Acids, Bases And Salts Unit Plan

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

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 .

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 = [H3 O ][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

CH3COOH

NaHCO3

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 ⇄

2- 7. HPO4 + H2O ⇄ (base)

- 2- - 8. HCO3 CO3 9. CH3COO CH3COOH 2- - 10. HPO4 11. IO3 - 12. H2O 13. NH2 - 3- 14. HS 15. C2H5SO7

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.

- 2- - H2O F NH3 SO3 HSO3 NaOH

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

ii) Write the Ka expression for the above.

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

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.

- 2- 1. HSO4 + HPO4 ⇄

2. HCN + H2O ⇄

- 3. HCO3 + H2S ⇄

2- + 4. HPO4 + NH4 ⇄

5. NH3 + H2O ⇄

1- 6. H2PO4 + 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 ⇄

- 2- HS03 + HPO4 ⇄

- 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.

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

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 ⇄ Ka= - F + H2O ⇄ Kb=

16. H2SO4 + NaOH →

17. Define conjugate pairs.

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

. +, - + - - 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:

-2 HClO4 Ka is very large HClO3 Ka=1.2x10 -5 -8 HClO2 Ka=8.0x10 HClO Ka=4.4x10

5. For an oxy acid what is the relationship between the number of O’s and acid strength? (Compare H2S04 and H2S03)

6. Which acid is stronger? HI03 or HIO2

+ - 7. Which produces more H30 ? H2CO3 or 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 ka, 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

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.

1. Formula of compound pH Classification

Reaction or reactions

Worksheet # 9 pH and pOH Calculations

Complete the chart:

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

+ - 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)2 when 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.

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 Kb For Weak Bases

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

- 1. NaNO2 (the basic ion is NO2 )

- -14 -11 Kb(NO2 ) = Kw = 1.0 x 10 = 2.2 x 10 -4 Ka(HNO2) 4.6 x 10

- 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.

[H+] = [OH-] = pH = pOH =

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

2- HPO4

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? 2- - - HPO4 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. SO3(g) + H2O(g) ⇄ H2SO4(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.

2SO2(g) + O2(g) ⇄ 2SO3(g)

Worksheet # 15 Buffers and indicators

Buffers

1. Definition

Acid Conjugate Base Salt

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.

[HInd]

[Ind-]

Color Change

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

[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

Volume of HCl added Volume of HCl added

Acids Unit Midterm Practice Test

1. Consider the following:

- - I H2CO3 + F ⇄ HCO3 + HF - - 2- II HCO3 + HC2O4 ⇄ H2CO3 + C2O4 - - 2- III HCO3 + H2C6H6O7 ⇄ H2CO3 + HC6H5O7 - 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 ⇄ 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

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

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 ⇄ 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 ⇄ 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

2- - - 14. Consider the following equilibrium: HC6H5O7 + HIO3 ⇄ 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) ⇄ 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

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

2- - - 19. Consider the following: SO4 + HNO2 ⇄ 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 ⇄ NaOH + H + + B. H2O + 2Na ⇄ Na2O + 2H 2- 2- C. H2O + CO3 ⇄ H2CO3 + O 2- - - D. H2O + CO3 ⇄ HCO3 + OH

+ - 21. Consider the following equilibrium: 2H2O(l) ⇄ 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 ⇄ 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 ⇄ CH3COOH + OH - 2- + B. CH3COO + H2O ⇄ CH3COO + H3O - + C. CH3COOH + H2O ⇄ CH3COO + H3O + - D. CH3COOH + H2O ⇄ CH3COOH2 + OH

+ - 27. Consider the following equilibrium: 2H2O + 57kJ ⇄ 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 ⇄ 2HClO B. Cl2O + H2O ⇄ 2ClO + H2 C. Cl2O + H2O ⇄ Cl2 + H2O2 D. Cl2O + H2O ⇄ Cl2 + O2 + H2

- 29. The conjugate acid of C6H50 is

- A. C6H4O B. C6H5OH 2- C. C6H4O + 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

-5 H3AsO4 Ka = 5.0 x 10 - -8 H2AsO4 Ka = 8.0 x 10 2- -10 HAsO4 Ka = 6.0 x 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

o + -6 - 34. In a solution at 25 C, the [H3O ] is 3.5 x 10 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

- 2- + A. HSO4 + H2O ⇄ SO4 + H3O - - B. HSO4 + H2O ⇄ H2SO4 + OH + 2- + C. KHSO4 + H2O ⇄ K + SO4 + H3O + - D. KHSO4 + H2O ⇄ 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 ] 2- - [HPO4 ] [H2PO4 ]

+ - + 2 2- C. [K ] [KHPO4 ] D. [K ] [HPO4 ] [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 ⇄ 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:

-5 Barbituric acid HC4H3N2O3 Ka = 9.8 x 10 -10 Sodium propanoate NaC3H5O2 Kb = 7.5 x 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?

2+ - A. Pb (aq) + 2Cl (aq) → PbCl2(s) B. HCl(aq) + NH3(aq) → NH4Cl(aq) C. BaI2(aq) + MgSO4(aq) → BaSO4(s) + MgI2(aq) - 2+ + 2+ 3+ D. MnO4 (aq) + 5Fe (aq) +8H (aq) → Mn (aq) + 5Fe (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 HClO4(aq) with LiOH(aq) is

+ - A. H (aq) + OH (aq) → H2O(l) - - B. HClO4(aq) + OH (aq) → ClO4 (aq) + H2O(l) C. HClO4(aq) + LiOH(aq) → LiClO4(aq) + H2O(l) + - + - D. H (aq) + ClO4 (aq) + Li (aq) + OH (aq) → LiClO4(aq) + H2O(l)

7. The equilibrium constant expression for a sulphurous acid is

+ - A Ka = [H ][HSO3 ]

+ - B. Ka = [H ][HSO3 ] [H2SO3]

+ 2- C. Ka = [2H ][SO3 ] [H2SO3] + 2- D. Ka = [H ][SO3 ] [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 - 2- 2- - HS + SO4 ⇄ S + HSO4

A. HS- and S2- 2- 2- B. SO4 and S - - C. HS and HSO4 2- - D. SO4 and HSO4

- 3. The conjugate acid of H2C6H5O 7 is

3- A. C6H5O7 2- B. HC6H5O7 C. H2C6H5O7 D. H3C6H5O7

4. Which one of the following substances is/are amphiprotic? - 2- (1) H3PO4 (2) H2PO4 (3) HPO4

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) 2+ - B. Ca (aq) + 2Br (aq) ⇄ CaBr2(s) C. 2HBr(aq) + Ca(OH)2(aq) ⇄ 2H2O(l) + CaBr2(s) + - 2+ - 2+ - D. 2H (aq) + 2Br (aq) + Ca (aq) + 2OH (aq) ⇄ 2H2O(l) + Ca (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) ⇄ NH3(aq) + H3O (aq) The stronger base is

