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E. Chemical Systems and Equilibrium · Acid-Base and Solubility Equilibria

Acid-Base and Solubility Equilibria: Unit Review

A1.2 A1.4 A1.5 A1.6 A1.8 A1.9 A1.13 A2.1 A2.2 E1.2 E2.1 E2.2 E2.4 E2.5 E3.4 E3.5 E3.6 E3.7 E3.8About 40 minutes
Learning goals
  • I can identify Brønsted-Lowry acids, bases and conjugate pairs, and predict which side a proton transfer favours.
  • I can calculate pH, pOH, Ka, Kb and percentage ionization for strong and weak acids, bases, salts and buffers.
  • I can calculate and interpret titration curves, choose indicators and solve titration problems.
  • I can use Ksp and Q to find solubility, predict precipitation and explain the common ion effect, and apply these ideas to real systems.

Words to know

conjugate acid-base pair
Two species that differ by one proton.
Ka × Kb = Kw
The link between the constants of a conjugate pair (1.0 × 10⁻¹⁴ at 25 °C).
buffer
A weak acid and its conjugate base (or weak base and conjugate acid) that resists pH change.
equivalence point
Where titrant and analyte have reacted in exact stoichiometric amounts.
Ksp
The equilibrium constant for a slightly soluble solid dissolving into its ions.

Let's review!

This unit took the general equilibrium ideas from Chemical Equilibrium (K, ICE tables, Q and Le Châtelier's principle) and applied them to two huge families of reactions. You redefined acids and bases as proton donors and acceptors, paired them into conjugate pairs, and used Kw to connect pH and pOH. You calculated pH for weak acids, weak bases and salts with Ka and Kb, explained and designed buffers, analysed titration curves and chose indicators. Then you turned to solubility equilibria: Ksp, molar solubility, predicting precipitates with Q, and the common ion effect. Finally, you used all of it to explain blood chemistry, tooth decay, acid lakes, ocean acidification and water treatment.

The next unit, Electrochemistry, studies a different kind of transfer: electrons instead of protons. This review has questions from every lesson. Try each one, then press Show answer to check.

How to use this review

Work one or two sections at a time with a scientific calculator, and keep the Ka and Ksp tables from the lessons beside you. For every calculation, write the equation first and decide what kind of solution you have (the second diagram below). Do stoichiometry before equilibrium whenever a strong acid or base is added. Round only at the end; pH values get as many decimal places as the concentration has significant digits. For each question you miss, note why (wrong K? forgot to dilute? pOH reported as pH?) and go back to that lesson. Finish with the 10-question quiz.

Brønsted-Lowryproton transfer (1, 2)Kw, pH, pOHstrong acids/bases (3)Ka, Kbweak acids/bases (4-6)Salt solutionshydrolysis (7)BufferspH = pKa + log ratio (8)Titration curvesindicators (9-11)Ksp, Q, common ionsolubility (12, 13)Applicationsblood, teeth, lakes, oceans, water (14)
How the lessons in this unit connect. The Brønsted-Lowry idea and equilibrium constants underlie everything else; the numbers in brackets are lesson numbers.
What is in the solution? Pick the tool.only a strong acid or strong base[H₃O⁺] or [OH⁻] = c (× 2 for M(OH)₂)only a weak acid or weak baseICE table with Ka or Kb; check 5%only a saltfind the ion that reacts; Kb = Kw ÷ Kaweak acid AND its conjugate basebuffer: pH = pKa + log([A⁻] ÷ [HA])acid and base mixedstoichiometry first, then one of the aboveions of a slightly soluble saltQ versus Ksp; common ion lowers s
A decision guide for pH and solubility problems. Identify what is in the solution, then use the matching method.
Just for fun

Acid-base chemistry in one sentence: everybody wants the proton, nobody gets to keep it forever, and Kw keeps score.

Brønsted-Lowry Acids and Bases: Proton Transfer

1

Identify the Brønsted-Lowry acid and base: HSO₄⁻(aq) + HCO₃⁻(aq) ⇌ SO₄²⁻(aq) + H₂CO₃(aq).

2

Write equations showing HPO₄²⁻ acting as an acid with OH⁻ and as a base with H₃O⁺.

