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Acids, Bases and Salts Unit Plan Period/Topic Worksheets Quiz

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

2. Arrhenius, Bronsted Bases, Kb and Strength 2

3. Acid & Base Reactions. Amphiprotic. Acid Chart. 3 2

4. Leveling effect, Anhydrides and Relationships. 4

5. Hydrolysis of Salts. Quiz. 5 3

6. Acid, Base & Salt Reactions. Hydrolysis. 6

7. Yamada’s Indicator Lab. Hydrolysis. 7 4

8. Ionization of Water, [H+] & [OH-], pH scale. 8

9. pH Calculations 9

10. pH Calculations for Weak Acids. 10 5

11. Ka from pH for Weak Acids. 11

12. Indicators Lab.

13. Kbs from Kas for Weak Bases. 12 6

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

15. Amphiprotic Ions- Kas and Kbs. 14 7

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

17. Titration Lab

18. Buffers & Indicators 15 8

19. Titration Curves. 16 9

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. List five properties of acids that are in your textbook. 2. List five properties of bases that are in your textbook. 3. Make some notes on the commercial acids: HCl and H2SO4 .

HCl

H2SO4

4. Make some notes on the commercial base NaOH. 5. Describe the difference between a concentrated and dilute acid (hint: concentration refers

to the molarity). Describe their relative conductivities. 6. Describe the difference between a strong and weak acid. Use two examples and write

equations to support your answer. Describe their relative conductivities.

7. Describe a situation where a strong acid would have the same conductivity as a weak acid

(hint: think about concentration)

Complete each acid reaction. Label each reactant and product as an acid or base. The first on is done for you. 1. HCN + H2O ⇄ H3O+ + CN-

Acid Base Acid Base 2. H3C6O7 + H2O ⇄

3. H3PO4 + H2O ⇄ 4. HF + H2O ⇄ 5. H2CO3 + H2O ⇄ 6. NH4

+ + H2O ⇄ 7. CH3COOH + H2O ⇄

8. HCl + H2O ⇄ 9. HNO3 + H2O ⇄ Write the equilibrium expression (Ka) for the first seven above reactions. The first one is done for you. 10. Ka = [H3O + ][CN - ] 14. Ka =

[HCN] 11. Ka = 15. Ka =

12. Ka = 16. Ka = 13. Ka = 17. Which acids are strong? 18. What does the term strong acid mean? 19. Why is it impossible to write an equilibrium expression for a strong acid? 20. Which acids are weak? 23. What does the term weak acid mean? 24. Explain the difference between a strong and weak acid in terms of electrical conductivity.

Acid Conjugate Base Base Conjugate Acid 14. HNO2 15. HCOO- 16. HSO3

- 17. IO3-

18. H2O2 19. NH3 20. HS- 21. CH3COO- 22. H2O 23. H2O Define: 22. Bronsted acid- 23. Bronsted base- 24. Arrhenius acid- 25. Arrhenius base- 26. List the six strong acids. 27. Rank the acids in order of decreasing strength. HCl H2S H3PO4 H2CO3 HF

HSO4

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

Making a Universal Indicator Lab Activity

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

Yamada’s Universal Indicator 5 drops thymol blue

6 drops methyl orange5 drops phenolphthalein

10 drops bromothymol blue20 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.

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 plateWash and return your eyedropper.Wash and return your beaker.Wash your hands. Compound Conductivity pH Classification HCl CH3COOH NaOH NH3 NaHCO3 H2CO3

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

NaCl Worksheet # 2 Conjugate Acid-Base Pairs Complete each reaction. Label each reactant and product as an acid or base.

1. HCN + H2O ⇄ H3O+ + CN-

Acid Base Acid Base

2. HCl + H2O ⇄

3. HF + H2O ⇄

4. F- + H2O ⇄

5. HSO4- + H2O ⇄

(acid)

6. NH4+ + H2O ⇄

7. HPO42- + H2O ⇄

(base)

Acid Conjugate Base Base Conjugate Acid 8. HCO3

- CO32- 9. CH3COO- CH3COOH

10. HPO42- 11. IO3

- 12. H2O 13. NH2

- 14. HS- 15. C2H5SO7

3- 16. Circle the strong bases.

Fe(OH)3 NaOH CsOH KOHZn(OH)2 Sr(OH)2 Ba(OH)2 Ca(OH)2

17. Rank the following acids from strongest to weakest.

H2S CH3COOH H2PO4- HI HCl HF

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

H2O F- NH3 SO32- 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 # 3 Using Acid Strength Tables

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

1. HSO4

- + HPO42- ⇄

2. HCN + H2O ⇄

3. HCO3- + H2S ⇄

4. HPO4

2- + NH4+ ⇄

5. NH3 + H2O ⇄ 6. H2PO4

1- + NH3 ⇄

7. HCO3- + HF ⇄

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

acid or base.

HSO4- + HCO3

- ⇄

H2PO4- + HC03

- ⇄

HS03- + HPO4

2- ⇄

NH3 + HC2O4-⇄

9. Explain why HF(aq) is a better conductor than HCN(aq).

10. Which is a stronger acid in water, HCl or HI? Explain! 11. State the important ion produced by an acid and a base. 12. Which is the stronger base? Which produces the least OH-? F- or CO3

-2

13. Define a Bronsted/Lowry acid and base.

14. Define an Arrhenius acid and base. 15. Complete each reaction and write the equilibrium expression.

HF + H2O ⇄ Ka= F- + H2O ⇄ Kb=

16. H2SO4 + NaOH → 17. Define conjugate pairs. 18. Give conjugate acids for: HS-, NH3, HPO4

-2, OH-, H2O, NH3, CO3-2

19. Give conjugate bases for: NH4

+, HF, H2PO4-, H3O+, OH-, HCO3

-, H2O Worksheet # 4 Acid and Basic Anhydrides 1. What is the strongest acid that can exist in water? Write an equation to show how a

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

stronger base would be reduced in strength by the leveling effect of water. 3. List three strong acids and three strong bases. 4. Rank the acids in decreasing strength:

HClO4 Ka is very large HClO3 Ka=1.2x10-2

HClO2 Ka=8.0x10-5 HClO Ka=4.4x10-8

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 # 5 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 parentacid-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 # 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 dissociation equation and hydrolysis equation (only for acidic and basic salts). For acids and bases write an equation to show how each reacts with water. 1. NH3

2. NaCl

3. HCl

4. NaCN

5. NaOH

6. FeCl3

7. HF

8. LiHCO3

9. Fe(NO3)3

10. MgCO3

11. H2S

12. HF

13. CaI2

14. Mg(OH)2

15. Ba(OH)2

16. Describe why Tums (CaCO3) neutralizes stomach acid. 17. Describe why Mg(OH)2 is used in Milk of Magnesia as an antacid instead of NaOH.

Worksheet # 7 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


pH Classification


3. Formula of compound pH Classification



pH Classification



pH Classification



pH Classification


7. Formula of compound pH Classification



pH Classification



pH Classification



pH Classification



pH Classification



pH Classification



pH Classification



pH Classification



pH Classification



pH Classification



pH Classification


Worksheet # 8 pH and pOH Calculations Complete the chart:

[H+] [OH-] pH pOH Acid/base/neutral1. 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 # 9 pH Calculations for Weak Acids 1. Calculate the [H+], [OH-], pH, and pOH for 0.20 M HCN.

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

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

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

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

[H+] = [OH-] = pH = pOH = 5. Calculate the pH of a saturated solution of Mg(OH)2. 6. Calculate the pH of a 0.200 M weak diprotic acid with a Ka = 1.8 x 10-6. 7. 350.0 mL of 0.20M Sr(OH)2 is diluted by adding 450.0 mL of water, calculate the pH of

the solution.

Worksheet # 10 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 # 11 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

-)

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

Ka(HNO2) 4.6 x 10-4

2. KCH3COO (the basic ion is CH3COO-) 3. NaHCO3

4. NH3

5. NaCN

6. Li2HPO4

7. KH2PO4

8. K2CO3

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

[H+] = [OH-] = pH = pOH = 10. The pH of 0.20 M H2CO3 is 3.53. Calculate the Ka for H2CO3. Compare your calculated

value with that in the table.

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

[H+] = [OH-] = pH = pOH = 12. The pH of 0.10 M CH3COOH is 2.87. Calculate the Ka for CH3COOH. Compare your

calculated value with that in the table. 13. 200.0 mL of 0.120 M H2SO4 reacts with 400.0 mL of 0.140 M NaOH. Calculate the pH

of the resulting solution.

Worksheet # 12 Acid and Base pH Calculations

For each weak bases calculate the [OH-], [H+], pOH and pH. Remember that you need to calculate Kb first. 1. 0.20 M CN-

2. 0.010 M NaHS (the basic ion is HS-) 3. 0.067 M KCH3COO 4. 0.40 M KHCO3

5. 0.60 M NH3

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

the acid using your acid chart.

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

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

[H+] = [OH-] = pH = pOH = 9. The pH of 1.65 M H3BO3 is 4.46. Calculate the Ka for H3BO3. Compare your calculated

value with that in the table. 10. The pH of 0.65 M NaX is 12.46. Calculate the Kb for NaX.

11. Consider the following reaction: 2HCl + Ba(OH)2 → BaCl2 + 2H2OWhen 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 # 13 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-

HPO42-

4. What is the strongest base in water? What is the strongest acid in water? Write equations

to explain your answer. 5. Match each equation:

Acid/base complete HCl + NaOH → NaCl + HOH Acid/base net ionic F- + HOH ⇄ HF + OH-

Solubility product H+ + OH- → HOH Hydrolysis AgCl(s) ⇄ Ag+ + Cl-

Acid/Base formula H20 ⇄ H+ + OH-

Ionization of water H+ + Cl- + Na+ + OH- → Na++ Cl- + H2O 6. HCl and HF. Describe each acid as: a) strong or weak b) high or low ionization c) large or small Ka d) good or poor conductor e) strong or weak electrolyte 7. Out of 0.2 M HCl and 1.0 M HF, which is the most concentrated?

