Acids Worksheet Questions

Power Point Lesson Notes- double click on the lesson number

Worksheets Quiz

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

2. Arrhenius, Bronsted Bases, Kb and Strength WS 2

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

4. Leveling effect, Anhydrides and Relationships. WS 4

5. Hydrolysis of Salts. WS 5 3

6. Acid, Base & Salt Reactions. Hydrolysis. WS 6 ClassifyingEverything Activity

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

8. Application of Hydrolysis and Ionization of Water Hydrolysis Quiz 1 Hydrolysis Quiz 1 Answers

Hydrolysis Quiz 2 Hydrolysis Quiz 2 Answers

9. Ionization of Water pH scale. WS 8 Hydrolysis Quiz 3 Hydrolysis Quiz 3 Answers
10. pH Calculations for Weak Acids. WS 9 5

11. Ka from pH for Weak Acids. WS 10 6

12. Indicators Lab. Lab handout 1 Lab handout 2 Data

13. Kbs from Kas for Weak Bases. WS 11

14. pH for Weak Bases pH [H+] [ OH- ] Relationships. WS 12 7

15.Amphiprotic Ions- Kas and Kbs. WS 13

16. Titration Lab. Primary Standards. Acids Midtermreview Test

17. Titration Lab

18. Buffers & Indicators WS 14 8

19. Titration Curves. WS 15 9     and    10

20. Review #1 Web Site Review Sheet Quizmebc

21. Review #2 Practice Test # 1 Practice Test 2

22. Test

WS #1 Conjugate Acid-Base Pairs

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

Acids conduct electricity, taste sour, neutralize bases, change the color of indicators, and react with some metals to produce hydrogen.

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

Bases conduct electricity, taste bitter, neutralize acids, change the color of indicators, and feel slippery.

3. Make some brief notes on the commercial acids: HCl and H₂SO₄ (p 112).

HCl - five common use

H₂SO₄- five common use

4. Make some brief notes on the commercial base NaOH - five common uses.

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

Concentrated means relatively high molarity and dilute means relatively low molarity.

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

A strong acid completely ionizes and a weak acid partially ionizes.

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

A weak acid could have a high molarity and the strong acid could have a low molarity.

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

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

1. HCN + H₂O ⇄ H₃O⁺ + CN^-

Acid Base Acid Base

2. H₃C₆O₇ + H₂O ⇄ H₂C₆O₇^- + H₃O⁺

acid base base acid

3. H₃PO₄ + H₂O ⇄ H₂PO₄^- + H₃O⁺

acid base base acid

4. HF + H₂O ⇄ F^- + H₃O⁺

acid base base acid

5. H₂CO₃ + H₂O ⇄ HCO₃^- + H₃O⁺

acid base base acid

6. NH₄⁺ + H₂O ⇄ NH₃ + H₃O⁺

acid base base acid

7. CH₃COOH + H₂O ⇄ CH₃COO^- + H₃O⁺

acid base base acid

8. HCl + H₂O → Cl^- + H₃O⁺

acid base base acid

9. HNO₃ + H₂O → H₃O⁺ + NO₃^-

acid base acid base

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

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

[HCN] [H₂CO₃]

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

[H₃C₆O₇] [NH₄⁺]

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

[H₃PO₄] [CH₃COOH]

13. Ka = [H₃O⁺] [F^-]

[HF]

17. Which acids are strong? The six on the top of the acid chart are strong.

18. What does the term strong acid mean? They complete ionization into ions. Such as: HCl + H₂O → Cl^- + H₃O⁺

19. Why is it impossible to write an equilibrium expression for a strong acid? Ka = [H₃O⁺] [Cl^-]

[HCl] is equal to zero and in math numbers divided by zero are undefined. [HCl]

20. Which acids are weak?

All acids listed on the acid chart below the top six.

23. What does the term weak acid mean?

Incomplete ionization. Such as: HF + H₂O ⇄ F^- + H₃O⁺

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

A strong acid is a good conductor. A weak acid conducts but not so good.

