The first ionization enthalpies of $\mathrm{Na}, \mathrm{Mg}$ and $\mathrm{Si}$ are respectively 496,737 and…
The first ionization enthalpies of $\mathrm{Na}, \mathrm{Mg}$ and $\mathrm{Si}$ are respectively 496,737 and $786 \mathrm{KJ} \mathrm{mol}^{-1}$. What would be the first ionization enthalpy of $\mathrm{Al}$ in $\mathrm{KJ} \mathrm{mol}^{-1}$ ?
$450$
$750$
$575$
$800$
Solution
$\mathrm{Na}, \mathrm{Mg}, \mathrm{Al}$ and $\mathrm{Si}$ lie in the same period $\left(3^{\text {rd }}\right.$ period) and therefore the effective nuclear charge increases as $\mathrm{Na} < \mathrm{Mg} < \mathrm{Al} < \mathrm{Si}$ because of which their ionization enthalpy (first) is also supposed to increase in this manner.
However, the configuration of $\mathrm{Al}$ is $1 s^2 2 s^2 \cdot 2 p^6 3 s^2 3 p^1$ in which the loss of one $3 \mathrm{p}$ - electron results in a stable, fully-filled configuration.
This stable configuration is not as stable as that of $\mathrm{Na}^{+}$ which is $1 s^2 2 s^2 2 p^6$ and therefore the first ionization energy of $\mathrm{Al}$ will be higher than that of $\mathrm{Na}$ but lower than that of $\mathrm{Mg}$.
$\Rightarrow$ First ionization energy of $\mathrm{Al}=575 \mathrm{~kJ} \mathrm{~mol}^{-1}$.