Acid-Base Theory

ACID-BASE THEORY

1. Give a balanced equation for the ionization of a typical Arrhenius acid such as H2S.

2. Give a balanced equation for the ionization of a typical Arrhenius base such as Ca(OH)2.

3. (a) Give a balanced equation for the proton transfer reaction of a typical Brönsted-Lowry acid such as

HCl with water.

(b) Identify the conjugate acid/base pairs in this system.

(d) Name the two Brönsted-Lowry acids present in this equilibrium system.

(e) Name the two Brönsted-Lowry bases present in this equilibrium system.

4. (a) Give a balanced equation for the proton transfer reaction of a typical Brönsted-Lowry base such as

NH3 with water.

(b) Identify the conjugate acid/base pairs in this system.

(d) Could the Arrhenius theory account for this basicity? Explain.

(e) Name the two Brönsted-Lowry acids present in this equilibrium system.

(f) Name the two Brönsted-Lowry bases present in this equilibrium system.

5. Consider the dissolving of pure H2SO4 in pure HC2H3O2.

(a) Identify the conjugate acid/base pairs.

(b) Would this reaction be considered an acid/base reaction under Arrhenius Theory? Explain.

6. Write an equation for each of the following reactions. Identify the acids and bases. Indicate one conjugate pair by drawing boxes around their formulas and the other conjugate pair by circling them.

(a) H3PO4 acting as a Brönsted acid in an aqueous solution.

(b) C2H3O21-acting as a Brönsted base in an aqueous solution.

(d) the nonaqueous reaction in which hydrogen nitrate acts as a base and sulphuric acid acts as an acid.

7. NaOH is a base and HClO3 is an acid. Each has an O-H covalent bond. Where the bonds break upon ionization in water accounts for whether or not they are acids or bases.

(a) Draw an electron dot diagram for each.

(b) Account for the bond breaking between the Na and the O-H in NaOH and between the O and the H in HClO3.

8. (a) Write equilibrium reaction equations for the dissociation of HNO3 and HNO2 in water.

(b) Identify the conjugate acid-base pairs for the forward and reverse reactions.

(c) Since both systems test as decidedly acidic with litmus, to which side of the reaction equation do each of the equilibria lie?

(d) Identify the stronger acid in each equilibrium.

(e) Identify the stronger base in each equilibrium.

(f) Give the Electron Dot Drawing for each of HNO3 and HNO2.

(g) On the basis of molecular structures, electronegativities and bond polarities, which of HNO3 and HNO2 will have the more polar O-H bond?

(h) Predict the relative acid strengths of HNO3 and HNO2. Explain.

9. In order for a molecule such as HF to ionize in water, the H-F bond must be broken and each of the ions created must be solvated. Each of these steps involves energy being gained or lost.

(a) In which hydrogen halide, HF or HI will the hydrogen halide bond require more energy to be broken? Explain.

(b) The I-1 ion will release more energy when solvated by a water molecule than will the F-1 ion. Discuss two factors that contribute to this energy release.

10. Some acids are polyprotic, i.e. they have more than one proton available for donation, eg. H2SO4. The loss of these protons may occur in successive steps.

(a) Write the ionization equation for the Brönsted-Lowry acid H2SO4 donating one proton.

(b) Identify the conjugate acid/base pairs.

(d) Write the ionization equation for the above product donating its remaining proton as a Brönsted-Lowry acid.

(e) Identify the conjugate acid/base pairs.

(f) Offer an explanation for why the reaction in (a) is more acidic than that in (d).

Solutions by Dilution

1. What is the concentration (molarity) of a solution of NaCl if 40. mL of a 2.5 M NaCl solution is diluted to a total volume of 500. mL? [0.20 M]

2. What is the concentration of a solution of Fe(NO3)3 if 80. mL of a 3.0 M Fe(NO3)3 solution is concentrated to a total volume of 15 mL? [16 M]

3. What volume of water is needed to create a 0.675M KOH solution from an original 75.0 mL of a 2.25 M KOH solution? [175 mL]

4. What is the concentration of a solution made by mixing 8.0 mL of a 1.0 M CaSO3 solution with 12.0 mL of a 2.0 M CaSO3 solution? [1.6 M]

5. How many mL of 12.0 M HCl is needed to make 250. mL of a 1.50 M HCl solution? [31.3 mL]

6. How would you prepare 250. mL of a 6.00 M NaNO3 solution from a 15.0 M NaNO3 solution? [To 150 mL of water, add 100 mL of 15 M NaNO3 solution]

Titration Practice Questions

1. If it takes 54.0 mL of 0.10 M NaOH to neutralize 125 mL of an HCl solution, what is the concentration of the acid? [0.043 M]