3

Explain why NH₃ is a base even though it has no OH⁻.

Conjugate Acid-Base Pairs and Relative Strength

1

Give the conjugate base of H₂S and the conjugate acid of HS⁻.

2

Label the conjugate pairs: HCOOH + CN⁻ ⇌ HCOO⁻ + HCN.

3

HCOOH (Ka 1.8 × 10⁻⁴) is a stronger acid than HCN (Ka 6.2 × 10⁻¹⁰). Which side does the reaction in the previous question favour?

The Ionization of Water: Kw, pH and pOH

1

Find [OH⁻], pH and pOH of 0.0030 mol/L HNO₃.

2

Find the pH of 0.0025 mol/L Ba(OH)₂.

3

20.0 mL of 0.150 mol/L HCl is mixed with 30.0 mL of 0.100 mol/L NaOH. Find the pH.

4

Is pure water at 37 °C (Kw = 2.4 × 10⁻¹⁴) acidic, basic or neutral? What is its pH?

Strong and Weak Acids and Bases: Ka and Kb

1

Write the Ka expression for HClO.

2

Rank by strength (strongest first): HCN (pKa 9.21), HF (pKa 3.18), CH₃COOH (pKa 4.74).

3

A 0.25 mol/L weak acid solution has [H₃O⁺] = 2.0 × 10⁻³ mol/L. Find the percentage ionization.

Weak Acid Calculations: pH, Ka and Percentage Ionization

1

Find the pH of 0.15 mol/L CH₃COOH (Ka 1.8 × 10⁻⁵).

2

A 0.050 mol/L solution of a weak acid has pH 3.10. Find Ka.

3

Find the pH of 0.010 mol/L HF (Ka 6.6 × 10⁻⁴). Is the approximation valid?

Weak Bases, Kb, and the Ka × Kb = Kw Link

1

Find the pH of 0.050 mol/L NH₃ (Kb 1.8 × 10⁻⁵).

2

Find Kb for the benzoate ion (Ka of benzoic acid = 6.3 × 10⁻⁵).

3

A base has pKb = 3.36. What is pKa of its conjugate acid?

Acid-Base Properties of Salt Solutions

1

Classify as acidic, basic or neutral: (a) NaNO₃ (b) KF (c) NH₄Br (d) Fe(NO₃)₃.

2

Find the pH of 0.20 mol/L CH₃COONa.

3

Is NaHSO₃ likely acidic or basic? For HSO₃⁻, Ka = 6.3 × 10⁻⁸ and the Ka of H₂SO₃ is 1.4 × 10⁻².

Buffer Solutions: How They Resist pH Change

1

Find the pH of a buffer that is 0.30 mol/L HCOOH and 0.20 mol/L HCOONa (Ka 1.8 × 10⁻⁴).

2

1.00 L of buffer contains 0.20 mol NH₃ and 0.20 mol NH₄Cl (pKa of NH₄⁺ 9.25). Find the pH after adding 0.050 mol HCl.

3

Name a conjugate pair suitable for a buffer at pH 7.0 and explain your choice.

Titration Curves I: Strong Acid and Strong Base

1

30.00 mL of 0.100 mol/L HCl is titrated with 0.150 mol/L NaOH. What is the equivalence volume?

2

For the titration in the previous question, find the pH after 10.00 mL of NaOH.

3

Why is the curve steepest at the equivalence point?

Titration Curves II: Weak Acids and Weak Bases

1

At the half-equivalence point of a weak acid titration, the pH is 4.87. Find Ka.

2

20.00 mL of 0.100 mol/L CH₃COOH is titrated with 0.100 mol/L NaOH. Find the pH at equivalence.

3

Sketch in words the curve for NH₃ titrated with HCl, giving the approximate pH at half-equivalence and at equivalence for 0.100 mol/L solutions.

Choosing Indicators and Solving Titration Problems

1

Choose an indicator for (a) HNO₃ with NaOH (b) HCOOH with KOH (c) CH₃NH₂ with HCl.

2

0.4084 g of KHP (204.22 g/mol, monoprotic) needs 19.80 mL of NaOH. Find [NaOH].

3

A student uses methyl orange to titrate acetic acid with NaOH. Will the calculated acid concentration be too high or too low? Why?

