Acids, Bases and Salts Unit Plan

Period/Topic Worksheets Quiz

1. Properties of Acids, Bases & Salts WS 1

2. Arrhenius, Bronsted Acids, Ka and Strength. WS 2 1

3. Arrhenius, Bronsted Bases, Kb and Strength WS 3

4. Acid & Base Reactions. Amphiprotic. Acid Chart. WS 4 2

5. Leveling effect, Anhydrides and Relationships. WS 5

6. Hydrolysis of Salts. Quiz. WS 6 3

7. Acid, Base & Salt Reactions. Hydrolysis. WS 7

8. Yamada’s Indicator Lab. Hydrolysis. WS 8 4

9. Ionization of Water, [H⁺] & [OH^-], pH scale. WS 9

10. pH Calculations for Weak Acids. WS 10 5

11. Ka from pH for Weak Acids. WS 11

12. Indicators Lab.

13. Kbs from Kas for Weak Bases. WS 12 6

14. pH for Weak Bases pH [H⁺] [OH^-] Relationships. WS 13

15. Amphiprotic Ions- Kas and Kbs. WS 14 7

16. Titration Lab. Primary Standards. Acids Midterm Practice Test

17. Titration Lab

18. Buffers & Indicators WS 15 8

19. Titration Curves. WS 16 9/10

20. Review # 1 Web Site Review Practice Test 1

21. Review # 2 Practice Test 2

22. Test

Worksheet # 1 Properties of Acids and Bases

1. Add 1 drop of each solution to 1 drop of the acid-base indicator in a spot plate. Record the colour in the data table below. Describe each solution as an acid or base in the space provided. Write the acid colour and base colour in the table below.

Indicator Phenolphthalein Litmus BromothymolBlue Acid/Base

Solution

HCl

NaOH

Vinegar

Ammonia

(NH₃)

Lemon Juice

Seven-up

Baking Soda

(NaHCO₃)

Indicator Acid Colour Base Colour

Phenolphthalein

Litmus

Bromothymol Blue

Wash and dry your spot plate before going on to step 2.

2. Wear safety goggles for this experiment. Pour approximately 50 mL of 1 M HCl into a fleaker. Add one level spoonful of Ca and cover with a plastic funnel. After 1 minute and not before light the top of the funnel using a match. Write the equation for the reaction below.

Wash and dry your fleaker before going on to step 3.

3. Taste a lemon and describe the taste in one word

4. Taste some baking soda and describe the taste in one word.

6. Test two drops of HCl for conductivity in a spot plate. Result:

Write an equation that accounts for the conductivity of HCl.

7. Test two drops of NaOH for conductivity in a spot plate. Result:

Write an equation that accounts for the conductivity of NaOH (dissociation).

Clean, dry and put away the spot plate

8. List five properties of acids that are in your textbook.

9. List five properties of bases that are in your textbook.

10. Make some notes on the commercial acids: HCl and H₂SO₄.

HCl

H₂SO₄

11. Make some notes on the commercial base NaOH.

12. Describe the difference between a concentrated and dilute acid (hint: concentration refers to the molarity). Describe their relative conductivities.

13. Describe the difference between a strong and weak acid. Use two examples and write equations to support your answer. Describe their relative conductivities.

14. Describe a situation where a strong acid would have the same conductivity as a weak acid (hint: think about concentration).

Worksheet # 2 Conjugate Acid-Base Pairs

Complete each acid reaction. Label each reactant and product as an acid or base. The first on is done for you.

1. HCN + H₂O ⇄ H₃O⁺ + CN^-

Acid Base Acid Base

2. H₃C₆O₇ + H₂O ⇄

3. H₃PO₄ + H₂O ⇄

4. HF + H₂O ⇄

5. H₂CO₃ + H₂O ⇄

6. NH₄⁺ + H₂O ⇄

7. CH₃COOH + H₂O ⇄

8. HCl + H₂O ⇄

9. HNO₃ + H₂O ⇄

Write the equilibrium expression (Ka) for the first seven above reactions. The first one is done for you.

10. Ka = [H₃O⁺][CN^-] 14. Ka =

[HCN]

11. Ka = 15. Ka =

12. Ka = 16. Ka =

13. Ka =

17. Which acids are strong?

18. What does the term strong acid mean?

19. Why is it impossible to write an equilibrium expression for a strong acid?

20. Which acids are weak?

23. What does the term weak acid mean?

24. Explain the difference between a strong and weak acid in terms of electrical conductivity.

Acid Conjugate Base Base Conjugate Acid

14. HNO₂ 15. HCOO^-

16. HSO₃^- 17. IO₃^-

18. H₂O₂ 19. NH₃

20. HS^- 21. CH₃COO^-

22. H₂O 23. H₂O

Define:

22. Bronsted acid-

23. Bronsted base-

24. Arrhenius acid-

25. Arrhenius base-

26. List the six strong acids.

27. Rank the acids in order of decreasing strength. HCl H₂S H₃PO₄ H₂CO₃ HF HSO₄

28. What would you rather drink vinegar or hydrochloric acid? Explain.

Making a Universal Indicator Lab Activity

Mix the following indicators in a 50 mL beaker. Stir with an eyedropper.

Yamada’s Universal Indicator

5 drops thymol blue

6 drops methyl orange

5 drops phenolphthalein

10 drops bromothymol blue

20 drops of water

Part 1. In a spot plate add two drops of each buffer solution to a cell. Add one drop of Yamada’s indicator to each. Record each colour on another lab sheet by colouring the cell the same colour. Make sure you are accurate because you will use this information for future labs and projects.

<---------- Acid Strength Increases ------ Neutral ----Base Strength Increases ------->

pH = 1	pH = 3	pH = 5	pH = 7	pH = 9	pH =11	pH = 13

Part 2. Test a drop of HCl, CH₃COOH, NaOH, NH₃, NaHCO₃, H₂CO₃ and NaCl solution for conductivity. Test with your Universal Indicator. Record the pH of each. Test with your Universal Indicator. Explain your results with what you know about acids and bases. Classify each as a strong or weak acid or base or neutral, acidic, or basic salt. Write an equation for each to show how they ionize in water using the Bronsted (Chemistry 12) definition of an acid.

Wash and dry your chem plate

Wash and return your eyedropper.

Wash and return your beaker.

Wash your hands.

Results

Compound Conductivity pH Classification

HCl

CH₃COOH

NaOH

NH₃

NaHCO₃

H₂CO₃

NaCl

Worksheet # 3 Conjugate Acid-Base Pairs

Complete each reaction. Label each reactant and product as an acid or base.

1. HCN + H₂O ⇄ H₃O⁺ + CN^-

Acid Base Acid Base

2. HCl + H₂O ⇄

3. HF + H₂O ⇄

4. F^- + H₂O ⇄

5. HSO₄^- + H₂O ⇄

(acid)

6. NH₄⁺ + H₂O ⇄

7. HPO₄^2- + H₂O ⇄

(base)

Acid Conjugate Base Base Conjugate Acid

8. HCO₃^- CO₃^2- 9. CH₃COO^- CH₃COOH

10. HPO₄^2- 11. IO₃^-

12. H₂O 13. NH₂^-

14. HS^- 15. C₂H₅SO₇^3-

16. Circle the strong bases.

Fe(OH)₃ NaOH CsOH KOH

Zn(OH)₂ Sr(OH)₂ Ba(OH)₂ Ca(OH)₂

17. Rank the following acids from strongest to weakest.

H₂S CH₃COOH H₂PO₄^- HI HCl HF

18. Rank the following bases from the strongest to weakest.

H₂O F^- NH₃ SO₃^2- HSO₃^- NaOH

19. i) Write the reaction of H₃BO₃ with water (remove one H⁺ only because it is a weak acid).

ii) Write the K_a expression for the above.

iii) What is the ionization constant for the acid (use your table). K_a =

20. List six strong acids.

21. List six strong bases.

22. List six weak acids in order of decreasing strength (use your acid/base table).

23. List six weak bases in order of decreasing strength (use your acid/base table).

Worksheet # 4 Using Acid Strength Tables

Acid-base reactions can be considered to be a competition for protons. A stronger acid can cause a weaker acid to act like a base. Label the acids and bases. Complete the reaction. State if the reactants or products are favoured.

1. HSO₄^- + HPO₄^2- ⇄

2. HCN + H₂O ⇄

3. HCO₃^- + H₂S ⇄

4. HPO₄^2- + NH₄⁺⇄

5. NH₃ + H₂O ⇄

6. H₂PO₄^1-+NH₃⇄

7. HCO₃^- + HF ⇄

8. Complete each equation and indicate if reactants or products are favoured. Label each acid or base.

HSO₄^- + HCO₃^- ⇄

H₂PO₄^-+ HC0₃^- ⇄

HS0₃^-+ HPO₄^2- ⇄

NH₃+ HC₂O₄^- ⇄

9.Explain why HF_(aq)is a better conductor than HCN_(aq).

10_.Which is a stronger acid in water, HCl or HI? Explain!

11. State the important ion produced by an acid and a base.

12. Which is the stronger base? Which produces the least OH^-? F^-or CO₃^-2

13. Define a Bronsted/Lowry acid and base.

14. Define an Arrhenius acid and base.

15. Complete each reaction and write the equilibrium expression.

HF + H₂O ⇄ K_a=

F^- + H₂O ⇄ K_b=

16. H₂SO₄ + NaOH →

17. Define conjugate pairs.

18. Give conjugate acids for: HS^-, NH₃, HPO₄^-2, OH^-, H₂O, NH₃, CO₃^-2

19^.Give conjugate bases for: NH₄^+, HF, H₂PO₄^-, H₃O⁺, OH^-, HCO₃^-, H₂O

Worksheet # 5 Acid and Basic Anhydrides

1. What is the strongest acid that can exist in water? Write an equation to show how a stronger acid would be reduced in strength by the leveling effect of water.

2. What is the strongest base that can exist in water? Write an equation to show how a stronger base would be reduced in strength by the leveling effect of water.

3. List three strong acids and three strong bases.

4. Rank the acids in decreasing strength:

HClO₄ K_a is very large HClO₃ K_a=1.2x10^-2

HClO₂ K_a=8.0x10^-5 HClO K_a=4.4x10^-8

5. For an oxy acid what is the relationship between the number of O’s and acid strength? (Compare H₂S0₄and H₂S0₃)

6. Which acid is stronger? HI0₃ or HIO₂

7. Which produces more H₃0⁺? H₂CO₃or HS0₄^-

8. Which produces more OH^-? F^-or HC0₃^-

9. Which conducts better NH₃ or NaOH (both .1M)? Why?

10. Which conducts better HF or HCN (both .1M)? Why?

11. Compare and contrast a strong and weak acid in terms of degree of ionization, size of k_a, conductivity, and concentration of H⁺.

Classify each formula as an acid anhydride, basic anhydride, strong acid, weak acid, strong, or weak base. For each formula write an equation to show how it reacts with water. For anhydrides write two equations.

Formula Classification Reaction

12. Na₂O

13. CaO

14. SO₃

15. CO₂

16. SO₂

17. HCl

18. NH₃

19. NaOH

20. HF

21. H₃PO₄

Worksheet # 6 Hydrolysis of Salts and Reactions of Acids and Bases

Describe each as an acid, base, neutral salt, acidic salt, or basic salt. For each salt write a parent acid-base formation equation, dissociation equation, and hydrolysis equation (only for acidic and basic salts). For acids and bases write an equation to show how each reacts with water.

1. NH₃

2. KCl

3. HNO₃

4. NaHCO₃

5. RbOH

6. AlCl₃

7. H₂C₂O₄

8. NaC₆H₅O

9. Co(NO₃)₃

10. Na₂CO₃

Worksheet # 7 Hydrolysis of Salts and Reactions of Acids and Bases

Describe each as an acid, base, neutral salt, acidic salt, or basic salt. For each salt write a dissociation equation and hydrolysis equation (only for acidic and basic salts). For acids and bases write an equation to show how each reacts with water.