Acid Conjugate Base Base Conjugate Acid

14. HNO₂ NO₂^- 15. HCOO^- HCOOH

16. HSO₃^- SO₃^2- 17. IO₃^- HIO₃

18. H₂O₂ HO₂^- 19. NH₃ NH₄⁺

20. HS^- S^2- 21. CH₃COO^- CH₃COOH

22. H₂O OH^- 23. H₂O H₃O⁺

Define:

22. Bronsted acid- a proton donor

23. Bronsted base- a proton acceptor

24. Arrhenius acid- a substance that ionizes in water to produce H⁺

25. Arrhenius base- a substance that ionizes in water to produce OH^-

26. List the six strong acids. HCLO₄ HI HBr HCl HNO₃ H₂SO₄

27. Rank the acids in order of decreasing strength.

HCl HSO₄^- H₃PO₄ HF H₂CO₃ H₂S

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

Vinegar. It is a weak acid and produces much less H₃0⁺ ion which is the corrosive part of an acid.

Making a Universal Indicator Lab Activity

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

Yamada’s Universal Indicator

5 drops thymol blue

8 drops methyl orange

5 drops phenolphthalein

10 drops bromothymol blue

20 drops of water

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

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

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

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

Wash and dry your acetate.

Wash and return your eyedropper.

Wash and return your beaker.

Wash your hands.

Results

Compound Conductivity pH Classification

HCl good 1 strong acid

CH₃COOH ok 3 weak acid

NaOH good 13 strong base

NH₃ ok 11 weak base

NaHCO₃ good 11 weak base

H₂CO₃ ok 3 weak acid

NaCl good 7 neutral salt

WS # 2 Conjugate Acid-Base Pairs

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

1. HCN + H₂O ⇄ H₃O⁺ + CN^-

2. HCl + H₂O ⇄ H₃O⁺ + Cl^-

3. HF + H₂O ⇄ H₃O⁺ + F^-

4. F^- + H₂O ⇄ HF + OH^-

5. HSO₄^- + H₂O ⇄ H₃O⁺ + SO₄^2-

(acid)

6. NH₄⁺ + H₂O ⇄ H₃O⁺ + NH₃

7. HPO₄^2- + H₂O ⇄ H₂PO₄^- + OH^-

(base)

Acid Conjugate Base Base Conjugate Acid

8. HCO₃^- CO₃^2- 9. CH₃COO^- CH₃COOH

10. HPO₄^-2 PO₄^3- 11. IO₃^- HIO₃

12. H₂O OH^- 13. NH₂^- NH₃

14. HS^- S^2- 15. C₂H₅SO₇^3- HC₂H₅SO₇^2-

16. Circle the strong bases.

Fe(OH)₃ NaOH CsOH KOH

Zn(OH)₂ Sr(OH)₂ Ba(OH)₂ Ca(OH)₂

17. Rank the following acids from strongest to weakest.

H₂S CH₃COOH H₂PO₄^- HI HCl HF

5 4 6 1 1 3

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

H₂O F^- NH₃ SO₃^2- HSO₃^- NaOH

6 4 2 3 5 1

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

H₃BO₃ + H₂O ⇄ H₂BO₃^- + H₃O⁺

ii) Write the K_a expression for the above.

         [H₃O⁺]         [H₂BO₃^-]

K_a = [H₃BO₃]

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

20. List six strong acids. HClO₄ HI HBr HCl HNO₃ H₂SO₄

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

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

HIO₃ H₂C₂O₄ H₂SO₃ HSO₄^- H₃PO₄ HNO₂

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

PO₄^3- CO₃^2- CN^- NH₃ H₂BO₃^- HS^-

WS # 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^- + HPO₄^2- ⇄ SO₄^2- + H₂PO₄^-

Acid Base Base Acid

Products are favoured as HSO₄^- is a stronger acid than H₂PO₄^-

2. HCN + H₂O ⇄ H₃O⁺ + CN^-

Acid Base Acid Base

Reactants are favoured as H₃O⁺ is a stronger acid than HCN.

