Q: A 0.085 M aqueous solution of a monoprotic acid (HA) has a pH of 3.30. What is the concentration of $\mathrm{H_3O^+}$ in the solution?

$5.02 \times 10^{-4}\ \mathrm{M}$

Q: A $0.155\ \mathrm{M}$ solution of $\mathrm{H_2CO_3}$ has $K_{a,1} = 4.3 \times 10^{-7}$ and $K_{a,2} = 5.6 \times 10^{-11}$. What is the concentration of $\mathrm{HCO_3^-}$ at equilibrium after the first dissociation?

$2.58 \times 10^{-4}\ \mathrm{M}$

Q: A solution contains $0.022\ \mathrm{M}$ Fe$^{2+}$ and $0.014\ \mathrm{M}$ Mg$^{2+}$. $K_{sp}$ for FeCO$_3$ is $3.0 \times 10^{-11}$ and for MgCO$_3$ is $6.82 \times 10^{-6}$. What is the minimum concentration of CO$_3^{2-}$ needed to begin precipitation of FeCO$_3$?

$1.4 \times 10^{-9}\ \mathrm{M}$

Q: What is the maximum concentration of CO$_3^{2-}$ that can be added before MgCO$_3$ begins to precipitate in the above solution?

$4.9 \times 10^{-4}\ \mathrm{M}$

Q: After FeCO$_3$ begins to precipitate, what is the remaining concentration of Fe$^{2+}$ when MgCO$_3$ just starts to precipitate?

$6.3 \times 10^{-8}\ \mathrm{M}$

Q: Which equation is used to calculate the pH of a buffer solution containing a weak acid and its conjugate base?

$\mathrm{pH = pK_a + \log \left( \frac{[\text{base}]}{[\text{acid}]} \right)}$

Question 1

A 0.085 M aqueous solution of a monoprotic acid (HA) has a pH of 3.30. What is the concentration of $\mathrm{H_3O^+}$ in the solution?

Accepted Answer

$5.02 	imes 10^{-4}\ \mathrm{M}$

Question 2

Given the reaction $\mathrm{HA\ (aq) + H_2O\ (l) ightarrow A^-\ (aq) + H_3O^+\ (aq)}$, and the initial concentration of HA is $0.085\ \mathrm{M}$, with $[H_3O^+] = 5.02 	imes 10^{-4}\ \mathrm{M}$ at equilibrium, what is the value of the acid dissociation constant $K_a$?

Accepted Answer

$3.0 	imes 10^{-6}$

Question 3

A solution contains $0.150\ \mathrm{M}$ HF. Given $K_a = 3.5 	imes 10^{-4}$, what is the pH of the solution?

Accepted Answer

$2.14$

Question 4

Why can the contribution of $\mathrm{HCOOH}$ to $[H_3O^+]$ be neglected in a mixture of $0.100\ \mathrm{M}$ HF and $0.100\ \mathrm{M}$ HCOOH?

Accepted Answer

Because $K_a$ for HCOOH is much smaller than that for HF, so its ionization is negligible compared to HF.

Question 5

A $0.155\ \mathrm{M}$ solution of $\mathrm{H_2CO_3}$ has $K_{a,1} = 4.3 	imes 10^{-7}$ and $K_{a,2} = 5.6 	imes 10^{-11}$. What is the concentration of $\mathrm{HCO_3^-}$ at equilibrium after the first dissociation?

Accepted Answer

$2.58 	imes 10^{-4}\ \mathrm{M}$

Question 6

For the reaction $\mathrm{C_5H_5N\ (aq) + H_2O\ (l) ightarrow C_5H_5NH^+\ (aq) + OH^-\ (aq)}$ with $K_b = 1.7 	imes 10^{-9}$ and $[C_5H_5N] = 0.125\ \mathrm{M}$, what is the pH of the solution?

Accepted Answer

$9.16$

Question 7

A buffer solution is $0.125\ \mathrm{M}$ in HNO$_2$ ($K_a = 4.6 	imes 10^{-4}$) and $0.145\ \mathrm{M}$ in NaNO$_2$. What is the pH of the buffer?

Accepted Answer

$3.40$

Question 8

If $1.5\ \mathrm{g}$ of HCl is added to the buffer solution described above, what is the new pH?

Accepted Answer

$3.13$

Question 9

A solution contains $0.022\ \mathrm{M}$ Fe$^{2+}$ and $0.014\ \mathrm{M}$ Mg$^{2+}$. $K_{sp}$ for FeCO$_3$ is $3.0 	imes 10^{-11}$ and for MgCO$_3$ is $6.82 	imes 10^{-6}$. What is the minimum concentration of CO$_3^{2-}$ needed to begin precipitation of FeCO$_3$?

Accepted Answer

$1.4 	imes 10^{-9}\ \mathrm{M}$

Question 10

What is the maximum concentration of CO$_3^{2-}$ that can be added before MgCO$_3$ begins to precipitate in the above solution?

Accepted Answer

$4.9 	imes 10^{-4}\ \mathrm{M}$

Question 11

After FeCO$_3$ begins to precipitate, what is the remaining concentration of Fe$^{2+}$ when MgCO$_3$ just starts to precipitate?

Accepted Answer

$6.3 	imes 10^{-8}\ \mathrm{M}$

Question 12

A solution contains $0.20\ \mathrm{M}$ Zn$^{2+}$ and excess Ba(OH)$_2$. $K_{sp}$ for Zn(OH)$_2$ is $3.0 	imes 10^{-15}$ and $K_f$ for Zn(OH)$_4^{2-}$ is $2.0 	imes 10^{5}$. What is the pH of the solution after Zn(OH)$_2$ dissolves as Zn(OH)$_4^{2-}$?

Accepted Answer

$12.97$

Question 13

Which of the following best describes the effect of adding a strong acid to a buffer solution?

Accepted Answer

The pH decreases slightly, but the buffer resists large changes in pH.

Question 14

In a titration of $0.125\ \mathrm{M}$ pyridine ($C_5H_5N$) with $0.100\ \mathrm{M}$ HCl, what is the volume of HCl required to reach the equivalence point if the initial volume of pyridine is $25.0\ \mathrm{mL}$?

Accepted Answer

$31.25\ \mathrm{mL}$

Question 15

Which equation is used to calculate the pH of a buffer solution containing a weak acid and its conjugate base?

Accepted Answer

$\mathrm{pH = pK_a + \log \left( \frac{[	ext{base}]}{[	ext{acid}]} ight)}$

Question 16

If $K_{sp}$ for FeCO$_3$ is $3.0 	imes 10^{-11}$, what is the equilibrium expression for the solubility product?

Accepted Answer

$K_{sp} = [Fe^{2+}][CO_3^{2-}]$

Question 17

A buffer is prepared by mixing $0.125\ \mathrm{M}$ HNO$_2$ and $0.145\ \mathrm{M}$ NaNO$_2$. If $1.5\ \mathrm{g}$ NaOH is added, what happens to the ratio $\frac{[	ext{base}]}{[	ext{acid}]}$?

Accepted Answer

The ratio increases because NaOH converts acid to base.

Question 18

Which of the following statements about the solubility of Zn(OH)$_2$ in the presence of excess OH$^-$ is correct?

Accepted Answer

The solubility increases due to formation of the complex ion Zn(OH)$_4^{2-}$.

General Chemistry II: Acid-Base Equilibria, Buffers, and Solubility Study Guide

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