1. NH₃

2. NaCl

3. HCl

4. NaCN

5. NaOH

6. FeCl₃

7. HF

8. LiHCO₃

9. Fe(NO₃)₃

10. MgCO₃

11. H₂S

12. HF

13. CaI₂

14. Mg(OH)₂

15. Ba(OH)₂

16. Describe why Tums (CaCO₃) neutralizes stomach acid.

17. Describe why Mg(OH)₂ is used in Milk of Magnesia as an antacid instead of NaOH.

Worksheet # 8 Yamada’s Indicator Activity

Acid, Base and Salt Lab

Purpose:

1) To use Yamada’s Indicator to determine the pH of various acids, bases and salts.

2) To classify compounds as strong acids, weak acids, strong bases, weak bases, neutral salts, acid anhydrides, and basic anhydrides.

3) To write reactions for each compound to show how each ionizes, hydrolyzes or reacts with water.

Procedure:

1) To a cell in a spot plate add one drop of solution or a very tiny amount of solid. Write the formula of the compound in the data table.

2) Add two drops of Yamada’s Indicator. Record the pH of the compound.

3) Classify the compound as a strong acid, weak acid, strong base, weak base, neutral salt, acid anhydride, or basic anhydride. Use the formula of the compound as well as the pH.

4) Write an equation to show the reaction of anhydrides with water, the hydrolysis of salts, or the ionization of acids or bases.

Data

1. Formula of compound

Classification

Reaction or reactions

2. Formula of compound

Classification

Reaction or reactions

3. Formula of compound

Classification

Reaction or reactions

4. Formula of compound

Classification

Reaction or reactions

5. Formula of compound

Classification

Reaction or reactions

6. Formula of compound

Classification

Reaction or reactions

7. Formula of compound

Classification

Reaction or reactions

8. Formula of compound

Classification

Reaction or reactions

9. Formula of compound

Classification

Reaction or reactions

10. Formula of compound

Classification

Reaction or reactions

11. Formula of compound

Classification

Reaction or reactions

12. Formula of compound

Classification

Reaction or reactions

13. Formula of compound

Classification

Reaction or reactions

14. Formula of compound

Classification

Reaction or reactions

15. Formula of compound

Classification

Reaction or reactions

16. Formula of compound

Classification

Reaction or reactions

17. Formula of compound

Classification

Reaction or reactions

Worksheet # 9 pH and pOH Calculations

Complete the chart:

	[H⁺]	[OH^-]	pH	pOH	Acid/base/neutral
1.	7.00 x 10^-3M
2.		8.75 x 10^-2M
3.			7.33
4.				4.00
5.					Neutral (2 sig figs)
6.				10.7
7.			2.553
8.	5.0 x 10^-10M
9.		4.7 x 10^-10M

10. Calculate the [H⁺], [OH^-], pH and pOH for a 0.20 M Ba(OH)₂ solution.

11. Calculate the [H⁺], [OH^-], pH and pOH for a 0.030 M HCl solution.

12. Calculate the [H⁺], [OH^-], pH and pOH for a 0.20 M NaOH solution.

13. 300.0 mL of 0.20 M HCl is added to 500.0 mL of water, calculate the pH of the solution.

14. 200.0 mL of 0.020 M HCl is diluted to a final volume of 500.0 mL with water, calculate the pH.

15. 150.0 mL of 0.40 M Ba(OH)₂ is placed in a 500.0 mL volumetric flask and filled to the mark with water, calculate the pH of the solution.

16. 250.0 mL of 0.20M Sr(OH)₂ is diluted by adding 350.0 mL of water, calculate the pH of the solution.

17. Calculate the pH of a 0.40 solution of Ba(OH)₂when 25.0 mL is added to 25.0 mL of water.

Worksheet # 10 pH Calculations for Weak Acids

1. Calculate the [H⁺], [OH^-], pH, and pOH for 0.20 M HCN.

[H⁺] = [OH^-] = pH = pOH =

2. Calculate the [H⁺], [OH^-], pH, and pOH for 2.20 M HF.

[H⁺] = [OH^-] = pH = pOH =

3. Calculate the [H⁺], [OH^-], pH, and pOH for 0.805 M CH₃COOH.

[H⁺] = [OH^-] = pH = pOH =

4. Calculate the [H⁺], [OH^-], pH, and pOH for 1.65 M H₃BO₃.

[H⁺] = [OH^-] = pH = pOH =

5. Calculate the pH of a saturated solution of Mg(OH)₂.

6. Calculate the pH of a 0.200 M weak diprotic acid with a Ka = 1.8 x 10^-6.

7. 350.0 mL of 0.20M Sr(OH)₂ is diluted by adding 450.0 mL of water, calculate the pH of the solution.

Worksheet # 11 pH Calculations for Weak Acids

1. The pH of 0.20 M HCN is 5.00. Calculate the Ka for HCN. Compare your calculated value with that in the table.

2. The pH of 2.20 M HF is 1.56. Calculate the Ka for HF. Compare your calculated value with that in the table.

3. The pH of 0.805 M CH₃COOH is 2.42. Calculate the Ka for CH₃COOH. Compare your calculated value with that in the table.

4. The pH of 1.65 M H₃BO₃ is 4.46. Calculate the Ka for H₃BO₃. Compare your calculated value with that in the table.

5. The pH of a 0.10 M diprotic acid is 3.683, calculate the Ka and identify the acid.

6. The pH of 0.20 M NH₃ is 11.227; calculate the Kb of the Base.

7. The pH of 0.40 M NaCN is 11.456; calculate the Kb for the basic salt. Start by writing an equation and an ICE chart.

8. The pH of a 0.10 M triprotic acid is 5.068, calculate the Ka and identify the acid.

9. How many grams of CH₃COOH are dissolved in 2.00 L of a solution with pH = 2.45?

Use questions 1 to 4 from last assignment to mark questions 1 to 4.

Worksheet # 12 K_b For Weak Bases

Determine the K_b for each weak base. Write the ionization reaction for each. Remember that K_w = K_a •K_b (the acid and base must be conjugates). Find the base on the right side of the acid table and use the K_a values that correspond. Be careful with amphiprotic anions! The first one is done for you.

1. NaNO₂ (the basic ion is NO₂^-)

Kb(NO₂^-) = Kw = 1.0 x 10^-14 = 2.2 x 10^-11

Ka(HNO₂) 4.6 x 10^-4

2. KCH₃COO (the basic ion is CH₃COO^-)

3. NaHCO₃

4. NH₃

5. NaCN

6. Li₂HPO₄

7. KH₂PO₄

8. K₂CO₃

9. Calculate the [H⁺], [OH^-], pH, and pOH for 0.20 M H₂CO₃.

[H⁺] = [OH^-] = pH = pOH =

10. The pH of 0.20 M H₂CO₃ is 3.53. Calculate the Ka for H₂CO₃. Compare your calculated value with that in the table.

11. Calculate the [H⁺], [OH^-], pH, and pOH for 0.10 M CH₃COOH.

[H⁺] = [OH^-] = pH = pOH =

12. The pH of 0.10 M CH₃COOH is 2.87. Calculate the Ka for CH₃COOH. Compare your calculated value with that in the table.

13. 200.0 mL of 0.120 M H₂SO₄ reacts with 400.0 mL of 0.140 M NaOH. Calculate the pH of the resulting solution.

Worksheet # 13 Acid and Base pH Calculations

For each weak bases calculate the [OH^-], [H⁺], pOH and pH. Remember that you need to calculate Kb first.

1. 0.20 M CN^-

2. 0.010 M NaHS (the basic ion is HS^-)

3. 0.067 M KCH₃COO

4. 0.40 M KHCO₃

5. 0.60 M NH₃

6. If the pH of a 0.10 M weak acid H₂X is 3.683, calculate the Ka for the acid and identify the acid using your acid chart.

7. Calculate the [H⁺], [OH^-], pH, and pOH for 0.80 M H₃BO₃.

[H⁺] = [OH^-] = pH = pOH =

8. Calculate the [H⁺], [OH^-], pH, and pOH for 0.25 M H₂CO₃.

[H⁺] = [OH^-] = pH = pOH =

9. The pH of 1.65 M H₃BO₃ is 4.46. Calculate the Ka for H₃BO₃. Compare your calculated value with that in the table.

10. The pH of 0.65 M NaX is 12.46. Calculate the Kb for NaX.

11. Consider the following reaction: 2HCl + Ba(OH)₂ → BaCl₂ + 2H₂O

When 3.16g samples of Ba(OH)₂ were titrated to the equivalence point with an HCl solution, the following data was recorded.

Trial Volume of HCl added

#1 37.80 mL

#2 35.49 mL

#3 35.51 mL Calculate the original [HCl]

12. Calculate the volume of 0.200M H₂SO₄ required to neutralize 25.0 ml of 0.100M NaOH.

13. 25.0 mL of .200 M HCl is mixed with 50.0 mL 0.100 M NaOH, calculate the pH of he resulting solution.

14. 10.0 mL of 0.200 M H₂SO₄ is mixed with 25.0 mL 0.200 M NaOH, calculate the pH of the resulting solution.

15. 125.0 mL of .200 M HCl is mixed with 350.0 mL 0.100 M NaOH, calculate the pH of the resulting solution.

16. Define standard solution and describe two ways to standardize a solution.

17. What is the [H₃O⁺] in a solution formed by adding 60.0 mL of water to 40.0 mL of 0.040 M KOH solution?

Worksheet # 14 Amphiprotic Ions and Review

1. List the properties of acids/bases.

2. Define the following:

Arhenius strong acid

Arhenius weak base

Bronsted strong acid

Bronsted weak base

Conjugate pair

Amphiprotic

Standard solution.

3. Show by calculation if the following amphiprotic ions are acids or bases:

HCO₃^-

H₂PO₄^-

HPO₄^2-

4. What is the strongest base in water? What is the strongest acid in water? Write equations to explain your answer.

5. Match each equation:

Acid/base complete HCl + NaOH → NaCl + HOH

Acid/base net ionic F^- + HOH → HF + OH^-

Solubility product H⁺ + OH^- → HOH

Hydrolysis AgCl_(s) → Ag⁺ + Cl^-

Acid/Base formula H₂0 → H⁺ + OH^-

Ionization of water H⁺ + Cl^-+ Na⁺ + OH^-→ Na⁺+ Cl^- + H₂O

6. HCl and HF. Describe each acid as:

a) strong or weak

b) high or low ionization

c) large or small Ka

d) good or poor conductor

e) strong or weak electrolyte

7. Out of 0.2 M HCl and 1.0 M HF, which is the most concentrated?

Which is the strongest acid?

8. Label the scale as strong/weak acid and strong/weak base.

|________________________|_________________________|__

pH 0 7 14

9. Which ions are amphiprotic?

HPO₄^2- HCl F^- HS^- H₂S H₂O

10. Write the net ionic equation between any strong acid and strong base.

11. Write the ionization equation for water.

12. Write the Kw expression.

13. H₂SO₃ + HS^- ⇄ H₂S + HSO₃^-

a) Are the reactants or products favoured?

b) Is the Keq large, small or about 1?

Determine the pH Write equations for each first!

14. .20M HCl pH=?

15. 0.20M Ba(OH)₂ pH=?

16. 0.20M H₂CO₃ pH=?

17. 0.40M KHCO₃ pH=?

18. The pH increases by 2 units. How does [H⁺] change?

19. The pH decreases by 1 unit. How does [H⁺] change?

20. a) For distilled water : pH= pOH= [H⁺]= [OH^-]=

b) For 1M HCl: pH= pOH= [H⁺]= [OH^-]=

c) For 1M NaOH: pH= pOH= [H⁺]= [OH^-]=

21. The pH of 0.20 M NaX is 12.50; calculate the Kb.

22. The pH of 0.2 M HX is 4.5; calculate the Ka.

23. 100.0 mL of 0.200 M NaOH is mixed with 100.0 mL of 0.180 M HCl. Calculate the pH of the resulting solution.

