Question

Among the following molecules H-bond is present in (a) \(\mathrm{NH}_{3}\) (b) \(\mathrm{PH}_{3}\) (c) \(\mathrm{H}_{2} \mathrm{~S}\) (d) \(\mathrm{CH}_{4}\)

Short answer

(a) \(\mathrm{NH}_{3}\) (b) \(\mathrm{PH}_{3}\) (c) \(\mathrm{H}_{2} \mathrm{~S}\) (d) \(\mathrm{CH}_{4}\) Answer: (a) \(\mathrm{NH}_{3}\)

Step-by-step solution

Understanding Hydrogen Bonding

Hydrogen bonding is a type of attractive interaction between a hydrogen atom bonded to a highly electronegative atom (such as nitrogen, oxygen, or fluorine) and another electronegative atom (N, O, or F). It is a relatively strong type of intermolecular force, responsible for many unique properties of substances, such as the high boiling point of water.

Analyzing \(\mathrm{NH}_{3}\)

In ammonia(\(\mathrm{NH}_{3}\)), nitrogen is the central atom, which is highly electronegative. It has three hydrogen atoms covalently bonded to it. The electronegativity difference between nitrogen and hydrogen creates polar covalent bonds, and the presence of a lone pair on nitrogen allows for hydrogen bonding with other molecules. So, hydrogen bonding is present in \(\mathrm{NH}_{3}\).

Analyzing \(\mathrm{PH}_{3}\)

Phosphine(\(\mathrm{PH}_{3}\)) has a phosphorous central atom bonded to three hydrogen atoms. The electronegativity difference between phosphorous and hydrogen is lesser than that of nitrogen and hydrogen; hence, the molecule is less polar. Moreover, phosphorus is not one of the highly electronegative atoms involved in hydrogen bonding (N, O, or F). Therefore, no hydrogen bonding is present in \(\mathrm{PH}_{3}\).

Analyzing \(\mathrm{H}_{2} \mathrm{~S}\)

Hydrogen sulfide(\(\mathrm{H}_{2}\mathrm{S}\)) consists of a sulfur atom bonded to two hydrogen atoms. Although the electronegativity difference between sulfur and hydrogen is significant, sulfur is not as electronegative as nitrogen, oxygen, or fluorine. So, \(\mathrm{H}_{2}\mathrm{S}\) does not have hydrogen bonding.

Analyzing \(\mathrm{CH}_{4}\)

Methane(\(\mathrm{CH}_{4}\)) contains a carbon atom bonded to four hydrogen atoms. Carbon does not belong to the highly electronegative atoms group (N, O, or F), so no hydrogen bonding is present in \(\mathrm{CH}_{4}\).

Conclusion

Among the given molecules, hydrogen bonding is present only in \(\boldsymbol{\mathrm{NH}_{3}}\). So, the correct option is (a) \(\mathrm{NH}_{3}\).

Key concepts

Electronegativity

Electronegativity can be thought of as the measure of an atom's ability to attract and hold onto electrons within a chemical bond. Atoms with higher electronegativity will tend to pull electrons closer towards themselves when bonded to atoms with lower electronegativity, creating an uneven distribution of electric charge.

For instance, in a water (H₂O) molecule, oxygen has a higher electronegativity compared to hydrogen. This causes the shared electrons to spend more time closer to the oxygen atom, resulting in a partial negative charge on the oxygen and a partial positive charge on each hydrogen atom. This electronegativity difference is foundational for the formation of polar covalent bonds, which are key in creating the conditions for hydrogen bonding.

Understanding the trend of electronegativity across periodic table is essential. Generally, electronegativity increases as you move from left to right across a period and decreases as you go down a group. Elements with high electronegativity such as flourine, oxygen, and nitrogen, are often involved in hydrogen bonding due to their strong attraction for electrons.

Intermolecular Forces

Intermolecular forces (IMFs) are forces of attraction or repulsion which act between neighboring particles (atoms, molecules, or ions). They are weaker than the strong covalent bonds within molecules, but are critical for dictating the physical properties of substances, such as boiling points, melting points, and solubilities.

There are several types of IMFs, including London dispersion forces, dipole-dipole interactions, and the strongest among them, hydrogen bonds. London dispersion forces occur in all molecular substances, whereas dipole-dipole interactions occur in polar molecules. Hydrogen bonds, a special case of IMFs, require a hydrogen atom to be covalently bonded to a highly electronegative atom (such as O, N, or F).

The unique properties of water, such as its high surface tension, can be attributed to the presence of hydrogen bonding. Ice floats in liquid water because hydrogen bonds hold water molecules in an open hexagonal structure, which is less dense than the liquid form. Understanding these interactions is crucial for predicting the behavior of molecules in both natural processes and industrial applications.

Polar Covalent Bonds

Polar covalent bonds are the result of two atoms with different electronegativities bonding together and unequally sharing electrons. One atom becomes partially negative (the more electronegative atom) because it pulls the electrons in the bond closer to itself, while the other atom becomes partially positive.

Consider hydrochloric acid (HCl) as an example. The chlorine atom is more electronegative than the hydrogen atom, which leads to a polar covalent bond between them. The uneven charge distribution enables these molecules to attract each other through dipole-dipole interactions, another common type of IMF.

Relevance to Hydrogen Bonding

For a molecule to form hydrogen bonds, it must feature polar covalent bonds between hydrogen and a highly electronegative atom. These polar covalent bonds are the prerequisite for the formation of hydrogen bonds; without them, the necessary unequal charge distribution that allows a hydrogen atom to be attracted to another electronegative atom in a neighboring molecule would be impossible. This condition explains why molecules like methane (CH₄) do not participate in hydrogen bonding, as the carbon-hydrogen bond is relatively nonpolar.