Which is the strongest acid? 8. Label the scale as strong/weak acid and strong/weak base.

|________________________|_________________________|__pH 0 7 14

9. Which ions are amphiprotic? HPO4

2- HCl F- HS- H2S H2O

10. Write the net ionic equation between any strong acid and strong base. 11. Write the ionization equation for water. 12. Write the Kw expression. 13. H2SO3 + HS- ⇄ H2S + HSO3

-

a) Are the reactants or products favoured?

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

Determine the pH Write equations for each first! 14. .20M HCl pH=? 15. 0.20M Ba(OH)2 pH=? 16. 0.20M H2CO3 pH=? 17. 0.40M KHCO3 pH=? 18. The pH increases by 2 units. How does [H+] change? 19. The pH decreases by 1 unit. How does [H+] change?

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

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

21. The pH of 0.20 M NaX is 12.50; calculate the Kb. 22. The pH of 0.2 M HX is 4.5; calculate the Ka. 23. 100.0 mL of 0.200 M NaOH is mixed with 100.0 mL of 0.180 M HCl. Calculate the pH

of the resulting solution. 24. How many grams of NaHCO3 are required to make 100.0 mL of 0.200M solution? 25. What volume of 0.200 M NaOH is required to neutralize 25.0 mL of 0.150 M H2SO4?

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

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

of water. 28. 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 # 14 Buffers and indicators Buffers 1. Definition 2. Acid Conjugate Base Salt HCN KHCO3

NH3 HF 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 13.00Final pH

a) 2 drops of 0.10 M HCl are added

b) 1 drop of 0.10 M NaOH is added c) 10 mL of 0.10 M HCl are added 6. Write an equation for the carbonic acid, sodium hydrogencarbonate buffer system. A few

drops of HCl are added. Describe the shift and each concentration change.

Equation:

Shift [H+] = [H2CO3] = [HCO3-] =

Indicators 1. Definition 2. Equilibrium equation 3. Colors for methyl orange HInd Ind- 4. Compare the relative sizes of [HInd] and [Ind-] at the following pH’s for methyl orange.

Color Relationship

pH = 2.0

pH = 3.7

pH = 5.0 5. HCl is added to methyl orange, describe if each increases or decreases.

[H+]

[HInd]

[Ind-]

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

[H+]

[HInd]

[Ind-]

Color Change

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

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

9. What is the Ka for methyl orange? 10. A solution is pink in phenolphthalein and colorless in thymolphthalein. What is the pH of

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

What is the pH of the solution? Worksheet # 15 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 pH

Volume of base added Volume of base added

14. Adding 100 ml 0.10 M HCl 15. Adding 100 ml .10 M HCl

to 50 mL 0.10M NH3 to 50 mL 0.10 M NaOH pH pH

Volume of HCl added Volume of HCl added

Acids Unit Midterm Practice Test 1. Consider the following:

I H2CO3 + F- ⇄ HCO3- + HF

II HCO3- + HC2O4

- ⇄ H2CO3 + C2O42-

III HCO3- + H2C6H6O7

- ⇄ H2CO3 + HC6H5O72-

The HCO3- is a base in

A. I onlyB. I and II onlyC. II and III onlyD. 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 favouredB. red and the reactants are favouredC. yellow and the products are favouredD. 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 mLB. 25.0 mL

C. 50.0 mLD. 100.0 mL

4. The pOH of 0.050 M HCl is

A. 0.050B. 1.30C. 12.70D. 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 mLB. 20.0 mLC. 40.0 mLD. 80.0 mL

6. Consider the followingI H3PO4 II H2PO4

- III HPO42- IV PO4

3-

Which of the above solutions are amphiprotic?A. I and II onlyB. II and III onlyC. I, II, and III onlyD. 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 H3BO3

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 shiftsA. left and the Kw increasesB. left and the Kw decreasesC. right and the Kw increasesD. right and the Kw decreases

9. Normal rainwater has a pH of approximately 6 as a result of dissolved A. oxygenB. carbon dioxideC. sulphur dioxideD. nitrogen dioxide

10. A 1.0 M solution of sodium dihydrogen phosphate is

A. acidic and the pH < 7.00B. acidic and the pH > 7.00C. basic and the pH < 7.00D. 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 isA. 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 thatA. accepts protonsB. neutralizes acidsC. donated electronsD. produces hydroxides ions in solution

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

A. 1.0 M HCNB. 1.0 M H2SO4

C. 1.0 M H3PO4

D. 1.0 M CH3COOH 14. Consider the following equilibrium: HC6H5O7

2- + HIO3 ⇄ H2C6H5O7- + IO3

-

The order of Bronsted-Lowry acids and bases isA. acid, base, acid, baseB. acid, base, base, acidC. base, acid, acid, baseD. 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 equilibriumA. shifts left and [H+] decreasesB. shifts left and [H+] increasesC. shifts right and [H+] decreasesD. 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 MB. 1.25 x 10-12 MC. 8.00 x 10-3 MD. 8.00 x 10-2 M

18. Consider the following:

I ammonium nitrate II calcium nitrate III iron III nitrateWhen dissolved in water, which of these salts would form a neutral solution? A. II onlyB. III onlyC. I and III onlyD. I, II, and III

19. Consider the following: SO4

2- + 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 isA. H2O + Na+ ⇄ NaOH + H+

B. H2O + 2Na+ ⇄ Na2O + 2H+

C. H2O + CO32- ⇄ H2CO3 + O2-

D. H2O + CO32- ⇄ 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, theA. [H3O+] has increased and the [OH-] has decreasedB. [H3O+] has increased and the [OH-] has increasedC. [H3O+] has decreased and the [OH-] has increasedD. [H3O+] has decreased and the [OH-] has decreased

22. A basic solution

A. tastes sourB. feels slipperyC. does not conduct electricityD. reacts with metals to release oxygen gas

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

A. H2SO4 + KOH → KSO4 + H2OB. H2SO4 + KOH → K2SO4 + H2OC. H2SO4 + 2KOH → K2SO4 + H2OD. H2SO4 + 2KOH → K2SO4 + 2H2O

24. An Arrhenius base is defined as a substance which

A. donates protonsB. donates electronsC. produces H+ in solutionD. produces OH- in solution

25. HS- + H3PO4 ⇄ H2S + H2PO4

- The order of Bronsted-Lowry acids and bases isA. acid, base, acid, base.B. acid, base, base, acidC. base, acid, acid, baseD. base, acid, base, acid

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

A. CH3COO- + H2O ⇄ CH3COOH + OH-

B. CH3COO- + H2O ⇄ CH3COO2- + 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 waterA. stays neutral and the [H3O+] increasesB. stays neutral and the [H3O+] decreasesC. becomes basic and [H3O+] decreasesD. 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. C6H5OHC. C6H4O2-

D. C6H5OH+

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

A. 1.0 M HClB. 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 HClB. a higher pOH than a solution of 1.0 M HClC. a higher [OH-] than a solution of 1.0 M HClD. a higher [H3O+] than a solution of 1.0 M HCl

32. Which of the following is the weakest acid

A. HIO3

B. HCNC. HNO3

D. C6H5COOH 33. Considering the following data

H3AsO4 Ka = 5.0 x 10-5

H2AsO4- Ka = 8.0 x 10-8

HAsO42- Ka = 6.0 x 10-10

The Kb value for H2AsO4- is

A. 2.0 x 10-10

B. 8.0 x 10-8

C. 1.2 x 10-7

D. 1.7 x 10-5

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

A. 3.5 x 10-20 MB. 2.9 x 10-9 MC. 1.0 x 10-7 MD. 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 MB. 6.0 x 10-5 MC. 6.3 x 10-1 MD. 1.7 x 104 M

36. An aqueous solution of NH4CN is

A. basic because Ka < KbB. basic because Ka > KbC. acidic because Ka < KbD. acidic because Ka > Kb

37. The net ionic equation for the predominant hydrolysis reaction of KHSO4 isA. HSO4

- + H2O ⇄ SO42- + H3O+

B. HSO4- + H2O ⇄ H2SO4 + OH-

C. KHSO4 + H2O ⇄ K+ + SO42- + 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 MB. 1.79 x 10-14 MC. 5.58 x 10-1 MD. 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 - ]

[HPO42-] [H2PO4

-]

C. [K + ] [KHPO 4- ] D. [K + ] 2 [HPO 4

2- ] [K2HPO4] [K2HPO4]

41. The solution with the highest pH is

A. 1.0 M NaClB. 1.0 M NaCNC. 1.0 M NaIO3

D. 1.0 M Na2SO4

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

A. 2.30B. 3.30C. 10.70D. 11.7

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:

Barbituric acid HC4H3N2O3 Ka = 9.8 x 10-5

Sodium propanoate NaC3H5O2 Kb = 7.5 x 10-10

Propanoic acid HC3H5O2 Ka = ?

Which is the stronger acid, propanoic acid or babituric acid? Explain using calculations. 8. A solution of 0.0100 M lactic acid, HC3H5O3, has a pH of 2.95. Calculate the Ka value. 9. a) Write equations showing the amphiprotic nature of water as it reacts with

HCO3-.

b) Calculate the Kb for HCO3-.