24. How many grams of NaHCO₃ are required to make 100.0 mL of 0.200M solution?

25. What volume of 0.200 M NaOH is required to neutralize 25.0 mL of 0.150 M H₂SO₄?

26. In a titration 25.0 mL of .200M H₂SO₄ is required to neutralize 10.0 mL NaOH. Calculate the concentration of the base.

27. Calculate the concentration of a solution of NaCl made by dissolving 50.0 g in 250.0 mL of water.

28. SO_3(g) + H₂O_(g) ⇄ H₂SO_4(l)

Equilibrium concentrations are found to be:

[SO₃] = 0.400 M [ H₂O] = 0.480 M [H₂SO₄] = 0.600 M

Calculate the value of the equilibrium constant.

29. 4.00 moles of SO₂ and 5.00 moles O₂ are placed in a 2.00 L container at 200º C and allowed to reach equilibrium. If the equilibrium concentration of O₂ is

2.00 M, calculate the Keq.

2SO_2(g) + O_2(g) ⇄ 2SO_3(g)

Worksheet # 15 Buffers and indicators

Buffers

1. Definition

Acid Conjugate Base Salt

HCN

KHCO₃

NH₃

NaCH₃COO

HC₂O₄^-

3. Write an equation for the first three buffer systems above.

4. Which buffer could have a pH of 4.0? Which buffer could have a pH of 10.0?

a) HCl & NaCl b) HF & NaF c) NH₃ & NH₄Cl

5. Predict how the buffer of pH = 9.00 will change. Your possible answers are 9.00, 8.98, 9.01, 2.00, and 13.00

Final pH

a) 2 drops of 0.10 M HCl are added

b) 1 drop of 0.10 M NaOH is added

c) 10 mL of 0.10 M HCl are added

6. Write an equation for the carbonic acid, sodium hydrogencarbonate buffer system. A few drops of HCl are added. Describe the shift and each concentration change.

Equation:

Shift [H⁺] = [H₂CO₃] = [HCO₃^-] =

Indicators

1. Definition

2. Equilibrium equation

3. Colors for methyl orange HInd Ind^-

4. Compare the relative sizes of [HInd] and [Ind^-] at the following pH’s for methyl orange.

Color Relationship

pH = 2.0

pH = 3.7

pH = 5.0

5. HCl is added to methyl orange, describe if each increases or decreases.

[H⁺]

[HInd]

[Ind^-]

Color Change

6. NaOH is added to methyl orange, describe if each increases or decreases.

[H⁺]

[HInd]

[Ind^-]

Color Change

7. State two equations that are true at the transition point of an indicator.

8. What indicator has a Ka = 4 x 10^-8

9. What is the Ka for methyl orange?

10. A solution is pink in phenolphthalein and colorless in thymolphthalein. What is the pH of the solution?

11. A solution is blue in bromothymol blue, red in phenol red, and yellow in thymol blue. What is the pH of the solution?

Worksheet # 16 Titration Curves

Choose an indicator and describe the approximate pH of the equivalence point for each titration. Complete each reaction.

pH Indicator

1. HCl + NaOH →

2. HF + RbOH →

3. HI + Ba(OH)₂ →

4. HCN + KOH →

5. HClO₄ + NH₃ →

6. CH₃COOH + LiOH →

7. Calculate the Ka of bromothymol blue.

8. An indicator has a Ka = 1 x 10^-10, determine the indicator.

9. Calculate the Ka of methyl orange.

10. An indicator has a Ka = 6.3 x 10^-13, determine the indicator.

11. Explain the difference between an equivalence point and a transition point.

Draw a titration curve for each of the following.

12. Adding 100 ml 1.0 M NaOH to 13. Adding 100 ml 1.0 M NaOH to

50 mL 1.0 M HCl 50 mL 1.0 M HCN

Volume of base added Volume of base added

14. Adding 100 ml 0.10 M HCl 15. Adding 100 ml .10 M HCl to 50 mL 0.10M NH₃ to 50 mL 0.10 M NaOH

pH pH

Volume of HCl added Volume of HCl added

Acids Unit Midterm Practice Test

1. Consider the following:

I H₂CO₃ + F^- D HCO₃^- + HF

II HCO₃^- + HC₂O₄^- D H₂CO₃ + C₂O₄^2-

III HCO₃^- + H₂C₆H₆O₇^- D H₂CO₃ + HC₆H₅O₇^2-

The HCO₃^- is a base in

A. I only

B. I and II only

C. II and III only

D. I, II, and III

2. Consider the following equilibrium for an indicator:

HInd + H₂O D Ind^- + H₃O⁺

When a few drops of indicator methyl red are added to 1.0 M HCl, the colour of the resulting solution is

A. red and the products are favoured

B. red and the reactants are favoured

C. yellow and the products are favoured

D. yellow and the reactants are favoured

3. The volume of 0.200 M Sr(OH)₂ needed to neutralize 50.0 mL of 0.200 M HI is

A. 10.0 mL

B. 25.0 mL

C. 50.0 mL

D. 100.0 mL

4. The pOH of 0.050 M HCl is

A. 0.050

B. 1.30

C. 12.70

D. 13.70

5. The volume of 0.450 M HCl needed to neutralize 40.0 mL of 0.450 M Sr(OH)₂ is

A. 18.0 mL

B. 20.0 mL

C. 40.0 mL

D. 80.0 mL

6. Consider the following

I H₃PO₄ II H₂PO₄^- III HPO₄^2- IV PO₄^3-

Which of the above solutions are amphiprotic?

A. I and II only

B. II and III only

C. I, II, and III only

D. II, III, and IV only

7. Which of the following solutions will have the largest [H₃O⁺]?

A. 1.0 M HNO₂

B. 1.0 M HBO₃

C. 1.0 M H₂C₂O₄

D. 1.0 M HCOOH

8. Consider the following: H₂O + 57 kJ D H₃O⁺ + OH^-

When the temperature of the system is increased, the equilibrium shifts

A. left and the Kw increases

B. left and the Kw decreases

C. right and the Kw increases

D. right and the Kw decreases

9. Normal rainwater has a pH of approximately 6 as a result of dissolved

A. oxygen

B. carbon dioxide

C. sulphur dioxide

D. nitrogen dioxide

10. A 1.0 M solution of sodium dihydrogen phosphate is

A. acidic and the pH < 7.00

B. acidic and the pH > 7.00

C. basic and the pH < 7.00

D. basic and the pH > 7.00

11. Consider the following equilibrium for an indicator:

HInd + H₂O D Ind^- + H₃O⁺

When a few drops of indicator chlorophenol red are added to a colourless solution of pH 4.0, the resulting solution is

A. red as [HInd] < [Ind^-]

B. red as [HInd] > [Ind^-]

C. yellow as [HInd] < [Ind^-]

D. yellow as [HInd] > [Ind^-]

12. A Bronsted-Lowry base is defined as a chemical species that

A. accepts protons

B. neutralizes acids

C. donated electrons

D. produces hydroxides ions in solution

13. Which of the following solutions will have the greatest electrical conductivity?

A. 1.0 M HCN

B. 1.0 M H₂SO₄

C. 1.0 M H₃PO₄

D. 1.0 M CH₃COOH

14. Consider the following equilibrium: HC₆H₅O₇^2- + HIO₃ D H₂C₆H₅O₇^- + IO₃^-

The order of Bronsted-Lowry acids and bases is

A. acid, base, acid, base

B. acid, base, base, acid

C. base, acid, acid, base

D. base, acid, base acid

15. Consider the following: H₂O_(l) D H⁺ + OH^-

When a small amount of 1.0 M KOH is added to the above system, the equilibrium

A. shifts left and [H⁺] decreases

B. shifts left and [H⁺] increases

C. shifts right and [H⁺] decreases

D. shifts right and [H⁺] increases

16. Which of the following has the highest pH?

A. 1.0 M NaIO₃

B. 1.0 M Na₂CO₃

C. 1.0 M Na₃PO₄

D. 1.0 M Na₂SO₄

17. In a 100.0 mL sample of 0.0800 M NaOH the [H₃O⁺] is

A. 1.25 x 10^-13 M

B. 1.25 x 10^-12 M

C. 8.00 x 10^-3 M

D. 8.00 x 10^-2 M

18. Consider the following:

I ammonium nitrate II calcium nitrate III iron III nitrate

When dissolved in water, which of these salts would form a neutral solution?