+ A. NH4 B. H2O C. NH3 + D. H3O

10. Consider the following equilibrium system: - - 3 OCl (aq) + HC7H5O2(aq) ⇄ HOCl(aq) + C7H5O2 (aq) Keq= 2.1 x 10

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 - - 2- H2BO3 (aq) + HCO3 (aq) ⇄ H2CO3(aq) + HBO3 (aq)

The two species acting as Bronsted-Lowry acids are

- A. HCO3 and H2CO3 - B. H2BO3 and H2CO3 - 2- C. HCO3 and HBO3 - 2- D. H2BO3 and HBO3

- - 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. O2- 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 - 2- + HSeO3 + H2O ⇄ SeO3 + 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

2- 13. The Kb expression for HPO4 is

3- + 2- - A. [PO4 ][H3 O ] B. [HPO4 ][OH ] 2- - [HPO4 ] [H2PO4 ]

- - 2- + C. [H2PO4 ][OH ] D. [HPO4 ][ H3 O ] 2- 3- [HPO4 ] [PO4 ]

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 + 2- A. Na2S ⇄ 2Na + S 2- - - B. S + H2O ⇄ OH + HS C. Na2S + 2H2O ⇄ 2NaOH + H2S + 2- + - D. 2Na + S + 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 2- B. SO3 - C. HSO3 D. H2SO3

9. Consider the following equilibrium expression - K= [H2 S][OH ] [HS-] This expression represents the

A. Kb for H2S B. Ka for H2S - C. Kb for HS - D. Ka 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) + NH3(aq) ⇄ CH3COO (aq) + NH4 (aq) The sequence of Bronsted-Lowry acids 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?

3- A. H3PO4 and PO4 - 3- B. H2PO4 and PO4 2- C. H3PO4 and HPO4 - 2- D. H2PO4 and HPO4

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 25oC, the equation representing the ionization of water is

A H2O + H2O ⇄ 2H2 + O2 B. H2O + H2O ⇄ H2O2 + H2 + 2- C. H2O + H2O ⇄ 4H + 2O + - 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. Kw B. Ka C. Kb D. Ksp

+ 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. NH3(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= pKw – [H3O ]

14. The pH of 0.20 M HNO3 is

A. 0.20 B. 0.63 C. 0.70 D. 1.58

- o 15. The [OH ] in 0.050 M HNO3 at 25 C 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? 3+ - A. Al(OH)3(s) ⇄ Al (aq) + 3OH (aq) + - B. HF(aq) + H2O(l) ⇄ H3O (aq) + F (aq) 2- - 2- C. Cr2O7 (aq) + 2OH (aq) ⇄ 2CrO4 (aq) + H2O(l) + - D. CH3NH2(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 ] Kw + - C. Increase such that [H3O ][OH ] = Kw + - D. Decrease such that [H3O ][OH ] = Kw

7. In a titration, 10.0 mL of H2SO4(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: + - -10 Hlnd ⇄ H + Ind Ka = 1.0 x 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

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

A. An acid because Ka < Kb B. An acid because Ka > Kb C. A base because Ka < Kb D. A base because Ka > Kb

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?

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 Kb is greater than Ka B. basic because Ka is greater than Kb C. acidic because Ka is greater than Kb D. acidic because Kb is greater than Ka

4. The concentration, Ka and pH values of four acids are given in the following table Acid pH ACID Concentration Ka pH X 4.0 Y 2.0 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?

-2 A. Ethyl red, Ka = 8 x 10 Solution 1.0 M HCl 1.0 M HAl 1.0 M HA2 -3 B. Congo red, Ka = 9.0 x 10 Colour Yellow Blue Yellow -8 C. Cresol red, Ka = 1.0 x 10 -11 D. Alizarin blue, Ka = 7.0 x 10

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:

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

A. HA2, HInd, HAl B. HInd, HAl, HA2 C. HA2, HAl, HInd D. HAl, 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. Ka 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

14- 12- 10- 8- 6- 4- 2-

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 2- D. H2C6H5O7 and HC6H5O7

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?

B A. .

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?

- 2- + A. H2PO4 + H2O ⇄ HPO4 + H3O - - B. H2PO4 + H2O ⇄ H3PO4 + OH 2- 3- + C. HPO4 + H2O ⇄ PO4 + H3O 2- - - D. HPO4 + 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.

- - - 2+ CN NH3 F OH Co(H2O)5(OH)

8. Determine the conjugates for each of the acids.

3+ + 2- HF HCN Al(H2O)6 NH4 HPO4

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

+ 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

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 -2 - -2 CO3 KOH H2PO4 HPO4

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 10-3 M 1.3 x 10-

5M 3.1 2.508 neutral (2sig figs)

2. What volume of 0.20 M H2SO4 is 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 10-6 M. 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).

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

- -3 18. A 0.50 M NH3 solution is found to have a OH concentration of 1.86 x 10 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 H2SO4 were 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 = [H3 O ][OH ] [H2O]

+ 2 B. Kw = [H3O ] [O2]

+ - C. Kw = [H3O ][OH ]

+ 2 2- D. Kw = [H3O ] [O ]

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

7 Volume HCl added I

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

-2 6. A sample containing 1.20 x 10 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: - - 2- HC2O4 + H2BO3 ⇄ H3BO3 + C2O4 Identify the Bronsted-Lowry bases in this equilibrium.

- A. H2BO3 and H3BO3 - B. HC2O4 and H3BO3 - 2- C. HC2O4 and C2O4 - 2- D. H2BO3 and C2O4

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

- - A. CH3COO + H2O ⇄ CH3COOH + OH - 2- B. CH3COO + H2O ⇄ H2O + CH2COO + - 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 2- B. CO3 - 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

+ 2- A. H2O ⇄ 2H + O 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

I. H2SO4 - II. HSO4 2- III. SO4

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

2- 23. The conjugate acid of HPO4 is

3- A. PO4 - B. H2PO4 2- C. H2PO4 3- D. H2PO4

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) 3+ A small amount of Fe(H2O)6 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

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

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 NH3(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 3- - 2- D CH3COOH + PO4 ⇄ CH3COO + HPO4

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 10-5 M 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 - + 2- C HCO3 + H2O ⇄ H3O + CO3 - + 2- - D 2HCO3 + H2O ⇄ H3O + CO3 + 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

+ - 2- - A [H3O ] > [HSO4 ] > [SO4 ] > [OH ] + 2- - - B [H3O ] > [SO4 ] > [HSO4 ] > [OH ] - 2- - + C [OH ] > [SO4 ] > [HSO4 ] > [H3O ] 2- - - + D [SO4 ] > [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

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

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 2- C HPO4 - 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

2- - - A H2O2 SO3 ⇄ HO2 HSO3 2- - - B H2O2 SO3 ⇄ HSO3 HO2 2- - - C SO3 H2O2 ⇄ HO2 HSO3 2- - - D SO3 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 NH3(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 NH3(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 NH3(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.