2. If it takes 50.0 mL of 0.50 M KOH to completely neutralize 125 mL of an H2SO4 solution, what is the concentration of the acid? [0.10 M]

3. What volume of 1.2 M HClO4 is needed to neutralize 5.8 mL of a 0.44 M Ba(OH)2 solution? [4.3 mL]

4. What is the pH of a solution that contains 0.0090 g of HCl in 100. mL of water? [pH = 2.61]

5. What is the concentration of H3O+ (or H+) in a flask if 20.0 mL of 0.10 M NaOH solution is added to 15.0 mL of 0.20 M HCl solution? [0.029 M]

Acid-Base Equilibrium Problems

NOTE: (sig. figs. calculated using Ka, Kb values where applicable)

1) Calculate the pH given the hydronium ion concentration for the following:

a) 3.5 x 10-8 M (7.46) b) 1.89 x 10-1 M (0.724)

2) Calculate the hydronium ion concentration for the following pH values:

a) 7.52 (3.0 x 10-8 M) b) 0.77 (1.7 x 10-1 M)

3) Calculate the pH of the following solutions with the following hydroxide ion concentrations:

a) 2.56 x 10-8 M (6.408) b) 7.9 x 10-2 M (12.90)

4) Calculate the hydroxide ion concentration for the following pH values:

a) 1.52 (3.3 x 10-13 M) b) 8.96 (9.1 x 10-6 M)

5) What is the pH of a 0.100 M HCN solution? (5.10)

6) Calculate the hydronium ion concentration in a 0.200 M acetic acid (CH3COOH) solution. (1.9 x 10-3 M)

7) What is the pH of a solution containing 0.200 M formic acid (HCOOH)? (2.22)

8) A pH meter is used to find the equilibrium pH for the following acids. Find the ionization constants (Ka or Kb) for each of the following systems:

a) 0.010 M CH3COOH has a pH = 3.37 (Ka = 1.9 x 10-5)

b) 0.020 M HCNO has a pH = 2.70 (Ka = 2.2 x 10-4)

c) 0.010 M (CH3)2NH has a pH = 11.34 (Kb = 6.1 x 10-4)

9) Calculate the percentage ionization of each of the following solutions

a) 0.0030 M HCN solution Ka = 4.9 x 10-10 (0.040 %)

b) 1.25 M CH3NH2 solution Kb = 4.4 x 10-4 (1.8 %)

10) A 5.00 g sample of CH3COOH is added to 500.0 mL of water. What is the pH of the solution? (2.77)

11) A 5.00 g sample of dimethylamine ((CH3)2NH) is added to 500.0 mL of water. Calculate the pH of the solution. (12.18)

12) Calculate the Ka for 0.200 M HAsO2 (aq) if 0.00550% of the HAsO2 solution is ionized. (Ka = 6.05 x 10-10)

ACID/BASE TITRATION CURVE PROBLEMS

For each of the following, 1st determine the equivalence point volume then calculate the pH at conditions a to d:

1. A 100. mL sample of a 0.0500M HNO3 solution is titrated with a 0.200M KOH solution.

(a) before any base is added (1.301)

(b) after 12.5 mL of base is added (1.654)

(d) after 37.5 mL of base is added (12.260)

2. A 15.0 mL sample of a 0.250M NaOH solution is titrated with a 0.125 M HBr solution.

(a) before any acid is added (13.398)

(b) after 15.0 mL of acid is added (12.796)

(d) after 45.0 mL of acid is added (1.505)

3. A 25.0 mL sample of a 0.200M HCOOH solution is titrated with a 0.100M NaOH solution. Ka = 1.8 x 10-4.

(a) before any base is added (2.22)

(b) after 25.0 mL of base is added (3.74)

(d) after 75.0 mL of base is added (12.398)

4. A 100. mL sample of a 0.100M NH3 solution is titrated with a 0.200M HCl solution.

Kb = 1.8 x 10-5.

(a) before any acid is added (11.13)

(b) after 25.0 mL of acid is added (9.26)

(d) after 75.0 mL of acid is added (1.544)

5. A 15.0 mL sample of a 0.250M HC2H3O2 solution is titrated with a 0.125 M LiOH solution. Ka = 1.8 x 10-5.

(a) before any base is added (2.67)

(b) after 15.0 mL of base is added (4.74)

(d) after 45.0 mL of base is added (12.496)

6. A 20.0 mL sample of a 0.250M CH3NH2 solution is titrated with a 0.125 M HCl solution. Kb = 4.4 x 10-4.

(a) before any acid is added (12.02)

(b) after 20.0 mL of acid is added (10.64)

(d) after 60.0 mL of acid is added (1.505)

BUFFER PROBLEMS

A 1.00 L sample of an aqueous solution contains 0.200 mol of acetic acid and 0.100 mol sodium acetate. Given that the Ka of acetic acid is 1.8 x 10-5, calculate the following;

(a) the pH of the solution [4.44]

(b) the pH of the solution after the addition of 1.00 mL of concentrated HCl (12.0 M). [4.36]

Calculate the pH of a 1.00 L solution which is 0.75 M lactic acid (Ka = 1.4 x 10-4) and 0.25 M sodium lactate before and after the addition of 2.00 mL of 16.0 M NaOH. [3.38] and [3.44].

