The table lists experimental molar enthalpies of vaporization ΔHvap at 298K for several non-polar alkanes and structurally similar polar aldehydes.
Substance
Molar mass (g mol−1)
ΔHvap (kJ mol−1)
Gas-phase dipole μ (D)
n-Butane
58.1
22.4
0.00
n-Pentane
72.2
25.8
0.00
n-Hexane
86.2
28.9
0.00
Propanal
58.1
29.7
2.70
Butanal
72.2
33.1
2.72
Pentanal
86.1
36.4
2.75
Assume that, for molecules of comparable size, London dispersion forces account for most of the cohesive energy in alkanes, while aldehydes experience an additional dipole–dipole contribution.
Use the data to analyze the relative magnitudes of London dispersion and dipole–dipole interactions.
(a) Using the three alkane data points, determine the best-fit linear relationship between ΔHvap and molar mass. Show your calculations and report the slope and intercept (both to two significant figures).
(b) Apply the linear relationship from part (a) to predict ΔHvap values for propanal, butanal, and pentanal based solely on London dispersion forces.
(c) For each aldehyde, calculate the difference between the experimental ΔHvap and the value predicted in part (b). Interpret this difference as an estimate of the dipole–dipole contribution per mole, and tabulate your results.
(d) Compare the estimated dipole–dipole contribution among the three aldehydes. Discuss whether the results are consistent with the nearly constant dipole moments listed, and explain any trend you observe.
(e) Identify one limitation of using enthalpy of vaporization data alone to separate London dispersion and dipole–dipole contributions, and propose an additional experiment or measurement that could reduce this uncertainty.