Acids and Bases Written Response January 1999 4. Consider the salt sodium oxalate, Na2C2O4 . a) Write the dissociation equation for sodium oxalate. (1 mark) b) A 1.0M solution of sodium oxalate turns pink when a few drops of the indicator phenolphthalein are added. Write a hydrolysis equation and explain why this salt causes the indicator to change colour. (2 marks) c) Calculate the equilibrium constant for the hydrolysis in b). (1 mark) 5. Calculate the pH of 0.50M H3BO3. (4 marks) 6. A 25.0mL sample of Sr(OH)2 is titrated with a standardized solution of HCl to the equivalence point. a) Write the formula equation for the neutralization. (1 mark) b) Write the net ionic equation for the neutralization. (1 mark) c) What is meant by the term “standardized” solution? (1 mark) d) Define equivalence point. (1 mark)

April 1999 6. Consider the following amphiprotic anions reacting with each other:

HC6H5O7 -2 + HC2O4− ?

a) Complete the Brönsted-Lowry acid-base equilibrium for the predominant reaction. (1 mark) b) Does the equilibrium above favour reactants or products? Explain. (1 mark) 7. Calculate the pH of a 1.5M H2S solution. (4 marks) 8. Consider the following reaction:

2HCl (aq) + Ba(OH)2 (s) → BaCl2 (aq) + 2H2O (l) When 3.16g samples of Ba(OH)2 were titrated to the equivalence point with an HCl solution, the following data were recorded:

Using the data above, calculate the original [HCl]. (4 marks) June 1999 6. In aqueous solutions, H3O+

is the strongest acid present. This phenomenon is called the levelling effect. Explain why this occurs. (2 marks)

7. A 1.00M OCl− solution has an [OH−] of 5.75 ×10−4 M.

a) Calculate Kb for OCl−

. (3 marks) b) Calculate Ka for HOCl. (1 mark) 8. Calculate the mass of NaOH needed to prepare 2.0 L of a solution with a pH of 12.00.

(3 marks) August 1999 6. Consider a Brönsted-Lowry acid-base equation, where HNO2 is a reactant and H2PO4− is a product. a) Complete the following equation. (1 mark)

HNO2 + ________ ________ + H2PO4−

b) Identify the weaker base in the equilibrium in part a). (1 mark)

7. A chemist prepares a solution by dissolving the salt NaIO3 in water. a) Write the equation for the dissociation reaction that occurs. (1 mark) b) Write the equation for the hydrolysis reaction that occurs. (1 mark) c) Calculate the value of the equilibrium constant for the hydrolysis in part b). (1 mark)

8. Calculate the pH of a solution prepared by adding 15.0mL of 0.500M H2SO4 to 35.0mL of 0.750M NaOH. (4 marks)

January 2000 7. Define the term amphiprotic. Give an example of an ion which is amphiprotic.

(2 marks) Definition: _______________________________________________________ _________________________________________________________________

Example: ________________________________________________________ 8. A 0.0200M solution of methylamine, CH3NH2, has a pH = 11.40.　 Calculate the Kb for methylamine. (4 marks) 9. A titration was performed by adding 0.115M NaOH to a 25.00mL sample of H2SO4. Calculate the [H2SO4] from the following data. (3 marks)

April 2000 7. A sample of a weak acid was found to conduct an electric current better than a sample of a strong acid. Explain these results in terms of ion concentration. (2 marks) 8. Calculate the [OH-] of 0.10M NH3. (4 marks)

9. A titration was performed by adding 0.175M H2C2O4 to a 25.00mL sample of NaOH. The following data was collected:

a) Calculate the [NaOH]. (3 marks) June 2000 6. A 0.100M solution of an unknown weak acid, HX, has a pH = 1.414.　 What is the Ka for HX? (4 marks) 7. Consider the salt ammonium acetate, NH4CH3COO. a) Write the equation for the dissociation of NH4CH3COO. (1 mark) b) Write equations for the hydrolysis reactions which occur. (2 marks) c) Explain why a solution of NH4CH3COO has a pH = 7.00.　 Support your answer with calculations. (2 marks) August 2000 5. a) Define the term Bronsted-Lowry conjugate acid-base pair. (1 mark) b) Give an example of a conjugate acid-base pair. (1 mark) Acid: Base:

