Question

Explain the following observations: (a) ${\mathrm{NH}}_3$ contains no ${\mathrm{OH}}^{-}$ ions, and yet its aqueous solutions are basic; (b) HF is called a weak acid, and yet it is very reactive; (c) although sulfuric acid is a strong electrolyte, an aqueous solution of ${\mathrm{H}}_2{\mathrm{SO}}_4$ contains more ${\mathrm{HSO}}_4^{-}$ ions than ${\mathrm{SO}}_4^{2-}$ ions.

Short answer

In summary: (a) Aqueous NH3 solutions are basic because NH3 acts as a Bronsted-Lowry base, accepting a proton from water molecules and producing OH- ions in the process. (b) HF is a weak acid due to its incomplete dissociation, but its reactivity is derived from the strong hydrogen-fluorine bond, enabling it to form stable compounds with other substances. (c) H2SO4 is a strong electrolyte, but its aqueous solutions contain more HSO4- ions than SO42- ions because its second dissociation step (from HSO4- to SO42-) is weak, favoring the formation of HSO4- ions.

Step-by-step solution

Explanation A: NH3 acts as a base without containing OH- ions

According to the Bronsted-Lowry theory of acids and bases, a substance can act as either an acid or a base by donating or accepting a proton (H+ ion), respectively. When ammonia (NH3) dissolves in water, it accepts a proton from water molecules, forming ammonium ions (NH4+) and hydroxide ions (OH-), as shown in the following equation: NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH– (aq) Thus, although NH3 molecules do not contain OH- ions, their interaction with water results in the production of OH- ions, making the solution basic.

Explanation B: HF being a weak acid and yet reactive

An acid is considered weak if it does not fully dissociate into its ions when dissolved in water. HF is a weak acid as it does not dissociate completely into H+ and F- ions: HF (aq) ⇌ H+ (aq) + F– (aq) However, HF is still very reactive because of the high strength of the hydrogen-fluorine bond. The high bond strength makes it difficult for HF to dissociate but, when it reacts with other substances, it can form stable compounds, such as metal fluorides and silicon tetrafluoride (SiF4), releasing a large amount of energy in the process.

Explanation C: H2SO4 being a strong electrolyte and aqueous solutions containing more HSO4- ions than SO42- ions

H2SO4 is a strong electrolyte as it dissociates completely into its ions when dissolved in water: H2SO4 (aq) → 2H+ (aq) + SO42- (aq) However, H2SO4 undergoes two dissociation steps: The first step is the dissociation of one H+ ion, forming HSO4- ions, which is a strong-acid dissociation. The second step is the dissociation of the remaining H+ ion from the HSO4- ion to form SO42- ions, which is a weak-acid dissociation: 1. H2SO4 (aq) → H+ (aq) + HSO4- (aq) (strong dissociation) 2. HSO4- (aq) ⇌ H+ (aq) + SO42- (aq) (weak dissociation) Since the second dissociation step is weak, the equilibrium favors the formation of HSO4- ions over SO42-. This means that the concentration of HSO4- ions will be higher than that of SO42- ions in an aqueous solution of H2SO4.

Key concepts

Bronsted-Lowry Theory

The Bronsted-Lowry theory fundamentally changed our understanding of acid-base reactions. It defines an acid as a proton (H+) donor and a base as a proton acceptor. This broader definition includes reactions where no hydrogen ions are apparent, such as when ammonia (NH₃) acts as a base in water. In the classroom problem, we observe that ammonia can make the solution basic, despite the absence of hydroxide ions in its molecular form. Ammonia readily accepts a proton from the water, leading to the formation of hydroxide ions (OH-), which are responsible for the basic nature of the solution.

The theory places emphasis not just on the presence of certain ions, but on their movement - the transfer of protons between molecules. Therefore, a substance can be an acid or a base depending on the reaction context, which opens up students to a more dynamic view of chemical interactions.

Weak Acids and Dissociation

Weak acids such as hydrofluoric acid (HF) can sometimes be confusing. Although they are termed 'weak' in the context of acid-base chemistry, this does not necessarily correlate with their reactivity in other types of chemical interactions. A weak acid only partially dissociates into its ions in an aqueous solution. This partial dissociation means that in a solution, there is an equilibrium between the undissociated molecules and the ions produced.

So, when we say HF is a weak acid, we are describing its behavior in water, where it exists considerably in its molecular form alongside fewer hydrated protons (H+) and fluoride ions (F-). The strength of the acid, in this case, is less about its reactive potential and more about its tendency not to release protons into the solution. However, what makes HF notably reactive is the energy released when it forms stable compounds, reflecting its chemical potential outside the scope of acid-dissociation.

Strong Electrolytes and Dissociation

Electrolytes are substances that dissolve in water to form free ions, thus becoming capable of conducting electricity. A strong electrolyte like sulfuric acid (H₂SO₄) dissociates completely in aqueous solution—except the dissociation occurs in multiple steps. In our textbook exercise, sulfuric acid's behavior in water is understood as a two-step process.

Initially, H₂SO₄ loses one proton, creating the bisulfate ion (HSO₄^-), a reaction which is virtually complete, thus making it a strong acid in this first dissociative step. The second proton removal, converting HSO₄^- to the sulfate ion (SO₄^2-), is less favorable, creating an equilibrium where the bisulfate ion is predominant in solution. This concept highlights not only the complexity of dissociation but also the graded nature of acid strength, even within a single compound like H₂SO₄.