Identify the kinds of intermolecular forces that might arise between molecules of each of the following substances: (a) \(\mathrm{NH}_{2} \mathrm{OH}\); (b) \(\mathrm{CBr}_{4}\); (c) \(\mathrm{H}_{2} \mathrm{SeO}_{4}\); (d) \(\mathrm{SO}_{2}\).

To fully understand why substances like NH₂OH and H₂SeO₄ exhibit hydrogen bonding, it's essential to dive deeper into the concept. Hydrogen bonding is a specific type of intermolecular force that occurs when a hydrogen atom, bonded to a highly electronegative element such as nitrogen, oxygen, or fluorine, is attracted to another electronegative atom in a nearby molecule. The result is a strong interaction relative to other types of intermolecular forces.

Hydrogen bonds are responsible for many of the unique properties of water, such as its high boiling point and surface tension. In our specific cases, NH₂OH exhibits this behavior due to nitrogen and oxygen having lone pairs of electrons that offer a site for hydrogen atoms of adjacent molecules to bond. H₂SeO₄, despite selenium being less electronegative than oxygen, can still participate in hydrogen bonding due to the presence of oxygen.

Dipole-dipole interactions are another significant type of force, arising between polar molecules where there is a partial positive charge on one end and a partial negative charge on the other. These are directional interactions, as positive ends of one molecule align with the negative ends of another. Molecules like SO₂ exemplify dipole-dipole interactions due to the substantial difference in electronegativity between sulfur and oxygen.

This polarity creates a dipole moment, where you can imagine a molecule having a 'north' and 'south' pole similar to a magnet. In a substance such as CBr₄, the symmetrical shape of the molecule makes the dipole moments cancel out, thus considerably weakening any dipole-dipole interactions.

The London dispersion forces are often thought of as the 'default' interaction because they can be found between all molecules, regardless of polarity. They are particularly significant in nonpolar compounds like CBr₄. The size and shape of molecules influence the magnitude of these forces; larger atoms or molecules with more electrons often have stronger London dispersion forces.

This happens because electrons in larger atoms can become more easily polarizable, creating temporary dipoles that induce corresponding dipoles in adjacent molecules, thus attracting them to each other. While these are the weakest of the intermolecular forces, they are omnipresent and can become quite strong in large, heavy atoms or molecules.

The term 'molecular polarity' refers to the distribution of electric charge over the atoms in a molecule. Compounds like SO₂ and H₂SeO₄ are polar due to an unequal sharing of electrons resulting from bonds between atoms with different electronegativities. This disparity leads to a net dipole moment where one end of the molecule is relatively positive, and the other is relatively negative.

Molecular polarity is crucial because it influences a substance's physical and chemical properties, including its solubility, melting point, and how it interacts with other molecules. It's also a determining factor in the kinds of intermolecular forces that can occur within a sample of the substance.

Electronegativity is a measure of how strongly an atom can attract and hold onto electrons in a chemical bond. Atoms like oxygen, nitrogen, and fluorine are among the most electronegative and can create polarized bonds leading to hydrogen bonding or dipole-dipole interactions. Understanding electronegativity helps us predict the kind of intermolecular forces that may arise in a substance.

For instance, in NH₂OH, the significant difference in electronegativity between hydrogen and the atoms it's bound to (nitrogen and oxygen) allows for robust hydrogen bonding. However, in symmetrical molecules like CBr₄, even with the difference in electronegativity between carbon and bromine, the symmetry results in a nonpolar molecule, minimizing the impact of electronegativity on intermolecular forces.