Acid-Base Titrations - AP Chemistry

1. Which option describes a buffer solution?

   Resists changes in pH upon addition of acid or base. Contains conjugate acid-base pair to maintain pH stability.

2. What is the equivalence point in a titration curve?

   Where moles of acid equal moles of base. Stoichiometric neutralization point on titration curve.

3. Calculate the pH of a 0.01 M HCl solution.

   pH = 2. Strong acid: pH = $-\log(0.01) = -\log(10^{-2}) = 2$.

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

   $pH = 7$. Equal concentrations of $H^+$ and $OH^-$ at 25°C.

5. Choose the correct pH range for phenolphthalein.

   8.2 to 10.0. Indicator's transition range for color change.

6. What is the pH of a 0.1 M NH₃ solution if $K_b = 1.8 \times 10^{-5}$?

   pH ≈ 11.1. Use $K_b$ expression: $[OH^-] = \sqrt{K_b \times C}$ then find pH.

7. Calculate the volume of 0.1 M NaOH needed to neutralize 50 mL of 0.1 M HCl.

   50 mL. Equal molarity and volume means 1:1 mole ratio requires equal volume.

8. Which indicator is commonly used for strong acid-strong base titrations?

   Phenolphthalein. Changes color at pH 8.2-10, ideal for strong acid-base endpoints.

9. State the formula for converting pH to hydrogen ion concentration.

   $[\text{H}^+] = 10^{-\text{pH}}$. Inverse relationship: pH is negative log of $[H^+]$.

10. In a titration, what is the titrant?

    The solution added from the burette. Standard solution delivered from burette during analysis.

11. Find and correct the error: 'The equivalence point is always at pH 7.'

    Not always; depends on the strength of acid and base. Only true for strong acid-strong base; weak acids/bases differ.

12. What is the color change of phenolphthalein in basic solutions?

    Colorless to pink. Phenolphthalein is colorless in acid, pink in base.

13. Identify the primary characteristic of a weak acid.

    Partially ionizes in solution. Incomplete dissociation results in equilibrium with undissociated acid.

14. What is the role of a pipette in a titration?

    To measure and transfer a specific volume of solution. Precision instrument for accurate volume transfer.

15. What is the purpose of a titration in acid-base chemistry?

    To determine the concentration of an unknown solution. Uses known titrant to find unknown concentration via stoichiometry.

16. Which equation relates pH and pOH?

    $\text{pH} + \text{pOH} = 14$. Fundamental relationship for aqueous solutions at 25°C.

17. What is the definition of a titrant?

    A solution of known concentration used in titration. Standard solution added from burette to determine analyte concentration.

18. Which option best describes a titration curve?

    Graph of pH versus volume of titrant added. Visual representation showing pH changes during titration process.

19. What is the effect of a strong base on the pH of a buffer?

    Slight increase in pH. Buffer components consume added base, minimizing pH change.

20. What is the effect of dilution on the pH of a strong acid?

    Increases the pH. Dilution decreases $[H^+]$, raising pH toward neutral.

21. What happens to pH when a strong acid is added to a buffer?

    Slight decrease in pH. Buffer components consume added acid, minimizing pH change.

22. Which term describes the solution whose concentration is to be determined during titration?

    Analyte. The unknown solution whose concentration is being determined.

23. State the formula for calculating molarity.

    $M = \frac{\text{moles of solute}}{\text{liters of solution}}$. Defines concentration as moles of solute per liter of solution.

24. Which acid-base pair has a higher pH at the equivalence point: strong acid-weak base or weak acid-strong base?

    Weak acid-strong base. Weak acid-strong base gives basic equivalence point (pH > 7).

25. Calculate the moles of HCl in 25 mL of 0.1 M solution.

    0.0025 moles. Convert volume to L (0.025 L), multiply by molarity (0.1 M).

26. What type of reaction occurs during a titration?

    Neutralization. Acid and base react to form salt and water.

27. What is the pH at the equivalence point for a strong acid-strong base titration?

    pH = 7. Equal amounts of strong acid and base neutralize to neutral pH.

28. Identify the endpoint in a titration.

    The point at which the indicator changes color. Visual signal used to determine when to stop adding titrant.

29. Find the pH of a solution with $[\text{OH}^-] = 10^{-3}$ M.

    pH = 11. pOH = $-\log(10^{-3}) = 3$, so pH = $14 - 3 = 11$.

30. What is the role of a burette in a titration?

    To deliver a precise volume of titrant. Graduated tube allows controlled addition and volume measurement.

31. What is the primary characteristic of a strong acid?

    Completely ionizes in solution. Dissociates nearly 100% in aqueous solution.

32. What is the endpoint of a titration?

    The point where the indicator changes color. Observable color change indicating approximate equivalence point.

33. What is a standard solution?

    A solution of known concentration. Precisely known concentration for quantitative analysis.

34. Determine the pH of a solution with $[\text{H}^+] = 10^{-6}$ M.

    pH = 6. pH is defined as $-\log[H^+] = -\log(10^{-6}) = 6$.

35. Identify the formula to convert pOH to hydroxide ion concentration.

    $[\text{OH}^-] = 10^{-\text{pOH}}$. Inverse relationship: pOH is negative log of $[OH^-]$.

36. What is the relationship between $K_a$, the acid dissociation constant, and acid strength?

    Higher $K_a$ indicates a stronger acid. Larger $K_a$ means greater ionization and stronger acid.

37. Identify the primary purpose of an indicator in a titration.

    To signal the endpoint of the titration. Color change marks approximate equivalence point.

38. State the formula to find pOH from hydroxide ion concentration.

    $\text{pOH} = -\text{log}[\text{OH}^-]$. Negative logarithm relationship for hydroxide concentration.

39. Calculate the pOH of a 0.01 M NaOH solution.

    pOH = 2. Strong base: pOH = $-\log(0.01) = 2$.

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