Identifying More Acidic Species

To determine the more acidic species in each pair, we need to consider the nature of the compounds and their behavior in terms of donating protons (Brønsted-Lowry acidity) or accepting electron pairs (Lewis acidity).

Pair (i): $NH_4Cl$ and $FeO$

• $NH_4Cl$ is a salt formed from a weak base ($NH_4OH$) and a strong acid ($HCl$). When dissolved in water, it undergoes hydrolysis, producing $H^+$ ions.
  $NH_4Cl(aq) \rightarrow NH_4^+(aq) + Cl^-(aq)$
  $NH_4^+(aq) + H_2O(l) \rightleftharpoons NH_4OH(aq) + H^+(aq)$
  This process releases protons, making the solution acidic.
• $FeO$ is an oxide of a metal (Iron). Metallic oxides are generally basic in nature. For example, many react with water to form bases or react with acids to form salt and water.

Comparing the two, $NH_4Cl$ behaves as an acidic salt in water, while $FeO$ is a basic oxide. Therefore, $NH_4Cl$ is the more acidic species.

Pair (ii): $H_3N \cdot BF_3$ and $BCl_3$

• $H_3N \cdot BF_3$ is an adduct formed by the reaction between the Lewis base ($NH_3$) and the Lewis acid ($BF_3$). The boron atom in $BF_3$ accepts an electron pair from the nitrogen atom in $NH_3$ to form a coordinate covalent bond. This adduct is a stable neutral molecule.
• $BCl_3$ is a compound where boron is bonded to three chlorine atoms. Boron has an incomplete octet and acts as a Lewis acid, readily accepting an electron pair.

We need to compare the Lewis acidity. The Lewis acidity of boron trihalides ($BX_3$) depends on the extent of $p\pi - p\pi$ backbonding from the halogen atom to the boron atom. This backbonding makes the boron less electron-deficient.
In $BF_3$, the small size and high electronegativity of Fluorine allows for effective $2p\pi - 2p\pi$ backbonding with boron, reducing its Lewis acidity.
In $BCl_3$, the larger size and lower electronegativity of Chlorine make $3p\pi - 2p\pi$ backbonding less effective compared to Fluorine in $BF_3$. This leaves the boron atom in $BCl_3$ more electron-deficient and thus a stronger Lewis acid than $BF_3$.
Since $BCl_3$ is a stronger Lewis acid than $BF_3$, and $H_3N \cdot BF_3$ is formed by the reaction of $BF_3$ (a weaker Lewis acid) with a base ($NH_3$), $BCl_3$ is the more acidic species (stronger Lewis acid) compared to the adduct $H_3N \cdot BF_3$.

Pair (iii): $HSO_3F$ and $HF$

• $HF$ is hydrofluoric acid, a weak acid in aqueous solution.
• $HSO_3F$ is fluorosulfuric acid, a superacid. Superacids are acids that are more acidic than pure sulfuric acid ($H_2SO_4$).

The acidity of an acid $HA$ is related to the stability of its conjugate base $A^-$. The more stable the conjugate base, the stronger the acid.
For $HF$, the conjugate base is $F^-$.
For $HSO_3F$, the conjugate base is $SO_3F^-$.
The $SO_3F^-$ ion is resonance stabilized, and the strong electron-withdrawing effect of the $-\text{SO}_3F$ group effectively delocalizes the negative charge, making it a very stable conjugate base. The $F^-$ ion from $HF$ does not have this extensive charge delocalization.
Due to the high stability of its conjugate base, $HSO_3F$ readily donates a proton and is significantly more acidic than $HF$. $HSO_3F$ is considered one of the strongest known Brønsted-Lowry acids.

Summary of More Acidic Species

Based on the analysis:

• In pair (i) $NH_4Cl$ and $FeO$, $NH_4Cl$ is more acidic.
• In pair (ii) $H_3N \cdot BF_3$ and $BCl_3$, $BCl_3$ is more acidic.
• In pair (iii) $HSO_3F$ and $HF$, $HSO_3F$ is more acidic.

Therefore, the more acidic species in (i), (ii) and (iii) are $NH_4Cl$, $BCl_3$, and $HSO_3F$ respectively.