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

Acids Quiz #1 Properties of Acids and BasesStrength

1. Drano®, a commercial product used to clean drains, contains small bits of aluminum metal andA. ammoniaB. acetic acidC. hydrochloric acidD. sodium hydroxide

2. A net ionic equation for the reaction between CH3COOH and KOH isA. CH3COOH(aq) + K+

(aq) ⇄ CH3COOK(aq)

B. CH3COOH(aq) + OH-(aq) ⇄ H2O(l) + CH3COO-

(aq)

C. CH3COOH(aq) + KOH(aq) ⇄ H2O(l) + CH3COOK(aq)

D. CH3COOH(aq) + K+(aq) + OH-

(aq)⇄ H2O(l) + KCH3COO(aq)

3. Which equation represents a neutralization reaction?A. Pb2+

(aq) + 2Cl-(aq) → PbCl2(s)

B. HCl(aq) + NH3(aq) → NH4Cl(aq)

C. BaI2(aq) + MgSO4(aq) → BaSO4(s) + MgI2(aq)

D. MnO4-(aq) + 5Fe2+

(aq) +8H+(aq) → Mn2+

(aq) + 5Fe3+(aq) + 4H2O(l)

4. An Arrhenius acid is a substance thatA. accepts a protonB. donates a protonC. produces H+ in solutionD. 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 NH4OH3 25 mL 0.1 M KOH4 50 mL 0.2 M NaOH

The beaker that requires the smallest volume of 1.0 M HCl for complete neutralization is:A. Beaker 1B. Beaker 2C. Beaker 3D. Beaker 4

6. The net ionic equation for the titration of HClO4(aq) with LiOH(aq) isA. 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 isA Ka = [H+][HSO3

-]B. Ka = [H + ][HSO 3

- ] [H2SO3]

C. Ka = [2H + ][SO 32- ]

[H2SO3]

D. Ka = [H + ][SO 32- ]

[H2SO3]8. To distinguish between a strong acid and a strong base, an experimenter could use

A. odorB. magnesiumC. a conductivity testD. 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. 2B. 3C. 4D. 5

10. There are two beakers on a laboratory desk. One beaker contains 1.0 M HCl andthe other contains tap water. To distinguish the acid solution from the water, onewould useA. a piece of copper.B. a piece of magnesiumC. phenolphthalein indicatorD. a piece or red litmus paper

11. Caustic soda, NaOH, is found inA. fertilizersB. beveragesC. toothpasteD. oven cleaners

12. Which of the following is the strongest acid?A. Acetic acidB. Oxalic acidC. Benzoic acidD. Carbonic acid

13. The acid used in the lead-acid storage battery isA. HClB. HNO3

C. H2SO4

D. CH3COOH

14. A definition for a Brønsted-Lowry acid is:A. the donation of H+

B. the donation of OH-

C. the acceptance of H+

D. the acceptance of OH-

15. What species will form when H+ ions are in the presence of H2O molecules?A. HO+

B. HO2+

C. H3O+

D. H2O2

16. What is the conjugate acid of the base HAsO42-?

A. AsO43-

B. H2AsO42-

C. H2AsO4-

D. H3AsO4

17. What species is the hydronium ion?A. HO+

B. HO2+

C. H2O2

D. H3O+

18. Which of the following is a characteristic of acidic solutions?A. They accept H+ ions.B. They have a pH more than 7.C. They react with Mg to produce H2 gas.D. They turn limus blue.

19. Which of the following is a common base found in drain cleaners?A. bleachB. vinegarC. milk of magnesiaD. sodium hydroxide

20. Which of the following is a characteristic of basic solutions?A. They donate H+ ions.B. They have a pH less than 7.C. They feel slippery.D. They turn limus red.

21. Which of the following is the weakest acid?A. Acetic acidB. Oxalic acidC. Benzoic acidD. Carbonic acid

22. What is the main difference between a strong acid and a weak acid?A. their degree of ionizationB. their reactivity with platinumC. their concentration in solutionD. their effect on an indicator

23. Which of the following best describes an acidic solution?Litmus Colour Reaction with Zn

A. red reactionB. red no reactionC. blue no reactionD. blue reaction

24. Identify the common acid found in the stomach.A. benzoic acidB. sulphuric acidC. perchloric acidD. hydrochloric acid

25. Which of the following best describes an basic solution?Litmus Colour pH

A. red less than 7B. red more than 7C. blue less than 7

D. blue more than 7

26. Which of the following is a general characteristic of Arrhenius acids?A. They produce H+ in solution.B. They accept an H+ from water.C. They turn bromothymol blue a blue colour.D. They react with H3O+ ions to produce H2.

27. Identify a conjugate pair from the equilibrium provided: PO4

3- + HCO3- ⇌ HPO4

2- + CO32-

A. CO32- and HPO4

2-

B. PO43- and HCO3

-

C. PO43- and HPO4

2-

D. HCO3- and HPO4

2-

28. Which of the following best describes a weak acid?

A. pH = 10B. It may be very soluble, but only partly ionized.C. It must be very soluble and completely ionized.D. It must be of low solubility and completely ionized.

29. Which of the following is a common property of acid solutions?A. They have a pH > 7.B. They turn red litmus blue.C. They have a slippery feeling.D. They turn pink phenolphthalein colourless.

30. What is a general characteristic of bases?A. They all donate H+.B. They all accept OH- .C. They will neutralize acidsD. They will react with acids to produce H2 gas.

Acids Quiz #2 Conjugates, Reactions, Amphiprotic, Arrhenius, Bronsted Bases, KB, & Strength

1. A test that could be safely used to distinguish a strong base from a weak base isA. tasteB. touchC. litmus paperD. electrical conductivity

2. Identify the two substances that act as Bronsted-Lowry bases in the equationHS- + SO4

2- ⇄ S2- + HSO4-

A. HS- and S2-

B. SO4 2-and S2-

C. HS- and HSO4-

D. SO42- and HSO4

-

3. The conjugate acid of H2C6H5O-7 is

A. C6H5O73-

B. HC6H5O72-

C. H2C6H5O7

D. H3C6H5O7

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

- (3) HPO42-

A. 2 onlyB. 3 onlyC. 1 and 2D. 2 and 3

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

(aq) + OH-(aq) ® H2O(l)

B. Ca2+(aq) + 2Br-

(aq) ® CaBr2(s)

C. 2HBr(aq) + Ca(OH)2(aq) ® 2H2O(l) + CaBr2(s)

D. 2H+(aq) + Ca(OH)2(aq) ® 2H2O(l) + Ca2+

(aq)

6. If reactants are favored in the following equilibrium, the stronger base must beHCN + HS - ⇄ H2S + CN -

A. H2SB. HS-

C. CN-

D. HCN

7. The hydronium ion, H3O+ is a water molecule that hasA. lost a protonB. gained a protonC. gained a neutronD. gained an electron

8. The complete ionic equation for the neutralization of acetic acid by sodium hydroxide isA. H+ + OH- ⇄ H2OB. CH3COOH + OH- ⇄ CH3COOH- + H2OC. CH3COOH + NaOH ⇄ NaCH3COOH + H2OD. 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 isA. NH4

+

B. H2OC. NH3

D. H3O+

10. Consider the following equilibrium system:OCl-

(aq) + HC7H5O2(aq) ⇄ HOCl(aq) + C7H5O2-(aq) Keq= 2.1 x 103

At EquilibriumA. products are favored and HOCl is the stronger acidB. reactants are favored and HOCl is the stronger acidC. products are favored and HC7H5O2 is the stronger acidD. reactants are favored and HC7H5O2 is the stronger acid

11. In the equilibrium systemH2BO3

- (aq) + HCO3

-(aq) ⇄ H2CO2(aq) + HBO3

2-(aq)

The two species acting as Bronsted-Lowry acids are

A. HCO3- and H2CO3

B. H2BO3- and H2CO3

C. HCO3- and HBO3

2-

D. H2BO3- and HBO3

2-

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

The stronger acid isA. HS-

B. H2C2O4

C. HC2O4-

D. H2S

13. Which species is not amphiprotic?A. H2OB. H3BO3

C. HPO42-

D. HC6H5O72-

14. Given the equilibrium: H2BO3- + H2PO4

- ⇌ H3BO3 + HPO42-

Which is the strongest acid?A. H2PO4

-

B. H3BO3

C. HPO42-

D. H2BO3-

15. Which equation represents the reaction of a Brønsted-Lowry base with water?A. 2Na + 2H2O ® 2NaOH + H2

B. P2H4 + H2O ⇌ P2H5+ + OH-

C. H2PO4- + H2O ⇌ H3O+ + HPO4

2-

D. H2C2O4 + H2O ⇌ H3O+ + HC2O4-

16. A definition for a Brønsted-Lowry acid is:A. the donation of H+

B. the donation of OH-

C. the acceptance of H+

D. the acceptance of OH-

17. Which of the following would be a typical pH of soda pop?A. 1.0B. 3.2C. 6.8D. 7.2

18. When three 25.0 mL samples of the strong acid H2SO4 were titrated with 0.25 M NaOH the following results were obtained: Volume of NaOH(aq) 47.2 mL 39.9 mL 40.1 mLWhat is the concentration of the H2SO4 sample?A. 0.20 MB. 0.21 MC. 0.40 MD. 0.42 M

19. Which of the following is a common base found in drain cleaners?A. baking sodaB. vinegarC. milk of magnesiaD. sodium hydroxide

20. Identify the common acid found in the stomach.A. benzoic acidB. sulphuric acidC. perchloric acidD. hydrochloric acid

21. Which of the following solutions would typically show the greatest electrical conductivity?A. 1.0 M weak acidB. 0.8 M weak baseC. 0.5 M strong acidD. 0.1 M strong base

22. Which of the following are amphiprotic in aqueous solutions?I. H3BO3 II. H2BO3

- III. HBO32- IV. BO3

3-

A. I onlyB. IV onlyC. I and II onlyD. II and III only

23. Which of the following best describes an acidic solution?Litmus Colour Reaction with Mg

A. red reactionB. red no reactionC. blue no reactionD. blue reaction

24. What is the Ka expression for H3PO4 ?

A. Ka = [H + ] 3 [PO4

3-]B. Ka = [H + ] 3

[H3PO4]

C. Ka = [H + ][H 2PO4-]

[H3PO4]

D. Ka = [H + ][H PO42-]

[H3PO4]

25. What volume of 0.250 M KOH is needed to titrate 0.00230 mol of the acid H2C2O4 ?A. 1.15 mLB. 4.60 mLC. 9.20 mLD. 18.4 mL

26. What is the net ionic equation for the titration of H3PO4(aq) with Sr(OH)2(aq) ? A. H+

(aq) + OH- (aq) ® H2O(l)

B. 6H+(aq) + 6OH-

(aq) ® 6H2O(l)

C. 2H3PO4(aq) + 3Sr2+(aq) + 6OH-

(aq) ® Sr3(PO4)2(s) + 6H2O(l)

D. 6H+(aq)

+ 2PO43-

(aq) + 3Sr2+(aq) + 6OH-

(aq) ® 3Sr2+ + 2PO43-

(aq) + 6H2O(l)

27. Which of the following is a piece of equipment used in acid-base titrations?A. buretteB. test tubeC. litmus paperD. graduated cylinder

28. Water has the greatest tendency to act as an acid with which of the following?A. Cl-

B. NO2-

C. H2PO4-

D. CH3COO-

29. Which net ionic equation best describes the reaction between NaOH and H2S?A. H+

(aq) + OH-(aq)

® H2O(l)

B. H2S(aq) + 2OH- ® 2H2O(l) + S2-(aq)

C. H2S(aq) + 2NaOH ® 2H2O(l) + Na2S(aq)

D. 2H+(aq) + S2-

(aq) + 2NaOH(aq) ® 2H2O(l) + 2Na+(aq) + S2-

(aq)

30. Which of the following is a general characteristic of Arrhenius acids?A. They produce H+ in solution.B. They accept an H+ from water.C. They donate a proton.D. They accept a proton.

