A researcher is studying a hypothetical four-step reaction A → E in which intermediates B, C, and D appear sequentially. The table lists the potential energy of each species relative to A (all energies in kJ mol⁻¹).
Reactants (A): 0
TS₁: +120 Intermediate (B): +10
TS₂: +95 Intermediate (C): –25
TS₃: +110 Intermediate (D): –40
TS₄: +80 Products (E): –90
To test the effect of solvent polarity, the researcher repeats the experiment in a more polar solvent. Only the energy of intermediate C is stabilized by an additional 35 kJ mol⁻¹; all other energies remain unchanged.
(a) Using the original energy data, calculate ΔH for each elementary step and the overall reaction.
(b) Identify the slowest step in the original mechanism and justify your answer numerically.
(c) For the reaction carried out in the more polar solvent, recalculate the ΔH for step 2 (B → C) and discuss how the change affects the heights of TS₂ and TS₃ relative to reactants if Hammond postulates are applied.
(d) Predict whether the rate-determining step changes in the more polar solvent. Explain fully, making specific reference to altered transition-state energies and any shifts in the reaction coordinate that you infer.
(e) Discuss how, if at all, the change in solvent polarity would influence the equilibrium constant for A ⇌ E. Support your prediction with enthalpic considerations drawn from the modified energy profile.