Acid-Base Reactions and Buffers - AP Chemistry

1. Which strong base completely dissociates in water?

   Sodium hydroxide ($NaOH$). Strong base with 100% dissociation in water.

2. What is the formula for calculating pH from hydrogen ion concentration?

   p$H = -\text{log}[H^+]$. Negative log transforms $H^+$ concentration to pH scale.

3. Identify the conjugate acid of the base $H_2O$.

   $H_3O^+$. Water gains a proton to form hydronium ion.

4. Which ion is the conjugate acid of $NH_2^-$?

   $NH_3$. Gains a proton to form ammonia from amide ion.

5. What happens to the pH of a buffer when a small amount of acid is added?

   pH decreases slightly. Buffer components neutralize added acid effectively.

6. Determine the pH if $[H^+] = 2.5 \times 10^{-4} M$.

   pH = 3.60. pH = $-\log(2.5 \times 10^{-4}) = 3.60$.

7. Identify the weak acid in the acetic acid-sodium acetate buffer system.

   Acetic acid ($CH_3COOH$). Weak acid component that partially ionizes.

8. What is the formula for calculating $pK_w$?

   $pK_w = pH + pOH$. Sum equals 14 at 25°C for aqueous solutions.

9. What is the effect of dilution on the pH of a buffer?

   Minimal change in pH. Buffer ratio remains constant during dilution process.

10. Calculate the pH if $[OH^-] = 1.0 \times 10^{-5} M$ at 25°C.

    pH = 9. pOH = 5, so pH = 14 - 5 = 9.

11. Identify the conjugate base of the acid $H_2O$.

    $OH^-$. Water loses a proton to form hydroxide ion.

12. Which strong acid completely dissociates in water?

    Hydrochloric acid ($HCl$). Strong acid with 100% ionization in water.

13. Identify the conjugate base of the acid $HCl$.

    $Cl^-$. Loses a proton from $HCl$ to form chloride ion.

14. What is the formula to calculate pOH from hydroxide ion concentration?

    p$OH = -\text{log}[OH^-]$. Negative log transforms $OH^-$ concentration to pOH scale.

15. What is the $K_a$ expression for acetic acid $CH_3COOH$?

    $K_a = \frac{[CH_3COO^-][H^+]}{[CH_3COOH]}$. Products over reactants for acid dissociation equilibrium.

16. Calculate the pH of a buffer with $pK_a = 4.76$ and $\frac{[A^-]}{[HA]} = 1$.

    pH = 4.76. When ratio equals 1, $\log(1) = 0$, so pH = $pK_a$.

17. Calculate the pH change when $0.01$ mol $HCl$ is added to a 1L buffer with $pK_a = 4.76$.

    Minimal pH change. Buffer capacity prevents large pH shifts effectively.

18. Calculate the pOH of a solution with $[OH^-] = 1.0 \times 10^{-4} M$.

    pOH = 4. pOH = $-\log(1.0 \times 10^{-4}) = 4$.

19. What is the primary function of a buffer?

    To maintain stable pH. Prevents large pH changes in chemical systems.

20. What is the relationship between $K_a$ and $K_b$ for a conjugate acid-base pair?

    $K_a \times K_b = K_w$. Product equals water's ionization constant at 25°C.

21. What is the value of $K_w$ at 25°C?

    $1.0 \times 10^{-14}$. Water's ionization constant at standard temperature.

22. Identify the weak base in the ammonia-ammonium chloride buffer system.

    Ammonia ($NH_3$). Weak base component that partially ionizes.

23. Identify the components of a buffer solution.

    Weak acid and its conjugate base. Both species neutralize added acids and bases.

24. State the Bronsted-Lowry definition of a base.

    Proton acceptor. Bases accept protons from other species in solution.

25. What defines a buffer solution?

    Resists changes in pH upon addition of acid or base. Weak acid-base pair maintains stable pH range.

26. What is the effect of adding a strong acid to a buffer solution?

    Buffer resists significant pH change. Buffer neutralizes acid through base component reaction.

27. Calculate the $pK_a$ for an acid with $K_a = 1.0 \times 10^{-5}$.

    $pK_a = 5$. $pK_a = -\log(1.0 \times 10^{-5}) = 5$.

28. Calculate the pH of a solution with $[H^+] = 1.0 \times 10^{-3} M$.

    pH = 3. pH = $-\log(1.0 \times 10^{-3}) = 3$.

29. What is the Henderson-Hasselbalch equation for pH?

    $\mathrm{pH} = \mathrm{p}K_a + \log \frac{[A^-]}{[HA]}$. Relates buffer pH to acid strength and base ratio.

30. What is the concentration of $H^+$ in a solution with pH 4?

    $1.0 \times 10^{-4} M$. $[H^+] = 10^{-4} = 1.0 \times 10^{-4}$ M.

31. Identify the conjugate acid of the base $NH_3$.

    $NH_4^+$. Gains a proton from $NH_3$ to form ammonium ion.

32. What is the pH of a neutral solution at 25°C?

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

33. Find the pH of a 0.1 M $HCl$ solution.

    pH = 1. Strong acid completely ionizes: $[H^+] = 0.1$ M.

34. Which ion is the conjugate base of $H_3PO_4$?

    $H_2PO_4^-$. Loses one proton from phosphoric acid.

35. What is the $K_b$ expression for ammonia $NH_3$?

    $K_b = \frac{[NH_4^+][OH^-]}{[NH_3]}$. Products over reactants for base ionization equilibrium.

36. What characteristic of a weak acid makes it suitable for a buffer?

    Partial ionization. Incomplete dissociation provides buffering capacity.

37. Which ion is responsible for the acidity of a solution?

    Hydrogen ion ($H^+$). Higher $H^+$ concentration means more acidic solution.

38. State the Bronsted-Lowry definition of an acid.

    Proton donor. Acids release protons to other species in solution.

Study these flashcards