31. Which of the following is a general characteristic of Bronstead acids?A. They produce H+ in solution.B. They accept an H+ from water.C. They donate a proton.D. They accept a proton.

32. Which of the following is a general characteristic of Arrhenius bases?A. They produce OH- in solution.B. They produce H3O+ in water.C. They donate a proton.D. They accept a proton.

33. Which of the following is a general characteristic of Bronstead base?A. They produce H+ in solution.B. They accept an H+ from water.C. They donate a proton.D. They accept a proton

Acids Quiz #3 Leveling effect, Anhydrides, Hydrolysis, Relationships

1. Which of the following oxides will form the most acidic solution?A. SO2

B. MgOC. Na2OD. Al2O3

2. Which one of the following salts will produce an acidic solution?A. KBrB. LiCNC. NH4ClD. NaCH3COO

3. The balanced equation for the reaction between sodium oxide and water isA. Na2O + H2O → 2NaOHB. 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. methaneB. carbon dioxideC. sulphur dioxideD. nitrogen dioxide

5. Which of the following oxides would hydrolyze to produce hydroxide ions?A. NOB. SO2

C. Cl2OD. Na2O

6. The approximate pH of “normal” rainwater isA. 0B. 6C. 7D. 8

7. Which of the following oxides would hydrolyze to produce hydronium ions?A. CaOB. SO2

C. MgOD. 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 solutionHSO3

- + HF ⇄ H2SO3 + F- The order of Bronsted-Lowry acids and bases in this equation is

A. acid + base ⇄ acid + baseB. acid + base ⇄ base + acid C. base + acid ⇄ base + acidD. base + acid ⇄ acid + base

10. The conjugate acid of OH- isA. H+

B. O2-

C. H2OD. H3O+

11. Which of the following 0.10 M solutions will have the greatest electricalconductivity?A. HFB. NH3

C. NaOHD. C6H5COOH

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

HSeO3- + H2O ⇄ H2SeO3 + OH- Kb

HSeO3- + H2O ⇄ SeO3

2-+ H3O+ Ka

An aqueous solution of HSeO3- is found to be acidic. HSeO3

- behaves mainly as aA. proton donor, and Kb is less than Ka

B. proton donor, and Kb is greater than Ka

C. proton acceptor, and Kb is less than Ka

D. proton acceptor, and Kb is greater than Ka

13. The Kb expression for HPO42- is

A. [PO43- ][H 3O + ] B. [HPO4

2- ][OH - ] [HPO4

2-] [H2PO4-]

C. [H2PO4- ][OH - ] D. [HPO4

2- ][ H 3O + ] [HPO4

2-] [PO43- ]

14. Given the equilibrium: H2BO3- + H2PO4

- ⇌ H3BO3 + HPO42-

Which is the strongest acid?A. HPO4

2-

B. H3BO3

C. H2PO4-

D. H2BO3-

15. Which species is not amphiprotic?A. H2OB. H3BO3

C. H2PO4-

D. H2C6H5O7-

16. What is produced when CH3NH2 acts as a base in water?A. CH3NH-

B. CH3NH3+

C. CH3NH2+

D. CH2NH2-

17. What species will form when H+ ions are in the presence of H2O molecules?A. HO+

B. H2O+

C. H3O+

D. H2O2+

18. What is the conjugate acid of the base HAsO42-?

A. AsO43-

B. H2AsO4-

C. H4AsO4+

D. H3AsO4

19. Which solution will have the greatest electrical conductivity?A. 0.50 M HClB. 0.10 M RbOHC. 0.50 M K3PO4

D. 2.0 M C6H12O6

20. The following equilibrium favours the formation of products:NH2OH + CH3NH2 ⇌ CH3NH- + NH3OH+

Which species is the strongest acid?A. NH3OH+

B. NH2OHC. CH3NH2

D. CH3NH-

21. Which of the following solutions would have the greatest [OH-]?A. 0.1 M HCO3

-

B. 0.1 M HPO42-

C. 0.1 M H2PO4-

D. 0.1 M CO32-

22. Which of the following amphiprotic ions will act predominantly as a base in solution?A. HSO3

-

B. HSO4-

C. H2PO4-

D. HPO42-

23. Which is the strongest base?A. Cl-

B. NO2-

C. HPO42-

D. CH3COO-

24. Water has the greatest tendency to act as an acid with which of the following?A. Cl-

B. NO2-

C. HPO42-

D. CH3COO-

25. Water has the greatest tendency to act as a base with which of the following?A. HFB. H2CO3

C. H3PO4

D. CH3COOH

26. Which of the following will have the smallest Ka value?A. HFB. H2CO3

C. H3PO4

D. CH3COOH

27. Which of the following will have the smallest Kb value?A. IO3

-

B. NH3

C. CN-

D. HPO42-

28. What volume of 0.500 M NaOH is required to neutralize 25.0 mL 0.250 M HBr ?A. 5.00 mLB. 12.5 mLC. 20.0 mLD. 25.0 mL

29. Which of the following Ka values represents the acid that is strongest?A. Ka = 2.8 x 10-18

B. Ka = 8.2 x 10-18

C. Ka = 4.4 x 10-22

D. Ka = 6.4 x 10-22

30. Which of the following Ka values represents the acid with the strongest conjugate base?A. Ka = 2.8 x 10-18

B. Ka = 8.2 x 10-18

C. Ka = 4.4 x 10-22

D. Ka = 6.4 x 10-22

Acids 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/anA. saltB. acidC. baseD. buffer

3. The net ionic equation for the hydrolysis of the salt, Na2S isA. Na2S ⇄ 2Na+ + S2-

B. S2- + H2O ⇄ OH- + HS+

C. Na2S + 2H2O ⇄ 2NaOH + H2SD. 2Na+ + S2- + 2H2O ⇄ 2Na+ + 2OH- + H2S

4. Which of the following solutions would be acidic?A. sodium acetateB. iron III chlorideC. sodium carbonateD. 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 onlyB. I and II onlyC. II and III onlyD. I, II and III

6. Which of the following, when dissolved in water, forms a basic solution?A. KClB. NaClO4

C. Na2CO3

D. NH4NO3

7. Which of the following oxides forms a basic solution?A. K2OB. CO2

C. SO3

D. NO2

8. Which of the following is amphiprotic in water?A. SO2

B. SO32-

C. HSO3-

D. H2SO3

9. Consider the following equilibrium expressionK= [H2S][OH - ]

[HS-] This expression represents the

A. Kb for H2SB. Ka for H2SC. Kb for HS-

D. Ka for HS-

10. The reaction of a strong acid with a strong base producesA. A salt and a waterB. A base and an acidC. A metallic oxide and waterD. 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 isA. acid, base, base, acidB. acid, base, acid, baseC. base, acid, base, acidD. base, acid, acid, base

12. The pH range of ‘acid rain’ is oftenA. 3 to 6B. 6 to 8C. 7 to 9D. 10 to 12

13. Water will act as a Bronsted-Lowry acid withA. NH3

B. H2SC. HCND. HNO3

14. Which of the following is a conjugate acid-base pair?A. H3PO4 and PO4

3-

B. H2PO4- and PO4

3-

C. H3PO4 and HPO42-

D. H2PO4- and HPO4

2-

15. What is the net ionic equation for the hydrolysis of NH4Cl ?A. NH4Cl ® NH4

+ + Cl-

B. Cl- + H2O ⇌ HCl + OH-

C. NH4+ + H2O ⇌ H3O+ + NH3

D. NH4+ + H2O ⇌ NH5

+ + OH-

16. What is the approximate pH of the salt NH4Cl?A. 0.0B. 5.0C. 7.0D. 9.0

17. Four samples of rain are collected from different geographic regions and the pH is measured for each sample.

Sample pH1 2.82 4.03 6.24 6.8

Which of the above samples would be classified as acid rain?A. 1 onlyB. 1 and 2C. 1, 2 and 3D. 1, 2, 3 and 4

18. Which of the following equations correctly represents the reaction of a metallic oxide with water?A. K2O + H2O ® 2KOHB. SO3 + H2O ® H2SO4

C. Na2O + H2O ® Na2O2 + H2

D. NaOH + HCl ® NaCl + HOH

19. Which of the following equations describes the dissociation of the salt, ammonium chloride, in water?A. NH4Cl(s) ® NH4Cl(aq)

B. NH4+

(aq) + Cl-(aq) ® NH4Cl(s)

C. NH4Cl(s) ® NH4+

(aq) + Cl-(aq)

D. NH4+

(aq) + H2O(l) ® H3O+(aq) + NH3(aq)

20. Which of the following salt solutions is acidic?A. NaBrB. AlCl3

C. Li2CO3

D. NaHCO3

21. What pH would most likely result when CO2 dissolves naturally in rainwater?A. 3.0B. 6.6C. 7.0D. 7.5

22. Which of the following is the net ionic equation that describes the hydrolysis that occurs in a K2CO3 solution?A. K2CO3 ® 2K+ + CO3

2-

B. 2K+ + CO32- ® K2CO3

C. CO32- + H2O ® HCO3

- + OH-

D. K2CO3 + 2H2O ® H2CO3 + 2KOH

23. What is the equilibrium expression for the predominant equilibrium in NaCN(aq)?A. Kb = [Na+][CN-]

[NaCN]B. Kb = [Na+][CN-]

[HCN]C. Kb = [HCN][OH-]

[CN-]D. Kb = [CN-][OH-]

[HCN]

24. Which of the following salts will be basic?A. KClB. NH4ClC. NaBrD. Na2HPO4

25. Which of the following salts will be acidic?A. KClB. NH4ClC. NaBrD. Na2HPO4

26. Which is an acid anhydride?A. Na2OB. NaClC. CO2

D. NH4Cl

27. Which formula has the highest pH when in water?A. Na2OB. NaClC. CO2

D. NH4Cl

28. Which formula has the lowest pH when in water?A. Na2OB. NaClC. FeCl3

D. NaHCO3

29. Which formula has the lowest pH when in water?A. NaCNB. Na2HPO4

C. NH4ClD. NaOH

30. Which of the following properties is true for a solution of NaNO3 ?A. It is neutral.B. It is very basic.C. It is slightly basic.D. It is slightly acidic.

