Acid-Base Properties of Salts

Two flasks containing aqueous solutions of the same concentration are available in the lab: one of sodium chloride, $\mathrm{NaCl}$, and the other of potassium nitrite, $\mathrm{KNO_2}$, at $25\,^{\circ}\mathrm{C}$.

1. a) Write the equations for the dissociation of both salts in water and identify the spectator ions that do not undergo hydrolysis.
2. b) Justify the acidic, basic, or neutral character of each of the two resulting solutions, given that $K_\mathrm{a}(\mathrm{HNO_2})=4.5\times 10^{-4}$ at $25\,^{\circ}\mathrm{C}$. Consider the solutions to be ideal.

Ammonium nitrate, $\mathrm{NH_4NO_3}$, is a compound widely used in fertilizer solutions at $25\,^{\circ}\mathrm{C}$.

1. a) Write the ionic dissociation and hydrolysis processes in water corresponding to this salt, indicating which species act as a Brønsted-Lowry acid or base.
2. b) Deduce, with justification, whether the pH of the resulting aqueous solution will be lower than, equal to or higher than $7$. Data at $25\,^{\circ}\mathrm{C}$: $K_\mathrm{b}(\mathrm{NH_3})=1.8\times 10^{-5}$. Consider the solutions to be ideal.

An aqueous solution is prepared by dissolving $4.10\text{ g}$ of sodium acetate, $\mathrm{CH_3COONa}$, in distilled water to a total volume of $500\text{ mL}$ at $25\,^{\circ}\mathrm{C}$.

1. a) Write the equation for the chemical hydrolysis equilibrium and determine the value of the hydrolysis constant $K_\mathrm{h}$.
2. b) Calculate the molar concentration of all the species present at equilibrium ($\mathrm{Na^+}$, $\mathrm{CH_3COO^-}$, $\mathrm{CH_3COOH}$, $\mathrm{OH^-}$ and $\mathrm{H_3O^+}$) and determine the final pH of the solution.

Data at $25\,^{\circ}\mathrm{C}$: Atomic molar masses in $\text{g}\times\text{mol}^{-1}$: $\mathrm{C}=12.0$; $\mathrm{H}=1.0$; $\mathrm{O}=16.0$; $\mathrm{Na}=23.0$. $K_\mathrm{a}(\mathrm{CH_3COOH})=1.8\times 10^{-5}$; $K_\mathrm{w}=1.0\times 10^{-14}$. Consider the solutions to be ideal.

An aqueous solution of ammonium fluoride, $\mathrm{NH_4F}$, is prepared at $25\,^{\circ}\mathrm{C}$. Consider the following equilibrium constant data at that temperature: $K_\mathrm{a}(\mathrm{HF})=6.6\times 10^{-4}$, $K_\mathrm{b}(\mathrm{NH_3})=1.8\times 10^{-5}$ and $K_\mathrm{w}=1.0\times 10^{-14}$.

1. a) Write the ionic reactions that take place when both ions from the salt interact with water.
2. b) Compare quantitatively the constants of the hydrolysis equilibria of the ions present and determine whether the resulting solution is acidic, basic or neutral. Consider the solutions to be ideal.

In the general chemistry lab, four beakers numbered 1 to 4 have been labeled, each containing a $0.10\text{ M}$ aqueous solution at $25\,^{\circ}\mathrm{C}$ of the following substances: $\mathrm{NaCl}$, $\mathrm{CH_3COONa}$, $\mathrm{NH_4Cl}$ and $\mathrm{KNO_3}$. The diagram of the experimental setup is shown below:

1. a) Classify each of the four salts, indicating whether the acid and the base from which it derives are strong or weak, and state which undergo hydrolysis.
2. b) Given that $K_\mathrm{a}(\mathrm{CH_3COOH})=1.8\times 10^{-5}$, $K_\mathrm{b}(\mathrm{NH_3})=1.8\times 10^{-5}$ and $K_\mathrm{w}=1.0\times 10^{-14}$ at $25\,^{\circ}\mathrm{C}$, arrange the four solutions in order of increasing pH, justifying your answer. Consider the solutions to be ideal.