Understand the trend of first ionization enthalpy

First ionization enthalpy is the energy required to remove the outermost electron from an isolated gaseous atom.

General periodic trends:

• It increases across a period from left to right due to increasing effective nuclear charge.
• It decreases down a group due to increase in atomic size and shielding.

Compare the given elements

The elements are:

• $\mathrm{C}$ : Period 2, Group 14
• $\mathrm{Al}$ : Period 3, Group 13
• $\mathrm{Si}$ : Period 3, Group 14
• $\mathrm{N}$ : Period 2, Group 15

Apply periodic trends

• Between period 2 and period 3, period 2 elements generally have higher ionization enthalpy than the corresponding period 3 elements.
• So, $\mathrm{C}$ and $\mathrm{N}$ should have higher values than $\mathrm{Al}$ and $\mathrm{Si}$.
• Within period 2, ionization enthalpy generally increases from $\mathrm{C}$ to $\mathrm{N}$.

Special stability of nitrogen

Nitrogen has electronic configuration: $1s^2\,2s^2\,2p^3$

The $2p^3$ configuration is half-filled, which is especially stable. Removing one electron from nitrogen disturbs this stable half-filled arrangement, so more energy is required.

Carbon has configuration: $1s^2\,2s^2\,2p^2$

This is less stable than nitrogen’s half-filled $2p^3$ configuration.

Conclusion

Therefore, among $\mathrm{C}$, $\mathrm{Al}$, $\mathrm{Si}$, and $\mathrm{N}$, the element with the highest first ionization enthalpy is: $\boxed{\mathrm{N}}$

Option check

• A: $\mathrm{C}$ — incorrect
• B: $\mathrm{Al}$ — incorrect
• C: $\mathrm{Si}$ — incorrect
• D: $\mathrm{N}$ — correct