A. II only

B. III only

C. I and III only

D. I, II, and III

19. Consider the following: SO₄^2- + HNO₂ D HSO₄^- + NO₂^-

Equilibrium would favour the

A. the products since HSO₄^- is a weaker acid than HNO₂

B. the reactants since HSO₄^- is a weaker acid than HNO₂

C. the products since HSO₄^- is a stronger acid than HNO₂

D. the reactants since HSO₄^- is a stronger acid than HNO₂

20. The net ionic equation for the hydrolysis of Na₂CO₃ is

A. H₂O + Na⁺ D NaOH + H⁺

B. H₂O + 2Na⁺ D Na₂O + 2H⁺

C. H₂O + CO₃^2- D H₂CO₃ + O^2-

D. H₂O + CO₃^2- D HCO₃^- + OH^-

21. Consider the following equilibrium: 2H₂O_(l) D H₃O⁺ + OH^-

A few drops of 1.0 M HCl are added to the above system. When equilibrium is

re-established, the

A. [H₃O⁺] has increased and the [OH^-] has decreased

B. [H₃O⁺] has increased and the [OH^-] has increased

C. [H₃O⁺] has decreased and the [OH^-] has increased

D. [H₃O⁺] has decreased and the [OH^-] has decreased

22. A basic solution

A. tastes sour

B. feels slippery

C. does not conduct electricity

D. reacts with metals to release oxygen gas

23. The balanced formula equation for the neutralization of H₂SO₄ by KOH is

A. H₂SO₄ + KOH → KSO₄ + H₂O

B. H₂SO₄ + KOH → K₂SO₄ + H₂O

C. H₂SO₄ + 2KOH → K₂SO₄ + H₂O

D. H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O

24. An Arrhenius base is defined as a substance which

A. donates protons

B. donates electrons

C. produces H⁺ in solution

D. produces OH^- in solution

25. Consider the following equilibrium: HS^- + H₃PO₄ D H₂S + H₂PO₄^-

The order of Bronsted-Lowry acids and bases is

A. acid, base, acid, base.

B. acid, base, base, acid

C. base, acid, acid, base

D. base, acid, base, acid

26. The equation representing the reaction of ethanoic acid with water is

A. CH₃COO^- + H₂O D CH₃COOH + OH^-

B. CH₃COO^- + H₂O D CH₃COO^2- + H₃O⁺

C. CH₃COOH + H₂O D CH₃COO^- + H₃O⁺

D. CH₃COOH + H₂O D CH₃COOH₂⁺ + OH^-

27. Consider the following equilibrium: 2H₂O + 57kJ D H₃O⁺ + OH^-

When the temperature is decreased, the water

A. stays neutral and the [H₃O⁺] increases

B. stays neutral and the [H₃O⁺] decreases

C. becomes basic and [H₃O⁺] decreases

D. becomes acidic and [H₃O⁺] increases

28. The equation for the reaction of Cl₂O with water is

A. Cl₂O + H₂O D 2HClO

B. Cl₂O + H₂O D 2ClO + H₂

C. Cl₂O + H₂O D Cl₂ + H₂O₂

D. Cl₂O + H₂O D Cl₂ + O₂ + H₂

29. The conjugate acid of C₆H₅0^- is

A. C₆H₄O^-

B. C₆H₅OH

C. C₆H₄O^2-

D. C₆H₅OH⁺

30. Which of the following solutions will have the greatest electrical conductivity?

A. 1.0 M HCl

B. 1.0 M HNO₂

C. 1.0 M H₃BO₃

D. 1.0 M HCOOH

31. A solution of 1.0 M HF has

A. a lower pH than a solution of 1.0 M HCl

B. a higher pOH than a solution of 1.0 M HCl

C. a higher [OH^-] than a solution of 1.0 M HCl

D. a higher [H₃O⁺] than a solution of 1.0 M HCl

32. Which of the following is the weakest acid

A. HIO₃

B. HCN

C. HNO₃

D. C₆H₅COOH

33. Considering the following data

H₃AsO₄ Ka = 5.0 x 10^-5

H₂AsO₄^- Ka = 8.0 x 10^-8

HAsO₄^2- Ka = 6.0 x 10^-10

The Kb value for H₂AsO₄^-is

A. 2.0 x 10^-10

B. 8.0 x 10^-8

C. 1.2 x 10^-7

D. 1.7 x 10^-5

34. In a solution at 25^oC, the [H₃O⁺] is 3.5 x 10^-6 M. The [OH^-] is

A. 3.5 x 10^-20 M

B. 2.9 x 10^-9 M

C. 1.0 x 10^-7 M

D. 3.5 x 10^-6 M

35. In a solution with a pOH of 4.22, the [OH-] is

A. 1.7 x 10^-10 M

B. 6.0 x 10^-5 M

C. 6.3 x 10^-1 M

D. 1.7 x 10⁴ M

36. An aqueous solution of NH₄CN is

A. basic because Ka < Kb

B. basic because Ka > Kb

C. acidic because Ka < Kb

D. acidic because Ka > Kb

37. The net ionic equation for the predominant hydrolysis reaction of KHSO₄ is

A. HSO₄^- + H₂O D SO₄^2- + H₃O⁺

B. HSO₄^- + H₂O D H₂SO₄ + OH^-

C. KHSO₄ + H₂O D K⁺ + SO₄^2- + H₃O⁺

D. KHSO₄ + H₂O D K⁺ + H₂SO₄ + OH^-

38. The [OH^-] in an aqueous solution always equals

A. Kw x [H₃O⁺]

B. Kw - [H₃O⁺]

C. Kw/[H₃O⁺]

D. [H₃O⁺]/Kw

39. The [H₃O⁺] in a solution with pOH of 0.253 is

A. 5.58 x 10^-15 M

B. 1.79 x 10^-14 M

C. 5.58 x 10^-1 M

D. 5.97 x 10^-1 M

40. The equilibrium expression for the hydrolysis reaction of 1.0 M K₂HPO₄ is

A. [H₂PO₄^-][OH^-] B. [H₃PO₄][OH^-]

[HPO₄^2-] [H₂PO₄^-]

C. [K⁺] [KHPO₄^-] D. [K⁺]² [HPO₄^2-]

[K₂HPO₄] [K₂HPO₄]

41. The solution with the highest pH is

A. 1.0 M NaCl

B. 1.0 M NaCN

C. 1.0 M NaIO₃

D. 1.0 M Na₂SO₄

42. The pH of 100.0 mL of 0.0050 M NaOH is

A. 2.30

B. 3.30

C. 10.70

D. 11.70

43. Consider the following equilibrium for an indicator: HInd + H₂O D Ind^- + H₃O⁺

At the transition point,

A. [HInd] > [Ind^-]

B. [HInd] = [Ind^-]

C. [HInd] < [Ind^-]

D. [HInd] = [H₃O⁺]

Acids Unit Midterm Practice Test Subjective

1. a) Write the net ionic equation for the reaction between NaHSO₃ and NaHC₂O₄.

b) Explain why the reactants are favoured in the above reaction.

2. What is the [H₃O⁺] in a solution formed by adding 60.0 mL of water to 40.0 mL

of 0.400 M KOH?

3. A solution of 0.100 M HOCN has a pH of 2.24. Calculate the Ka value for the acid.

4. Calculate the pH in 100.0 mL 0.400 M H₃BO₃.

5. Calculate the pH of the solution formed by mixing 20.0 mL of 0.500 M HCl with 30.0 mL of 0.300 M NaOH.

6. a) Write the balanced equation representing the reaction of HF with H₂O.

b) Identify the Bronsted-Lowry bases in the above equation.

7. Consider the following data:

Barbituric acid HC₄H₃N₂O₃ Ka = 9.8 x 10^-5

Sodium propanoate NaC₃H₅O₂ Kb = 7.5 x 10^-10

Propanoic acid HC₃H₅O₂ Ka = ?

Which is the stronger acid, propanoic acid or babituric acid? Explain using calculations.

8. A solution of 0.0100 M lactic acid, HC₃H₅O₃, has a pH of 2.95. Calculate the Ka value.

9. a) Write equations showing the amphiprotic nature of water as it reacts with HCO₃^-.

b) Calculate the Kb for HCO₃^-.

10. Calculate the [H₃O⁺] in 0.550 M C₆H₅COOH.

Quiz #1 Properties of Acids, Bases, Salts, Arrhenius Bronsted Acids, Ka, Strength

1. Drano®, a commercial product used to clean drains, contains small bits of aluminum metal and

A. ammonia

B. acetic acid

C. hydrochloric acid

D. sodium hydroxide

2. A net ionic equation for the reaction between CH₃COOH and KOH is

A. CH₃COOH_(aq) + K⁺_(aq)⇄ CH₃COOK_(aq)

B. CH₃COOH_(aq) + OH^-_(aq) ⇄ H₂O_(l) + CH₃COO^-_(aq)

C. CH₃COOH_(aq) + KOH_(aq) ⇄ H₂O_(l) + CH₃COOK_(aq)

D. CH₃COOH_(aq) + K⁺_(aq) + OH^-_(aq)⇄ H₂O_(l) + KCH₃COO_(aq)

3. Which equation represents a neutralization reaction?

A. Pb²⁺_(aq)+ 2Cl^-_(aq)→ PbCl_2(s)

B. HCl_(aq) + NH_3(aq) → NH₄Cl_(aq)

C. BaI_2(aq) + MgSO_4(aq) → BaSO_4(s) + MgI_2(aq)

D. MnO₄^-_(aq) + 5Fe²⁺_(aq)+8H⁺_(aq)→ Mn²⁺_(aq) + 5Fe³⁺_(aq)+ 4H₂O_(l)

4. An Arrhenius acid is a substance that

A. accepts a proton

B. donates a proton

C. produces H⁺ in solution

D. produces OH^- in solution

5. Consider the following data table:

Breaker	Volume	Contents
1	15 mL	0.1 M Sr(OH)₂
2	20 mL	0.2 M NH₄OH
3	25 mL	0.1 M KOH
4	50 mL	0.2 M NaOH

Identify the beaker that requires the smallest volume of 1.0 M HCl for complete neutralization

A. Beaker 1

B. Beaker 2

C. Beaker 3

D. Beaker 4

6. The net ionic equation for the titration of HClO_4(aq) with LiOH_(aq) is

A. H⁺_(aq)+ OH^-_(aq)→ H₂O_(l)

B. HClO_4(aq)+ OH^-_(aq)→ ClO₄^-_(aq)+ H₂O_(l)

C. HClO_4(aq) + LiOH_(aq)→ LiClO_4(aq)+ H₂O_(l)

D. H⁺_(aq)+ ClO₄^-_(aq)+ Li⁺_(aq)+ OH^-_(aq)→ LiClO_4(aq)+ H₂O_(l)

7. The equilibrium constant expression for a sulphurous acid is

A K_a = [H⁺][HSO₃^-]

B. K_a = [H⁺][HSO₃^-]

[H₂SO₃]

C. K_a = [2H⁺][SO₃^2-]

[H₂SO₃]

D. K_a = [H⁺][SO₃^2-]

[H₂SO₃]

8. To distinguish between a strong acid and a strong base, an experimenter could use

A. odor

B. magnesium

C. a conductivity test

D. the common ion test

9. How many acids from the list below are known to be weak acids?

HCl, HF, H₂SO₃, H₂SO₄, HNO₃, HNO₂

A. 2

B. 3

C. 4

D. 5

10. There are two beakers on a laboratory desk. One beaker contains 1.0 M HCl and the other contains tap water. To distinguish the acid solution from the water, one would use

A. a piece of copper.

B. a piece of magnesium

C. phenolphthalein indicator

D. a piece or red litmus paper

11. Caustic soda, NaOH, is found in

A. fertilizers

B. beverages

C. toothpaste

D. oven cleaners

12. Which of the following is the strongest acid?

A. Acetic acid

B. Oxalic acid

C. Benzoic acid

D. Carbonic acid

13. The acid used in the lead-acid storage battery is

A. HCl

B. HNO₃

C. H₂SO₄

D. CH₃COOH

Quiz #2 Conjugates, Amphiprotic, Arrhenius, Bronsted Bases, Kb, & Strength

1. A test that could be safely used to distinguish a strong base from a weak base is

A. taste

B. touch

C. litmus paper

D. electrical conductivity

2. Identify the two substances that act as Bronsted-Lowry bases in the equation

HS^- + SO₄^2- ⇄ S^2-+ HSO₄^-

A. HS^- and S^2-

B. SO₄^2-and S^2-

C. HS^- and HSO₄^-

D. SO₄^2- and HSO₄^-

3. The conjugate acid of H₂C₆H₅O^-₇ is

A. C₆H₅O₇^3-

B. HC₆H₅O₇^2-

C. H₂C₆H₅O₇

D. H₃C₆H₅O₇

4. Which one of the following substances is/are amphiprotic?

(1) H₃PO₄ (2) H₂PO₄^- (3) HPO₄^2-

A. 2 only

B. 3 only

C. 1 and 2

D. 2 and 3

5. The net ionic equation for the neutralization of HBr by Ca(OH)₂ is

A. H⁺_(aq)+ OH^-_(aq)⇄ H₂O_(l)

B. Ca²⁺_(aq)+ 2Br^-_(aq)⇄ CaBr_2(s)

C. 2HBr_(aq) + Ca(OH)_2(aq)⇄ 2H₂O_(l) + CaBr_2(s)

D. 2H⁺_(aq)+ 2Br^-_(aq)+ Ca²⁺_(aq) + 2OH^-_(aq) ⇄ 2H₂O_(l) + Ca²⁺_(aq) + 2Br ^-_(aq)

6. If reactants are favored in the following equilibrium, the stronger base must be

HCN + HS ^- ⇄ H₂S + CN ^-

A. H₂S

B. HS^-

C. CN^-

D. HCN

7. The hydronium ion, H₃O⁺ is a water molecule that has

A. lost a proton

B. gained a proton

C. gained a neutron

D. gained an electron

8. The complete ionic equation for the neutralization of acetic acid by sodium hydroxide is

A. H⁺ + OH^-⇄ H₂O

B. CH₃COOH + OH^- ⇄ CH₃COO^- + H₂O

C. CH₃COOH + NaOH ⇄ NaCH₃COOH + H₂O

D. CH₃COOH + Na⁺ + OH^- ⇄ Na⁺ + CH₃COO^- + H₂O

9. In the following Bronsted – Lowry acid-base equation:

NH₄⁺_(aq) + H₂O_(l) ⇄ NH_3(aq)+ H₃O⁺_(aq)

The stronger base is

A. NH₄⁺

B. H₂O

C. NH₃

D. H₃O⁺

10. Consider the following equilibrium system:

OCl^-_(aq)+ HC₇H₅O_2(aq)⇄ HOCl_(aq)+ C₇H₅O₂^-_(aq) K_eq= 2.1 x 10³

At Equilibrium

A. products are favored and HOCl is the stronger acid

B. reactants are favored and HOCl is the stronger acid

C. products are favored and HC₇H₅O₂ is the stronger acid

D. reactants are favored and HC₇H₅O₂ is the stronger acid

11. In the equilibrium system

H₂BO₃^-_(aq) + HCO₃^-_(aq)⇄ H₂CO_3(aq) + HBO₃^2-_(aq)

