Animated Solution for Chemistry - Periodicity in Properties: The correct order of first ionisation enthalpy is

Why the Periodic Table Isn't Always Predictable

When you first learn about the periodic table, it feels like a beautifully organized grid where everything follows strict rules. One of the most fundamental rules is the trend of first ionization enthalpy—the energy required to strip the most loosely bound electron from an isolated gaseous atom.

As we move from left to right across a period, the atomic number increases. This means more protons are added to the nucleus, increasing the effective nuclear charge ($Z_{\text{eff}}$). This stronger nuclear pull holds the outermost electrons more tightly, making it harder to remove them. Therefore, the general rule of thumb is: Ionization enthalpy increases across a period.

If we apply this logic blindly to the third-period elements Magnesium (Mg), Aluminum (Al), Phosphorus (P), and Sulfur (S), we would expect the order to be:

But chemistry loves exceptions, and these exceptions are born from the beautiful quantum mechanical structure of atoms.

Magnesium vs. Aluminum

Let's zoom in on Magnesium ($Z=12$) and Aluminum ($Z=13$).

The electronic configuration of Magnesium is $[\text{Ne}]3s^2$. Notice that the $3s$ orbital is fully filled. A fully filled subshell is a state of high stability. The electrons are perfectly paired, and the $s$-orbital penetrates closer to the nucleus, feeling a strong attractive force.

Now look at Aluminum: $[\text{Ne}]3s^{2}3p^1$. The outermost electron resides alone in the $3p$ orbital. This single $p$-electron is shielded by the inner $3s$ electrons and is slightly further from the nucleus. It is much easier to pluck away this lone $3p$ electron from Aluminum than to break the stable, fully filled $3s$ pair of Magnesium.

Because of this extra stability in Magnesium, it requires more energy to ionize than Aluminum. Thus, the trend reverses:

Phosphorus vs. Sulfur

We see a similar quantum drama unfold between Phosphorus ($Z=15$) and Sulfur ($Z=16$).

The electronic configuration of Phosphorus is $[\text{Ne}]3s^{2}3p^3$. According to Hund's Rule, these three $p$-electrons occupy the $p_x$, $p_y$, and $p_z$ orbitals singly. This creates a half-filled subshell. Half-filled subshells possess extra stability due to a highly symmetrical distribution of electron density and maximum exchange energy.

Sulfur, on the other hand, has the configuration $[\text{Ne}]3s^{2}3p^4$. Here, one of the $p$-orbitals must contain a pair of electrons. These two electrons occupying the same small region of space repel each other strongly (inter-electronic repulsion). This repulsion makes it relatively easier to remove one of the paired electrons to achieve the stable half-filled state of $3p^3$.

Consequently, breaking the stable half-filled configuration of Phosphorus requires significantly more energy than removing the repelling electron from Sulfur. Once again, the trend reverses:

Putting It All Together

Now that we have uncovered the quantum secrets of these elements, we can construct the true, experimentally verified order of their first ionization enthalpies.

Aluminum is the easiest to ionize because of its lone $3p$ electron. Magnesium comes next, guarding its fully filled $3s$ orbital. Then comes Sulfur, eager to lose an electron to relieve repulsion. Finally, Phosphorus stands as the hardest to ionize among the four, fiercely protecting its perfectly symmetrical half-filled $3p$ subshell.

The final correct order is:

This problem is a classic reminder that in chemistry, general trends only tell half the story. The real magic lies in the quantum mechanical details of electronic configurations!