Acids, Bases and Salts Unit Plan

Notes- double click on the lesson number and download Power Point Viewer if you do not have it.

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 ClassifyingEverything Activity 8. Yamada’s Indicator Lab. Hydrolysis. WS 8 4 9. Ionization of Water, [H+] & [ OH- ], pH scale. Hydrolysis Quiz 1 Hydrolysis Quiz 2 Hydrolysis Quiz 3 10. pH for Strong Acids and Bases. Hydrolysis Quiz WS 9 11. pH Calculations for Weak Acids. WS 10 5 12. Ka from pH for Weak Acids. WS 11 13. Indicators Lab. 14. Kbs from Kas for Weak Bases. WS 12 6 15. pH for Weak Bases pH [H+] [ OH-] Relationships. WS 13 16.Amphiprotic Ions- Kas and Kbs. WS 14 7 17. Titration Lab. Primary Standards. Acids Midtermreview Test 18. Titration Lab 19. Buffers & Indicators WS 15 8 20. Titration Curves. WS 16 9/ 10 21. Review #1 Web Site Review Sheet Quizmebc 22. Review #2 Practice Test # 1 Practice Test 2 23. Test

Text book Hebden Read Unit IV

WS #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 Bromothymol Blue Acid or Base

Solution: HCl Clear Red Yellow Acid

NaOH Pink Blue Blue Base

Vinegar Clear Red Yellow Acid

Ammonia Pink Blue Blue Base (NH3) Lemon Juice Clear Red Yellow Acid

Seven-up Clear Red Yellow Acid

Baking Soda Pink Blue Blue Base (NaHCO3)

Indicator Acid Colour Base Colour

Phenolphthalein Clear Pink

Litmus Red Blue

Bromothymol Blue Yellow 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 Mg 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. 2HCl(aq) + Mg(s) → H2(g) + MgCl2(aq) Wash and dry your fleaker before going on to step 3. 3. Taste a lemon and describe the taste in one word Sour 4. Taste some baking soda and describe the taste in one word. Bitter 6. Test two drops of HCl for conductivity in a spot plate. Result: Good Conductor Write an equation that accounts for the conductivity of HCl.

HCl → H+ + Cl-

7. Test two drops of NaOH for conductivity in a spot plate. Result: Good Conductor Write an equation that accounts for the conductivity of NaOH (dissociation). NaOH → Na+ + OH- Clean, dry and put away the spot plate 8. List five properties of acids that are in your textbook. Acids conduct electricity, taste sour, neutralize bases, change the color of indicators, and react with some metals to produce hydrogen.

9. List five properties of bases that are in your textbook. Bases conduct electricity, taste bitter, neutralize acids, change the color of indicators, and feel slippery. 10. Make some brief notes on the commercial acids: HCl and H2SO4 (p 112). HCl

11. Make some brief notes on the commercial base NaOH (p 114).

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

13. Describe the difference between a strong and weak acid (p 121-124). Use two examples and write equations to support your answer. Describe their relative conductivities. A strong acid completely ionizes and a weak acid partially ionizes.

14. Describe a situation where a strong acid would have the same conductivity as a weak acid (hint: think about concentration). A weak acid could have a high molarity and the strong acid could have a low molarity.

Complete this worksheet for next period. Read pages 107-126 for homework.

WS #2 Conjugate Acid-Base Pairs

Complete each acid reaction. Label each reactant or 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 ⇄ H2C6O7 + H3O acid base base acid

- + 3. H3PO4 + H2O ⇄ H2PO4 + H3O acid base base acid

- + 4. HF + H2O ⇄ F + H3O acid base base acid

- + 5. H2CO3 + H2O ⇄ HCO3 + H3O acid base base acid

+ + 6. NH4 + H2O ⇄ NH3 + H3O acid base base acid

- + 7. CH3COOH + H2O ⇄ CH3COO + H3O acid base base acid

- + 8. HCl + H2O → Cl + H3O acid base base acid

+ - 9. HNO3 + H2O → H3O + NO3 acid base acid base

Write the equilibrium expression (Ka) for the first seven above reactions.

+ - + - 10. Ka = [H3O ] [ CN ] 14. Ka = [H3O ] [HCO3 ] [HCN] [H2CO3]

+ - + 11. Ka = [H3O ] [H2C6O7 ] 15. Ka = [H3O ] [NH3] + [H3C6O7] [NH4 ]

+ - + - 12. Ka = [H3O ] [H2PO4 ] 16. Ka = [H3O ] [CH3COO ] [H3PO4] [CH3COOH]

+ - 13. Ka = [H3O ] [F ] [HF] 17. Which acids are strong? The six on the top of the acid chart are strong. 18. What does the term strong acid mean? They complete ionization into ions. Such as: HCl + H2O - + → Cl + H3O

+ - 19. Why is it impossible to write an equilibrium expression for a strong acid? Ka = [H3O ] [Cl ] [HCl] is equal to zero and in math numbers divided by zero are undefined. [HCl] 20. Which acids are weak?

All acids listed on the acid chart below the top six. 23. What does the term weak acid mean? - + Incomplete ionization. Such as: HF + H2O ⇄ F + H3O 24. Explain the difference between a strong and weak acid in terms of electrical conductivity. A strong acid is a good conductor. A weak acid conducts but not so good.

- - 14. HNO2 NO2 15. HCOO HCOOH - 2- - 16. HSO3 SO3 17. IO3 HIO3 - + 18. H2O2 HO2 19. NH3 NH4 - 2- - 20. HS S 21. CH3COO CH3COOH - + 22. H2O OH 23. H2O H3O

Define:

22. Bronsted acid- a proton donor 23. Bronsted base- a proton acceptor 24. Arrhenius acid- a substance that ionizes in water to produce H+

25. Arrhenius base- a substance that ionizes in water to produce OH- 26. List the six strong acids. HCLO4 HI HBr HCl HNO3 H2SO4 27. Rank the acids in order of decreasing strength. - HCl HSO4 H3PO4 HF H2CO3 H2S 28. What would you rather drink vinegar or hydrochloric acid? Explain. + Vinegar. It is a weak acid and produces much less H30 ion which is the corrosive part of an acid. 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 8 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 acetate. Wash and return your eyedropper. Wash and return your beaker.

Wash your hands.