The two species acting as Bronsted-Lowry acids are

A. HCO₃^- and H₂CO₃

B. H₂BO₃^- and H₂CO₃

C. HCO₃^- and HBO₃^2-

D. H₂BO₃^-and HBO₃^2-

12. Consider the following equilibrium HS^- + H₂C₂O₄ ⇄ HC₂O₄^- + H₂S

The stronger acid is

A. HS^-

B. H₂C₂O₄

C. HC₂O₄^-

D. H₂S

Quiz #3 Leveling effect, Anhydrides, Hydrolysis, Relationships

1. Which of the following oxides will form the most acidic solution?

A. SO₂

B. MgO

C. Na₂O

D. Al₂O₃

2. Which one of the following salts will produce an acidic solution?

A. KBr

B. LiCN

C. NH₄Cl

D. NaCH₃COO

3. The balanced equation for the reaction between sodium oxide and water is

A. Na₂O + H₂O → 2NaOH

B. Na₂O + H₂O → 2NaH + O₂

C. Na₂O + H₂O → 2Na + H₂O₂

D. Na₂O + H₂O → 2Na + H₂⁺O₂

4. ‘Normal’ rainwater is slightly acidic due to the presence of dissolved

A. methane

B. carbon dioxide

C. sulphur dioxide

D. nitrogen dioxide

5. Which of the following oxides would hydrolyze to produce hydroxide ions?

A. NO

B. SO₂

C. Cl₂O

D. Na₂O

6. The approximate pH of “normal” rainwater is

A. 0

B. 6

C. 7

D. 8

7. Which of the following oxides would hydrolyze to produce hydronium ions?

A. CaO

B. SO₂

C. MgO

D. Na₂O

8. Which of the following gasses results in the formation of acid rain?

A. H₂

B. O₃

C. SO₂

D. NH₃

9. Consider the following acid base solution

HSO₃^- + HF ⇄ H₂SO₃ + F^-

The order of Bronsted-Lowry acids and bases in this equation is

A. acid + base ⇄ acid + base

B. acid + base ⇄ base + acid

C. base + acid ⇄ base + acid

D. base + acid ⇄ acid + base

10. The conjugate acid of OH^- is

A. H⁺

B. O^2-

C. H₂O

D. H₃O⁺

11. Which of the following 0.10 M solutions will have the greatest electrical conductivity?

A. HF

B. NH₃

C. NaOH

D. C₆H₅COOH

12. The amphiprotic ion HSeO₃^- can undergo hydrolysis according to the following equations

HSeO₃^- + H₂O ⇄ H₂SeO₃ + OH^-	K₁
HSeO₃^- + H₂O ⇄ SeO₃^2-+ H₃O⁺	K₂

An aqueous solution of HSeO₃^- is found to be acidic. This observation indicates that when it is added to water, HSeO₃^- behaves mainly as a

A. proton donor, and K₁ is less than K₂

B. proton donor, and K₁ is greater than K₂

C. proton acceptor, and K₁ is less than K₂

D. proton acceptor, and K₁ is greater than K₂

13. The K_bexpression for HPO₄^2-is

A. [PO₄^3-][H₃O⁺] B. [HPO₄^2-][OH^-]

[HPO₄^2-] [H₂PO₄^-]

C. [H₂PO₄^-][OH^-] D. [HPO₄^2-][ H₃O⁺]

[HPO₄^2-] [PO₄^3-]

Quiz #4 Anhydrides, Hydrolysis

1. Which of the following pairs of gases are primarily responsible for producing “acid rain”?

A. O₂ and O₃

B. N₂ and O₂

C. CO and CO₂

D. NO₂ and SO₂

2. Sodium potassium tartrate (NaKC₄H₄O₆) is used to raise the pH of fruit during processing. In this process, sodium potassium tartrate is being used as a/an

A. salt

B. acid

C. base

D. buffer

3. The net ionic equation for the hydrolysis of the salt, Na₂S is

A. Na₂S ⇄ 2Na⁺ + S^2-

B. S^2-+ H₂O ⇄ OH^-+ HS^-

C. Na₂S + 2H₂O ⇄ 2NaOH + H₂S

D. 2Na⁺ + S^2-+ 2H₂O ⇄ 2Na⁺ + 2OH^- + H₂S

4. Which of the following solutions would be acidic?

A. sodium acetate

B. iron III chloride

C. sodium carbonate

D. potassium chloride

5. Consider the following salts: I. NaF II. NaClO₄ III. NaHSO₄

Which of these salts, when dissolved in water, would form a basic solution?

A. I only

B. I and II only

C. II and III only

D. I, II and III

6. Which of the following, when dissolved in water, forms a basic solution?

A. KCl

B. NaClO₄

C. Na₂CO₃

D. NH₄NO₃

7. Which of the following oxides forms a basic solution?

A. K₂O

B. CO₂

C. SO₃

D. NO₂

8. Which of the following is amphiprotic in water?

A. SO₂

B. SO₃^2-

C. HSO₃^-

D. H₂SO₃

9. Consider the following equilibrium expression

K= [H₂S][OH^-]

[HS^-]

This expression represents the

A. K_b for H₂S

B. K_a for H₂S

C. K_b for HS^-

D. K_a for HS^-

10. The reaction of a strong acid with a strong base produces

A. A salt and a water

B. A base and an acid

C. A metallic oxide and water

D. A non-metallic oxide and water

11. Consider the following equilibrium:

CH₃COOH_(aq) + NH_3(aq)⇄ CH₃COO^- _(aq)+ NH₄⁺_(aq)

The sequence of Bronsted-Lowryacids and bases in the above equilibrium equation is

A. acid, base, base, acid

B. acid, base, acid, base

C. base, acid, base, acid

D. base, acid, acid, base

12. The pH range of ‘acid rain’ is often

A. 3 to 6

B. 6 to 8

C. 7 to 9

D. 10 to 12

13. Water will act as a Bronsted-Lowry acid with

A. NH₃

B. H₂S

C. HCN

D. HNO₃

14. Which of the following is a conjugate acid-base pair?

A. H₃PO₄ and PO₄^3-

B. H₂PO₄^- and PO₄^3-

C. H₃PO₄ and HPO₄^2-

D. H₂PO₄^- and HPO₄^2-

Quiz #5 pH calculations for Strong and Weak Acids

1. The 1.0 M acidic solution with the highest pH is

A. H₂S

B HNO₂

C. HNO₃

D. H₃BO₃

2. At 25^oC, the equation representing the ionization of water is

A H₂O + H₂O ⇄ 2H₂ + O₂

B. H₂O + H₂O ⇄ H₂O₂ + H₂

C. H₂O + H₂O ⇄ 4H⁺ + 2O^2-

D. H₂O + H₂O ⇄ H₃O⁺ +OH^-

3. The pH of a 0.3 M solution of NH₃ is approximately

A. 14.0

B. 11.0

C. 6.0

D. 3.0

4. The pH of an aqueous solution is 4.32. The [OH^-] is

A. 6.4 x 10^-1 M

B. 4.8 x 10^-5 M

C. 2.1 x 10^-10 M

D. 1.6 x 10^-14 M

5. The pH of an aqueous solution is 10.32. The [OH^-] is

A. 5.0 x 10^-12 M

B. 2.0 x 10^-11 M

C. 4.8 x 10^-11 M

D. 2.1 x 10^-4 M

6. The pH of a 0.025 M HClO₄ solution is

A. 0.94

B. 1.60

C. 12.40

D. 13.06

7. Consider the following equilibrium: H₂O_(l) + H₂O_(l) ⇄ H₃O⁺_(aq)+ OH^-_(aq)

The equilibrium constant for this system is referred to as

A. K_w

B. K_a

C. K_b

D. K_sp

8. The [H₃O⁺] in a solution of pH = 0.60 is

A. 4.0 x 10^-14 M

B. 2.2 x 10^-1 M

C. 2.5 x 10^-1 M

D. 6.0 x 10^-1 M

9. A solution is prepared by adding 100 mL of 10 M of HCl to a 1 litre volumetric flask and filling it to the mark with water. The pH of this solution is

A. -1

B. 0

C. 1

D. 7

10. The approximate pH of a 0.06 M solution of CH₃COOH is

A. 1

B. 3

C. 11

D. 13

11. The [OH^-] is greater than the [H₃O⁺] in

A. HCl_(aq)

B. NH_3(aq)

C. H₂O_(aq)

D. CH₃COOH_(aq)

12. The pH of 0.15 M HCl is

A. 0.15

B. 0.71

C. 0.82

D. 13.18

13. Which of the following equations correctly relates pH and [H₃O⁺]?

A. pH= log [H₃O⁺]

B. pH= 14 - [H₃O⁺]

C. pH= -log [H₃O⁺]

D. pH= pK_w – [H₃O⁺]

14. The pH of 0.20 M HNO₃ is

A. 0.20

B. 0.63

C. 0.70

D. 1.58

15. The [OH^-] in 0.050 M HNO₃ at 25^oC is

A. 5.0 x 10^-16 M

B. 1.0 x 10^-14 M

C. 2.0 x 10^-13 M

D. 5.0 x 10^-2 M

Quiz #6 Ka’s from pH Kb’s from Ka’s

1. The Kb for the dihydrogen phosphate ion is

A. 1.3 x 10^-12

B. 6.3 x 10^-8

C. 1.6 x 10^-7

D. 7.1 x 10^-3

2. What volume of 0.100 M NaOH is required to neutralize a 10.0 mL sample of 0.200 M H₂SO₄?

A. 5.0 mL

B. 10.0 mL

C. 20.0 mL

D. 40.0 mL

3. Consider the following equilibriums:

I	HCO₃^-+ H₂O ⇄ H₂CO₃ + OH^-
II	NH₄⁺ + H₂O ⇄ H₃O⁺ + NH₃
III	HSO₃^- + H₃O⁺ ⇄ H₂O + H₂SO₃

Water acts as a Bronsted-Lowry base in

A. III only

B. I and II only

C. II and III only

D. I, II, and III

4. Which of the following is represented by a Kb expression?

A. Al(OH)_3(s) ⇄ Al³⁺_(aq) + 3OH^-_(aq)

B. HF_(aq) + H₂O_(l) ⇄ H₃O⁺_(aq)+ F^-_(aq)

C. Cr₂O₇^2-_(aq)+ 2OH^-_(aq)⇄ 2CrO₄^2-_(aq)+ H₂O_(l)

D. CH₃NH_2(aq)+ H₂O_(l) ⇄ CH₃NH₃⁺_(aq) + OH^-_(aq)

5. A student combines 0.25 mol of NaOH and 0.20 mol of HCl in water to make 2.0 liters of solution. The pH of the solution is

A. 1.3

B. 1.6

C. 12.4

D. 12.7

6. If OH^- is added to a solution, the [H₃O⁺] will

A. Remain constant

B. Adjust such that [H₃O⁺]= [OH^-]

K_w

C. Increase such that [H₃O⁺][OH^-] = K_w

D. Decrease such that [H₃O⁺][OH^-] = K_w

7. In a titration, 10.0 mL of H₂SO_4(aq)is required to neutralize 0.0400 mol of NaOH.

From this data, the [H₂SO₄] is

A. 0.0200 M

B. 2.00 M

C. 4.00 M

D. 8.00 M

8. Consider the following equilibrium for an acid-base indicator:

Hlnd ⇄ H⁺ + Ind^- K_a = 1.0 x 10^-10

Which of the following statements is correct at pH 7.0?