A buffered solution contains 0.25 M NH3 (Kb = 1.8 x 10-5) and 0.40 M NH4Cl. Calculate the pH of this solution. [9.05]

How many grams of sodium acetate must be dissolved in a 0.500 L solution of 0.50 M acetic acid to obtain a buffer solution with a pH of 5.00? [37 g]

A chemist needs a solution buffered at pH 4.30 and can choose from the following acids and their sodium salts.

Chloroacetic Acid Ka = 1.35 x 10-3

Propanoic Acid Ka = 1.3 x 10-5

Benzoic Acid Ka = 6.4 x 10-5

Hypochlorous Acid Ka = 3.5 x 10-8

Calculate the ratio of [Acid]/[Base] required for each system to yield the desired pH. Which system will work best?

Lewis Acid/Base

1. (a) Give the Lewis definition of an acid.

(b) Give the Lewis definition of a base.

Cu2+ + 4 NH3 « Cu(NH3)42+

(d) Identify the Lewis acids and bases in the following list:

CN-1 AlCl3 Ag+ S2O3-2 Ni2+ Fe3+ BeF2 PCl3

2. Write equations for the following and label the Lewis acid and base.

(a) the hydration of calcium ions when calcium chloride dissolves in water

(b) the formation of a complex ion when silver chloride is dissolved in aqueous ammonia

REVIEW: ACID/BASE PROBLEMS

1. Predict and explain the relative pH of a solution of sodium acetate. Include all equations showing the reactions involved in your explanation.

2. Predict and explain the relative pH of a solution of ammonium bisulphite, NH4HSO3. Include all equations showing the reactions involved in your explanation.

3. Calculate the pH of a 0.175M HC2H3O2 solution. ( 2.75 )

4. Calculate the pH of a 0.175M NH3 solution. ( 11.25 )

5. A 100. mL sample of a 0.0500M strong acid, HX, is titrated with a 0.200M strong base, MOH, solution. Calculate the equivalence point and the pH at:

(a) before any base is added ( 1.301 )

(b) after 1/2 the volume to equivalence point is added ( 1.654 )

(d) after 1 1/2 the volume to equivalence point is added ( 12.260 )

6. A 100. mL sample of a 0.0500M strong base, MOH, is titrated with a 0.200M strong acid, HX, solution. Calculate the equivalence point and the pH at:

(a) before any acid is added ( 12.699 )

(b) after 1/2 the volume to equivalence point is added ( 12.346 )

(d) after 1 1/2 the volume to equivalence point is added ( 1.740 )

7. A titration of 60.0 mL of a 0.200M weak acid, HA, solution using a 0.250M strong base, MOH, solution was performed. Given that the Ka = 1.2 x 10-6 calculate the equivalence point and the pH at:

(a) before any base is added ( 3.31 )

(b) after 1/2 the volume to equivalence point is added ( 5.92 )

(d) after 1 1/2 the volume to equivalence point is added ( 12.658 )

8. A titration of 60.0 mL of a 0.200M weak base, B:, solution using a 0.250M strong acid, HX, solution was performed. Given that the Kb= 1.2 x 10-6 calculate the equivalence point and the pH at:

(a) before any acid is added ( 10.69 )

(b) after 1/2 the volume to equivalence point is added ( 8.08 )

(d) after 1 1/2 the volume to equivalence point is added ( 1.342 )

9. a) A 1.00 L sample of an aqueous solution contains 0.200 mol of hypochlorous acid and 0.100 mol of sodium hypochlorite. Given that the Ka of hypochlorous acid is 3.0 x 10-8, calculate the pH of the solution. ( 7.22 )

b) Calculate the pH after the addition of 1.5 mL of 5.0M NaOH. ( 7.29 )

10. a) Calculate the pH of a 1.00 L solution which is 0.50 M lactic acid (Ka = 1.4 x 10-4) and 0.25 M sodium lactate. ( 3.55 )

b) Calculate the pH after the addition of 1.5 mL of 12M HCl. ( 3.50 )