Acids Quiz #5 pH calculations for Strong and Weak Acids

1. The 1.0 M acidic solution with the highest pH isA. H2SB HNO2

C. HNO3

D. H3BO3

2. At 25 oC, the equation representing the ionization of water isA H2O + H2O ⇄ 2H2 + O2

B. H2O + H2O ⇄ H2O2 + H2

C. H2O + H2O ⇄ 4H+ + 2O2- D. H2O + H2O ⇄ H3O+ +OH-

3. The pH of a 0.3 M solution of NH3 is approximatelyA. 14.0B. 11.0C. 6.0D. 3.0

4. The pH of an aqueous solution is 4.32. The [OH-] isA. 6.4 x 10-1 MB. 4.8 x 10-5 MC. 2.1 x 10-10 MD. 1.6 x 10-14 M

5. The pH of an aqueous solution is 10.32. The [OH-] isA. 5.0 x 10-12 MB. 2.0 x 10-11 MC. 4.8 x 10-11 MD. 2.1 x 10-4 M

6. The pH of a 0.025 M HClO4 solution isA. 0.94B. 1.60C. 12.40D. 13.06

7. Consider the following equilibriumH2O(l) + H2O(l) ⇄ H3O+

(aq) + OH-(aq)

The equilibrium constant for this system is referred to asA. Kw B. Ka C. Kb D. Ksp

8. The [H3O-] in a solution of pH 0.60 isA. 4.0 x 10-14 MB. 2.2 x 10-1 MC. 2.5 x 10-1 MD. 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 isA. -1

B. 0C. 1D. 7

10. The approximate pH of a 0.06 M solution of CH3COOH isA. 1B. 3C. 11D. 13

11. The [OH-] is greater than the [H3O+] inA. HCl(aq)

B. NH3(aq)

C. H2O(aq)

D. CH3COOH(aq)

12. The pH of 0.15 M HCl isA. 0.15B. 0.71C. 0.82D. 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 isA. 0.20B. 0.63C. 0.70D. 1.58

15. The [OH-] in 0.050 M HNO3 at 25oC isA. 5.0 x 10-16 MB. 1.0 x 10-14 MC. 2.0 x 10-13 MD. 5.0 x 10-2 M

16. What is the approximate pH of a 1.0 M solution of the salt NH4Cl?A. 0.0B. 5.0C. 7.0D. 10.0

17. What is the approximate pH of a 1.0 M solution of the salt NH3?A. 2.0B. 5.0C. 7.0D. 10.0

18. What is the approximate pH of a 1.0 M solution of the salt NaHCO3?A. 2.0B. 5.0C. 7.0D. 9.0

19. What is the approximate pH of a 1.0 M solution of the salt AlCl3?A. 0.0B. 3.0C. 7.0D. 9.0

20. What is the approximate pH of a 1.0 M solution of the salt NaHSO4?A. 0.0B. 3.0C. 7.0D. 9.0

21. Which of the following is a definition of pH?A. pH = +Log[H3O+]B. pH = -Log[OH-]C. pH = -Log[H3O+]D. pH = pOH + pKw

22. What is the mass of NaOH required to prepare 100.0 mL of NaOH(aq) that has a pH = 13.62 ?A. 0.38 gB. 0.42 gC. 1.67 gD. 0.14 g

23. What is the [KOH] in a KOH solution that has a pH =12.00?A. 0.010 MB. 0.56 MC. 2.0 MD. 2.0 x 10-12 M

24. What is the [H3O+] in 0.70 M HCN ?A. 0.70 MB. 1.9 x 10-5 MC. 1.0 x 10-6 MD. 2.4 x 10-10 M

25. What is the value of pKw for water at 25 0C ?A. 10 x 10-14

B. 10 x 10-7

C. 7.00D. 14.00

26. What pH would most likely result when CO2 dissolves naturally in rainwater?A. 3.5B. 6.5C. 7.2D. 7.8

27. What is the pOH of 0.05 M Ba(OH)2 ?A. 1.0B. 1.3C. 12.7D. 13.0

28. What is the pH of 0.5 M Sr(OH)2 ?A. 0.0B. 0.3C. 12.7D. 14.0

29. Which of the following statements is true for an acidic solution at 25 0C ?A. pH > 7.0B. pOH < 7.0C. [H3O+] < [OH-]D. [H3O+] > [OH-]

30. What is the [OH-] in 0.025 M HCl ?A. 2.5 x 10-16 MB. 4.0 x 10-13 MC. 2.5 x 10-2 MD. 1.6 M

Acids Quiz #6 Ka’s from pH Kb’s from Ka’s

1. The Kb for the dihydrogen phosphate ion isA. 1.4 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 mLB. 10.0 mLC. 20.0 mLD. 40.0 mL

3. Consider the following equilibria:

I HCO3- + H2O ⇄ H2CO3 + OH-

II NH4+ + H2O ⇄ H3O+ + NH3

III HSO3- + H3O ⇄ H2O + H2SO3

Water acts as a Bronsted-Lowry base inA. III onlyB. I and II onlyC. II and III onlyD. I, II, and III

4. Which of the following is represented by a Kb expression?A. Al(OH)3(s) ⇄ Al3+

(aq) + 3OH-(aq)

B. HF(aq) + H2O(l) ⇄ H3O+(aq) + F-

(aq)

C. Cr2O72-

(aq) + 2OH-(aq) ⇄ 2CrO4

2-(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 isA. 1.3B. 1.6C. 12.4D. 12.7

6. If OH- is added to a solution, the [H3O+] willA. Remain constantB. 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 MB. 2.00 MC. 4.00 MD. 8.00 M

8. Consider the following equilibrium for an acid-base indicator:Hlnd ⇄ H+ + Ind- Ka = 1.0 x 10-10

Which of the following statements is correct at pH 7.0?

A. [Ind-] < [HInd]B. [Ind-] = [HInd]C. [Ind-] > [HInd]D. [Ind-] = [H+] = [HInd]

9. Which of the following indicators would be yellow at pH 7 and blue at pH 10?A. thymol blueB. methyl violetC. bromthymol blueD. 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 isA. 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 withA. 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. 0.20B. 0.70C. 13.30D. 13.80

13. The 1.0 M acid solution with the largest [H3O+] isA. HNO2

B. H2SO3

C. H2CO3

D. H3BO3

14. Consider the following indicator equilibrium:HInd(aq) + H2O(l) ⇌ H3O+ + Ind-

(aq)

Colourless blueWhat is the effect of adding HCl to a blue sample of this indicator?

Equilibrium Shift Colour ChangeA. left less blueB. left more blueC. right less blueD. right more blue


(aq)

Colourless blueWhat is the effect of adding NaOH to a light blue sample of this indicator?

Equilibrium Shift Colour ChangeA. left less blueB. left more blueC. right less blueD. right more blue


(aq)

yellow blueWhat is the effect of adding FeCl3 to a green sample of this indicator?

Equilibrium Shift Colour ChangeA. left yellowB. left blueC. right less blueD. right more blue

17. An indicator has a Ka = 4 x 10-6. Which of the following is true for this indicator? pH at Transition Point Indicator

A. 4.0 methyl orangeB. 4.0 bromcresol green C. 5.4 methyl redD. 5.4 bromcresol green

18. An indicator has a Ka = 6.3 x 10-13. Which of the following is true for this indicator? pH at Transition Point Indicator

A. 4.0 methyl orangeB. 4.0 bromcresol green C. 12.2 alizarin yellowD. 12.2 indigo carmine

19. A salt forms in the reaction between HF(aq) and NaOH(aq). What is the net ionic equation for the hydrolysis of this salt?A. NaF(aq) ® Na+

(aq) + F-(aq)

B. HF(aq) + H2O(l) ⇌ H3O+(aq) + F-

(aq)

C. F-(aq) + H2O(aq) ⇌ HF(aq) + OH-

(aq)

D. HF(aq) + NaOH(aq) ® NaF(aq) + H2O(l)

20. What is the Ka for the indicator that is yellow in its basic form and blue in its acid form?A. 6 x 10-13

B. 2 x 10-9

C. 2 x 10-7

D. 3 x 10-5

21. What term is used to describe the point at which a chemical indicator changes colour?A. titration pointB. transition pointC. equivalence pointD. stoichiometric point

22. What is the net ionic equation in the reaction between HF(aq) and NaOH(aq)?A. NaF(aq) ® Na+

(aq) + F-(aq)

B. HF(aq) + OH-(aq) ⇌ F-

(aq) + H2O(l)

C. F-(aq) + H2O(aq) ⇌ HF(aq) + OH-

(aq)

D. HF(aq) + NaOH(aq) ® NaF(aq) + H2O(l)

23. Consider the ionization of water: 2H2O ⇌ H3O+ + OH-

What happens to the pH when KOH is added to water?A. pH increases since [H3O+] increases.B. pH increases since [H3O+] decreases.C. pH decreases since [H3O+] increases.D. pH decreases since [H3O+] decreases.