Results Compound Conductivity pH Classification

HCl good 1 strong acid

CH3COOH ok 3 weak acid

NaOH good 13 strong base

NH3 ok 11 weak base

NaHCO3 good 11 weak base

H2CO3 ok 3 weak acid

NaCl good 7 neutral salt

WS #3 Conjugate Acid-Base Pairs Complete each reaction. Label each reactant or product as an acid or base.

+ - 1. HCN + H2O ⇄ H3O + CN

+ - 2. HCl + H2O ⇄ H3O + Cl

+ - 3. HF + H2O ⇄ H3O + F

- - 4. F + H2O ⇄ HF + OH

- + 2- 5. HSO4 + H2O ⇄ H3O + SO4 (acid)

+ + 6. NH4 + H2O ⇄ H3O + NH3

2- - - 7. HPO4 + H2O ⇄ H2PO4 + OH (base)

- 2- - 8. HCO3 CO3 9. CH3COO CH3COOH -2 3- - 10. HPO4 PO4 11. IO3 HIO3 - - 12. H2O OH 13. NH2 NH3 - 2- 3- 2- 14. HS S 15. C2H5SO7 HC2H5SO7

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 5 4 6 1 1 3

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

- 2- - H2O F NH3 SO3 HSO3 NaOH 6 4 2 3 5 1

+ 19. i) Write the reaction of H3BO3 with water (remove one H only because it is a weak acid). - + H3BO3 + H2O ⇄ H2BO3 + H3O

ii) Write the Ka expression for the above. + - [H3O ] [H2BO3 ] Ka = [H3BO3]

-10 iii) What is the ionization constant for the acid (use your table). Ka = 7.3 x 10

20. List six strong acids. HClO4 HI HBr HCl HNO3

21. List six strong bases. NaOH KOH LiOH RbOH CsOH Ba(OH)2

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

- HIO3 H2C2O4 H2SO3 HSO4 H3PO4 HNO2

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

3- 2- - - - PO4 CO3 CN NH3 H2BO3 HS

WS #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. - 2- 2- - 1. HSO4 + HPO4 ⇄ SO4 + H2PO4 Acid Base Base Acid - - Products are favoured as HSO4 is a stronger acid than H2PO4 + - 2. HCN + H2O ⇄ H3O + CN Acid Base Acid Base + Reactants are favoured as H3O is a stronger acid than HCN. - - 3. HCO3 + H2S ⇄ H2CO3 + HS Base Acid Acid Base

Reactants are favoured as H2CO3 is a stronger acid than H2S 2- + - 4. HPO4 + NH4 ⇄ H2PO4 + NH3 Base Acid Acid Base - + Reactants are favoured as H2PO4 is a stronger acid than NH4 + - 5. NH3 + H2O ⇄ NH4 + OH Base Acid Acid Base - Reactants are favoured as OH is a stronger base than NH3 1- 2- + 6. H2PO4 + NH3 ⇄ HPO4 + NH4 Acid Base Base Acid

2- Products are favoured as NH3 is a stronger base than HPO4 - - 7. HCO3 + HF ⇄ H2CO3 + F Base Acid Acid Base Products are favoured as HF is a stronger acid than H2CO3

8. Complete each equation and indicate if reactants or products are favoured. Label each acid or base. - - 2- HSO4 + HCO3 ⇄ H2CO3 + SO4 products are favoured

- 2- H2PO4 + HC03 ⇄ HPO4 + H2CO3 reactants are favoured

- 2- - 2- HS03 + HPO4 ⇄ H2PO4 + SO3 products are favoured

- + 2- NH3 + HC2O4 ⇄ NH4 + C2O4 products are favoured

9. Explain why HF(aq) is a better conductor than HCN(aq). HF is a stronger acid and creates more ions. 10. Which is a stronger acid in water, HCl or HI? Explain! Both are strong acids and have the same strength as both completely ionize to from H+. 11. State the important ion produced by an acid and a base. + + - Acid: H or H3O Base: OH 12. Which is the stronger base? Which produces the least OH-? - -2 F is the weaker base and produces the least OH- CO3 is the stronger base 13. Define a Bronsted/Lowry acid and base. An acid is a proton donor and a base is a proton acceptor.

14. Define an Arrhenius acid and base. An acid ionizes in water to produce H+ and a base ionizes in water to produce OH-. 15. Complete each reaction and write the equilibrium expression. + - + - HF + H2O ⇄ H3O + F Ka= [ H3O ][ F ] Kb= [HF] [ OH-] - - - F + H2O ⇄ HF + OH [HF] [F ]

16. H2SO4 + 2NaOH → Na2SO4 + 2HOH

17. Define conjugate pairs. Acid base pairs that differ by one proton. - -2 - 18. Give conjugate acids for: HS , NH3, HPO4 , OH , H2O, NH3, -2 CO3 + - + + - H2S NH4 H2PO4 HOH H3O NH4 HCO3 . +, - + - - 19 Give conjugate bases for: NH4 HF, H2PO4 , H3O , OH , HCO3 , H2O - -2 2- -2 NH3 F HPO4 HOH O CO3 OH-

WS #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. + + - H3O HCl + H2O → H3O + Cl 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. OH- NaOH → Na+ + OH-

3. List three strong acids and three strong bases. HCl HI HClO4 NaOH KOH LiOH

4. Rank the acids in decreasing strength: -2 HClO4 1 Ka is very large HClO3 2 Ka=1.2x10 -5 -8 HClO2 3 Ka=8.0x10 HClO 4 Ka=4.4x10

5. For an oxy acid what is the relationship between the number of O’s and acid strength? (Compare H2S04 and H2S03) The more O’s the stronger the acid.

6.Which acid is stronger? HI03 or HIO2

+ - 7.Which produces more H30 ? H2CO3 or HS04

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

9.Which conducts better NH3 or NaOH (both .1M)? Why? NaOH is a strong acid.

10.Which conducts better HF or HCN (both .1M)? Why? HF is a stronger acid. 11. Compare and contrast a strong and weak acid in terms of degree of ionization, size of ka, conductivity, and concentration of H+. Strong acid: complete ionization, very large Ka, good conductor, high [H+]. Weak acid: partial ionization, small Ka, OK conductor, low [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 basic anhydride Na2O + H2O → 2NaOH

