Question

Without referring to Table 6 , predict which compound in each of the following pairs has the higher heat of vaporization (cal/mol): (a) \(\mathrm{H}_{2} \mathrm{O}\) or \(\mathrm{H}_{2} \mathrm{Se}\) (b) \(\mathrm{H}_{2} \mathrm{~S}\) or \(\mathrm{H}_{2} \mathrm{Te}\)

Short answer

(a) H₂O has a higher heat of vaporization; (b) H₂Te has a higher heat of vaporization.

Step-by-step solution

Understanding Heat of Vaporization

The heat of vaporization is the amount of energy required to transform a given quantity of a substance from a liquid into a gas at constant pressure. Molecules with stronger intermolecular forces generally have higher heats of vaporization because more energy is needed to overcome these forces.

Compare Water (H₂O) and Hydrogen Selenide (H₂Se)

Water (HiO) has strong hydrogen bonds due to the high electronegativity of oxygen, leading to strong intermolecular interactions. Hydrogen selenide (HSe), with selenium, a less electronegative element than oxygen, experiences weaker dipole-dipole interactions. Thus, compound water (HO) likely has a higher heat of vaporization.

Compare Hydrogen Sulfide (H₂S) and Hydrogen Telluride (H₂Te)

Hydrogen sulfide (HS) and hydrogen telluride (HTe) both primarily experience dipole-dipole interactions, though these interactions are weaker than hydrogen bonding. Considering size and polarizability, hydrogen telluride being larger and more polarizable, it can form stronger London dispersion forces than hydrogen sulfide, suggesting HTe might have a higher heat of vaporization.

Key concepts

Intermolecular Forces

Intermolecular forces are the attractions between molecules that determine many physical properties of substances, such as boiling points, melting points, and heat of vaporization. These forces are generally weaker than the intramolecular forces, which are the bonds within a molecule, such as covalent bonds. The different types of intermolecular forces include:

• London Dispersion Forces: These are the weakest type of intermolecular forces that occur between all molecules, whether they are polar or nonpolar. They arise from temporary fluctuations in the electron distribution around atoms, leading to short-lived dipoles that induce dipoles in neighboring molecules.
• Dipole-Dipole Interactions: These occur between polar molecules due to their permanent dipoles. The positive end of one molecule is attracted to the negative end of another, leading to these relatively strong attractions.
• Hydrogen Bonding: This is a special case of dipole-dipole interaction. It occurs in molecules where hydrogen is bonded to highly electronegative atoms such as oxygen, nitrogen, or fluorine, leading to very strong dipolar interactions.

Intermolecular forces play a crucial role in determining the heat of vaporization, as stronger forces require more energy to break during the phase change from liquid to gas.

Hydrogen Bonding

Hydrogen bonding is a significant intermolecular force that significantly affects the properties of substances. It occurs when hydrogen is covalently bonded to fluorine, oxygen, or nitrogen - atoms known for their high electronegativities. The result is a strong dipole, where hydrogen bears a partial positive charge, and the electronegative atom bears a partial negative charge. These configurations allow hydrogen atoms to form strong interactions with lone pairs of electrons on nearby molecules.
Consider water ( H
,O), a prime example of hydrogen bonding. The oxygen atom, being more electronegative, pulls electrons closer, leaving the hydrogen with a positive partial charge. This allows water molecules to form an extensive network of hydrogen bonds, leading to water’s relatively high heat of vaporization and boiling point compared to other small molecules. These bonds require substantial energy to break as they denote a higher degree of intermolecular forces.
Thus, hydrogen bonding is not only responsible for the liquid state of water but also gives rise to its many unique properties, such as high surface tension and a solid form less dense than its liquid counterpart.

Dipole-Dipole Interactions

Dipole-dipole interactions occur in polar molecules that have permanent dipoles due to uneven electron distribution, resulting in a molecule with a positive end and a negative end. These interactions are stronger than London dispersion forces but weaker than hydrogen bonds. The presence of polar bonds and geometrical arrangement that does not cancel out the dipoles leads to such interactions.
Take, for instance, hydrogen sulfide ( H
,S), which is a polar molecule because of its bent shape and difference in electronegativity between hydrogen and sulfur. The individual dipoles do not cancel out, resulting in a net dipole moment. The dipole-dipole interactions in hydrogen sulfide, while significant, cannot compete with hydrogen bonding as seen in water. This notable difference means substances with only dipole-dipole interactions generally have a lower heat of vaporization than those engaging in hydrogen bonding.
Understanding dipole-dipole interactions helps predict properties such as the relative heat of vaporization in compounds and can indicate why certain molecules demonstrate higher energy requirements for phase changes compared to others. Such insights are essential for making accurate predictions in scenarios described in chemical exercises.