Final answer:
Straight-chain compounds have higher boiling points than branched compounds because their larger surface area allows for more extensive London dispersion forces between molecules.
The more extended structure of straight-chain molecules also results in better molecular packing in the liquid phase, requiring more energy to transition to a gas.
Step-by-step explanation:
When we compare the boiling points of straight-chain compounds to those of branched compounds, a clear pattern emerges.
Straight-chain compounds generally exhibit higher boiling points than their branched counterparts due to several interrelated factors concerning their molecular structure and intermolecular forces.
The primary force at play here is the London dispersion force, a weak intermolecular force that is nonetheless the driving force behind the boiling points of nonpolar organic molecules like alkanes.
The larger surface area of a straight-chain alkane allows for more extensive contact between molecules, resulting in greater London dispersion forces which in turn leads to higher boiling points.
A branched compound, in contrast, tends to be more compact, presenting a smaller surface area for intermolecular interactions and therefore exhibits a lower boiling point.
Additionally, straight-chain molecules tend to 'pack' together more efficiently in the liquid state, creating a higher demand for energy to change into the gaseous state.
For any given molar mass, this difference in boiling point can be quite significant. For example, n-pentane has an extended structure and a higher boiling point compared to its isomer neopentane, which is almost spherical in shape and has a lower boiling point.
The degree of branching and molecular shape are critical in determining the strength of intermolecular attractions and therefore boiling points.