Q.The first ionization enthalpy values (in kJ mol$^{-1}$) of group 13 elements are: B 801, Al 577, Ga 579, In 558, Tl 589. How would you explain this deviation from the general trend?

The key idea is that the general trend of decreasing ionization enthalpy down a group is disrupted by the d-block and f-block contraction effects, which increase the effective nuclear charge experienced by the outermost electrons.

Reasoning:

1. From B to Al, ionization enthalpy drops sharply (801 → 577 kJ/mol) as expected, because the atomic radius increases and shielding improves.
2. From Al to Ga, the value barely changes (577 → 579). This is due to the d-block contraction: the 3d electrons in Ga (electronic configuration: $[\text{Ar}] 3d^{10} 4s^2 4p^1$) shield the 4p electron poorly, so the effective nuclear charge is higher than expected, nearly cancelling the size increase. …

The first ionization enthalpy of an element is the energy required to remove the most loosely bound electron from a gaseous atom. In group 13 (boron family), the general trend expected is a steady decrease as we move down the group. This is because atomic size increases, and the outermost electron is farther from the nucleus, making it easier to remove.

However, the given values show a clear deviation:

B (801) > Al (577) ≈ Ga (579) > In (558) < Tl (589).

Notice that gallium has a slightly higher value than aluminium, and thallium has a higher value than indium. This is not a simple monotonic drop.

Let’s break down the reasons step by step.

1. The expected trend and the role of atomic size

   Down the group, each successive element has an additional electron shell. This increases the atomic radius, so the outermost electron is, on average, farther from the nucleus. The force of attraction between the nucleus and the electron weakens, so ionization enthalpy should decrease. This holds true from B to Al (801 → 577) and from Al to Ga (577 → 579 is a tiny increase, but nearly constant), and from Ga to In (579 → 558). But the jump from In to Tl (558 → 589) is an increase, which is opposite to expectation.

2. The key factor: shielding by inner electrons

   The effective nuclear charge ($Z_{\text{eff}}$) felt by the outermost electron depends not only on the actual nuclear charge but also on how well inner electrons shield it. Electrons in different subshells shield differently.

   • s and p electrons shield relatively well.
   • d electrons (in the penultimate shell) shield poorly because their orbitals are more diffuse and penetrate less.
   • f electrons (in the antepenultimate shell) shield even more poorly.

3. Gallium’s anomaly: the d-block contraction

   Gallium (atomic number 31) comes after the first transition series (Sc to Zn). Its electron configuration is $[\text{Ar}] 3d^{10} 4s^2 4p^1$. The 3d electrons are in the inner shell but do not shield the 4p electron effectively. As a result, the 4p electron experiences a higher $Z_{\text{eff}}$ than expected. This increases the ionization enthalpy of Ga, making it nearly equal to that of Al, despite being one period lower.

   Watch out

   A common mistake is to think that Ga should have a much lower ionization enthalpy than Al simply because it is below. But the poor shielding by the 3d electrons counteracts the size increase, so the value barely changes.

4. Indium’s drop: normal trend resumes

   Indium (atomic number 49) has configuration $[\text{Kr}] 4d^{10} 5s^2 5p^1$. The 4d electrons also shield poorly, but the atomic radius increase from Ga to In is significant enough that the ionization enthalpy still drops (579 → 558). The d-block contraction effect is present but not strong enough to reverse the trend.

5. Thallium’s anomaly: the f-block contraction (lanthanoid contraction) …