Which of the following moiecules are likely to form hydrogen bonds: (a) \(\mathrm{HF}\); (b) \(\mathrm{CH}_{4}\); (c) \(\mathrm{NH}_{3}\); (d) \(\mathrm{CH}_{3} \mathrm{OH}\) ?

Electronegativity is a chemical property describing an atom's ability to attract and hold onto electrons. This property plays a pivotal role in forming various types of chemical bonds and intermolecular forces.

Atoms with high electronegativity, such as nitrogen, oxygen, and fluorine, have a greater pull on the shared electrons in a bond. This tug-of-war can result in the electron density being more skewed towards these atoms, creating a polar effect within the molecule. As a result, one end of the molecule becomes slightly negative, while the other end becomes slightly positive, leading to what is known as a dipole.

Understanding the concept of electronegativity can shed light on the behavior of molecules in various environments. For instance, in the exercise provided, the electronegativity of fluorine, oxygen, and nitrogen explains why molecules like HF, NH3, and CH3OH can form hydrogen bonds – they contain hydrogen atoms that are bound to these highly electronegative atoms.

Dipole-dipole interactions are intermolecular forces that occur between molecules that possess a permanent dipole moment, due to differences in electronegativity between the atoms involved in a bond. When you have two polar molecules, the positive end of one molecule is attracted to the negative end of another molecule.

This attraction is what keeps the molecules in close proximity, affecting properties like the boiling point, melting point, and solubility of substances. The stronger the dipole-dipole interactions, typically, the higher the boiling and melting points. This is critical to the understanding of hydrogen bonding because it is a specific, unusually strong type of dipole-dipole interaction that occurs whenever hydrogen is bonded to atoms like nitrogen, oxygen, or fluorine, which are seen in the molecules of the exercise HF and NH3, and less directly with CH3OH.

The molecular structure dictates how atoms are arranged within a molecule and determines the molecule's geometry and physical properties. It is insightful to examine a molecule's shape to predict its ability to participate in hydrogen bonding and other intermolecular forces.

Linear, bent, tetrahedral, and trigonal planar are examples of molecular geometries that can influence a molecule's polarity and reactivity. A molecule like CH4, which has a tetrahedral structure and is symmetric, does not produce a permanent dipole and therefore cannot participate in hydrogen bonding. On the other hand, the unique structures of HF, NH3, and CH3OH, which include polar bonds with highly electronegative atoms, enable these molecules to engage in hydrogen bonding.

Intermolecular forces are the forces of attraction or repulsion between neighboring particles (atoms, molecules, or ions). They are weaker than chemical bonds but essential in explaining physical properties such as boiling and melting points, vapor pressure, and viscosity.

Examples of intermolecular forces include dipole-dipole interactions, London dispersion forces, and hydrogen bonding. Hydrogen bonding is particularly strong among these, and it occurs in molecules where a hydrogen atom is directly bonded to an electronegative element which also has lone pairs to facilitate the bond, as explained in the exercise. This particular bond is a key player in determining the properties of many compounds, such as water's high boiling point or the structure of DNA. Molecules like HF and NH3, which have hydrogen bonds, exhibit higher boiling points compared to similar sized molecules that do not have such bonds.