6. Consider the acids HCl and HF. a) Only one of the following reactions occurs. Complete the equation for the reaction which does occur. (1 mark)

i) HCl + F- ____________________________　

ii) HF + Cl- ____________________________ b) For the reaction that occurs, are reactants or products favoured? Explain. (1 mark) c) Explain why the other reaction will not occur. (1 mark) 7. Calculate the [H3O+] of 0.10M HNO2. (3 marks) 8. Write the formula equation and the net ionic equation for the reaction between 0.10M H2SO4 and 0.10M Sr(OH)2. (3 marks) Formula equation: Net ionic equation: January 2001 6. Consider the following Bronsted-Lowry equilibrium:

H2SO3 (aq) + HPO4 -2(aq) H2PO4- (aq) + HSO3- (aq)

a) Identify the two Bronsted-Lowry acids in the above equilibrium. (1 mark) b) Define the term conjugate acid. (1 mark)

7. A 250.0mL sample of HCl with a pH of 2.000 is completely neutralized with 0.200M NaOH. a) What volume of NaOH is required to reach the stoichiometric point? (4 marks) b) Write the net ionic equation for the above neutralization reaction. (1 mark) c) If the HCl were titrated with a 0.200M NH3(aq) instead of 0.200M NaOH, how would the volume of base required to reach the equivalence point compare with the volume calculated in part a) ? Explain your answer. (1 mark) 8. Consider the following equilibrium:

energy + 2H2O 　　 H　 3O+　 + OH-

a) Explain how pure water can have a pH = 7.30. 　 (2 marks) b) Calculate the value of Kw for the sample of water with a pH = 7.30. 　 (2 marks) April 2001 6.a) Write the equation for the predominant reaction of HC2O4

- with HSO3　　. (1 mark)

HC2O4

- + HSO3　　　　 __________ + ___________

B Identify a Br¿nsted-Lowry conjugate acid base pair from the above reaction. (1 mark) Acid:__________________________ Base:______________________________

7.a) In the space below, sketch the titration curve for the reaction when 0.10M HCl is added to 10.0mL of 0.10M NaOH. (3 marks)

b) Describe two changes in the titration curve that would result from using 0.10M CH3COOH in place of the HCl. (2 marks) i)___________________________________________________________ ii)___________________________________________________________ 8. Calculate the [OH-] in 0.50M NH3(aq). (5 marks) June 2001 6. Write the equation for the predominant reaction between HSO4　　and H2PO4

　　. (2 marks)

7. A 0.20M solution of a weak acid, HA, has a pH 1.32.　　 Use calculations and the table of “Relative Strengths of Brønsted-Lowry Acids and Bases” from the Data Booklet to determine the identity of the acid. (5 marks) The Acid is:________________________.

8. Calculate the pH of a solution prepared by mixing 15.0mL of 0.50M HCl with 35.0mL of 1.0M NaOH. (4 marks) 9. Identify a gas which causes acid rain, and write an equation showing this gas reacting with water. (2 marks) August 2001 6. Consider the 0.10M solutions of the following two acids:

a) What can you conclude about the acids that will explain these different pH values? (1 mark)

You can conclude that acid HA:_____________________________ You can conclude that acid HB:_____________________________ b) Compare the volume of 0.10M NaOH needed to neutralize equal volumes of

each of these acid samples. (1 mark)

7. Consider a 1.0M solution of NH4F.

a) Write both hydrolysis reactions that occur when NH4F is dissolved in water. (2 marks)

b) Will the above NH4F solution be acidic, basic, or neutral? Support your answer with calculations. (2 marks) 8. An indicator is often used during acid-base titrations. a) Define the term transition point for an indicator. (1 mark) b) Calculate the Ka value for methyl red. (1 mark) c)A mixture of indicators is made by combining equal amounts of methyl orange and bromthymol blue. Complete the following table, showing the colour of each indicator and the mixture at pH of 5 and pH of 9. (2 marks)

9. Will HC2O4 act predominantly as an acid or as a base in solution? Support your answer with calculations. (3 marks)

January 2002 7) The two reactants in an acid-base reaction are HNO2(aq) and HCO3

-(aq).

