Practice Questions

Justify why the pH of a solution of sodium acetate ($\text{CH}_3\text{COONa}$) is greater than 7.

What is a buffer solution?

For the reaction $2 N_2O_5(g) \rightleftharpoons 4 NO_2(g) + O_2(g)$, the equilibrium constant $K_c$ is $1.5 \times 10^{-5}$ at a certain temperature. Calculate the equilibrium constant $K_c'$ for the reverse reaction, $4 NO_2(g) + O_2(g) \rightleftharpoons 2 N_2O_5(g)$, at the same temperature.

Examine the following Brønsted-Lowry acid-base reaction and identify the two conjugate acid-base pairs. $HSO_4^-(aq) + H_2O(l) \rightleftharpoons SO_4^{2-}(aq) + H_3O^+(aq)$

Explain three key characteristics of chemical equilibrium.

Write the expression for the equilibrium constant, $K_c$, for the following heterogeneous equilibrium: $CaCO_3(s) + 2 H^+(aq) \rightleftharpoons Ca^{2+}(aq) + H_2O(l) + CO_2(g)$

State the law of mass action.

Define dynamic equilibrium.

Define the ionic product of water ($K_w$).

Propose an experimental design to distinguish between a strong acid (like $0.1 \text{ M HCl}$) and a weak acid (like $0.1 \text{ M CH}_3\text{COOH}$) using electrical conductivity measurements. Justify your reasoning.

Justify the classification of the ammonia molecule ($\text{NH}_3$) and the boron trifluoride molecule ($\text{BF}_3$) as a Lewis base and a Lewis acid, respectively. Create a diagram showing the formation of the adduct between them.

Describe the Arrhenius concept of acids and bases with one example for each.

Critique the statement: "Adding a catalyst to a reaction at equilibrium increases the amount of product formed."

Write the expression for the equilibrium constant, $K_c$, for the following reactions: (i) $4\text{NH}_3(\text{g}) + 5\text{O}_2(\text{g}) \rightleftharpoons 4\text{NO}(\text{g}) + 6\text{H}_2\text{O}(\text{g})$ (ii) $\text{Ag}_2\text{O}(\text{s}) + 2\text{HNO}_3(\text{aq}) \rightleftharpoons 2\text{AgNO}_3(\text{aq}) + \text{H}_2\text{O}(\text{l})$ (iii) $\text{CaCO}_3(\text{s}) \rightleftharpoons \text{CaO}(\text{s}) + \text{CO}_2(\text{g})$

Define pH. What is the pH of a neutral solution at 298 K? Explain the relationship between pH and pOH.

Describe the Brönsted-Lowry concept of acids and bases. For each species below, identify its conjugate base. For species that can also act as a base, identify their conjugate acid. (i) HF (ii) H₂SO₄ (iii) NH₃ (iv) H₂O (v) HCO₃⁻

For the reaction $H_2(g) + I_2(g) \rightleftharpoons 2HI(g)$, the equilibrium constant $K_c$ is 54.8 at 700 K. If 1.00 mol of $H_2$ and 1.00 mol of $I_2$ are placed in a 2.00 L vessel and allowed to reach equilibrium, calculate the equilibrium concentrations of $H_2$, $I_2$, and HI.

State Le Chatelier's principle. Explain the effect of adding an inert gas at constant volume on the equilibrium of a gaseous reaction.

Propose a reason why it is more difficult to remove the second proton from sulphuric acid ($\text{H}_2\text{SO}_4$) compared to the first.

Calculate the pOH of a solution at 298 K that has a hydrogen ion concentration of $2.5 \times 10^{-4}$ M.

What is the common ion effect? Explain with the example of the ionization of acetic acid.

Analyze the species $Ag^+$ and determine if it acts as a Lewis acid or a Lewis base. Justify your answer.

The solubility product constant ($K_{sp}$) for lead(II) chloride ($PbCl_2$) is $1.6 \times 10^{-5}$ at 298 K. Calculate the molar solubility of $PbCl_2$ in pure water.

Analyze the effect on the equilibrium position of the reaction $CO(g) + 2H_2(g) \rightleftharpoons CH_3OH(g)$ when the pressure is increased by decreasing the volume at a constant temperature. Explain your reasoning using Le Chatelier's principle.

For the reaction $N_2(g) + O_2(g) \rightleftharpoons 2 NO(g)$, the value of $K_c$ is $4.1 \times 10^{-4}$ at 2000 K. If the equilibrium concentrations of $N_2$ and $O_2$ are $0.50$ M and $0.25$ M respectively, calculate the equilibrium concentration of NO.

For the synthesis of ammonia, $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, the equilibrium constant $K_c$ is 0.061 at 500 K. A reaction mixture at 500 K is found to contain $[N_2] = 1.0$ M, $[H_2] = 2.0$ M, and $[NH_3] = 0.5$ M. Analyze the state of the reaction and predict the direction in which the net reaction will proceed.

A student prepares a solution by mixing equal volumes of $0.2 \text{ M}$ acetic acid ($\text{CH}_3\text{COOH}$) and $0.2 \text{ M}$ sodium acetate ($\text{CH}_3\text{COONa}$). The student predicts the pH will be exactly equal to the $\text{p}K_a$ of acetic acid. Critique this prediction and calculate the actual pH of the resulting buffer solution. (Given: $\text{p}K_a$ of $\text{CH}_3\text{COOH} = 4.76$)

A buffer solution is prepared by mixing 100 mL of 0.20 M acetic acid ($CH_3COOH$) with 100 mL of 0.30 M sodium acetate ($CH_3COONa$). Calculate the pH of this buffer solution. The $pKa$ of acetic acid is 4.76.