A. [Ind^-] < [HInd]

B. [Ind^-] = [HInd]

C. [Ind^-] > [HInd]

D. [Ind^-] = [H⁺] = [HInd]

9. Which of the following indicators would be yellow at pH 7 and blue at pH 10?

A. thymol blue

B. methyl violet

C. bromthymol blue

D. bromcresol green

10. Consider the following equilibrium for phenolphthalein: HInd ⇄ H⁺ + Ind^-

When phenolphthalein is added to 1.0 M NaOH, the color of the resulting solution is

A. pink as [HInd] is less than [Ind^-]

B. pink as [HInd] is greater than [Ind^-]

C. colorless as [HInd] is less than [Ind^-]

D. colorless as [HInd] is greater than [Ind^-]

11. Water acts as a base when it reacts with

A. CN^-

B. NH₃

C. NO₂^-

D. NH₄⁺

12. What is the pH of a solution prepared by adding 0.50 mol KOH to 1.0 L of

0.30 M HNO₃?

A. 0.20

B. 0.70

C. 13.30

D. 13.80

13. The 1.0 M acid solution with the largest [H₃O⁺] is

A. HNO₂

B. H₂SO₃

C. H₂CO₃

D. H₃BO₃

Quiz #7 pH for Weak bases, pH Relationships, Amphiprotic Calculations

1. In water, the hydrogen sulphide ion, HPO₄^2-, will act as

A. An acid because K_a< K_b

B. An acid because K_a > K_b

C. A base because K_a< K_b

D. A base because K_a > K_b

2. A student records the pH of 1.0 M solution of two acids. Which of the following statements can be concluded from the above data?

Acid	pH
X	4.0
Y	2.0

A. Acid X is stronger than acid Y

B. Acid X and acid Y are both weak

C. Acid X is diprotic while acid Y is monoprotic

D. Acid X is 100 times more concentrated than acid Y

3. When added to water, the hydrogen carbonate ion, HCO₃^-, produces a solution, which is

A. basic because K_b is greater than K_a

B. basic because K_a is greater than K_b

C. acidic because K_a is greater than K_b

D. acidic because K_b is greater than K_a

4. The concentration, Ka and pH values of four acids are given in the following table

ACID	Concentration	K_a	pH
HA	3.0 M	2.0 x 10^-5	2.1
HB	0.7 M	4.0 x 10^-5	2.3
HC	4.0 M	1.0 x 10^-5	2.2
HD	1.5 M	1.3 x 10^-5	2.4

Based on this data, the strongest acid is

A. HA

B. HB

C. HC

D. HD

5. Which of the following 0.10 M solutions is the most acidic?

A. AlCl₃

B. FeCl₃

C. CrCl₃

D. NH₄Cl

6. Which of the following acid-base indicators has a transition point between pH 7 and pH 9?

A. Ethyl red, K_a = 8 x 10^-2

B. Congo red, K_a = 9.0 x 10^-3

C. Cresol red, K_a = 1.0 x 10^-8

D. Alizarin blue, K_a = 7.0 x 10^-11

Quiz #8 Buffers and Indicators

1. Consider the following acid solutions:

I. H₂CO₃ II. HClO₄ III. HF

Which of the above acids would form a buffer solution when its conjugate base is added?

A. I only

B. II only

C. I and III only

D. I, II, and III only

2. Consider the following base indicator:

HInd ⇄ H⁺ + Ind^-

When the indicator is added to different solutions, the following data are obtained:

Solution	1.0 M HCl	1.0 M HA_l	1.0 M HA₂
Colour	Yellow	Blue	Yellow

The acids HA_l, HA₂, and HInd listed in the order of decreasing acid strength is

A. HA₂, HInd, HA_l

B. HInd, HA_l, HA₂

C. HA₂, HA_l, HInd

D. HA_l, HInd, HA₂

3. Which of the following compounds, when added to a solution of ammonium nitrate, will result in the formation of a buffer solution?

A. Ammonia

B. Nitric acid

C. Sodium nitrate

D. Ammonium chloride

4. Which of the following represents a buffer equilibrium?

A. HI + H₂O ⇄ H₃O⁺ + I^-

B. HCl + H₂O ⇄ H₃O + Cl^-

C. HCN + H₂O ⇄ H₃O⁺ + CN^-

D. HClO₄ + H₂O ⇄ H₃O + ClO₄^-

5. Consider the following equilibrium:

HF_(aq) + H₂O_(l) ⇄ H₃O⁺_(aq)+ F^-_(aq)

The above system will behave as a buffer when the [F^-] is approximately equal to

A. K_a

B. [HF]

C. [H₂O]

D. [H₃O⁺]

6. A basic buffer solution can be prepared by mixing equal numbers of moles of

A. NH₄CL and HCl

B. NaCl and NaOH

C. Na₂CO₃ and NaHCO₃

D. NaCH₃COOH and CH₃COOH

Quiz #9 Titrations and Titration Curves

1. Which of the following indicators would be used when titrating a weak acid with a strong base?

A. Methyl red

B. Methyl violet

C. Indigo carmine

D. Phenolphthalein

2. Which of the following acid-base pairs would result in an equivalence point with pH greater than 7.0?

A. HCl and LiOH

B. HNO₃ and NH₃

C. HClO₄ and NaOH

D. CH₃COOH and KOH

3. Which of the following standardized solutions should a chemist select when titrating a 25.00 mL sample of 0.1 M NH₃, using methyl red as an indicator?

A. 0.114 M HCl

B. 6.00 M HNO₃

C. 0.105 M NaOH

D. 0.100 M CH₃COOH

4. Consider the following 0.100 M solutions I. H₂SO₄ II. HCl III. HF

The equivalence point is reached when 10.00 mL of 0.100 NaOH has been added to 10.00 mL of solution(s)

A. II only

B. I and II only

C. II and III only

D. I, II and III

Which pair of 0.10 M solutions would result in the above titration curve?

A. HF and KOH

B. HCl and NH₃

C. H₂S and NaOH

D. HNO₃ and KOH

6. A suitable indicator for the above titration is

A. Methyl violet

B. Alizarin yellow

C. Thymolphthalein

D. Bromcresol green

7. The pH scale is

A. direct

B. inverse

C. logarithmic

D. exponential

Which of the following indicators should be used in the titration represented by the above titration curve?

A. Orange IV

B. Methyl red

C. Phenolphthalein

D. Alizarin yellow

9. Which of the following indicators should be used when 1.0 M HNO₂ is titrated with NaOH_(aq)?

A. Methyl red

B. Thymol blue

C. Methyl orange

D. Indigo carmine

10. Which of the following solutions should be used when titrating a 25.00 mL sample of CH₃COOH that is approximately 0.1 M?

A. 0.150 M NaOH

B. 0.001 M NaOH

C. 3.00 M NaOH

D. 6.00 M NaOH

11. What volume of 0.250 M H₂SO₄ is required to neutralize 25.00 mL of 2.50 M KOH?

A. 125 mL

B. 150 mL

C. 250 mL

D. 500mL

12. Which of the following pairs of substances form a buffer system for human blood?

A. HCl and Cl^-

B. NH₃ and NH₂^-

C. H₂CO₃ and HCO₃^-

D. H₂C₆H₅O₇ and HC₆H₅O₇^2-

Quiz #10 Review

1. How many moles of Mg(OH)₂ are required to neutralize 30.00 mL of 0.150 M HCl?

A. 2.25 x 10^-3 mol

B. 4.50 x 10^-3 mol

C. 5.00 x 10^-3 mol

D. 9.00 x 10^-3 mol

2. The approximate Ka for the indicator phenolphthalein is

A. 6 x 10^-19

B. 8 x 10^-10

C. 6 x 10^-8

D. 2 x 10^-2

3. A new indicator, “B.C. Blue (HInd),” is red in bases and blue in acids. Describe the shift in equilibrium and the resulting color change if 1.0 M HIO₃ is added to a neutral, purple solution of this indicator: HInd + H₂O ⇄ H₃O⁺ + Ind^-

A. Equilibrium shifts left, and colour becomes red

B Equilibrium shifts left, and colour becomes blue

C. Equilibrium shifts right, and colour becomes red

D. Equilibrium shifts right, and colour becomes blue

4. Which one of the following combinations would act as an acid buffer?

A. HCl and NaOH

B. KOH and KBr

C. NH₃ and NH₄Cl

D. CH₃COOH and NaCH₃COO

5. What is the pH at the transition point of an indicator if its Ka is 7.9 x 10^-3?

A. 0.98

B. 2.10

C. 7.00

D. 11.90

6. Which of the following pH curves best represents the titration of sodium hydroxide with hydrochloric acid?

7. A student prepares a buffer by placing ammonium chloride in a solution of ammonia. Equilibrium is established according to the equation: NH₃ + H₂O ⇄ NH₄⁺ + OH^-

When a small amount of base is added to the buffer, the base reacts with

A. NH₃ and the pH decreases

B. NH₄⁺ and the pH decreases

C. NH₃ to keep the pH relatively constant

D. NH₄⁺ to keep the pH relatively constant

8. At the equivalence point in a titration involving 1.0 M solutions, which of the following combinations would have the lowest conductivity?

A. Nitric acid and barium hydroxide

B. Acetic acid and sodium hydroxide

C. Sulphuric acid and barium hydroxide

D. Hydrochloric acid and sodium hydroxide

9. An indicator HInd produces a yellow colour in 0.1 M HCl solution and a red colour in 0.1 M HCN solution. Therefore, the following equilibrium:

HCN + Ind^- ⇄ HInd + CN^-

A. Products are favored and the stronger acid is HInd

B. Products are favored and the stronger acid is HCN

C. Reactants are favored and the stronger acid is HInd

D. Reactants are favored and the stronger acid is HCN

10. The indicator methyl red is red in a solution of NaH₂PO₄. Which of the following equations is consistent with this observation?

A. H₂PO₄^- + H₂O ⇄ HPO₄^2-+ H₃O⁺

B. H₂PO₄^- + H₂O ⇄ H₃PO₄+ OH^-

C. HPO₄^2-+ H₂O ⇄ PO₄^3-+ H₃O⁺

D. HPO₄^2-+ H₂O ⇄ H₂PO₄^- + OH^-

11. Consider the following acid-base indicator equilibrium:

HInd_(aq) + H₂O_(l) ⇄ H₃O⁺_(aq)⇄ Ind^-_(aq)

Which of the following statements describes the conditions that exist in an indicator equilibrium system at its transition point?

A. [HInd] = [Ind^-]

B. [Ind^-] = [H₃O⁺]

C. [HInd] = [H₃O⁺]

D. [H₃O⁺] = [H₂O]

12. Which of the following titrations would have an equivalence point less that pH 7?

A. NH₃ and HCl

B. NaOH and HNO₃

C. Ba(OH)₂ and H₂SO₄

D. KOH and CH₃COOH

Web Review of Acids

1. List five properties of:

a) acids

b) bases

2. What ion is produced when an acid reacts with water? A base?

3. Define:

Conjugate

Arhenius strong acid

Bronsted weak acid

Bronsted strong base

Ionization of water

Equivalence point

Transition point

Buffer

Hydrolysis.

4. Identify the acids or bases in the following equation. Are the reactants or products favoured?

HC0₃^- + HF ⇄ H₂CO₃ + F^-

5. Classify each compound as a strong or weak acid or base; acidic or basic anhydride; acidic, basic, or neutral salt; or buffer system. Write an equation to show how each reacts with water.

NH₃	AlCl₃	H₂CO₃	HClO₄	KCN	NH₄Cl	KOH
SO₂	NaF	HCl	NaI	K₂O	NaOH	CO₂

NH₃ and NH₄Cl

NaCH₃COO and CH₃COOH

6. H⁺is short for _______.

7. Determine the conjugates for each of the bases.

CN^-

NH₃

F^-

OH^-

Co(H₂O)₅(OH)²⁺

8. Determine the conjugates for each of the acids.

HCN

Al(H₂O)₆³⁺

NH₄⁺

HPO₄^2-

9. Describe a strong and weak acid as well as a strong and weak base in terms of each of the following:

strong acid weak acid strong base weak base

Conductivity

Size of Ka

Size of Kb

Degree of Ionization

pH.

10. Why is the strongest acid in water H₃O⁺? Explain!

11. Why is the strongest base in water OH^-? Explain!

12. Which has the higher pH H₂S0₃ or H₃BO₃? Explain!

13. Which has the higher pH NaCN or NaF? Explain!

14. A buffer has a pH of 9.00. 2 drops of a dilute strong acid are added. Estimate how the pH changes?

15. a) Complete the chart by indicating the pairs required to make buffer solutions. For example HCN (weak acid) and NaCN (salt containing the conjugate of the weak acid) will make a buffer solution. b) Write an equation the describes the equilibrium for each buffer. c) Circle the formulas that have high concentrations.

