Chemistry - Complete P-block Question with Solution | TestHub

A disproportionation reaction involves an element in an intermediate oxidation state simultaneously undergoing oxidation and reduction. For phosphorus oxoacids, the oxidation state of phosphorus determines if disproportionation can occur.

Let's analyze the oxidation states of phosphorus in each option:

1. H₃PO₂ (Hypophosphorous acid):

Let the oxidation state of P be $x$.

$3(+1)+x+2(-2)=0$

$3+x-4=0$

$x=+1$

Phosphorus in H₃PO₂ is in the +1 oxidation state, which is an intermediate state (P can go to 0, -3, +3, +5). Thus, it can disproportionate.

2. H₃PO₄ (Orthophosphoric acid):

Let the oxidation state of P be $x$.

$3(+1)+x+4(-2)=0$

$3+x-8=0$

$x=+5$

Phosphorus in H₃PO₄ is in its highest possible oxidation state (+5). Therefore, it can only be reduced and cannot undergo further oxidation, meaning it cannot disproportionate.

3. H₃PO₃ (Phosphorous acid):

Let the oxidation state of P be $x$.

$3(+1)+x+3(-2)=0$

$3+x-6=0$

$x=+3$

Phosphorus in H₃PO₃ is in the +3 oxidation state, which is an intermediate state. Thus, it can disproportionate.

Therefore, H₃PO₄ does not disproportionate because phosphorus is already in its highest oxidation state (+5).

The final answer is $\boxed{H_3{\mathrm{PO}}_4}$.