Identify the conjugate acid-base pairs in the following reaction: $\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$

Explain the three classifications of reversible reactions based on the extent to which they proceed. Provide a general description for each category in terms of the value of the equilibrium constant, $K_c$.

The equilibrium constant $K_c$ for the reaction $2 SO_2(g) + O_2(g) \rightleftharpoons 2 SO_3(g)$ is 280 at 1000 K. Calculate the value of $K_p$ for this reaction at the same temperature. (Use $R = 0.0831$ L bar mol$^{-1}$ K$^{-1}$)

For the synthesis of methanol, $\text{CO}(\text{g}) + 2\text{H}_2(\text{g}) \rightleftharpoons \text{CH}_3\text{OH}(\text{g})$, the equilibrium constant $K_c$ is $10.5$ at $500 \text{ K}$. A reaction mixture at $500 \text{ K}$ in a $2.0 \text{ L}$ flask contains $0.20 \text{ mol}$ of $\text{CO}$, $0.30 \text{ mol}$ of $\text{H}_2$, and $0.05 \text{ mol}$ of $\text{CH}_3\text{OH}$. Evaluate whether the system is at equilibrium. If not, formulate the direction in which the net reaction will proceed.

Formulate the relationship between $K_p$ and $K_c$. Using this relationship, evaluate the condition under which $K_p = K_c$ for a gaseous reaction. Provide a balanced chemical equation for a reaction that satisfies this condition.

The value of $\Delta G^\ominus$ for the reaction $\text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g})$ is $-33.2 \text{ kJ mol}^{-1}$ at $298 \text{ K}$. Evaluate the spontaneity of the reaction under standard conditions. Then, formulate the expression relating $\Delta G^\ominus$ to the equilibrium constant $K_p$ and calculate the value of $K_p$. (Given: $R = 8.314 \text{ J K}^{-1} \text{mol}^{-1}$).

A researcher is studying the equilibrium $\text{N}_2\text{O}_4(\text{g}) \rightleftharpoons 2\text{NO}_2(\text{g})$. Initially, $0.50 \text{ mol}$ of $\text{N}_2\text{O}_4$ is placed in a $2.0 \text{ L}$ sealed vessel at $350 \text{ K}$. At equilibrium, it is found that $20\%$ of the $\text{N}_2\text{O}_4$ has dissociated. Evaluate the system to determine the equilibrium concentrations of both gases, and then calculate the equilibrium constants $K_c$ and $K_p$. (Given: $R = 0.0831 \text{ L bar mol}^{-1} \text{K}^{-1}$).

The solubility product, $K_{sp}$, for lead(II) chloride ($\text{PbCl}_2$) is $1.6 \times 10^{-5}$. Evaluate what happens when $100 \text{ mL}$ of $0.05 \text{ M Pb(NO}_3)_2$ is mixed with $100 \text{ mL}$ of $0.05 \text{ M NaCl}$. Formulate a conclusion based on a calculation of the ionic product, $Q_{sp}$.

Calculate the pH of a 0.05 M solution of pyridine ($C_5H_5N$), a weak base. The base ionization constant ($K_b$) for pyridine is $1.77 \times 10^{-9}$.

You are tasked with determining which of two sparingly soluble salts, Calcium Carbonate ($\text{CaCO}_3$, $K_{sp} = 2.8 \times 10^{-9}$) or Calcium Fluoride ($\text{CaF}_2$, $K_{sp} = 5.3 \times 10^{-9}$), is more soluble in pure water. A colleague suggests that since the $K_{sp}$ of $\text{CaF}_2$ is larger, it must be more soluble. Critique this reasoning and determine the molar solubility of each salt to justify your conclusion.

Design a buffer solution with a pH of $9.00$ using ammonia ($\text{NH}_3$) and ammonium chloride ($\text{NH}_4\text{Cl}$). Justify the ratio of the concentrations of the base and its conjugate acid required. Propose a reason why this buffer would be ineffective at maintaining a pH of $5.00$. (Given: $\text{p}K_b$ for $\text{NH}_3 = 4.75$).

Evaluate whether a solution can be simultaneously saturated with AgCl ($K_{sp} = 1.8 \times 10^{-10}$) and AgBr ($K_{sp} = 5.0 \times 10^{-13}$). Justify your conclusion.

The solubility product ($K_{sp}$) of magnesium hydroxide, $Mg(OH)_2$, is $1.8 \times 10^{-11}$. Calculate its molar solubility in a solution that is buffered at a pH of 9.0.

Formulate a prediction about the direction of equilibrium shift for the reaction $\text{PCl}_5(\text{g}) \rightleftharpoons \text{PCl}_3(\text{g}) + \text{Cl}_2(\text{g})$ if an inert gas is added at constant pressure. Justify your answer.

Define solubility product constant ($K_{sp}$). Write the expression for the solubility product of Calcium Phosphate, $\text{Ca}_3(\text{PO}_4)_2$.

List five important features of the equilibrium constant. Explain how the equilibrium constant for a reverse reaction is related to the equilibrium constant for the forward reaction, using a general example.

Calculate the pH of a 0.10 M solution of sodium cyanide (NaCN). The acid dissociation constant ($K_a$) for hydrocyanic acid (HCN) is $4.9 \times 10^{-10}$.