Solubility Equilibria and the Solubility Product, Ksp

1

Write the Ksp expression for Ca₃(PO₄)₂.

2

Find the molar solubility of PbSO₄ (Ksp 2.5 × 10⁻⁸).

3

A saturated solution of Ag₂SO₄ has [Ag⁺] = 0.030 mol/L. Find Ksp.

Predicting Precipitation and the Common Ion Effect

1

20.0 mL of 0.0020 mol/L CaCl₂ is mixed with 30.0 mL of 0.0010 mol/L NaF. Will CaF₂ (Ksp 3.5 × 10⁻¹¹) precipitate?

2

Find the molar solubility of BaSO₄ (Ksp 1.1 × 10⁻¹⁰) in 0.0010 mol/L BaCl₂.

3

Explain why Mg(OH)₂ dissolves in stomach acid but would dissolve even less in a solution of NaOH.

Acid-Base and Solubility Equilibria in Our Bodies, Environment and Industry

1

If dissolved CO₂ in blood rises while HCO₃⁻ stays the same, what happens to blood pH? How does the body respond quickly?

2

Explain why Canadian Shield lakes are sensitive to acid rain.

3

Give one way solubility equilibria are used to clean water, with an equation.

Check your understanding

Quiz

1.

In H₂PO₄⁻(aq) + NH₃(aq) ⇌ HPO₄²⁻(aq) + NH₄⁺(aq), what is the conjugate base of H₂PO₄⁻?

2.

What is the pH of 0.010 mol/L KOH at 25 °C?

3.

A 0.10 mol/L solution of a weak acid has pH 3.50. What is its Ka?

4.

What is Kb for the conjugate base of an acid with pKa 4.00?

5.

Which 0.10 mol/L solution has the highest pH?

6.

A buffer contains equal moles of HA and A⁻. A small amount of NaOH is added. What happens?

7.

At the equivalence point of a weak acid titrated with a strong base, the pH is

8.

Which indicator best detects the equivalence point when NH₃ is titrated with HCl?

9.

What is the molar solubility of a salt MX₂ with Ksp = 4.0 × 10⁻¹²?

10.

Adding NaCl to a saturated AgCl solution causes

For parents and teachers

Use the review to find what to revisit:

  • Mislabelling acids and bases, or changing a formula by the wrong number of protons → Lesson 1 (Brønsted-Lowry Acids and Bases).
  • Pairing two reactants as conjugates, or predicting the wrong side → Lesson 2 (Conjugate Acid-Base Pairs and Relative Strength).
  • pH/pOH conversions, the factor of 2 for Ba(OH)₂, or forgetting total volume after mixing → Lesson 3 (Kw, pH and pOH).
  • Including water in Ka, or ranking strength backwards → Lesson 4 (Ka and Kb).
  • Treating a weak acid as strong, skipping the 5% check → Lesson 5 (Weak Acid Calculations).
  • Reporting pOH as pH, or using the wrong Ka to find Kb → Lesson 6 (Weak Bases).
  • Calling all salts neutral, or not comparing Ka and Kb for amphiprotic ions → Lesson 7 (Salt Solutions).
  • Putting the buffer ratio upside down, or skipping the stoichiometry step → Lesson 8 (Buffers).
  • Errors in pH before and after equivalence, or mole ratios for diprotic acids → Lesson 9 (Titration Curves I).
  • Confusing half-equivalence with equivalence, or using the wrong method in each region → Lesson 10 (Titration Curves II).
  • Choosing an indicator by name, or back-titration slips → Lesson 11 (Indicators and Titration Problems).
  • Ksp in terms of s (4s³, 27s⁴), or comparing Ksp across formula types → Lesson 12 (Ksp).
  • Forgetting to dilute after mixing, or common-ion setup → Lesson 13 (Precipitation and the Common Ion Effect).
  • Explaining real systems (blood, teeth, lakes, oceans) → Lesson 14 (Our Bodies, Environment and Industry).

Looking ahead: the next unit, Electrochemistry, studies electron transfer (redox) reactions, batteries and electrolysis. The idea of comparing the strength of two competitors, which you used for acids and bases, returns there as comparing oxidizing and reducing agents.