Henderson-Hasselbalch Equation - AP Chemistry

1. What does the Henderson-Hasselbalch equation assume about ionic strength?

   It is constant. Simplifies calculations by ignoring activity coefficients.

2. For a buffer with $pK_a = 4.5$, what is pH when $[A^-] = 0.2$ M and $[HA] = 0.1$ M?

   pH = 4.8. $\text{log}(2) = 0.3$, so pH = 4.5 + 0.3.

3. What is the relation between pH and $pK_a$ when pH < $pK_a$?

   The solution has more weak acid than conjugate base. Lower pH indicates more acid than base present.

4. How does dilution affect the pH of a buffer according to the Henderson-Hasselbalch equation?

   pH remains relatively constant. Ratio stays constant when both components dilute equally.

5. If $[A^-]/[HA] = 2$, find pH given $pK_a = 7.2$.

   pH = 7.5. $\text{log}(2) = 0.3$, so pH = 7.2 + 0.3.

6. What is the primary use of the Henderson-Hasselbalch equation?

   To calculate the pH of buffer solutions. Combines weak acid-base pairs to resist pH changes.

7. What is the effect of temperature on $pK_a$ in the Henderson-Hasselbalch equation?

   $pK_a$ can change with temperature. Temperature affects equilibrium constant values.

8. What is the relation between pH and $pK_a$ when pH > $pK_a$?

   The solution has more conjugate base than acid. Higher pH indicates more base than acid present.

9. What is the impact on pH if $[A^-]$ decreases in the buffer solution?

   pH decreases. Less conjugate base shifts equilibrium toward lower pH.

10. How is the Henderson-Hasselbalch equation derived?

    From the acid dissociation constant expression. Taking negative log of the acid dissociation expression.

11. Find pH for $pK_a = 3.5$, $[A^-] = 0.03$ M, $[HA] = 0.01$ M.

    pH = 4.0. $\text{log}(3) ≈ 0.5$, so pH = 3.5 + 0.5.

12. What is the significance of the Henderson-Hasselbalch equation in titrations?

    It helps estimate pH at various stages of a titration. Predicts pH changes during weak acid-base reactions.

13. How does the Henderson-Hasselbalch equation help in buffer preparation?

    Determines component ratios for desired pH. Calculate required acid-to-base ratio for target pH.

14. If $[HA] = 0.1$ M and $[A^-] = 0.01$ M, how does pH compare to $pK_a$?

    pH < $pK_a$. More acid than base means pH below $pK_a$.

15. What is the effect on pH if $[A^-]$ increases relative to $[HA]$?

    pH increases. Larger numerator in ratio increases log term value.

16. Calculate pH if $pK_a = 9.3$, $[A^-] = 0.05$ M, $[HA] = 0.1$ M.

    pH = 9.0. $\text{log}(0.5) = -0.3$, so pH = 9.3 - 0.3.

17. What does $pK_a$ represent in the Henderson-Hasselbalch equation?

    The negative logarithm of the acid dissociation constant. Higher $pK_a$ means weaker acid dissociation.

18. State the logarithmic property used in the Henderson-Hasselbalch equation.

    $\text{log}(a/b) = \text{log}(a) - \text{log}(b)$. Allows separation of the concentration ratio term.

19. For a buffer with $pK_a = 6.3$, what is pH when $[A^-] = 0.5$ M, $[HA] = 0.25$ M?

    pH = 6.6. $\text{log}(2) = 0.3$, so pH = 6.3 + 0.3.

20. What happens to pH if $[A^-] = [HA]$ in the Henderson-Hasselbalch equation?

    pH equals $pK_a$. When the ratio equals 1, the log term becomes zero.

21. In the Henderson-Hasselbalch equation, what does a higher $pK_a$ indicate?

    A weaker acid. Larger $pK_a$ corresponds to smaller $K_a$ value.

22. If $pK_a = 6.1$ and $[A^-] = 0.25$ M, $[HA] = 0.25$ M, what is the pH?

    pH = 6.1. Equal concentrations make log term zero.

23. What does the ratio $[A^-]/[HA]$ signify in the Henderson-Hasselbalch equation?

    Relative amounts of conjugate base and weak acid. Determines whether solution is acidic or basic.

24. How does the Henderson-Hasselbalch equation relate to buffers?

    It calculates the pH of a buffer given $pK_a$, $[A^-]$, and $[HA]$. Uses component concentrations to predict buffer pH.

25. Which condition is assumed in the Henderson-Hasselbalch equation?

    The solution is a buffer solution. Contains both weak acid and conjugate base in equilibrium.

26. Identify the meaning of $[A^-]$ in the Henderson-Hasselbalch equation.

    The concentration of the conjugate base. The deprotonated form of the weak acid.

27. What is the logarithmic form of the Henderson-Hasselbalch equation?

    $pH = pK_a + \text{log}(\frac{[A^-]}{[HA]})$. Standard form using base-10 logarithm.

28. What assumption is made about $[A^-]$ and $[HA]$ in the Henderson-Hasselbalch equation?

    They are equilibrium concentrations. Assumes no significant change from initial values.

29. Identify the meaning of $[HA]$ in the Henderson-Hasselbalch equation.

    The concentration of the weak acid. The protonated form that can donate hydrogen ions.

30. Find the pH if $pK_a = 8.0$, $[A^-] = 0.1$ M, $[HA] = 0.05$ M.

    pH = 8.3. $\text{log}(2) = 0.3$, so pH = 8.0 + 0.3.

31. Calculate the pH if $pK_a = 5.4$, $[A^-] = 0.2$ M, $[HA] = 0.2$ M.

    pH = 5.4. Equal concentrations make the log term zero.

32. How is the Henderson-Hasselbalch equation used in biological systems?

    To maintain pH in physiological systems. Blood pH regulation uses bicarbonate buffer system.

33. What does the term 'buffer capacity' refer to?

    The amount of acid or base the buffer can neutralize. Measures resistance to pH change upon acid/base addition.

34. Find the pH if $pK_a = 4.75$, $[A^-] = 0.1$ M, $[HA] = 0.1$ M.

    pH = 4.75. Equal concentrations make the log term zero.

35. Determine the pH if $pK_a = 7.8$, $[A^-] = 0.3$ M, $[HA] = 0.1$ M.

    pH = 8.3. $\text{log}(3) ≈ 0.5$, so pH = 7.8 + 0.5.

36. State the formula for the Henderson-Hasselbalch equation.

    $pH = pK_a + \text{log}(\frac{[A^-]}{[HA]})$. Relates buffer pH to acid strength and component ratio.

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