Unit 8: Acids and Bases.

8.1 Introduction to Acids and Bases.

Calculate the value of pH, pOH, [H⁺], or [OH⁻] given one of the values.

  1. A solution has a pH of 8.00 at 25oC. What is the pOH?
  2. The [OH-] of a solution is 0.300 M. Determine the [H3O+].

8.2 pH and pOH of Strong Acids and Bases.

Differentiate strong and weak acids.

  1. Identify which of the following are strong acids: HCl, HF, HNO3, H2SO4, HCN, and HClO4.
  2. List the 6 strong acids.

Differentiate strong and weak bases

  1. Identify which of the following are strong bases: NH3, CH3NH2, NaOH, Ca(OH)2, Sr(OH)2, Ba(OH)2, and LiOH.
  2. Explain why NH3 is classified as a weak base even though it reacts with water to produce OH−.

Calculate the pH, pOH, [H3O+], or [OH-] of a strong acid or base from its concentration.

  1. What is the pH of 2.5 M HCl?
  2. What are the [H3O+] and [OH-] concentrations of a 1.0 M Ca(OH)2 solution?

8.3 Weak Acid and Base Equilibria.

Compare strong vs. weak acids and bases.

  1. Two monoprotic acids, HA and HB, have the same initial concentration. The pH of HA is 2.00 and the pH of HB is 3.00. Which acid is stronger? Explain your reasoning.
  2. Explain the difference between the strength of an acid and the concentration of an acid.

Use a Ka expression to calculate the pH of a weak acid (or Kb for a weak base).

  1. A 0.100 M solution of a weak acid, HA, has a Ka of 1.8 × 10-5. Calculate the pH of the solution.
  2. A 0.250 M solution of NH3 has a Kb of 1.8 × 10-5. Calculate the pH of the solution.

Calculate the pKa of a weak acid or pKb of a weak base.

  1. A weak acid has a Ka of 6.3 × 10-5. Calculate its pKa.
  2. A weak base has a Kb of 4.0 × 10-6. Calculate its pKb.

Determine the % ionization of a weak acid or base given its pKa (or pKb) and starting concentration.

  1. A 0.100 M solution of a weak acid has a pKa of 4.75. Calculate the percent ionization of the acid.
  2. A 0.0100 M weak acid is 2.0% ionized. What is the concentration of H3O+ at equilibrium?

8.4 Acid-Base Reactions and Buffers.

Determine the pH of a solution after complete neutralization of a strong acid and strong base.

  1. 25.0 mL of 0.100 M HCl is mixed with 30.0 mL of 0.100 M NaOH. Determine the pH of the resulting solution.

Write the equation for acid-base reactions (strong-strong, weak-strong, or weak-weak).

  1. Write the net ionic equation for the reaction between HCl(aq) and NaOH(aq).
  2. Write the net ionic equation for the reaction between CH3COOH(aq) and NaOH(aq).
  3. Write the net ionic equation for the reaction between NH3(aq) and HCl(aq).

Determine the pH of a solution after neutralization.

  1. 25.0 mL of 0.100 M HCl is mixed with 15.0 mL of 0.100 M NaOH. Determine the pH of the resulting solution.
  2. 25.0 mL of 0.100 M CH3COOH is titrated with 25.0 mL of 0.100 M NaOH. At the equivalence point, is the solution acidic, neutral, or basic? Explain why.

8.5 Acid-Base Titrations.

Produce a titration curve.

  1. Sketch the titration curve for a weak monoprotic acid being titrated with a strong base. Label the initial pH, buffer region, half-equivalence point, equivalence point, and region after the equivalence point.

Identify key areas on a titration curve: equivalence point, half-equivalence point, buffering region.

  1. On a titration curve for a weak acid titrated with a strong base, identify the equivalence point and half-equivalence point. What is the relationship between pH and pKa at the half-equivalence point?

Use the shape of a titration curve to determine if an acid is mono-, di-, or triprotic.

  1. A titration curve contains two distinct equivalence points. What can you conclude about the number of acidic protons in the original acid? Explain your reasoning.
  2. How many equivalence points would you expect for the titration of a triprotic acid with a strong base?

Identify the major species present at any point on a titration curve.

  1. A weak acid HA is being titrated with NaOH. Identify the major acid-base species present before any NaOH is added, at the half-equivalence point, at the equivalence point, and after the equivalence point.

8.6 Molecular Structure of Acids and Bases.

Identify the acidic proton in a carboxylic acid.

  1. In CH3COOH, identify the hydrogen atom that is most likely to be donated as H+. Explain why the other hydrogen atoms are much less acidic.

Justify the strength of an acid by examining the electronegativity of the anion.

  1. HCl and HI are both strong acids. Using the electronegativity of the conjugate base and the H-X bond, explain why HI is a stronger acid than HCl.
  2. Compare the acidity of H2O and NH3 by considering the relative electronegativities of O and N.

8.7 pH and pKa.

Determine if a weak acid or conjugate base is in higher concentration by comparing pH and pKa.

  1. A solution containing a weak acid HA has a pH of 4.20 and a pKa of 3.75. Which species, HA or A−, is present at the higher concentration? Explain your reasoning.
  2. If pH = pKa for a weak acid, what can you conclude about the relative concentrations of HA and A−?

Select an appropriate indicator for a titration.

  1. A weak acid is being titrated with a strong base. Should an indicator that changes color between pH 3 and 5 or one that changes color between pH 8 and 10 be selected? Explain your reasoning.
  2. Why is it important for the color-change range of an indicator to overlap with the steep portion of the titration curve?

8.8 Properties of Buffers.

Determine if a solution is a buffer.

  1. Which of the following solutions would function as a buffer? Explain your reasoning.
    (a) HCl and NaCl
    (b) CH3COOH and CH3COONa
    (c) NaOH and KCl
    (d) HNO3 and KNO3

Explain how a buffer solution stabilizes the pH of a system.

  1. Explain how a buffer containing CH3COOH and CH3COO− responds when a small amount of H3O+ is added.
  2. Explain how the same buffer responds when a small amount of OH− is added.

8.9 Henderson-Hasselbalch Equation.

Use the Henderson-Hasselbalch equation to determine the pH of a buffer solution.

  1. A buffer contains 0.20 M CH3COOH and 0.30 M CH3COO−. The pKa of CH3COOH is 4.76. Calculate the pH of the buffer.

Qualitatively describe how a buffer resists pH changes when acid or base is added.

  1. A buffer contains significant concentrations of HA and A−. Predict what happens to the pH when a small amount of HCl is added. Explain your reasoning at the particle level.
  2. Predict what happens to the pH of the same buffer when a small amount of NaOH is added.

8.10 Buffer Capacity.

Qualitatively describe whether a given buffer can accommodate additional acid or base.

  1. Two buffers have the same pH. Buffer A contains 0.010 M HA and 0.010 M A−, while Buffer B contains 0.50 M HA and 0.50 M A−. Which buffer has the greater capacity to resist a large addition of strong acid or base? Explain why.

8.11 pH and Solubility.

Determine if the solubility of a salt will be impacted by the pH of the solvent.

  1. Predict whether lowering the pH of a solution will increase the solubility of CaCO3. Explain your reasoning.

Explain the impact of pH on solubility using Le Châtelier’s principle.

  1. Carbonate ions react with H3O+ to form H2CO3. Explain, using Le Châtelier’s principle, why adding acid increases the solubility of CaCO3.
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