24. Which of the following salt solutions is acidic?A. NaBrB. FeCl3

C. LiCND. NaHCO3

25. Which of the following salt solutions is neutral?A. NaBrB. FeCl3

C. LiCND. NaHCO3

26. What happens to the [ion] in water when a small amount of HBr(aq) is added?A. [H3O+] = [OH-] = 1.0 x 10-7 MB. [H3O+] and [OH-] both increaseC. [H3O+] increases and [OH-] decreasesD. [H3O+] increases and [OH-] is unchanged

27. What is the Kb value for H2PO4- ?

A. 1.3 x 10-12

B. 6.2 x 10-8

C. 1.6 x 10-7

D. 2.5 x 103

28. What is one of the Ka values for thymol blue?A. 2 x 10-9

B. 2 x 10-7

C. 1 x 10-7

D. 5 x 102

29. Consider the indicator equilibrium: HInd(aq) + H2O(l) ⇌ H3O+ + Ind-

(aq)

yellow blueWhich of the following is true about the transition point of this indicator?A. pH = 7.0B. [HInd] = [Ind-]C. [HInd] > [Ind-]D. moles of H3O+ moles of Ind-

30. What is the [OH-] in 0.25 M HCl ?A. 2.5 x 10-15 MB. 4.0 x 10-14 MC. 2.5 x 10-1 MD. 2.5 x 10-13 M

Acids Quiz #7 pH Relationships, Amphiprotic Calculations

1. In water, the hydrogen phosphate ion, HPO42-, 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 acidsAcid pH

X 4.0Y 2.0Which of the following statements can be concluded from the above data?A. Acid X is stronger than acid YB. Acid X and acid Y are both weakC. Acid X is diprotic while acid Y is monoproticD. Acid X is 100 times more concentrated than acid Y

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

which isA. 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 Concentration Ka pH

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

Based on this data, the strongest acid isA. HAB. HBC. HCD. HD

5. Which of the following 0.10 M solutions is the most acidic?A. AlCl3

B. FeCl3

C. CrCl3

D. NH4Cl

6. Which of the following acid-base indicators has a transition point between pH 7 and pH 9?A. Ethyl red, Ka = 8 x 10-2

B. Congo red, Ka = 9.0 x 10-3

C. Cresol red, Ka = 1.0 x 10-8

D. Alizarin blue, Ka = 7.0 x 10-11

7. Which of the following would be a typical pH of lemon juice?A. 1.0B. 3.0C. 6.8D. 7.2

8. What is the [KOH] in a KOH solution that has a pH = 13.00?A. 0.10 MB. 0.56 MC. 1.0 MD. 1.0 x 10-13 M


10. The S2- ion is a base with an equilibrium constant of 0.77, what is the Ka value for HS-?A. 1.3 x 10-14

B. 9.1 x 10-8

C. 1.1 x 10-7

D. 7.7 x 1013

11. Which of the following is true as a result of the predominant hydrolysis of NaHCO3?

Solution ReasonA. basic Ka > KbB. basic Ka < KbC. acidic Ka > KbD. acidic Ka < Kb

12. Consider the ionization of water: H2O + H2O ⇌ H3O+ + OH-

What happens to the pH when 0.1 M NaOH is added to water?A. pH increases since [H3O+] increases.B. pH increases since [H3O+] decreases.C. pH decreases since [H3O+] increases.D. pH decreases since [H3O+] decreases.

13. Which of the following amphiprotic ions will act predominantly as a base in solution?A. HSO3

-

B. HSO4-

C. H2PO4-

D. HPO42-

14. Which of the following statements is true for a alkaline solution at 25 0C ?A. pH < 7.0B. pOH > 7.0C. [H3O+] > [OH-]D. [H3O+] < [OH-]

15. What is the [OH-] in 0.025 M HCl ?A. 2.5 x 10-16 MB. 4.0 x 10-13 MC. 2.5 x 10-2 MD. 2.5 x 10-12 M

16. Which of the following statements is true for an acidic solution at 25 0C ?A. pH > 7.0B. pOH < 7.0C. [H3O+] < [OH-]D. [H3O+] > [OH-]

17. What is the [H+] in 0.025 M NaOH?A. 2.5 x 10-16 MB. 4.0 x 10-13 MC. 2.5 x 10-2 MD. 2.5 x 10-12 M

18. What [H3O+] results when 25 mL of 1.0 M HCl is mixed with 10 mL of 2.0 M KOH ?A. 0.70 MB. 0.00 MC. 0.10 MD. 0.14 M

19. What is the pH of 0.70 M HCN ?A. 0.70 MB. 1.9 x 10-5 MC. 5.32D. 4.73

20. Consider the following indicator equilibrium: HInd + H2O ⇌ H3O+ + Ind-

yellow blueWhat is the result of adding FeCl3 to this indicator?

Equilibrium Shift ColourA. left blueB. left yellowC. right blueD. right yellow

21. An indicator changes colour when 1.0 M HCl is added. If the indicator has a Ka = 1.0 x 10-10, determine the indicator and the pH at its transition point.

Indicator pHA. phenolphthalein 4.0B. phenolphthalein 10.0C. thymolphthalein 4.0D. thymolphthalein 10.0

22. What is the Kb value for HC6H5O72- ?

A. 1.0 x 10-14

B. 5.9 x 10-10

C. 2.4 x 10-8

D. 4.1 x 107

23. What is produced when MgO is added to water?A. the metal MgB. the acid HMgOC. the base Mg(OH)2

D. the amphiprotic species HMgO2

24. Which of the following is a major source of NO2(g), which contributes to the problem of acid rain?A. a smelterB. an air conditionerC. a nuclear power plantD. the automobile engine

25. The indicator bromothymol blue can be described by the following equilibrium equation:

HInd + H2O ⇌ H3O+ + Ind-

yellow blue A base is added to the above indicator. After equilibrium has been re-established,

how do the H3O+ and the colour of the solution compare with the original equilibrium?

H3O+ Colour of SolutionA. decreases turns blueB. decreases turns yellowC. increases turns blueD. increases turns yellow

26. The indicator bromothymol blue can be described by the following equilibrium equation:

HInd + H2O ⇌ H3O+ + Ind-

yellow blue An acid is added to the above indicator. After equilibrium has been re-established,

how do the H3O+ and the colour of the solution compare with the original equilibrium?

H3O+ Colour of SolutionA. decreases turns blueB. decreases turns yellowC. increases turns blueD. increases turns yellow

27. What is the Ka value for the indicator neutral red? A. 1 x 10-14

B. 4 x 10-8

C. 7.4D. 14.0

28. What is a common source of SO2(g) ? A. automobilesB. a car batteryC. a lead smelterD. oxidation of iron


30. What is the value of pKw for water at 25 0C ?A. 1.0 x 10-14

B. 10 x 10-7

C. 7.00D. 14.00

Acids Quiz #8 Buffers and Indicators

1. Consider the following acid solutions:I. H2CO3II. HClO4 III. HF

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

A. I onlyB. II onlyC. I and III onlyD. I, II, and III only

2. Consider the following base indicator:HInd ⇄ H+ + Ind-

When the indicator is added to different solutions, the following data are obtained: Solution 1.0 M HCl 1.0 M HA1 1.0 M HA2

Colour Yellow Blue YellowThe acids HAl, HA2, and HInd listed in the order of decreasing acid strength isA. HA2, HInd, HAl

B. HInd, HAl, HA2

C. HA2, HAl, HIndD. 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. AmmoniaB. Nitric acidC. Sodium nitrateD. 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 toA. Ka

B. [HF]C. [H2O]D. [H3O+]

6. A basic buffer solution can be prepared by mixing equal numbers of moles ofA. NH4Cl and HClB. NaCl and NaOHC. Na2CO3 and NaHCO3

D. NaCH3COO and CH3COOH

7. Which of the following would be used to prepare an acidic buffer solution?A. HF and H3O+

B. NaHS and H2SC. NH3 and NH4ClD. HCl and NaCl

8. HCN(aq) + H2O ⇌ H3O+(aq) + CN-

(aq)

1.0 M 0.0001 M 0.0001 M Why is the solution above not considered to be a true buffer solution?

A. excessive [HCN]B. excessive [H3O+]C. insufficient [CN-]D. insufficient [HCN]

9. What could be added to 1.0 L of this solution in order for it to behave as a true buffer?A. 1.0 mol HClB. 1.0 mol HCNC. 1.0 mol H3O+

D. 1.0 mol NaCN

10. Which of the following could typically be used to prepare a buffer solution?A. NaHS and H2SB. H2S and Na2SC. HNO3 and NaNO3

D. HNO3 and NaNO2

11. What happens to the pH of a buffer solution if a small amount of acid is added?A. The pH remains constant.B. The pH increases slightly.C. The pH decreases slightly.D. The pH decreases significantly.