a) Write the equation for the above reaction. b) Define the term conjugate acid-base pair. c) Write the formulas for a conjugate acid-base pair for the above reaction. 8) At 10.0℃, Kw = 2.95 × 10-15 for pure water. a) Calculate the pH of water at 10.0℃. b) A mixture of the indicators phenolphthalein and bromcresol green is added to the water. What is the resulting colour of the mixture? Explain. 9) At a particular temperature a 1.0M H2S solution has a pH = 3.75. Calculate the value of Ka at this temperature. 10) What is the main function of a buffer solution? April 2002 6) The ion H2PO4

- is an amphiprotic anion. a) Define the term amphiprotic. b) Write the balanced equation for the reaction when H2PO4

- reacts with HF. 7) Write an equation for a reaction in which H2O acts only as a Bronsted-Lowry base. 8) Calculate the pH of 0.25M Sr(OH)2. 9) Calculate the pH of 0.25 M NH4Cl.

10) A 0.1M unknown acid is titrated with 0.10M NaOH and the following titration curve results: a) Choose a suitable indicator (other than phenolphthalein) and give a reason for your choice. b) Is the unknown acid weak or strong? Explain. June 2002 7) The cyanide ion, CN-, is a Bronsted-Lowry base. a) Define Bronsted-Lowry base. b) Write the equation representing the reaction of CN- with water. c) Identify a conjugate pair in b) above. 8) Write an equation to show the ionization of water. 9) Calculate the pH of 1.50M NH3.

10) Consider the following buffer equilibrium:

Using Le Châtelier’s Principle, explain what happens to the pH of the buffer solution When a small amount of NaOH is added. August 2002 7) Consider the following equilibria:

І

CH3COOH + OCN- ⇄ HOCN + CH3COO-

І І.

CH3COOH + ClO- ⇄ HClO + CH3COO-

a) In equation І above, the reactants are favoured. Identify the stronger acid. b) In equation І І above, the products are favoured. Identify the stronger acid. c) Consider the following reaction: HOCN + ClO- ⇄ OCN- + HClO Does this reaction favour reactants or products? Explain. 8) At 60℃, the pH = 6.51 for pure water. Determine the value of Kw at this temperature. 9) Calculate the pH of 0.35M H2CO3.

10) A strong acid – strong base titration has a pH = 7.0 at the equivalence point. A weak acid-strong base titration has a pH > 7.0 at the equivalence point. a) What causes the difference in this pH values? b) Select one indicator which could be used for both titrations. January 2003 7) Write a chemical reaction showing an amphiprotic anion reacting as a base in water. 8) Calculate the pOH of 0.25M Sr(OH)2. 9) A 2.00M diprotic acid has a pH of 0.50. Calculate its Ka value. 10) The following two experiments were conducted: Titration A : A strong acid was titrated with a strong base. Titration B : A weak acid was titrated with a strong base. a) How does the pH at the equivalence point of Titration B compare with the pH at the equivalence point of Titration A? b) Explain your answer to a). April 2003 6) An acid-base reaction occurs between HSO3

- and IO3-.

a) Write the equation for the equilibrium that results. b) Identify one conjugate acid-base pair in the reaction. c) State whether reactants or products are favoured, and explain how you arrived at your answer.

7) At 10℃, Kw = 2.95 × 10-15. a) Determine the pH of water at 10℃. b) State whether water at this temperature is acidic, basic or neutral, and explain. 8) Calculate the pH of 0.50M H2S. June 2003 6) a) Write an equation to represent the predominant reaction when HC2O4

- is mixed with HSO4-.

b) Justify your statement by comparing Ka values. c) Identify a conjugate acid-base pair. d) Predict whether the equilibrium will favour the formation of reactants or products. Explain. 7) Write an equation representing the ionization of water and state both ion concentrations that exist for pure water to have a pH = 7.20. 8) Calculate the pH of 0.25M NaHCO3, a basic salt. 9) Explain why the action of a buffer solution is limited. August 2003 6) Write the net ionic equation for the acid-base reaction that occurs between NaCN(aq) and NH4Cl(aq). 7) Define the term amphiprotic and give an example of an amphiprotic anion.

8) At 20℃, the ionization constant of water (Kw) is 6.76 × 10-15. Calculate the [H3O+] of water at 20℃. 9) Calculate the pH of 0.50M NaF. 10) Outline a procedure to prepare a buffer solution.

Acids and Bases Written Response Answers

January 1999

April 1999

June 1999

8.

