Which of the following molecules are likely to form hydrogen bonds: (a) \(\mathrm{D}_{2} \mathrm{O}\); (b) \(\mathrm{CH}_{3} \mathrm{COOH}\); (c) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\); (d) \(\mathrm{H}_{3} \mathrm{PO}_{4}\) ?

Chemical bonding is the force that holds atoms together in molecules. Atoms bond to achieve greater stability, which often comes in the form of complete electron shells. When looking at the molecules capable of forming hydrogen bonds, we are dealing with covalent bonds, where two atoms share electrons. Covalent bonds occur between nonmetal atoms that have similar electronegativities. Not all covalent bonds are equal, however; some are polar, with a significant difference in electronegativity between the bonded atoms.
For hydrogen bonding to occur, a special case of polarity is necessary where hydrogen is covalently bonded to one of the highly electronegative elements such as nitrogen, oxygen, or fluorine. This creates a significant positive charge on the hydrogen, allowing it to 'reach out' and interact with a lone pair of electrons on another electronegative atom in a neighboring molecule. This interaction is the hydrogen bond which, while weaker than a covalent bond, is a strong type of intermolecular force.

Intermolecular forces are the forces of attraction or repulsion which act between neighboring particles (atoms, molecules, or ions). They are weaker than the intramolecular forces that keep a molecule intact but are crucial in predicting the physical properties of substances.
There are different types of intermolecular forces such as London dispersion forces, dipole-dipole interactions, and hydrogen bonds. Hydrogen bonds are particularly interesting due to their strong influence on properties such as boiling point, surface tension, and viscosity. Substances that can form hydrogen bonds, like water, typically have higher boiling points than those that do not. This is because it requires more energy to break the additional hydrogen bonds in the process of changing from liquid to gas phase.
When analyzing molecules for the potential to form hydrogen bonds, one must consider the presence of a highly electronegative atom with lone pairs (commonly N, O, or F) and a hydrogen atom bonded to one of these electronegative atoms.

Electronegativity is a measure of how strongly an atom can attract bonding electrons to itself. It's a relative scale, with fluorine being the most electronegative element, set at a value of around 4.0 on the Pauling scale. As a rule of thumb, the higher the electronegativity, the more an atom pulls electrons towards itself when forming a chemical bond.
In the context of hydrogen bonding, the difference in electronegativity between hydrogen and its covalently bonded partner (N, O, or F) creates a polar bond. This means that the electrons in the bond are not shared equally, leading to partial charges on the atoms. It is this polarity that allows hydrogen bonds to form; the positive end of one polar molecule is attracted to the negative end of another (the lone pairs on the electronegative atom), creating a strong intermolecular bond. Understanding electronegativity is crucial in predicting and explaining the behavior of molecules with respect to their interactions and the types of bonds they can form.

Molecular structure refers to the arrangement of atoms within a molecule, including the geometric shape, bond lengths, angles, and the three-dimensional configuration. This structure is determined by the principles of quantum mechanics, electronegativity, and the repulsion between electron pairs.
In hydrogen bonding, molecular structure plays a pivotal role. For hydrogen bonds to form, there must be an optimal orientation where the hydrogen atom—the positive dipole—is in the vicinity of a lone pair on an electronegative atom of another molecule. Therefore, the physical shape of a molecule can enhance or inhibit its ability to participate in hydrogen bonding. For instance, in water (H2O), the bent shape of the molecule is perfect for forming hydrogen bonds, which result in many of water's unique properties. Molecules with linear, angular, or compact shapes containing the right atoms can also engage in hydrogen bonding, effectively influencing the physical properties of the substance.