Final answer:
The vapor pressures of acetone and chloroform are influenced by their specific intermolecular forces; acetone having weaker forces and therefore a higher vapor pressure, while chloroform's stronger forces lead to a lower vapor pressure. Acetone's polar carbonyl is less able to hydrogen bond than chloroform's chlorines, affecting the boiling points and vaporization energies.
Step-by-step explanation:
The vapor pressures of pure acetone and pure chloroform can be explained by their intermolecular forces (IMFs) and internal degrees of freedom. Acetone (CH3)2CO has a polar carbonyl group which allows for dipole-dipole interactions. However, these interactions are not as strong as the hydrogen bonding seen in molecules that have -OH groups, leading to a relatively high vapor pressure because the molecules can escape more readily into the gas phase.
Chloroform (CHCl3) experiences stronger IMFs due to its larger mass and its ability to participate in hydrogen bonding, albeit less extensively than alcohols. This result in lower vapor pressure compared to acetone at the same temperature since more energy is required to overcome these interactions for chloroform molecules to vaporize.
Acetone's overall molecular structure, with its combination of polar and nonpolar characteristics, leads to less restrictive IMFs than chloroform, which has heavier atoms and a slight ability to form hydrogen bonds due to the presence of the chlorine atoms. These differences delineate why acetone has a lower boiling point and higher vapor pressure than chloroform.