Weak Acid or Base	Salt
HF
	NaCH₃COO
NH₃
	NaCN
H₂CO₃
	KH₂PO₄
HCH₃COO

16. Match each equation with its type:

Acid/base formula equation	F^- + HOH_(l) ⇄ HF + OH^-
Acid/base net ionic equation	HCl + NaOH →NaCl + HOH_(l)
Solubility product ionization equation	H⁺ + OH^- → HOH_(l)
Hydrolysis of a weak acid	AgCl_(s)⇄ Ag⁺ + Cl^-
Hydrolysis of a weak base	H₂0_(l)⇄ H⁺+ OH^-
Ionization of water	NH₄⁺ + H₂0 ⇄ NH₃ + H₃O⁺

17. Write the equilibrium expressions for each of the above equations except for the second and third reactions.

18. A student tested the electrical conductivity of two acid solutions. One solution was a strong acid and the other a weak acid. Both solutions had the same conductivity. Explain how this could be possible.

19. Describe in terms of hydrolysis how NaCH₃COO can be added to potato chips in order to produce the vinegar flavour.

20. Describe what happens to the H⁺ and the OH^-when the pH increases by 2 units.

21. Describe two gases responsible for acid rain. Write equations to show how they react with water. What gas naturally lowers the pH of normal rain?

22. Complete the following reaction using a formula equation, complete ionic quation and net ionic equation. H₂C₂O₄+ NaOH →

23. Write the equilibrium expression for phosphoric acid in water.

24.

a) An indicator HInd is red in acid and blue in base. Write the equation for the indicator and explain the colours.
b) What is true at the transition point?
c)What is the color of this indicator in a solution of AlCl₃?
d) In the above solution, what is larger [HInd] or [Ind^-] ?
e) Calculate the Ka for Phenolphthalein.
f) What indicator has a Ka of approximately 1.0 X 10^-10?

25. Give an example of a monoprotic, diprotic, and triprotic acid. Write an equation for each to show how they ionize in water.

26. Give the approximate pH of the equivalence for each titration. Choose an appropriate indicator.

Acid	Base	pH of Equivalence Point	Indicator
HCl	NaOH
H₂SO₄	NH₃
HF	KOH

27. Which of the following will have the lowest pH?

HCLO

HClO₂

HClO₃

HClO₄

28. Pick the formulae that are amphiprotic.

H₂SO₄	H₂0	F^-	HCO₃^-
CO₃^-2	KOH	H₂PO₄^-	HPO₄^-2

CALCULATIONS

1. Calculate the quantities required to complete the table. In the first row write the general equations for each quantity. Watch your significant figures.

[H⁺] =	[OH^-] =	pH =	pOH =
[H⁺]	[OH^-]	pH	pOH	Acid/base/neutral
5.0 x 10^-3M
	1.3 x 10^-5M
		3.1
			2.508
				neutral (2sig figs)

2. What volume of 0.20 M H₂SO₄is needed to neutralize 50.00 ml 0.30M NaOH?

3. What mass of NaF is required to prepare 100.0 ml of 0.300 M solution?

4. 35.5 mL of 0.300 M NaOH is required to neutralize 10.0 mL of H₂SO₄. What is the acid concentration?

5. 100.0 mL of .200 M HCl is mixed with 120.0 mL of 0.200M NaOH. Calculate the pH of the resulting solution.

6. Calculate the Ka for phenolphthalein.

7. The Ksp of AgOH is 6.8 x 10^-12. Calculate the pH.

8. The OH^- concentration in 0.10M NaCN is 2.7 x 10^-6M. Calculate the Kb from this information only.

9. Calculate the pH for 0.20 M HCl.

10. Calculate the pH for 0.10M Ba(OH)₂.

11. Calculate the pH for 0.40 M HCN.

12. Calculate the pH for 0.40 M Na₂CO₃.

13. What is the pH for 0.30 M NaCl?

14. Calculate the pH of 0.20 M NH₃.

15. Calculate the pH of 0.20 M NH₄Cl.

16. Show by calculation if H₂PO₄^-is an acid or base (compare the Kb and Ka).

17. Calculate the pH of a saturated solution of Mg(OH)₂ if the Ksp is 1.2 X 10^-11.

18. A 0.50 M NH₃ solution is found to have a OH^- concentration of 1.86 x 10^-3 M. Using this data only calculate the Kb.

19. A 0.18 M acid HX has a pH of 2.40. What is the Ka?

20. The following data were recorded when 25.00 mL of H₂SO₄were titrated with 0.1030 M NaOH. The volumes of NaOH used in three runs were: 46.06 mL, 44.52 mL, 44.54 mL. Calculate the acid concentration.

21. The following data were recorded when 10.00ml of NaOH were titrated with 0.1030M H₂SO₄. The volumes of H₂SO₄ used in three runs were: 12.55 mL, 12.55 mL, 12.10 mL. Calculate the base concentration.

Acids Practice Test # 1

1. An equation representing the reaction of a weak acid with water is

A. HCl + H₂O ⇄ H₃O⁺ + Cl^-

B. NH₃ + H₂O ⇄ NH₄⁺ + OH^-

C. HCO₃^- H₂O ⇄ H₂CO₃ + OH^-

D. HCOOH + H₂O ⇄ H₃O⁺ + HCOO^-

2. The equilibrium expression for the ion product constant of water is

A. Kw = [H₃O⁺][OH^-]

[H₂O]

B. Kw = [H₃O⁺]²[O₂]

C. Kw = [H₃O⁺][OH^-]

D. Kw = [H₃O⁺]²[O^2-]

3. Consider the following graph for the titration of 0.1 M NH₃ with 1.0 M HCl.

pH

Volume HCl added

14

7

0

I

II

III

IV

A buffer solution is present at point

A. I

B. II

C. III

D. IV

4. Consider the following equilibrium system for an indicator: HInd + H₂O ⇄ H₃O⁺ + Ind^-

Which two species must be of two different colours in order to be used as an indicator?

A. HInd and H₂O

B. HInd and Ind^-

C. H₃O⁺ and Ind^-

D. Hind and H₃O⁺

5. Which of the following indicators is yellow at pH 10.0?

A. methyl red

B. phenol red

C. thymol blue

D. methyl violet

6. A sample containing 1.20 x 10^-2 mole HCl is completely neutralized by 100.0 mL of Sr(OH)₂. What is the [Sr(OH)₂]?

A. 6.00 x 10^-3 M

B. 6.00 x 10^-2 M

C. 1.20 x 10^-1 M

D. 2.4 x 10^-1 M

7. Which of the following titrations will have the highest pH at the equivalence point?

A. HBr with NH₃

B. HNO₂ with KOH

C. HCl with Na₂CO₃

D. HNO₃ with NaOH

8. An Arrhenius acid is defined as a chemical species that

A. is a proton donor.

B. is a proton acceptor.

C. produces hydrogen ions in solution.

D. produces hydroxide ions in solution.

9. Consider the following acid-base equilibrium system:

HC₂O₄^-+ H₂BO₃^- ⇄ H₃BO₃ + C₂O₄^2-

Identify the Bronsted-Lowry bases in this equilibrium.

A. H₂BO₃^- and H₃BO₃

B. HC₂O₄^- and H₃BO₃

C. HC₂O₄^- and C₂O₄^2-

D. H₂BO₃^- and C₂O₄^2-

10. The equation representing the predominant reaction between NaCH₃COO with water is

A. CH₃COO^- + H₂O ⇄ CH₃COOH + OH^-

B. CH₃COO^- + H₂O ⇄ H₂O + CH₂COO^2-

C. CH₃COOH + H₂O ⇄ H₃O⁺ + CH₃COO^-

D. CH₃COOH + H₂O ⇄ CH₃COOH₂⁺ + OH^-

11. Which of the following solutions will have the lowest electrical conductivity?

A. 0.10 M HF

B. 0.10 M NaF

C. 0.10 M H₂SO₃

D. 0.10 M NaHSO₃

12. Which of the following is the strongest Bronsted-Lowry base?

A. NH₃

B. CO₃^2-

C. HSO₃^-

D. H₂BO₃^-

13. A 1.0 x 10^-4 M solution has a pH of 10.00. The solute is a

A. weak acid

B. weak base

C. strong acid]

D. strong base

14. The ionization of water at room temperature is represented by

A. H₂O ⇄ 2H⁺ + O^2-

B. 2H₂O ⇄ 2H₂ + O₂

C. 2H₂O ⇄ H₂+ 2OH^-

D. 2H₂O ⇄ H₃O⁺ + OH^-

15. Addition of HCl to water causes

A. both [H₃O⁺] and [OH^-] to increase

B. both [H₃O⁺] and [OH^-] to decrease

C. [H₃O⁺] to increase and [OH^-] to decrease

D. [H₃O⁺] to decrease and [OH^-] to increase

16. Consider the following:

I. H₂SO₄

II. HSO₄^-

III. SO₄^2-

Which of the above is/are present in a reagent bottle labeled 1.0 M H₂SO₄?

A. I only

B. I and II only

C. II and III only

D. I, II, and III

17. The pH of 0.10 M KOH solution is

A. 0.10

B. 1.00

C. 13.00

D. 14.10

18. An indicator changes colour in the pH range 9.0 to 11.0. What is the value of the Ka for the indicator?

A. 1 x 10^-13

B. 1 x 10^-10

C. 1 x 10^-7

D. 1 x 10^-1

19. Which of the following are amphiprotic in aqueous solution?

I. HBr

II. H₂O

III. HCO₃^-

IV. H₂C₆H₅O₇^-

A I and II only

B. II and IV only

C. II, III, and IV only

D. I, II, III, and IV

20. Which of the following always applies at the transition point for the indicator Hind?

A. [Ind^-] = [OH^-]

B. [HInd] = [Ind^-]

C. [Ind^-] = [H₃O⁺]

D. [HInd] = [H₃O⁺]

21. Calculate the [H₃O⁺] of a solution prepared by adding 10.0 mL of 2.0 M HCl to 10.0 mL of 1.0 M NaOH.

A. 0.20 M

B. 0.50 M

C. 1.0 M

D. 2.0 M

22. Both acidic and basic solutions

A. taste sour

B. feel slippery

C. conduct electricity

D. turn blue litmus red

23. The conjugate acid of HPO₄^2- is

A. PO₄^3-

B. H₂PO₄^-

C. H₂PO₄^2-

D. H₂PO₄^3-

24. What is the value of the Kw at 25 ^oC?

A., 1.0 x 10^-14

B. 1.0 x 10^-7

C. 7

D. 14

25. Consider the following equilibrium: 2H₂O_(l) ⇄ H₃O⁺_(aq)+ OH^-_(aq)

A small amount of Fe(H₂O)₆³⁺is added to water and equilibrium is re-established. Which of the following represents the changes in ion concentrations?