August 1999 6a)

b)

7a)

b)

c)

8)

January 2000 7)

8)

9)

April 2000 7)

8)

9a)

b)

June 2000 6)

7a)

b)

c)

August 2000 5a)

b)

6a)

b)

c)

7)

8)

January 2001 6a)

b)

7a)

b)

c)

8a)

b)

April 2001 6a)

b)

7a)

b)

8)

June 2001 6)

7)

8)

9)

August 2001 6a)

b)

7a)

b)

8a)

b)

c)

9)

January 2002 7) a) HNO2(aq) + HCO3

-(aq) ⇄ H2CO3(aq) + NO2

-(aq)

b) A conjugate acid-base pair are two species whose formulas differ by a proton. c) HNO2 and NO2

- OR H2CO3 and HCO3-

8) a) b) Resulting colour : Blue at pH = 7.265. Explanation : Bromcresol green is blue at pH 7.265; phenolphthalein is colourless. 9)

10) The main function of a buffer solution is to resist changes in pH. April 2002 6) a) Amphiprotic describes the ability to act as an acid in one reaction and as a base in a different reaction. b) H2PO4

- + HF ⇄ H3PO4 + F-

7) HSO3

- + H2O ⇄ H3O+ + SO32-

8) [OH-] = 0.50M pOH = 0.30 pH = 13.70 9) 10) a) Suitable Indicator: Thymol blue Reason : The transition point is close to the equivalence point. b) Unknown Acid: Weak *Explanation : - the shape of the curve is characteristic of a weak acid being titrarted with a strong base - the initial pH is greater than 1.0 - the pH at the equivalence point is greater than 7.0

June 2002 7) a) Bronsted-Lowry base is a proton acceptor. b) CN- + H2O ⇄ HCN + OH-

c) CN- and HCN OR H2O and OH-

8) 2H2O(l) ⇄ H3O+

(aq) + OH-(aq)

9) 10) NaOH causes [H3O+] to decrease. The equilibrium shifts right. The pH almost returns to the former value.

August 2002 7) a) HOCN b) CH3COOH c) Products are favoured because HOCN is a stronger acid than HClO. 8) 9) 10) a) A strong acid-strong base titration produces a neutral salt while a weak acid-strong base titration produces a basic salt. b) Phenolphthalein

January 2003 7) HPO4

2- + H2O ⇄ H2PO4- + OH-

8) [OH-] = 2(0.25M) = 0.50M pOH = - log(0.50) = 0.30 9) 10) a) The pH at the equivalence point of Titration A = 7.0. The pH at the equivalence point of Titration B > 7.0. b) Neutral salt formed in titration A, a basic salt is formed in titration B. April 2003 6) a) HSO3

- + IO3- ⇄ SO3

2- + HIO3 b) HSO3

- and SO32- OR IO3

- and HIO3 c) Reactants are favoured. HSO3

- is a weaker acid than HIO3 OR IO3- is a weaker base than SO3

2-

7) a) Kw = 2.95 × 10-15 = [H3O+][OH-] Since [H3O+] = [OH-],

[H3O+]2 = 2.95 × 10-15

[H3O+] = 5.43 × 10-8

pH = 7.265

b) Since [H3O+] = [OH-], the water is neutral.

8)

June 2003 6) a) HC2O4

- + HSO4- ⇄ H2C2O4 + SO4

2-

b) KaHSO4

- > KaHC2O4-

c) HSO4

- and SO42- OR HC2O4

- and H2C2O4 d) Reactants are favuored since KaHSO4

- < KaH2C2O4 7) H2O(l) + H2O(l) ⇄ H3O+

(aq) + OH-(aq)

Since pH = 7.20, [H2O+] = 6.3 × 10-8 M

[H3O+] = [OH-] = 6.3 × 10-8 M 8) 9) Buffer action depends on the presence of - sufficient amounts - of weak acid and conjugate base in the buffer solution. August 2003 6) CN-

(aq) + NH4+

(aq) ⇄ HCN(aq) + NH3(aq) 7) Definition : Amphiprotic describes a substance that can act as either an acid or a base. Example : HCO3

-

8) Kw = [H3O+][OH-] = 6.76 × 10-15

Since [H3O+] = [OH-], [H3O+]2 = 6.76 ×10-15

[H3O+] = 8.22 ×10-8 M 9) 10) Prepare an aqueous mixture that contains: 1. a weak acid 2. a salt of its conjugate base 3. the acid and salt in sufficient concentrations