13. CaO basic anhydride CaO + H2O → Ca(OH)2

14. SO3 acid anhydride SO3 + H2O → H2SO4

15. CO2 acid anhydride CO2 + H2O → H2CO3

16. SO2 acid anhydride SO2 + H2O → H2SO3

+ - 17. HCl strong acid HCl + H2O → H3O + Cl

+ - 18. NH3 weak base NH3 + H2O ⇄ NH4 + OH

19. NaOH strong base NaOH → Na+ + OH-

+ - 20. HF weak acid HF + H2O ⇄ H3O + F

+ - 21. H3PO4 weak acid H3PO4 + H2O ⇄ H3O + H2PO4

WS # 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 weak base

+ - NH3 + H2O ⇄ NH4 + OH

2. KCl neutral salt

HCl + KOH → KCl + H2O

KCl → K+ + Cl-

3. HNO3 strong acid

+ - HNO3 + H2O → H3O + NO3

4. NaHCO3 basic salt

H2CO3 + NaOH → NaHCO3 + H2O

+ - NaHCO → Na + HCO3 - - HCO3 + H2O ⇄ H2CO3 + OH

5. RbOH strong base

RbOH → Rb+ + OH-

6. AlCl3 acid salt

3HCl + Al(OH)3 → AlCl3 + 3H2O

+3 - AlCl3 → Al + 3Cl

3+ 2+ + Al(H2O)6 ⇄ Al(H2O)5(OH) + H

7. H2C2O4 weak acid

+ - H2C2O4 + H2O ⇄ H3O + HC2O4

8. NaC6H5O basic salt

C6H5OH + NaOH → NaC6H5O + H2O

+ - NaC6H5O → Na + C6H5O

- - C6H5O + H2O ⇄ C6H5OH + OH

9. Co(NO3)3 acid salt

3HNO3 + Co(OH)3 → Co(NO3)3 + 3H2O

+3 - Co(NO3)3 → Co + 3NO3

3+ 2+ + Co(H2O)6 ⇄ Co(H2O)5(OH) + H

10. Na2CO3 basic salt

H2CO3 + 2NaOH → Na2CO3 + 2H2O

+ -2 Na2CO3 → 2Na + CO3

-2 - - CO3 + H2O ⇄ HCO3 + OH

WS # 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 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 weak base

+ - NH3 + H2O ⇄ NH4 + OH

2. NaCl neutral salt

NaCl → Na+ + Cl-

3. HCl strong acid

+ - HCl + H2O → H3O + Cl

4. NaCN basic salt

NaCN → Na+ + CN- - - CN + H2O ⇄ HCN + OH

5. NaOH strong base

NaOH → Na+ + OH-

6. FeCl3 acid salt

+3 - FeCl3 → Fe + 3Cl

3+ 2+ + Fe(H2O)6 ⇄ Fe(H2O)5(OH) + H

7. HF weak acid

+ - HF + H2O ⇄ H3O + F

8. LiHCO3 basic salt + - LiHCO3 → Li + HCO3 - - HCO3 + H2O ⇄ H2CO3 + OH

9. Fe(NO3)3 acid salt

+3 - Fe(NO3)3 → Fe + 3NO3

3+ 2+ + Fe(H2O)6 ⇄ Fe(H2O)5(OH) + H

10. MgCO3 basic salt

+2 -2 MgCO3 → Mg + CO3 -2 - - CO3 + H2O ⇄ HCO3 + OH

11. H2S weak acid

+ - H2S + H2O ⇄ H3O + HS

12. HF weak acid

+ - HF + H2O ⇄ H3O + F

13. CaI2 neutral salt

+2 - CaI2 → Ca + 2I

14. Mg(OH)2 weak base

+2 - Mg(OH)2 ⇄ Mg + 2OH

15. Ba(OH)2 strong base

+2 - Ba(OH)2 → Ba + 2OH

16. Describe why Tums (CaCO3) neutralizes stomach acid. It is a weak base and will neutralize acid. basic salt +2 -2 CaCO3 → Ca + CO3 -2 - - CO3 + H2O ⇄ HCO3 + OH

17. Describe why Mg(OH)2 is used in Milk of Magnesia as an antacid instead of NaOH. - - Mg(OH)2 is weak base and releases OH slowly, whereas NaOH is a strong base which releases OH in high concentrations which is corrosive. +2 - Mg(OH)2 ⇄ Mg + 2OH NaOH → Na+ + OH-

WS #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 Fe(NO3)3 pH 2 Classification acid salt

+3 - Reaction or reactions Fe(NO3)3 → Fe + 3NO3

3+ 2+ + Fe(H2O)6 ⇄ Fe(H2O)5(OH) + H

2. Formula of compound NaCH3COO pH 10 Classification basic salt

+ - Reaction or reactions NaCH3COO → Na + CH3COO

- - CH3COO + H2O ⇄ CH3COOH + OH 3. Formula of compound K2HPO4 pH 10 Classification basic salt

+ -2 Reaction or reactions K2HPO4 → 2K + HPO4

-2 - - HPO4 + H2O ⇄ H2PO4 + OH

4. Formula of compound HCL pH 0 Classification strong acid

Reaction or reactions HCl → H+ + Cl-

5. Formula of compound Al2(SO4)3 pH 3 Classification acid salt

+3 -2 Reaction or reactions Al2(SO4)3 → 2Al + 3SO4

3+ 2+ + Al(H2O)6 ⇄ Al(H2O)5(OH) + H

6. Formula of compound Na2CO3 pH 12 Classification basic salt

+ -2 Reaction or reactions Na2CO3 → 2Na + CO3

-2 - - CO3 + H2O ⇄ HCO3 + OH

7. Formula of compound P2O5 pH 2 Classification acid anhydride

Reaction or reactions P2O5 + H2O → H2P2O6

+ - H2P2O6 ⇄ H + HP2O6

8. Formula of compound Cu(NO3)2 pH 4 Classification acid salt

2+ - Reaction or reactions Cu(NO3)2 → Cu + 2NO3 2+ Cu(H2O)6 <⇄ + + Cu(H2O)5(OH) + H

9. Formula of compound Fe2(SO4)3 pH 3 Classification acid salt

+3 -2 Reaction or reactions Fe2(SO4)3 → 2Fe + 3SO4

3+ 2+ + Fe(H2O)6 ⇄ Fe(H2O)5(OH) + H

10. Formula of compound N2O5 pH 0 Classification acid anhydride

Reaction or reactions N2O5 + H2O → H2N2O6 → 2HNO3

+ - HNO3 → H + NO3

11. Formula of compound Ca(OH)2 pH 12 Classification weak base

+2 - Reaction or reactions Ca(OH)2 ⇄ Ca + 2OH

12. Formula of compound KHSO4 pH 2 Classification acid salt

+ - Reaction or reactions KHSO4 → K + HSO4

- + 2- HSO4 ⇄ H + SO4

13. Formula of compound NaHCO3 pH 10 Classification basic salt

+ - Reaction or reactions NaHCO3 → Na + HCO3

- - HCO3 + H2O ⇄ H2CO3 + OH 14. Formula of compound CaCO3 pH 10 Classification basic salt

