Question

In which of the following substances would you expect to find hydrogen bonding? (a) \(\mathrm{HI}\) (d) \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\) (b) \(\mathrm{NH}_{3}\) (e) \(\mathrm{H}_{2} \mathrm{O}\) (c) \(\mathrm{CH}_{2} \mathrm{~F}_{2}\)

Short answer

\text{C}_{2}\text{H}_{5}\text{OH}, \text{NH}_{3}, \text{H}_{2}\text{O} exhibit hydrogen bonding.

Step-by-step solution

Identify the Requirements for Hydrogen Bonding

Hydrogen bonding occurs when hydrogen is bonded directly to a highly electronegative atom such as nitrogen (N), oxygen (O), or fluorine (F). These atoms must also have available lone pairs of electrons.

Evaluate \(\text{HI}\) for Hydrogen Bonding

In \(\text{HI}\), the hydrogen is bonded to iodine, which is less electronegative and lacks lone pairs that are strongly attracting hydrogen. Therefore, \(\text{HI}\) does not exhibit hydrogen bonding.

Evaluate \(\text{C}_{2}\text{H}_{5}\text{OH}\) for Hydrogen Bonding

\text{Ethanol} (\(\text{C}_{2}\text{H}_{5}\text{OH}\)) has a hydrogen atom directly bonded to oxygen, which is a highly electronegative atom with lone pairs. Therefore, ethanol exhibits hydrogen bonding.

Evaluate \(\text{NH}_{3}\) for Hydrogen Bonding

In \( \text{NH}_{3} \), the hydrogen is bonded to nitrogen, which is highly electronegative and has lone pairs. Therefore, ammonia (\( \text{NH}_{3} \)) exhibits hydrogen bonding.

Evaluate \(\text{H}_{2}\text{O}\) for Hydrogen Bonding

Water (\(\text{H}_{2}\text{O}\)) has hydrogen atoms directly bonded to oxygen, which is a highly electronegative atom with lone pairs. Therefore, water exhibits hydrogen bonding.

Evaluate \(\text{CH}_{2}\text{F}_{2}\) for Hydrogen Bonding

In \(\text{CH}_{2}\text{F}_{2}\), the hydrogen atoms are not bonded to fluorine directly but to carbon. Due to this bonding configuration, it does not exhibit hydrogen bonding.

Key concepts

Intermolecular Forces

Intermolecular forces are the forces of attraction between molecules. They are essential because they determine the physical properties of substances, such as boiling point, melting point, and solubility. There are several types of intermolecular forces, including van der Waals forces, dipole-dipole interactions, and hydrogen bonding.

Hydrogen bonding is a special type of dipole-dipole interaction and is generally stronger than other dipole-dipole interactions. This type of bond occurs when hydrogen is bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine.
For instance, in water (H₂O), the hydrogen atoms are attracted to the oxygen atoms of nearby water molecules, forming a hydrogen bond.

Understanding intermolecular forces is crucial because it helps explain why substances behave differently under various conditions.

Electronegativity

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. The higher the electronegativity, the more strongly an atom can attract bonding electrons.

In the context of hydrogen bonding, the atoms involved must be highly electronegative. For instance, nitrogen (N), oxygen (O), and fluorine (F) are highly electronegative elements essential in enabling hydrogen bonding.

A useful way to remember this concept is by looking at the periodic table: electronegativity increases across a period from left to right and decreases down a group. Thus, hydrogen bonding is generally observed in molecules where hydrogen is directly bonded to N, O, or F.

This increase in electronegativity across the periodic table helps in predicting the strength and behavior of bonding between different atoms.

Lone Pairs

Lone pairs are pairs of valence electrons that are not shared with another atom in a covalent bond. These electrons are found in the outermost shell of an atom and play a critical role in determining the geometry and reactivity of molecules.

In the context of hydrogen bonding, the presence of lone pairs on highly electronegative atoms (N, O, F) is essential. These lone pairs can form hydrogen bonds with hydrogen atoms in neighboring molecules. For example, in water (H₂O), oxygen has two lone pairs of electrons that participate in hydrogen bonding, contributing to water's unique properties like high boiling point and surface tension.

Lone pairs also affect the shape of molecules. The repulsion between lone pairs and bonding pairs leads to specific geometric arrangements, as explained by the VSEPR (Valence Shell Electron Pair Repulsion) theory.

Understanding lone pairs is crucial for predicting molecule behavior and interactions, as well as how molecules will bond and react chemically.