Can you get this USNCO question?

Atomic structure and periodicityNational 2022

Which diatomic molecule has the greatest bond dissociation enthalpy?

25% of national qualifiers got this right

Explanation

draft

Bond dissociation enthalpies (approximate, kJ/mol):
- $\ce{F-F}$ : 159
- $\ce{Cl-Cl}$ : 243
- $\ce{Br-Br}$ : 193
- $\ce{I-I}$ : 151

At first glance one might expect $\ce{F-F}$ to be strongest since F is the smallest halogen and bond strength generally decreases with increasing atomic size (as seen going from Cl to Br to I). However, the $\ce{F-F}$ bond is anomalously weak because fluorine's small size causes severe repulsion between the lone pairs on each atom, destabilizing the bond.

So the trend is not monotonic: $\ce{Cl-Cl} > \ce{Br-Br} > \ce{F-F} > \ce{I-I}$ , making $\ce{Cl-Cl}$ the strongest diatomic halogen bond.

Assumes bond strength simply increases as atomic radius decreases, ignoring the lone-pair repulsion effect that weakens the small $\ce{F-F}$ bond.

Correct: $\ce{Cl-Cl}$ has the highest bond dissociation enthalpy (~243 kJ/mol) among the halogens because it is small enough for strong orbital overlap but large enough to avoid the severe lone-pair repulsion that weakens $\ce{F-F}$ .

Overestimates bromine's bond strength by assuming a smooth decrease from chlorine, without recognizing chlorine has the true maximum in the series.

Incorrectly assumes bond strength increases with atomic size (larger polarizability), when in fact larger atoms have weaker orbital overlap, giving iodine the weakest bond, not the strongest.

Concept · Trends in halogen bond dissociation enthalpy (lone-pair repulsion in F2)