+2 -2 Reaction or reactions CaCO3 → Ca + CO3

-2 - - CO3 + H2O <⇄ HCO3 + OH

15. Formula of compound CaO pH 12 Classification basic anhydride

Reaction or reactions CaO + H2O → Ca(OH)2

+2 - Ca(OH)2 ⇄ Ca + 2OH

16. Formula of compound Al2(SO4)3 pH 3 Classification acidic salt

+3 -2 Reaction or reactions Al2(SO4)3 → 2Al + 3SO4

3+ 2+ + Al(H2O)6 ⇄ Al(H2O)5(OH) + H 17. Formula of compound NaCl pH 7 Classification neutral salt

Reaction or reactions NaCl → Na+ + Cl-

WS # 9 - pH and pOH Calculations Complete the chart:

[H+] [OH-] pH pOH Acid/base/neutral 1. 7.00 x 10-3 M 1.43 x 10- 2.155 11.845 acid 12M 2. 1.14 x 10- 8.75 x 10-2 12.942 1.058 base 13M M 3. 4.7x 10-8M 2.1 x 10-7M 7.33 6.67 base 4. 1.0 x 10-10M 1.0 x 10-4M 10.00 4.00 base 5. 1.0 x 10-7M 1.0 x 10-7M 7.00 7.00 Neutral (2sig figs) 6. 5 x 10-4M 2 x 10-11M 3.3 10.7 acid 7. 2.80 x 10-3M 3.57 x 10- 2.553 11.447 acid 12M 8. 5.0 x 10-10 M 2.0 x 10-5M 9.30 4.70 base 9. 2.1 x 10-5M 4.7 x 10-10 M 4.67 9.33 acid

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

+2 - Ba(OH)2 ⇄ Ba + 2OH 0.20M 0.20M 0.40M

[OH-] = 0.40 M [H+] = 2.5 x 10-14 M pH = 13.60 pOH = 0.40

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

HCl → H+ + Cl- 0.030M 0.030M

[H+] = 0.030M [OH-] = 3.3 x 10-13 M pH = 1.52 pOH = 12.48

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

+ - NaOH → Na + OH 0.20M 0.20M 0.20M

[OH-] = 0.20 M [H+] = 5.0 x 10-14 M pH = 13.30 pOH = 0.70

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

HCl → H+ + Cl- 300.0 x 0.20 M = 0.075 M 0.075 M pH = -Log[H+] = 1.12 800.0

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

HCl → H+ + Cl- 200.0 x 0.020 M = 0.0080 M 0.0080 M pH = -Log[H+] = 2.10 500.0

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.

2+ - Ba(OH)2 → Ba + 2OH

150.0 x 0.40 M = 0.12 M 0.12 M 0.24 M 500.0

pOH = -Log[OH-] = 0.62 pH = 14.00 - pOH = 13.38

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

2+ - Sr(OH)2 → Sr + 2OH

250.0 x 0.20 M = 0.083 M 0.083 M 0.1667 M 600.0

pOH = -Log[OH-] = 0.78 pH = 14.00 - pOH = 13.22

17. Calculate the pH of a saturated solution of 0.40M Ba(OH)2 when 25 mL was added 25.0 mL of water.

2+ - Ba(OH)2 ⇄ Ba + 2OH (25)0.40 M 0.20 M0.40 M (50)

[OH-] = 0.40

pOH = 0.40

pH = 13.60

WS # 10 pH Calculations for Weak Acids 1. Calculate the [H+], [OH-], pH, and pOH for 0.20 M HCN.

HCN ⇄ H+ + CN-

I 0.20 M 0 0

C x x x

E 0.20 - x x x

x2 = 4.9 x 10-10 0.20 - x

x = 9.9 x 10-6 M

[H+] = 9.9 x 10-6 M [OH-] = 1.0 x 10-9 M pH = 5.00 pOH = 9.00

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

[H+] = 2.8 x 10-2 M [OH-] = 3.6 x 10-13 M pH = 1.56 pOH = 12.44 + - 3. Calculate the [H ], [OH ], pH, and pOH for 0.805 M CH3COOH.

[H+] = 3.8 x 10-3 M [OH-] = 2.6 x 10-12 M pH = 2.42 pOH = 11.58 + - 4. Calculate the [H ], [OH ], pH, and pOH for 1.65 M H3BO3.

[H+] = 3.5 x 10-5 M [OH-] = 2.9 x 10-10 M pH = 4.46 pOH = 9.54

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

2+ - Mg(OH)2 ⇄ Mg + 2OH x x 2x

Ksp = [Mg2+][OH-]2

5.6 x 10-12 = 4x3

[OH-] = 2x = 2.22 x 10-4 M

pH = 10.35

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

+ - H2X ⇄ H + X Note- only lose one proton for any weak acid!!

I 0.200 M 0 0

C x x x

E 0.20 - x x x

Small Ka approximation x = 0

x 2 = 1.8 x 10-6 0.20

x = 6.0 x 10-4 M

[H+] = 6.0 x 10-4 M [OH-] = 1.7 x 10-11 M pH = 3.22 pOH = 10.78

2+ - Sr(OH)2 → Sr + 2OH

350.0 x 0.20 M = 0.0875 M 0.0875 M 0.175 M 800.0

pOH = -Log[OH-] = 0.76 pH = 14.00 - pOH = 13.24

WS # 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.

[H+] = 10-pH = 10-5.00 = 0.0000100 M

HCN ⇄ H+ + CN-

I 0.20 M 0 0

C 0.0000100 M 0.0000100 M 0.0000100 M

E 0.19999 0.0000100 M 0.0000100 M

Ka = (0.0000100)2 = 5.0 x 10-10

0.19999

Ka = 5.0 x 10-10 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.

Ka = 3.5 x 10-4 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.

Ka = 1.8 x 10-5 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.

Ka = 7.3 x 10-10

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

[H+] = 10-pH = 10-3.683 = 0.0002075 M

+ - H2X ⇄ H + HX Note a diprotic weak acid only loses one proton.

I 0.10 M 0 0

C 0.0002075 M 0.0002075 M 0.0002075 M

E 0.09979 0.0002075 M 0.0002075 M

Ka = (0.0002075)2 = 4.3 x 10-7 0.09979

-7 Ka = 4.3 x 10 Carbonic acid H2CO3 Look up on Ka Table.

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

pOH = 14.00 - pH = 2.773

[OH-] = 10-pOH = 0.001686 M

+ - NH3 + H2O ⇄ NH4 + OH

I 0.20 M 0 0

C 0.001686 M 0.001686 M 0.001686 M

E 0.1983 M 0.001686 M 0.001686 M

Kb= (0.001686)2 = 1.4 x 10-5 0.1983

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

pOH = 14.00 - pH = 2.544

[OH-] = 10-pOH = 0.002858 M

- - CN + H2O ⇄ HCN + OH

I 0.40 M 0 0

C 0.002858 M 0.002858 M 0.002858 M

E 0.3971 M 0.002858 M 0.002858 M

Kb= (0.002858)2 = 2.0 x 10-5 0.3971

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

[H+] = 10-pH = 10-5.068 = 8.55 x 10-6 M

+ - H3X ⇄ H + H2X Note a triprotic weak acid only loses one proton.