3. HCO₃^- + H₂S ⇄ H₂CO₃ + HS^-

Base Acid Acid Base

Reactants are favoured as H₂CO₃ is a stronger acid than H₂S

4. HPO₄^2- + NH₄⁺ ⇄ H₂PO₄^- + NH₃

Base Acid Acid Base

Reactants are favoured as H₂PO₄^- is a stronger acid than NH₄⁺

5. NH₃ + H₂O ⇄ NH₄⁺ + OH^-

Base Acid Acid Base

Reactants are favoured as OH^- is a stronger base than NH₃

6. H₂PO₄^1- + NH₃ ⇄ HPO₄^2- + NH₄⁺

Acid Base Base Acid

Products are favoured as NH₃ is a stronger base than HPO₄^2-

7. HCO₃^- + HF ⇄ H₂CO₃ + F^-

Base Acid Acid Base

Products are favoured as HF is a stronger acid than H₂CO₃

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

HSO₄^- + HCO₃^- ⇄ H₂CO₃ + SO₄^2- products are favoured

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

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

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

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

HF is a stronger acid and creates more ions.

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

Both are strong acids and have the same strength as both completely ionize to from H⁺.

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

Acid: H⁺ or H₃O⁺ Base: OH^-

12. Which is the stronger base? Which produces the least OH^-?

F^- is the weaker base and produces the least OH- CO₃^-2 is the stronger base

13. Define a Bronsted/Lowry acid and base.

An acid is a proton donor and a base is a proton acceptor.

14. Define an Arrhenius acid and base.

An acid ionizes in water to produce H⁺ and a base ionizes in water to produce OH^-.

15. Complete each reaction and write the equilibrium expression.

HF + H₂O ⇄ H₃O⁺ + F^- K_a= [ H₃O⁺][ F^-] K_b= [HF][ OH^-]

F^- + H₂O ⇄ HF + OH^- [HF] [F^-]

16. H₂SO₄ + 2NaOH → Na₂SO₄ + 2HOH

17. Define conjugate pairs.

Acid base pairs that differ by one proton.

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

H₂S NH₄⁺ H₂PO₄^- HOH H₃O⁺ NH₄⁺ HCO₃^-

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

NH₃ F^- HPO₄^-2 HOH O^2- CO₃^-2 OH^-

WS # 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.

H₃O⁺ HCl + H₂O → H₃O⁺ + Cl^-

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

OH- NaOH → Na⁺ + OH^-

3. List three strong acids and three strong bases.

HCl HI HClO₄ NaOH KOH LiOH

4. Rank the acids in decreasing strength:

HClO₄ 1 K_a is very large HClO₃ 2 K_a=1.2x10^-2

HClO₂ 3 K_a=8.0x10^-5 HClO 4 K_a=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) The more O’s the stronger the acid.

6.Which acid is stronger? HI0₃ or HIO₂

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

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

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

NaOH is a strong base.

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

HF is a stronger acid.

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

Strong acid: complete ionization, very large Ka, good conductor, high [H⁺].

Weak acid: partial ionization, small Ka, OK conductor, low [H⁺].

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

Formula Classification Reaction

12. Na₂O basic anhydride Na₂O + H₂O → 2NaOH

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

14. SO₃ acid anhydride SO₃ + H₂O → H₂SO₄

15. CO₂ acid anhydride CO₂ + H₂O → H₂CO₃

16. SO₂ acid anhydride SO₂ + H₂O → H₂SO₃

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

18. NH₃ weak base NH₃ + H₂O D NH₄⁺ + OH^-

19. NaOH strong base NaOH → Na⁺ + OH^-

20. HF weak acid HF + H₂O D H₃O⁺ + F^-

21. H₃PO₄ weak acid H₃PO₄ + H₂O D H₃O⁺ + H₂PO₄^-

WS # 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 parent acid-base formation equation, dissociation equation, and hydrolysis equation (only for acidic and basic salts). For acids and bases write an equation to show how each reacts with water.

