$, and ionisation enthalpy climbs accordingly.

2. Atomic Radius

The farther an electron sits from the nucleus, the weaker the electrostatic attraction it experiences (Coulomb's law: force $\propto \frac{1}{r^2}$). Larger atoms have their valence electrons at greater distances, making them easier to remove.

This is why ionisation enthalpy generally decreases as you move down a group: each successive element adds a new shell, dramatically increasing atomic size.

3. Shielding (Screening) Effect

Inner-shell electrons repel outer-shell electrons, partially canceling the nuclear attraction. This "shielding" reduces the effective nuclear charge ($Z_{\text{eff}}$) felt by valence electrons:

where $\sigma$ is the shielding constant. More inner electrons mean greater shielding and lower ionisation enthalpy. Down a group, the addition of complete inner shells significantly increases shielding, contributing to the decrease in ionisation enthalpy.

4. Electronic Configuration (Subshell Stability)

Certain electron configurations are exceptionally stable and resist ionisation:

• Half-filled subshells ($p^3, d^5, f^7$): Electrons occupy separate orbitals with parallel spins, minimizing electron-electron repulsion (exchange energy stabilization).
• Fully-filled subshells ($s^2, p^6, d^{10}, f^{14}$): Maximum symmetry and stability.

Removing an electron from these stable arrangements requires extra energy, creating anomalies in otherwise smooth trends.

Periodic Trends in Ionisation Enthalpy

Across a Period (Left to Right)

Ionisation enthalpy generally increases across a period.

Why? Moving left to right:

• Nuclear charge increases (more protons).
• Atomic radius decreases (electrons added to the same shell, pulled closer).
• Shielding remains roughly constant (no new inner shells).

The net effect is a stronger hold on valence electrons.

Key exceptions arise from subshell stability:

These dips interrupt the overall upward trend but are entirely predictable from electronic structure.

Down a Group (Top to Bottom) …