I 0.10 M 0 0

C 8.55 x 10-6 M8.55 x 10-6 M8.55 x 10-6 M

E 0.10 M 8.55 x 10-6 M8.55 x 10-6 M

Ka = (8.55 x 10-6)2 = 7.3 x 10-10 0.10

-10 Ka = 7.3 x 10 Boric acid H3BO3 Look up on Ka Table.

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

[H+] = 10-2.45 = 0.003548 M

+ - CH3COOH ⇄ H + CH3COO

I x 0 0

C 0.003548 M 0.003548 M 0.003548 M

E x - 0.003548 M 0.003548 M 0.003548 M

+ - Keq = [H ][CH3 COO ] [CH3COOH] 1.8 x 10-5 = (0.003548)(0.003548) [CH3COOH]

[CH3COOH] = 0.6994 M 2.00 L x 0.6994 moles x 60.0 g = 84 g 1 L 1 mole

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

WS # 12 Kb For Weak Bases Determine the Kb for each weak base. Write the ionization reaction for each. Remember that Kw = Ka • Kb (the acid and base must be conjugates). Find the base on the right side of the acid table and use the Ka values that correspond. Be careful with amphiprotic anions! - 1. 1. NaNO2 (the basic ion is NO2 ) 2. - -14 Kb(NO2 ) = Kw = 1.0 x 10 -4 Ka(HNO2) 4.6 x 10 3. Kb = 2.2 x 10-11

- -10 2. 2. KCH3COO (the basic ion is CH3COO ) Kb = 5.6 x 10

-8 3. 3. NaHCO3 Kb = 2.3 x 10

-5 4. NH3 Kb = 1.8 x 10

5. NaCN Kb = 2.0 x 10-5

-7 6. Li2HPO4 Kb = 1.6 x 10

-12 7. KH2PO4 Kb = 1.3 x 10

-4 8. K2CO3 Kb = 1.8 x 10

[H+] = 2.9 x 10-4 M [OH-] = 3.4 x 10-11 M pH = 3.53 pOH = 10.47

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.

Ka = 4.4 x 10-7

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

[H+] = 1.3 x 10-3 M [OH-] = 7.5 x 10-12 M pH = 2.87 pOH = 11.13

12. The pH of 0.10 M CH3COOH is 2.87. Calculate the Ka.

[H+] = 10-2.87 = 0.001349 M

+ - CH3COOH ⇄ H + CH3COO

I 0.10 M 0 0

C 0.001349 M 0.001349 M 0.001349 M

E 0.09865 M 0.001349 M 0.001349 M

+ - Ka = [H ][CH3 COO ] [CH3COOH] Ka = (0.001349)( 0.001349) (0.09865)

Ka = 1.8 x 10-5

 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. 

 H2SO4 + 2NaOH ® Na2SO4 + 2HOH   0.200 L x 0.120 mol = 0.0240 mol 0.400 L x 0.140 mol = 0.0560 mol  L L   I 0.0240 mole 0.0560 mole Note the loss of sig figs!  Beware subtraction!  C 0.0240 mole 0.0480 mole   E 0 0.0080 mole    [OH-] = 0.0080 mole = 0.013 M  0.6000 L    Note the final volume is used as the  two solutions are mixed together.  200.0 mL + 400.0 mL = 600.0 mL   pOH = 1.88   pH = 12.12 WS # 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-

Kb(CN-) = Kw = 1.0 x 10-14 = 2.0408 x 10-5

Ka(HCN) 4.9 x 10-10

- - CN + H2O ⇄ HCN + OH I 0.20 0 0 C x x x E 0.20 - x x x

x2 = 2.0408 x 10-5

0.20 - x

x = [OH-] = 2.0 x 10-3 M

[OH-] = 2.0 x 10-3 M pOH = 2.69 pH = 11.31 [H+] = 4.9 x 10-12 M

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

Kb = 1.1 x 10-7 [OH-] = 3.3 x 10-5 M pOH = 4.48 pH = 9.52 [H+] = 3.0 x 10-10 M

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 HX is 3.683, calculate the Ka for the acid and identify the acid using your acid chart.

HX ⇄ H+ X-

I 0.100 M 0 0

C - 0.0002075 0.0002075 0.0002075

E 0.09979 0.0002075 0.0002075

Ka = (0.0002075)2 = 4.3 x 10-7 Carbonic acid (0.09979)

[H+] = 2.4 x 10-5 M [OH-] = 4.1 x 10-10 M pH = 4.62 pOH = 9.38

[H+] = 3.3 x 10-4 M [OH-] = 3.0 x 10-11 M pH = 3.48 pOH = 10.52

Ka = 7.3 x 10-10 [OH-] = 2.88 x 10-10 M pH = 4.46 [H+] = 3.47 x 10-5 pOH = 9.54

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

pOH = 14.00 - 12.46 = 1.54 [OH-] = 10-1.54 = 0.02884 M

- - CN + H2O ⇄ HCN + OH I 0.65 M 0 0 C 0.02884 M 0.02884 M 0.02884 M E 0.6212 M 0.02884 M 0.02884 M

(0.02884)2 Kb = (0.6212)

Kb = 1.3 x 10-3

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 Reject #2 35.49 mL #3 35.51 mL Calculate the original [HCl] = 1.04M

35.50 mL Average

2HCl + Ba(OH)2 → BaCl2 + 2H2O 0.03550 L 3.16 g 3.16 g Ba(OH)2 x 1 mole x 2 moles HCl Molarity = 171.3g 1 mole Ba(OH)2

0.03550 L

[HCl] = 1.04M

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

0.00625 L

13. 25.0 ml of .200M HCl is mixed with 50.0 ml .100M NaOH, calculate the pH of the resulting solution.

