Chemistry 112A First Exam Key

Consider the following balanced chemical reaction: $2 ClO (g) \rightleftharpoons Cl (g) + ClO_2 (g))$. The equilibrium constant, K, for this reaction at 385 K is $4.26 \times 10^{-4}$. If 1.240 atm of $ClO (g)$ and 0.330 atm of $Cl (g)$ are sealed in a container at 385 K, and no $ClO_2 (g)$ is present initially, what is the partial pressure of $ClO (g)$ (in atm) after equilibrium is established?

Consider the following balanced chemical reaction: $BrO (g) + ClO (g) \rightleftharpoons ClO_2 (g) + Br (g))$. The equilibrium constant, K, for this reaction at 385 K is 0.040 and the reaction is endothermic, as written. Which of the following statements is TRUE regarding Plot A?

Consider the generic, balanced chemical reaction: $A(g) + 2 B(g) \rightleftharpoons 2 C(g)$ which is endothermic. If the reaction is allowed to reach equilibrium and is then heated from the original temperature, $T_1$, to a new, higher temperature, $T_2$, in which direction does the reaction spontaneously shift to re-establish equilibrium?

Consider the generic, balanced chemical reaction: $A(g) + 2 B(g) \rightleftharpoons C(g) + D(g)$ which is exothermic. If the reaction is allowed to reach equilibrium and is then heated from the original temperature, $T_1$, to a new, higher temperature, $T_2$, in which direction does the reaction spontaneously shift to re-establish equilibrium?

Consider the generic, balanced chemical reaction: $A(g) + B(g) \rightleftharpoons C(s) + 2 D(g)$ which is endothermic (Note that product C is a pure solid.). If the reaction is allowed to reach equilibrium and is then heated from the original temperature, $T_1$, to a new, higher temperature, $T_2$, in which direction does the reaction spontaneously shift to re-establish equilibrium?

Consider the balanced gas phase reaction shown below: $4 NO (g) + 6 H_2O (g) \rightleftharpoons 4 NH_3 (g) + 5 O_2 (g)$. K = $8.32 \times 10^{-19}$ at temperature T. If 1.25 atm of $NH_3(g)$ and 1.25 atm of $O_2(g)$ are placed in a sealed container at temperature T, and no NO (g) or $H_2O (g)$ are initially present, what is the partial pressure of $NH_3 (g)$ (in atm) after equilibrium has been established?

At high temperatures, the solid compound $(NH_4)_2CO_3$ decomposes according to the following balanced chemical reaction: $(NH_4)_2CO_3 (s) \rightleftharpoons 2 NH_3 (g) + CO_2 (g) + H_2O (g)$. When excess $(NH_4)_2CO_3 (s)$ is sealed in an evacuated container at 750 K, the total pressure is 3.52 atm after equilibrium is established. What is the equilibrium constant, K, for this reaction at 750 K?

At high temperatures, the solid compound $(NH_4)_2CO_3$ decomposes according to the following balanced chemical reaction: $(NH_4)_2CO_3 (s) \rightleftharpoons 2 NH_3 (g) + CO_2 (g) + H_2O (g)$. If the equilibrium constant K at 775 K is equal to 5.47. Explain why when 1.40 atm of each gas are initially sealed in a container at 775 K, no solid is observed to form. Explain your reasoning in 1-2 sentences

At high temperatures, the solid compound $(NH_4)_2CO_3$ decomposes according to the following balanced chemical reaction: $(NH_4)_2CO_3 (s) \rightleftharpoons 2 NH_3 (g) + CO_2 (g) + H_2O (g)$. If the equilibrium constant K at 775 K is equal to 5.47. Determine the minimum amount of $NH_3 (g)$ that should be added to the container described in part (b, i) above in order for solid to form at 775 K. Report your answer in terms of the minimum amount of pressure (in atm) that should be added to the 1.40 atm of $NH_3 (g) already present.

A 0.40 M aqueous solution of HF is prepared at 25°C. Given: $K_a (HF, 25°C) = 6.6 \times 10^{-4}$. What is the percent dissociation of HF after equilibrium is established?

Using your work from part (a), determine the number of moles of $HClO_4$ that should be added to 600.0 mL of pure water at 25°C so that this solution would have the same final pH as the 0.40 M aqueous solution of HF described above.

Comparing a 600.0 mL sample of the 0.40 M aqueous solution of HF and the 600.0 mL sample of the $HClO_4$ solution, is this statement TRUE or FALSE: After equilibrium is established for each solution, the $[H_3O^+]$ is the same for both solutions.

Comparing a 600.0 mL sample of the 0.40 M aqueous solution of HF and the 600.0 mL sample of the $HClO_4$ solution, is this statement TRUE or FALSE: The same number of moles of $OH^−$ would be required to completely neutralize each solution.

Comparing a 600.0 mL sample of the 0.40 M aqueous solution of HF and the 600.0 mL sample of the $HClO_4$ solution, is this statement TRUE or FALSE: Both solutions would have a pH of 7.00 after completely neutralizing the acids with OH.

At 25°C, $K_a$ for the weak acid, hypochlorous acid, HClO, is $3.0 \times 10^{-8}$ ($pK_a$=7.53). You may assume a temperature of 25°C throughout this problem. Determine the mass (in g) of solid potassium hypochlorite, $K^+ClO^−$, that must be added to a 500 mL aqueous solution of 0.10 M HClO to prepare a buffer solution with pH = 8.00.

Write the balanced chemical equation that represents the governing equilibrium in this buffer: $K^+ClO^−$ + HClO

50 mL of a 0.20 M aqueous HCl solution is then added to the buffer described in part (a). Determine the pH of the buffer after this addition of HCl.

Using your work from part (c), calculate the number of moles of KClO that must be added to the buffer solution in part (c) to bring it back to the original pH of 8.00.

Volume of 0.50 M NaOH Added (mL). Determine the number of moles of HA in the original 200.0 mL solution (i.e., before any NaOH solution was added).

Calculate the $K_b$ for the conjugate weak base of the weak acid HA.

Which of the following is the monoprotic weak acid, HA, in the titration curve shown above?

When 500.0 mL of a 0.50 M aqueous solution of $MgCl_2$ is added to 250.0 mL of a 0.50 M aqueous solution of NaF at 25° C, a precipitate of the sparingly-soluble salt, $MgF_2$, is formed. What is the mass of $MgF_2 (in grams) that precipitates immediately upon mixing the two solutions, assuming complete reaction?

In order to establish equilibrium, a small amount of the $MgF_2$ precipitate determined in part (a) re-dissolves back into solution. What is the mass of $MgF_2 (in grams) that re-dissolves?

Equilibrium Constant (K)

Dynamic Equilibrium

Partial Pressure

Method of Successive Approximations (MOSA)

Assumption in Equilibrium Calc

Reversible Reaction

Equilibrium State

Initial Pressure

Change in Pressure

Equilibrium Pressures

Simplifying the Quadratic

K Expression Setup

Solving for 'x'

Verify the approximation

Equilibrium Pressures Calculation

Reaction Quotient (Q)

Temp and Equilibrium

Inert gas and equilibrium

Catalyst and Equilibrium

Le Chatelier's Principle

Pressure and equilibrium

Reactants and equilibrium

Products and equilibrium

Chemical Kinetics

Reaction Rate

Catalyst

Homogeneous Catalyst

Heterogeneous Catalyst

Activation Energy (Ea)

Reaction Intermediate