12. Consider the following buffer equilibrium:HF + H2O ⇌ H3O+ + F-

What concentration would limit the buffering action if acid were added?A. [F-]B. [HF]C. [H2O]D. [H3O+]

13. Consider the following buffer equilibrium:HF + H2O ⇌ H3O+ + F-

What concentration would limit the buffering action if base were added?A. [F-]B. [HF]C. [H2O]D. [H3O+]

14. Which term does the following statement best describe? A mixture of a weak acid and its conjugate base, each with distinguishing colours. A. bufferB. titrationC. indicatorD. conjugate

15. The HC2O4-(aq) ion will act as

A. a base since Ka > KbB. a base since Ka < KbC. an acid since Ka > KbD. an acid since Ka < Kb

16. What do a chemical indicator and a buffer solution typically both contain? A. a strong acid and its conjugate acid

B. a strong acid and its conjugate baseC. a weak acid and its conjugate acid

D. a weak acid and its conjugate base

17. What is the approximate pH and Ka at the transition point for phenol red?

pH KaA. 6.0 6.3 x 10-7

B. 7.3 3.1 x 10-14

C. 7.3 5.0 x 10-8

D. 8.0 1.0 x 10-8

18. Which of the following acids could not be present in a buffer solution?A. HFB. HNO2

C. H2SO3

D. HCl

19. The indicator phenol red will be red in which of the following solutions?A. 1.0 M HFB. 1.0 M HBrC. 1.0 M NH4ClD. 1.0 M NaHCO3

20. Which of the following chemical indicators has a Ka = 1.6 x 10-4 ?A. methyl orangeB. phenolphthaleinC. thymolphthaleinD. bromcresol green

21. Which of the following chemical indicators has a Ka = 7.9 x 10-10 ?A. methyl orangeB. phenolphthaleinC. thymolphthaleinD. bromcresol green

22. What is the Ka for phenolphthalein?A. 1 x 10-14

A. 8 x 10-10

A. 1 x 109

A. 1.0 x 10-7

23. What volume of 0.250 M KOH is required to titrate 0.00230 mol of the weak acid H2CO4 ?A. 1.15 mLB. 4.60 mLC. 9.20 mLD. 18.4 mL

24. What is the net ionic equation for the titration of H3PO4(aq) with Sr(OH)2(aq) ? A. H+

(aq) + OH- (aq) ® H2O(l)

B. 6H+(aq) + 6OH-

(aq) ® 6H2O(l)

C. 2H3PO4(aq) + 3Sr2+(aq) + 6OH-

(aq) ® Sr3(PO4)2(s) + 6H2O(l)

D. 6H+(aq)

+ 2PO43-

(aq) + 3Sr2+(aq) + 6OH-

(aq) ® 3Sr2+ + 2PO43-

(aq) + 6H2O(l)

25. Which of the following will have the smallest Ka value?A. HFB. H2CO3

C. H3PO4

D. CH3COOH

26. Which of the following will have the smallest Kb value?A. IO3

-

B. NH3

C. CN-

D. HPO42-

27. What volume of 0.500 M NaOH is required to neutralize 25.0 mL 0.250 M HBr ?A. 5.00 mLB. 12.5 mLC. 20.0 mLD. 25.0 mL

28. Which of the following equations correctly represents the reaction of a metallic oxide with water?A. K2O + H2O ® 2KOHB. SO3 + H2O ® H2SO4

C. Na2O + H2O ® Na2O2 + H2

D. NaOH + HCl ® NaCl + HOH

29. Which of the following equations describes the dissociation of the salt, ammonium chloride, in water?A. NH4Cl(s) ® NH4Cl(aq)

B. NH4+

(aq) + Cl-(aq) ® NH4Cl(s)

C. NH4Cl(s) ® NH4+

(aq) + Cl-(aq)

D. NH4+

(aq) + H2O(l) ® H3O+(aq) + NH3(aq)

30. Which of the following salt solutions is acidic?A. NaBrB. AlCl3

C. Li2CO3

D. NaHCO3

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 redB. Methyl violetC. Indigo carmineD. Phenolphthalein

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

A. HCl and LiOHB. HNO3 and NH3

C. HClO4 and NaOHD. 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 HClB. 6.00 M HNO3

C. 0.105 M NaOHD. 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 onlyB. I and II onlyC. II and III onlyD. I, II and III

5.

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

A. HF and KOHB. HCl and NH3

C. H2S and NaOHD. HNO3 and KOH

6. A suitable indicator for the above titration is

A. Methyl violetB. Alizarin yellowC. Thymolphthalein D. Bromcresol green

7. The pH scale is

A. directB. inverseC. logarithmicD. exponential

8.

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

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

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

A. Orange IVB. Methyl redC. 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 redB. Thymol blueC. Methyl orangeD. 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 NaOHB. 0.001 M NaOHC. 3.00 M NaOHD. 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 mLB. 150 mLC. 250 mLD. 500mL

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

A. HCl and Cl-

B. NH3 and NH2-

C. H2CO3 and HCO3-

D. H2C6H5O7 and HC6H5O72-

Acids Quiz #10 Review Answers

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 molC. 5.00 x 10-3 molD. 9.00 x 10-3 mol

2. The approximate Ka for the indicator phenolphthalein isA. 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 redB Equilibrium shifts left, and colour becomes blueC. Equilibrium shifts right, and colour becomes redD. Equilibrium shifts right, and colour becomes blue

4. Which one of the following combinations would act as an acid buffer?A. HCl and NaOHB. KOH and KBrC. NH3 and NH4ClD. 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. 0.98B. 2.10C. 7.00D. 11.90

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

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 withA. NH3 and the pH decreases

pH pH

B.

pH

C.

pH

D.

B. NH4+ and the pH decreases

C. NH3 to keep the pH relatively constantD. 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 hydroxideB. Acetic acid and sodium hydroxideC. Sulphuric acid and barium hydroxideD. 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 HIndB. Products are favored and the stronger acid is HCNC. Reactants are favored and the stronger acid is HIndD. Reactants are favored and the stronger acid is HCN

10. The indicator methyl red is red in a solution of NaH2PO4. Which of the following equations is consistent with this observation?A. H2PO4

- + H2O ⇄ HPO42- + H3O+

B. H2PO4- + H2O ⇄ H3PO4

+ OH-

C. HPO42- + H2O ⇄ PO4

3- + H3O+

D. HPO42- + 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 HClB. NaOH and HNO3

C. Ba(OH)2 and H2SO4

D. KOH and CH3COOH

13. Oxalic acid dihydrate is a pure, stable, crystalline substance. Which of the following describes one of its uses in acid-base titrations? A. bufferB. primary standardC. indicatorD. acid anhydride

14. What is the complete ionic equation that describes the reaction of HCl(aq) with Pb(OH)2(s)?A. H+

(aq) + OH-(aq) ® H2O(l)

B. 2HCl(aq) + Pb(OH)2(s) ® PbCl2(s) + 2H2O(l)

C. 2H+(aq) + 2Cl-

(aq) + Pb(OH)2(s) ® PbCl2(s) + 2H2O (l)

D. 2H+(aq) + 2Cl-

(aq) + Pb2+(aq) + 2OH-

®Pb2+(aq) + 2Cl-

(aq) + 2H2O (l)

15. What is the formula equation that describes the reaction of HCl(aq) with Pb(OH)2(s)?

A. H+(aq) + OH-

(aq) ® H2O(l)

B. 2HCl(aq) + Pb(OH)2(s) ® PbCl2(s) + 2H2O(l)

C. 2H+(aq) + 2Cl-

(aq) + Pb(OH)2(s) ® PbCl2(s) + 2H2O (l)

D. 2H+(aq) + 2Cl-

(aq) + Pb2+(aq) + 2OH-

®Pb2+(aq) + 2Cl-

(aq) + 2H2O (l)

16. What term is used to describe the point at which a chemical indicator changes colour?A. titration pointB. transition pointC. equivalence pointD. neutralization point

17. What term is used to describe the point in a titration where the acid has completely reacted with the base.A. titration pointB. transition pointC. equivalence pointD. neutralization point

18. Which of the following is a piece of equipment typically used in acid-base titrations?A. buretteB. cuvetteC. litmus paperD. graduated cylinder

19. Identify an environmental problem associated with acid rain.A. increasing the pH of lakesB. the green house effectC. chemical decomposition of rainwaterD. metals leeching from river rocks accumulating in lakes

20. A buffer solution is prepared using sufficient amounts of H2S and NaHS. What limits this buffer’s effectiveness when NaOH is added?A. [H2S]B. [HS-]C. [OH-]D. [H3O+]

21. Which of the following is not a good use for an acid-base titration curve?A. to determine the concentration of the acidB. to select a suitable indicator for the titrationC. to determine whether the acid is strong or weakD. to select a suitable primary standard for the titration

22. A substance which completely produces hydroxide ions in solution is a definition of which of the following?A. a strong Arrhenius acidB. a strong Arrhenius baseC. a weak Arrhenius baseD. a strong Brønsted-Lowry base

23. When a strong acid is titrated with a strong base, what will the pH value be at the equivalence point?

A. 0.0B. 5.0C. 7.0D. 9.0

24. At a certain point in a strong acid-strong base titration, the moles of H+ are equal to the moles of OH-. This is a definition of which of the following?A. the end pointB. the titration pointC. the transition pointD. the equivalence point

25. A 25.0 mL sample of H2SO4(aq) is titrated with 15.5 mL of 0.50 M NaOH(aq). What is the concentration of the H2SO4(aq) ?A. 0.078 MB. 0.16 MC. 0.31 MD. 0.62 M

26. When a weak acid is titrated with a strong base, what will the pH value be at the equivalence point?A. 0.0B. 6.7C. 7.0D. 8.8

27. A substance which completely produces accepts a proton from an acid is a definition of which of the following?A. a strong Arrhenius acidB. a strong Arrhenius baseC. a weak Arrhenius baseD. a strong Brønsted-Lowry base

28. Which of the following is a piece of equipment typically used in acid-base titrations?A. pipetteB. test tubeC. litmus paperD. graduated cylinder

29. Water has the greatest tendency to act as an acid with which of the following?A. NO3

-

B. NO2-

C. H2PO4-

D. CH3COO-

30. Which net ionic equation best describes the reaction between NaOH and H2S?A. H+

(aq) + OH-(aq)

® H2O(l)

B. H2S(aq) + 2OH- ® 2H2O(l) + S2-(aq)

C. H2S(aq) + 2NaOH ® 2H2O(l) + Na2S(aq)

D. 2H+(aq) + S2-

(aq) + 2NaOH(aq) ® 2H2O(l) + 2Na+(aq) + S2-

(aq)

Web Review of Acids

1. List five properties of:

a) acids

b) bases

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

3. Define:

Conjugate

Arhenius strong acid

Bronsted weak acid

Bronsted strong base

Ionization of water

Equivalence point

Transition point

Buffer

Hydrolysis.

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

HC03- + HF ⇄ H2CO3 + F-

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

NH3 AlCl3 H2CO3 HClO4 KCN NH4Cl KOH

SO2 NaF HCl NaI K2O NaOH CO2

NH3 and NH4Cl NaCH3COO and CH3COOH

6. H+is short for _______.

7. Determine the conjugates for each of the bases.

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

8. Determine the conjugates for each of the acids.

HF HCN Al(H2O)63+ NH4

+ HPO42-

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

strong acid weak acid strong base weak base

Conductivity

Size of Ka

Size of Kb

Degree of Ionization

pH.