No excess pH = 7.000

14. 10 ml of 0.10M H2SO4 is mixed with 25 ml 0.20M NaOH, calculate the pH of the resulting solution.

pH = 12.456

15. 125.0 ml of .200M HCl is mixed with 350.0 ml .100M NaOH, calculate the pH of the resulting solution.

pH = 12.323

16. Define standard solution and describe two ways to standardize a solution. A standard solution is one of known molarity. If you make the solution from a weighed amount of solid and dilute it to a final volume in a volumetric flask it is a standard solution. If you titrate a solution to determine its concentration it is a standard 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? [H+] = 6.3 x 10-13 M

WS # 14 Review

1. List the properties of acids/bases. Acids- conduct electricity, taste sour, change the color of indicators, neutralize bases, react with active metals like Mg to produce H2 gas. Bases- conduct electricity, taste bitter, change the color of indicators, neutralize acids, feel slippery. 2. Define the following: Arhenius strong acid- completely ionizes to form H+ Arhenius weak base- partially ionizes to form OH- Bronsted strong acid- completely donates a proton to a base Bronsted weak base- partially accepts a proton to an acid Conjugate pair – an acid base pair that differs by one proton Amphiprotic- a chemical species that can be an acid or base Standard solution- a solution of known molarity

3. Show by calculation if the following amphiprotic ions are acids or bases: a) - -11 -8 HCO3 Base Ka = 5.6 x 10 Kb = 2.3 x 10 b) - -8 -12 H2PO4 Acid Ka = 6.2 x 10 Kb = 1.3 x 10 c) 2- -13 -7 HPO4 Base Ka = 2.2 x 10 Kb = 1.6 x 10

4. What is the strongest base in water? What is the strongest acid in water? Write equations to explain your answer. Base OH- NaOH → Na+ + OH- Acid H+ HCl → H+ + Cl-

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/weak b) high/low ionization c) large or small Ka d) good/poor conductor e) strong or weak electrolyte 7. 0.2M HCl and 1.0M 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 SA WA WB SB

9. Which ions are amphiprotic? -2 - - HPO4 HCl F HS H2S H2O

10. Write the net ionic equation between any acid and base. H+ + OH- → HOH

+ - 11. Write the ionization equation for water. H20 → H + OH

12. Write the Kw expression. Kw = [H+][OH-] = 1.0 x 10-14

- - 13. H2SO3 + HS <====> H2S + HSO3 a) Are the reactants or products favoured? b) Are the Keq large, small or about 1?

14. 0.20M HCl pH = 0.70

15. 0.20M Ba(OH)2 pH = 13.60

16. 0.20M H2CO3 pH = 3.53

17. 0.40M KHCO3 pH = 9.98

18. The pH increases by 2 units. How does [H+] change? Decreases by a factor of 100

19. The pH decreases by 1 unit. How does [H+] change? Increases by a factor of 10

20. a) For distilled water : pH = 7.00 pOH =7.00 [H+] = 1.0 x 10-7 M [OH-] = 1.0 x 10-7 M b) For 1M HCl: pH = 0.0 pOH =14.0 [H+] = 1 M [OH-] = 1.0 x 10-14 M c) For 1M NaOH pH = 14.0 pOH =0.0 [H+] = 1.0 x 10-14 M [OH-] = 1 M

21. The pH of .20M NaX is 12.50, calculate the Kb.

Kb = 5.9 x 10-3

22. The pH of .2M HX is 4.5, calculate the Ka.

Ka = 5 x 10-9

23. 100mL of .200M NaOH is mixed with 100ml of .180M HCl. Calculate the pH of the resulting solution.

pH = 12.00 24. How many grams of NaHCO3 are required to make 100mL of .200M solution?

1.68 g

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

0.0375 L

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

1.00 M

27. Calculate the concentration of a solution of NaCl made by dissolving 50.0g in 250mL of water.

3.42 M

28. SO3(g) + H2O(g) ⇄ H2SO4(l) Equilibrium concentrations are found to be: [SO3] = 0.400M [ H2O] = 0.480M [H2SO4] = 0.600M Calculate the value of the equilibrium constant.

29. 2SO2(g) + O2(g) ⇄ 2SO3(g) 4.00 moles of SO2 and 5.00 moles O2 are placed in a 2.00 L container at 200ºC and allowed to reach equilibrium. If the equilibrium concentration of O2 is 2.00M, calculate the Keq.

Keq = 0.500 Ws # 15 Buffers 1. Definition (buffer) A solution that is made by mixing a weak acid or base with a salt containing the conjugate which maintains a relatively constant pH.

Acid Conjugate Base Salt

HCN CN- NaCN

- H2CO3 HCO3 KHCO3

+ NH4 NH3 NH4Cl

HF F- NaF

- CH3COOH CH3COO NaCH3COO

- H2C2O4 HC2O4 Na HC2O4

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

HCN ⇄ H+ + CN-

+ - H2CO3 ⇄ H + HCO3

+ - NH3 + H2O ⇄ NH4 + OH

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 answers are 9.00, 8.98, 9.01, 2.00, and 13.00

Final pH a) 2 drops of 0.10M HCl are added 8.98 b) 1 drop of 0.10M NaOH is added 9.01 c) 10 mL of 1.0 M HCl are added 2.00

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: H2CO3 ⇄ H + HCO3

+ - Shift left [H ] = increases [H2CO3] = increases [HCO3 ] = decreases

Indicators 1. Definition (Indicator) A weak acid whose conjugate base is a different color. 2. Equilibrium equation HInd ⇄ H+ + Ind-

3. Colors for methyl orange HInd red Ind- yellow

4. Compare the relative sizes of [HInd] and [Ind-] at the following pH’s. Color Relationship pH = 2.0 red [Hind] > [Ind-] pH = 3.7 orange [Hind] = [Ind-] pH = 5.0 yellow [Hind] < [Ind-]

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

[H+] increases

[HInd] increases

[Ind-] decreases

Color Change yellow to red

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

[H+] decreases

[HInd] decreases

[Ind-] increases

Color Change red to yellow

7. State two equations that are true at the transition point of an indicator. [Hind] = [Ind-] Ka = [H+]

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

9. What is the Ka for methyl orange. 2 x 10-4

10. A solution is pink in phenolphthalein and colorless in thymolphthalein. What is the pH of the solution? pH = 10 11. A solution is blue in bromothymol blue, red in phenol red, and yellow in thymol blue. What is the pH of the solution? pH = 8 Ws # 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 ------> 7 bromothymol blue

2. HF + RbOH ------> 9 phenolphthalein

3. HI + Ba(OH)2 ------> 7 bromothymol blue

4. HCN + KOH ------> 9 phenolphthalein

5. HClO4 + NH3 ------> 5 bromocresol green

6. CH3COOH + LiOH ------> 9 phenolphthalein

7. Calculate the Ka of bromothymol blue. Ka = 2 x 10-7

8. An indicator has a ka = 1 x 10-10, determine the indicator. Thymolphthalein

9. Calculate the Ka of methyl orange. Ka = 2 x 10-4

10. An indicator has a ka = 6.3 x 10-13, determine the indicator. Indigo Carmine

11. Explain the difference between an equivalence point and a transition point. The equivalence point refers to endpoint of a titration (moles acid = moles base) and a transition point refers to when an indicator changes color.

Draw14 a titration curve for each of the following. 12. Adding 100 ml 1.0 M NaOH to 50 mL 1.0 M HCl 13. Adding 100 ml 1.0 M NaOH to 50 mL 1.0 M HCN

0 50 100 0 50 100

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

1 4 pH pH

0 50 100 0 50 100

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