[H₃O⁺] [OH^-]

A. increases increases

B. increases decreases

C. decreases decreases

D. decreases increases

26. Consider the following equilibrium for an indicator: HInd + H₂O ⇄ H₃O⁺ + Ind^-

In a solution of pH of 6.8, the colour of bromthymol blue is

A. blue because [HInd] = [Ind^-]

B. green because [HInd] = [Ind^-]

C. green because [HInd] < [Ind^-]

D. yellow because [HInd] > [Ind^-]

27. The indicator with Ka = 4 x 10-8 is

A. neutral red

B. methyl red

C. indigo carmine

D. phenolphthalein

28. In a titration a 25.00 mL sample of Sr(OH)₂ is completely neutralized by 28.60 mL of 0.100 M HCl. The concentration of the Sr(OH)₂ is

A. 1.43 x 10^-3 M

B. 2.86 x 10^-3 M

C. 5.72 x 10^-2 M

D. 1.14 x 10^-1 M

29. A student mixes 15.0 mL of 0.100 M NaOH with 10.0 mL of 0.200 M HCl. The resulting solution is

A. basic

B. acidic

C. neutral

D. amphiprotic

30. Which of the following salts will dissolve in water to produce a neutral solution?

A. LiF

B. CrCl₃

C. KNO₃

D. NH₄Cl

31. What is the value of the Kb for HC₆H₅O₇^2-?

A. 5.9 x 10^-10

B. 2.4 x 10^-8

C. 4.1 x 10^-7

D. 1.7 x 10^-5

32. The pOH of 0.015 M HCl solution is

A. 0.97

B. 1.80

C. 12.18

D. 13.03

33. Which of the following will produce an acidic solution?

A. NaCl

B. NH₄NO₃

C. Ca(NO₃)₂

D. Ba(NO₃)₂

34. Which of the following salts will dissolve in water to produce an acid solution?

A. LiF

B. CrCl₃

C. KNO₃

D. NaCl

35. Which of the following salts will dissolve in water to produce a basic solution?

A. LiF

B. CrCl₃

C. KNO₃

D. NH₄Cl

36. A student mixes 400 mL of 0.100 M NaOH with 100 mL of 0.200 M H₂SO₄. The resulting solution has a pH of

A. 14.000

B. 0.000

C. 13.800

D. 7.000

37. A student mixes 500 mL of 0.400 M NaOH with 500 mL of 0.100 M H₂SO₄. The resulting solution has a pH of

A. 14.000

B. 0.000

C. 13.000

D. 7.000

38. The strongest acid in water is

A. HClO₄

B. HI

C. HF

D. H₃O⁺

39. The formula that has the highest pH in water is

A. HF

B. H₂CO₃

C. H₂C₂O₄

D. HCN

39. The formula that has the highest pH in water is

A. NaF

B. NaCl

C. H₂C₂O₄

D. NaCN

Subjective

1. A chemist prepares a solution by dissolving the salt NaCN in water.

a) Write the equation for the dissociation reaction that occurs.

b) Write the equation for the hydrolysis reaction that occurs.

c) Calculate the value of the equilibrium constant for the hydrolysis

2. A 3.50 x 10^-3 M sample of unknown acid, HA has a pH of 2.90. Calculate the value of the Ka and identify this acid.

3. Calculate the mass of NaOH needed to prepare 2.0 L of a solution with a pH of 12.00.

4. A 1.00 M OCl^- solution has an [OH-] of 5.75 x 10^-4 M. Calculate the Kb for OCl^-.

5. Calculate the pH of a solution prepared by adding 15.0 mL of 0.500 M H₂SO₄ to 35.0 mL of 0.750 M NaOH.

6. Determine the pH of a 0.10 M solution of hydrogen cyanide.

7. Determine the pH of 0.100 M NH₃.

8. Determine the pH of a saturated solution of Mg(OH)₂.

Acids Practice Test # 2

1. What colour would 1.0 M HCl be in an indicator mixture consisting of phenol red and thymolphthalein?

A red

B blue

C yellow

D colourless

2. During a titration, what volume of 0.500 M KOH is necessary to completely neutralize 10.0 mL of 2.00 M CH₃COOH?

A 10.0 mL

B 20.0 mL

C 25.0 mL

D 40.0 mL

3. Which indicator has a Ka = 1.0 x 10^-6?

A neutral red

B thymol blue

C thymolpthalein

D chlorophenol red

4. Acid is added to a buffer solution. When equilibrium is reestablished the buffering effect has resulted in [H₃O⁺]

A increasing slightly

B decreasing slightly

C increasing considerably

D decreasing considerably

5. A buffer solution will form when 0.10 M NaF is mixed with an equal volume of

A 0.10 M HF

B 0.10 M HCl

C 0.10 M NaCl

D 0.10 M NaOH

6. Which of the following statements applies to 1.0 M NH_3(aq)but not to 1.0 M NaOH_(aq)?

A partially ionizes

B neutralizes an acid

C has a pH greater than 7

D turns bromocresol green from yellow to blue

7. In which of the following are the reactants favoured?

A HNO₂ + CN^- ⇄ NO₂^- + HCN

B H₂S + HCO₃^- ⇄ HS^- + H₂CO₃

C H₃PO₄ + NH₃ ⇄ H₂PO₄^- + NH₄⁺

D CH₃COOH + PO₄^3- ⇄ CH₃COO^- + HPO₄^2-

8. What is the pOH of a solution prepared by adding 0.50 moles of NaOH to prepare 0.50 L of solution?

A 0.00

B 0.30

C 14.00

D 13.70

9. What is the [H₃O⁺] in a solution with a pH = 5.20?

A 1.4 x 10^-14

B 1.6 x 10^-9

C 6.3 x 10^-6

D 7.1 x 10^-1

10. Consider the following equilibrium: 2H₂O_(l) + energy ⇄ H₃O⁺_(aq)+ OH^-_(aq)

What will cause the pH to increase and the Kw to decrease?

A adding a strong acid

B adding a strong base

C increasing the temperature

D decreasing the temperature

11. The complete neutralization of 15.0 mL of KOH requires 0.0250 moles H₂SO₄. The [KOH] was

A 1.50 M

B 1.67 M

C 3.33 M

D 6.67 M

12. What is the [H₃O⁺] at the equivalence point for the titration between HBr and KOH?

A 1.0 x 10^-9 M

B 1.0 x 10^-7 M

C 1.0 x 10^-5M

D 0.0 M

13. Which of the following would form a buffer solution when equal moles are mixed together?

A HCl and NaCl

B HCN and NaCN

C KNO₃ and KOH

D Na₂SO₄ and NaOH

14. Which of the following acids has the weakest conjugate base?

A HIO₃

B HNO₂

C H₃PO₄

D CH₃COOH

15. When 10.0 ml of 0.10 M HCl is added to 10.0 mL of water, the concentration of H₃O⁺ in the final solution is

A 0.010 M

B 0.050 M

C 0.10 M

D 0.20 M

16. The conjugate base of an acid is produced by

A adding a proton to the acid

B adding an electron to the acid

C removing a proton from the acid

D removing an electron from the acid

17. Which of the following represents the predominant reaction between HCO₃^- and water?

A 2HCO₃^- ⇄ H₂O + 2CO₂

B HCO₃^- + H₂O ⇄ H₂CO₃ + OH^-

C HCO₃^- + H₂O ⇄ H₃O⁺ + CO₃^2-

D 2HCO₃^- + H₂O ⇄ H₃O⁺ + CO₃^2- + OH^- + CO₂

18. Water acts as an acid when it reacts with

I CN^-

II NH₃

III HClO₄

IV CH₃COO^-

A I and IV only

B II and III only

C I, II, and IV

D II, III, and IV

19. In a solution of 0.10 M H₂SO₄, the ions present in order of decreasing concentration are

A [H₃O⁺] > [HSO₄^-] > [SO₄^2-] > [OH^-]

B [H₃O⁺] > [SO₄^2-] > [HSO₄^-] > [OH^-]

C [OH^-] > [SO₄^2-] > [HSO₄^-] > [H₃O⁺]

D [SO₄^2-] > [HSO₄^-] > [OH^-] > [H₃O⁺]

20. Which of the following will dissolve in water to produce an acidic solution?

A CO₂

B CaO

C MgO

D Na₂O

21. Which of the following solutions will have a pH = 1.00?

I 0.10 M HCl

II 0.10 M HNO₂

III 0.10 M NaOH

A I only

B II only

C I and II only

D I, II, and III

22. Ka for the acid H₂AsO₄^- is 5.6 x 10^-8. What is the value of the Kb for HAsO₄^2-?

A 5.6 x 10^-22

B 3.2 x 10^-14

C 1.8 x 10^-7

B 2.4 x 10^-4

23. In a titration, which of the following has a pH = 7.00 at the equivalence point?

A NH₃ and HNO₃

B KOH and HCl

C NaF and HCl

D Ca(OH)₂ and CH₃COOH

24. Which of the following salts dissolves to produce a basic solution?

A KCl

B NH₄Br

C Fe(NO₃)₃

D LiCH₃COO

25. Calculate the pH in a 0.200 M solution of Sr(OH)₂.

A 1.40

B 1.70

C 13.30

D 13.60

26. Which of the following solutions has a pH less than 7.00?

A NaCl

B LiOH

C NH₄NO₃

D KCH₃COO

27. Which of the following will form a basic aqueous solution?

A HSO₃^-

B HSO₄^-

C HPO₄^2-

D HC₂O₄^-

28. What is the approximate Ka value for the indicator chlorophenol red?

A 1 x 10^-14

B 1 x 10^-8

C 1 x 10^-6

D 1 x 10^-3

29. What is the approximate pH of the solution formed when 0.040 mol NaOH is added to 2.00 L of 0.020 M HCl?

A 0.00

B 1.40

C 1.70

D 7.00

30. In which one of the following equations are the Bronsted acids and bases all correctly identified?

Acid + Base ⇄ Base + Acid

A H₂O₂ SO₃^2- ⇄ HO₂^- HSO₃^-

B H₂O₂ SO₃^2- ⇄ HSO₃^- HO₂^-

C SO₃^2- H₂O₂ ⇄ HO₂^- HSO₃^-

D SO₃^2- H₂O₂⇄HSO₃^- HO₂^-

31. Which of the following titrations will always have an equivalence point at a

pH > 7.00?

A weak acid with a weak base

B strong acid with a weak base

C weak acid with a strong base

D strong acid with a strong base

32. A buffer solution may contain equal moles of

A weak acid and strong base

B strong acid and strong base

C weak acid and its conjugate base

D strong acid and its conjugate base

33. A gas which is produced by burning coal and also contributes to the formation of acid rain is

A H₂

B O₃

C SO₂

D C₃H₈

34. Which of the following 1.0 M salt solutions is acidic?

A BaS

B NH₄Cl

C Ca(NO₃)₂

D NaCH₃COO

35. Which of the following statements applies to 1.0 M NH_3(aq) but not to 1.0 M NaOH(aq)?

A partially ionizes

B neutralizes and acid

C has a pH greater than 7

D turns bromcresol green from yellow to blue

36. When the indicator thymol blue is added to 0.10 M solution of an unknown acid, the solution is red. The acid could be

A HF

B H₂S

C HCN

D HNO₃

Subjective

1. Calculate the pH of the solution prepared by mixing 15.0 mL of 0.50 M HCl with 35.0 mL 0.50 M NaOH.

2. Calculate the [OH^-] in 0.50 M NH_3(aq).

3. A titration was performed by adding 0.175 M H₂C₂O₄ to a 25.00 mL sample of NaOH.

The following data was collected.

Trial 1 Trial 2 Trial 3

Final volume of H₂C₂O₄ from burette (mL) 23.00 39.05 20.95

Initial volume of H₂C₂O₄ from burette (mL) 4.85 23.00 5.00

Calculate the [NaOH]

4. A 250.0 mL sample of HCl with a pH of 2.000 is completely neutralized with 0.200 M NaOH. What volume of NaOH is required to reach the stoichiometric point.

5. If the HCl were titrated with 0.200 M NH_3(aq) instead of 0.200 M NaOH, how would the volume of base required to reach the equivalence point compare with the volume calculated in the last question? Explain your answer.

6. Consider the following salt ammonium acetate, NH₄CH₃COO.

a) Write the equation for the dissociation of NH₄CH₃COO.

b) Write the equations for the hydrolysis reactions that occur.

c) Explain why a solution of NH₄CH₃COO has a pH =7.00. Support your answer with a calculation.

7. Consider the following equilibrium: energy + 2H₂O ⇄ H₃O⁺ + OH^-

a) Explain how pure water can have a pH = 7.30.

b) Calculate the value of the Kw for the sample of water with a pH = 7.30.

Worksheet # 1 Properties of Acids and Bases

Acid Base Acid Base

pH = 5

pH = 7

pH =11

Data

Worksheet # 9 pH and pOH Calculations

Worksheet # 14 Amphiprotic Ions and Review

pH pH

Volume of HCl added Volume of HCl added