Question

Nitrogen, \(\mathrm{N}_{2}\), combines with hydrogen, \(\mathrm{H}_{2}\), to give ammonia, \(\mathrm{NH}_{3}\). Nitrogen and hydrogen are odorless gases that are not very soluble in water. Ammonia is also a gas, but it has an extremely powerful odor and is very soluble in water. How can you explain this great difference in properties?

Short answer

The great difference in properties between Nitrogen (N2), Hydrogen (H2), and Ammonia (NH3) can be explained by their molecular structures, bond types, and intermolecular forces. N2 and H2 are nonpolar molecules with weak van der Waals forces, leading to their odorlessness and low solubility in water. In contrast, NH3 has polar covalent bonds and strong intermolecular forces (hydrogen bonding) due to the electronegativity difference between nitrogen and hydrogen atoms, resulting in high solubility in water and a pungent odor.

Step-by-step solution

Understanding the molecular structure of Nitrogen (N2), Hydrogen (H2), and Ammonia (NH3)

Nitrogen gas (N2) contains two nitrogen atoms bonded with a triple bond. Hydrogen gas (H2) has two hydrogen atoms bonded with a single bond. Ammonia (NH3) is composed of one nitrogen atom bonded with three hydrogen atoms. In ammonia, nitrogen contributes with its unshared electron pair (lone pair).

Comparing the bond types and polarity in N2, H2, and NH3

In Nitrogen (N2) and Hydrogen (H2), the atoms are bonded through covalent bonds. Due to the identical electronegativities of the atoms in N2 and H2, these bonds are nonpolar. In contrast, Ammonia (NH3) possesses polar covalent bonds between the nitrogen and hydrogen atoms. Nitrogen atom, being more electronegative than hydrogen, attracts the shared electrons towards itself, resulting in a partial negative charge on the nitrogen atom and partial positive charges on the hydrogen atoms.

Examining the intermolecular forces

In N2 and H2, as the bonds are nonpolar, the intermolecular forces between their molecules are weak van der Waals forces (London dispersion forces). On the other hand, Ammonia (NH3) has hydrogen bonding, a strong intermolecular force, due to the presence of polar bonds between nitrogen and hydrogen atoms.

Relating the intermolecular forces to solubility and odor

In general, polar molecules with strong intermolecular forces like hydrogen bonding are more soluble in polar solvents such as water due to the ability to form strong interactions with the solvent molecules. In this case, Ammonia (NH3), being a polar molecule with hydrogen bonding, dissolves easily in water. In contrast, Nitrogen (N2) and Hydrogen (H2) are nonpolar with weak van der Waals forces and do not dissolve well in water. As for the odor, the strong polarity and intermolecular forces in ammonia lead to a higher interaction with olfactory receptors, resulting in its pungent smell. Nitrogen and hydrogen, being nonpolar and having weak intermolecular forces, exhibit no odor. In conclusion, the differences in properties of Nitrogen (N2), Hydrogen (H2), and Ammonia (NH3) can be primarily attributed to their molecular structures, bond types, and intermolecular forces.

Key concepts

Chemical Polarity

Chemical polarity is a fundamental concept in chemistry that describes how unequal sharing of electrons in a molecule leads to a distribution of electric charge across the molecule. This imbalance results in a molecule having a positive end and a negative end, much like a magnet with its north and south poles.

For example, in a water molecule (H_2O), oxygen is more electronegative than hydrogen, which means oxygen draws electrons toward itself more strongly. Consequently, the oxygen end of the molecule carries a slight negative charge, whereas the hydrogen end carries a slight positive charge, making water a polar molecule.

In the ammonia molecule (NH_3), the nitrogen atom is more electronegative than the hydrogen atoms. This creates a scenario where the shared electrons are closer to the nitrogen atom, giving it a partial negative charge and the hydrogen atoms a partial positive charge. This polarity is crucial because it gives rise to hydrogen bonding, which significantly affects properties like solubility and reactivity.

Hydrogen Bonding

Hydrogen bonding is a strong type of intermolecular force that occurs in polar molecules where hydrogen is directly bonded to a highly electronegative atom such as nitrogen, oxygen, or fluorine. Interestingly, despite being called a 'bond', it's not a chemical bond in the traditional sense but rather a strong attraction between a hydrogen atom carrying a partial positive charge and another electronegative atom carrying a partial negative charge in a different molecule.

In an NH_3 molecule, the nitrogen atom has a lone electron pair that forms a relatively strong hydrogen bond with the hydrogen atom of another ammonia molecule, or with water when dissolved. These bonds are instrumental in giving ammonia its high solubility in water and its sharp odor because they enhance the molecule's interaction with our olfactory receptors, unlike nonpolar molecules like N_2 and H_2, which lack this bonding and are less interactive.

Solubility in Water

Solubility in water is heavily influenced by the intermolecular forces present between solute and solvent molecules. Polar substances tend to dissolve well in polar solvents like water because the positive and negative ends of the molecules can attract and stabilize each other, leading to what's known as 'like dissolves like'. In contrast, nonpolar substances do not mix well with polar solvents because there is a lack of attraction between the molecules.

Ammonia's solubility in water is a classic illustration of this principle. The polar NH_3 molecules readily form hydrogen bonds with the polar water molecules, allowing it to dissolve. However, the nonpolar N_2 and H_2 molecules have little affinity for water and thus do not dissolve to any significant extent. This aspect is crucial for understanding how solutes interact with solvents and is a key factor in predicting the behavior of substances when mixed with water.