Weak Acid and Base Equilibria - AP Chemistry

1. What is the $pH$ of a neutral solution at $25^\text{o}$C?

   $pH = 7$. Equal concentrations of $H^+$ and $OH^-$ at equilibrium.

2. Identify the relationship between $K_w$, $K_a$, and $K_b$ for conjugate acid-base pairs.

   $K_w = K_a \times K_b$. For conjugate acid-base pairs, their equilibrium constants multiply to give $K_w$.

3. What is the expression for the equilibrium constant $K_b$ of a weak base?

   $K_b = \frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}$. Shows equilibrium between weak base and its protonation products.

4. What does a smaller $K_b$ value indicate about the strength of a base?

   The base is weaker. Lower $K_b$ means less dissociation and weaker base.

5. What does a larger $K_a$ value indicate about the strength of an acid?

   The acid is stronger. Higher $K_a$ means greater dissociation and stronger acid.

6. How do you find $[OH^-]$ from $pOH$?

   $[OH^-] = 10^{-pOH}$. Inverse logarithm converts pOH back to concentration.

7. What is the Bronsted-Lowry definition of a base?

   A base is a proton acceptor. Accepts protons from acids in chemical reactions.

8. What is the general expression for the equilibrium constant $K_a$ of a weak acid?

   $K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}$. Shows equilibrium between weak acid and its dissociation products.

9. What is the expression for the percent ionization of a weak acid?

   Percent ionization = \frac{[\text{H}^+]}{[\text{HA}]_0} \times 100\text{%}. Ratio of ionized acid to initial concentration times 100%.

10. What is the formula for the autoionization of water?

    $2 \text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{OH}^-$. Water molecules exchange protons to form ions.

11. Calculate the $pH$ for $[OH^-] = 4 \times 10^{-5}$ M.

    $pH = 9.6$. Find $pOH = 4.4$, then $pH = 14 - 4.4 = 9.6$.

12. Which is stronger: an acid with $K_a = 1.0 \times 10^{-4}$ or $K_a = 1.0 \times 10^{-6}$?

    The acid with $K_a = 1.0 \times 10^{-4}$ is stronger. Larger $K_a$ value indicates greater acid dissociation.

13. What is the $pH$ of a $0.01$ M sulfuric acid solution assuming complete dissociation?

    $pH = 2$. Diprotic acid: $[H^+] = 0.02$ M, so $pH = -\log(0.02) = 2$.

14. What is the relationship between $pH$ and $pOH$ at $25^\text{o}$C?

    $pH + pOH = 14$. Sum of pH and pOH equals 14 at standard temperature.

15. What is the value of $K_w$ at $25^\text{o}$C?

    $K_w = 1.0 \times 10^{-14}$. Ion product of water at standard temperature.

16. What is the effect of temperature on the value of $K_w$?

    $K_w$ increases with temperature. Higher temperature increases water dissociation.

17. What is the conjugate base of the bicarbonate ion, $HCO_3^{-}$?

    $CO_3^{2-}$. Loses another proton to form carbonate ion.

18. What is the conjugate acid of ammonia, NH₃?

    NH₄^+. Gains a proton to form the conjugate acid.

19. Find the $pK_a$ given $K_a = 1.8 \times 10^{-5}$.

    $pK_a = 4.74$. Use $pK_a = -\log(K_a) = -\log(1.8 \times 10^{-5}) = 4.74$.

20. Identify the Lewis definition of an acid.

    An acid is an electron pair acceptor. Accepts electron pairs from Lewis bases.

21. How does $pH$ change when a weak acid is diluted?

    $pH$ increases slightly. Dilution reduces concentration but increases degree of ionization.

22. Calculate the $pOH$ of a solution with $[OH^-] = 2 \times 10^{-3}$ M.

    $pOH = 2.7$. Use $pOH = -\log(2 \times 10^{-3}) = 2.7$.

23. How is $pK_b$ calculated from $K_b$?

    $pK_b = -\text{log}(K_b)$. Negative logarithm converts base constant to pK scale.

24. Find $K_a$ given $pK_a = 4.76$.

    $K_a = 1.74 \times 10^{-5}$. Use $K_a = 10^{-pK_a} = 10^{-4.76} = 1.74 \times 10^{-5}$.

25. Identify the effect of dilution on the $pH$ of a weak acid solution.

    Dilution increases $pH$ slightly. Equilibrium shifts right, increasing ionization and pH.

26. How is $[H^+]$ calculated from $pH$?

    $[H^+] = 10^{-pH}$. Inverse logarithm converts pH back to concentration.

27. What is the $pH$ of a $0.05$ M hydrochloric acid solution?

    $pH = 1.3$. Strong acid completely dissociates: $pH = -\log(0.05) = 1.3$.

28. Determine the $pK_b$ given $K_b = 4.5 \times 10^{-4}$.

    $pK_b = 3.35$. Use $pK_b = -\log(4.5 \times 10^{-4}) = 3.35$.

29. Determine the $pH$ of a $0.1$ M NaOH solution.

    $pH = 13$. Strong base: $[OH^-] = 0.1$ M, so $pOH = 1$, $pH = 13$.

30. What does a $pH$ greater than $7$ indicate?

    The solution is basic. pH above 7 means hydroxide concentration exceeds hydronium.

31. Determine the $pH$ of a solution if $[H^+] = 1 \times 10^{-3}$ M.

    $pH = 3$. Use $pH = -\log(1 \times 10^{-3}) = 3$.

32. State the formula for calculating $pH$ from $[H^+]$ concentration.

    $pH = -\text{log}([H^+])$. Negative logarithm converts concentration to pH scale.

33. What is the relationship between $pK_a$ and the strength of an acid?

    Lower $pK_a$ means stronger acid. Lower $pK_a$ corresponds to larger $K_a$ and stronger acid.

34. What is the conjugate base of acetic acid, CH₃COOH?

    CH₃COO⁻. Loses a proton to form the conjugate base.

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