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

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

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

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

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

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

Weak Acid or Base Salt

HF

NaCH3COO

NH3

NaCN

H2CO3

KH2PO4

HCH3COO

16. Match each equation with its type:

Acid/base formula equation F- + HOH(l) ⇄ HF + OH-

Acid/base net ionic equation HCl + NaOH →NaCl + HOH(l)

Solubility product ionization equation H+ + OH- → HOH(l)

Hydrolysis of a weak acid AgCl(s) ⇄ Ag+ + Cl-

Hydrolysis of a weak base H20(l) ⇄ H+ + OH-

Ionization of water NH4+ + H20 ⇄ NH3 + H3O+

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

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

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

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

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

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

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

24.

a) An indicator HInd is red in acid and blue in base. Write the equation for the indicator and explain the colours.

b) What is true at the transition point?

c)What is the color of this indicator in a solution of AlCl3?

d) In the above solution, what is larger [HInd] or [Ind-] ?

e) Calculate the Ka for Phenolphthalein.

f) What indicator has a Ka of approximately 1.0 X 10-10?

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

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

Acid Base pH of Equivalence Point Indicator

HCl NaOH

H2SO4 NH3

HF KOH

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

HCLO HClO2 HClO3 HClO4

28. Pick the formulae that are amphiprotic.

H2SO4 H20 F- HCO3-

CO3-2 KOH H2PO4

- HPO4-2

CALCULATIONS

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

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

[H+] [OH-] pH pOH Acid/base/neutral

5.0 x 10-3 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).

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

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

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

20. The following data were recorded when 25.00 mL of 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 = [H3O + ][OH - ] [H2O]

B. Kw = [H3O+]2[O2]C. Kw = [H3O+][OH-]D. Kw = [H3O+]2[O2-]

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

A buffer solution is present at pointA. IB. IIC. IIID. 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 H2OB. 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 redB. phenol redC. thymol blueD. methyl violet

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

[Sr(OH)2]?A. 6.00 x 10-3 MB. 6.00 x 10-2 MC. 1.20 x 10-1 MD. 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 KOHC. HCl with Na2CO3

D. HNO3 with NaOH

Volume HCl added

pH14

7

0

III

IIIIV

8. An Arrhenius acid is defined as a chemical species thatA. is a proton donor.B. is a proton acceptor.C. produces hydrogen ions in solution.D. produces hydroxide ions in solution.

9. Consider the following acid-base equilibrium system: HC2O4- + H2BO3

- ⇄ H3BO3 + C2O42-

Identify the Bronsted-Lowry bases in this equilibrium.A. H2BO3

- and H3BO3

B. HC2O4- and H3BO3

C. HC2O4- and C2O4

2-

D. H2BO3- and C2O4

2-

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

A. CH3COO- + H2O ⇄ CH3COOH + OH-

B. CH3COO- + H2O ⇄ H2O + CH2COO2- 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 HFB. 0.10 M NaFC. 0.10 M H2SO3

D. 0.10 M NaHSO3

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

A. NH3

B. CO32-

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 acidB. weak baseC. strong acid]D. strong base

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

A. H2O ⇄ 2H+ + O2-

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 increaseB. both [H3O+] and [OH-] to decreaseC. [H3O+] to increase and [OH-] to decreaseD. [H3O+] to decrease and [OH-] to increase

16. Consider the following:I. H2SO4

II. HSO4-

III. SO42-

Which of the above is/are present in a reagent bottle labeled 1.0 M H2SO4?A. I onlyB. I and II onlyC. II and III only

D. I, II, and III

17. The pH of 0.10 M KOH solution isA. 0.10B. 1.00C. 13.00D. 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. HBrII. H2OIII. HCO3

-

IV. H2C6H5O7-

A I and II onlyB. II and IV onlyC. II, III, and IV onlyD. 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 MB. 0.50 MC. 1.0 MD. 2.0 M

22. Both acidic and basic solutions

A. taste sourB. feel slipperyC. conduct electricityD. turn blue litmus red

23. The conjugate acid of HPO4

2- isA. PO4

3-

B. H2PO4-

C. H2PO42-

D. H2PO43-

24. What is the value of the Kw at 25 oC?A., 1.0 x 10-14

B. 1.0 x 10-7

C. 7D. 14

25. Consider the following equilibrium: 2H2O(l) ⇄ H3O+

(aq) + OH-(aq)

A small amount of Fe(H2O)63+ is added to water and equilibrium is re-established. Which of the following

represents the changes in ion concentrations?[H3O+] [OH-]

A. increases increasesB. increases decreasesC. decreases decreasesD. 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 isA. 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 redB. methyl redC. indigo carmineD. 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 isA. 1.43 x 10-3 MB. 2.86 x 10-3 MC. 5.72 x 10-2 MD. 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. basicB. acidicC. neutralD. amphiprotic

30. Which of the following salts will dissolve in water to produce a neutral solution?A. LiFB. CrCl3

C. KNO3

D. NH4Cl 31. What is the value of the Kb for HC6H5O7

2-?A. 5.9 x 10-10

B. 2.4 x 10-8

C. 4.1 x 10-7

D. 1.7 x 10-5

32. The pOH of 0.015 M HCl solution isA. 0.97B. 1.80C. 12.18D. 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. LiFB. CrCl3

C. KNO3

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

A. LiFB. CrCl3

C. KNO3

D. NH4Cl36. A student mixes 400 mL of 0.100 M NaOH with 100 mL of 0.200 M H2SO4. The resulting solution has a

pH ofA. 14.000B. 0.000C. 13.800D. 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 ofA. 14.000B. 0.000C. 13.000D. 7.000

38. The strongest acid in water isA. HClO4

B. HIC. HFD. H3O+

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

A. HFB. H2CO3

C. H2C2O4

D. HCN39. The formula that has the highest pH in water is

A. NaFB. NaClC. 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 redB blueC yellowD 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 mLB 20.0 mLC 25.0 mLD 40.0 mL

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

A neutral redB thymol blueC thymolpthaleinD chlorophenol red

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

[H3O+]A increasing slightlyB decreasing slightlyC increasing considerablyD decreasing considerably

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

A 0.10 M HFB 0.10 M HClC 0.10 M NaClD 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 ionizesB neutralizes an acidC has a pH greater than 7D turns bromocresol green from yellow to blue

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

A HNO2 + CN- ⇄ NO2- + HCN

B H2S + HCO3- ⇄ HS- + H2CO3

C H3PO4 + NH3 ⇄ H2PO4- + NH4

+

D CH3COOH + PO43- ⇄ CH3COO- + HPO4

2-

8. What is the pOH of a solution prepared by adding 0.50 moles of NaOH to prepare 0.50 L of solution?A 0.00B 0.30C 14.00D 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 acidB adding a strong baseC increasing the temperatureD decreasing the temperature

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

A 1.50 MB 1.67 MC 3.33 MD 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 MB 1.0 x 10-7 MC 1.0 x 10-5 MD 0.0 M

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

A HCl and NaClB HCN and NaCNC KNO3 and KOHD 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

isA 0.010 MB 0.050 MC 0.10 MD 0.20 M

16. The conjugate base of an acid is produced by

A adding a proton to the acidB adding an electron to the acidC removing a proton from the acidD removing an electron from the acid

17. Which of the following represents the predominant reaction between HCO3

- and water?A 2HCO3

- ⇄ H2O + 2CO2

B HCO3- + H2O ⇄ H2CO3 + OH-

C HCO3- + H2O ⇄ H3O+ + CO3

2-

D 2HCO3- + H2O ⇄ H3O+ + CO3

2- + OH- + CO2

18. Water acts as an acid when it reacts with

I CN-

II NH3

III HClO4

IV CH3COO-

A I and IV onlyB II and III onlyC I, II, and IVD II, III, and IV

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

A [H3O+] > [HSO4-] > [SO4

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

2-] > [HSO4-] > [OH-]

C [OH-] > [SO42-] > [HSO4

-] > [H3O+]D [SO4

2-] > [HSO4-] > [OH-] > [H3O+]

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

A CO2

B CaOC MgOD Na2O

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

I 0.10 M HClII 0.10 M HNO2

III 0.10 M NaOHA I onlyB II onlyC I and II onlyD I, II, and III

22. Ka for the acid H2AsO4

- is 5.6 x 10-8. What is the value of the Kb for HAsO42-?

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 HClC NaF and HClD Ca(OH)2 and CH3COOH

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

A KClB NH4BrC Fe(NO3)3

D LiCH3COO 25. Calculate the pH in a 0.200 M solution of Sr(OH)2.

A 1.40B 1.70C 13.30D 13.60

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

A NaClB LiOHC NH4NO3

D KCH3COO 27. Which of the following will form a basic aqueous solution?

A HSO3-

B HSO4-

C HPO42-

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 0.00B 1.40C 1.70D 7.00

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

Acid + Base ⇄ Base + AcidA H2O2 SO3

2- ⇄ HO2- HSO3

-

B H2O2 SO32- ⇄ HSO3

- HO2-

C SO32- H2O2 ⇄ HO2

- HSO3-

D SO32- 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 baseB strong acid with a weak baseC weak acid with a strong baseD strong acid with a strong base

32. A buffer solution may contain equal moles of

A weak acid and strong baseB strong acid and strong baseC weak acid and its conjugate baseD 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 BaSB NH4ClC 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 ionizesB neutralizes and acidC has a pH greater than 7D 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 beA HFB H2SC HCND 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 3Final volume of H2C2O4 from burette (mL) 23.00 39.05 20.95Initial volume of H2C2O4 from burette (mL) 4.85 23.00 5.00Calculate 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.

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€¦ · Web view · 2009-04-2320. HS- 21. CH3COO- ... C. toothpaste. D. oven cleaners. 12. Which of the following is the strongest acid? ... 19. The indicator phenol red will