1. NaHCO₃ basic salt

H₂CO₃ + NaOH → NaHCO₃ + H₂O

NaHCO₃ → Na⁺ + HCO₃^-

HCO₃^- + H₂O ⇄ H₂CO₃ + OH^-

2. AlCl₃ acid salt

3HCl + Al(OH)₃ → AlCl₃ + 3H₂O

AlCl₃ → Al⁺³ + 3Cl^-

Al(H₂O)₆³⁺ ⇄ Al(H₂O)₅(OH)²⁺ + H⁺

3. NaC₆H₅O basic salt

C₆H₅OH + NaOH → NaC₆H₅O + H₂O

NaC₆H₅O → Na⁺ + C₆H₅O^-

C₆H₅O^- + H₂O ⇄ C₆H₅OH + OH^-

4. Co(NO₃)₃ acid salt

3HNO₃ + Co(OH)₃ → Co(NO₃)₃ + 3H₂O

Co(NO₃)₃ → Co⁺³ + 3NO₃^-

Co(H₂O)₆³⁺ ⇄ Co(H₂O)₅(OH)²⁺ + H⁺

5. Na₂CO₃ basic salt

H₂CO₃ + 2NaOH → Na₂CO₃ + 2H₂O

Na₂CO₃ → 2Na⁺ + CO₃^-2

CO₃^-2 + H₂O ⇄ HCO₃^- + OH^-

6. H₂C₂O₄ weak acid

H₂C₂O₄ + H₂O ⇄ H₃O⁺ + HC₂O₄^-

7. NH₃ weak base

NH₃ + H₂O ⇄ NH₄⁺ + OH^-

8. KCl neutral salt

HCl + KOH → KCl + H₂O

KCl → K⁺ + Cl^-

9. HNO₃ strong acid

HNO₃ + H₂O → H₃O⁺ + NO₃^-

10. RbOH strong base

RbOH → Rb⁺ + OH^-

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

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

1. NH₃ weak base

NH₃ + H₂O ⇄ NH₄⁺ + OH^-

2. NaCl neutral salt

NaCl → Na⁺ + Cl^-

3. HCl strong acid

HCl + H₂O → H₃O⁺ + Cl^-

4. NaCN basic salt

NaCN → Na⁺ + CN^-

CN^- + H₂O ⇄ HCN + OH^-

5. NaOH strong base

NaOH → Na⁺ + OH^-

6. FeCl₃ acid salt

FeCl₃ → Fe⁺³ + 3Cl^-

Fe(H₂O)₆³⁺ ⇄ Fe(H₂O)₅(OH)²⁺ + H⁺

7. HF weak acid

HF + H₂O ⇄ H₃O⁺ + F^-

8. LiHCO₃ basic salt

LiHCO₃ → Li⁺ + HCO₃^-

HCO₃^- + H₂O ⇄ H₂CO₃ + OH^-

9. Fe(NO₃)₃ acid salt

Fe(NO₃)₃ → Fe⁺³ + 3NO₃^-

Fe(H₂O)₆³⁺ ⇄ Fe(H₂O)₅(OH)²⁺ + H⁺

10. MgCO₃ basic salt

MgCO₃ → Mg⁺² + CO₃^-2

CO₃^-2 + H₂O ⇄ HCO₃^- + OH^-

11. H₂S weak acid

H₂S + H₂O ⇄ H₃O⁺ + HS^-

12. HF weak acid

HF + H₂O ⇄ H₃O⁺ + F^-

13. CaI₂ neutral salt

CaI₂ → Ca⁺² + 2I^-

14. Be(OH)₂ weak base

Be(OH)₂ ⇄ Mg⁺² + 2OH^-

15. Ba(OH)₂ strong base

Ba(OH)₂ → Ba⁺² + 2OH^-

16. Describe why Tums (CaCO₃) neutralizes stomach acid. It is a weak base and will neutralize acid.

basic salt

CaCO₃ → Ca⁺² + CO₃^-2

CO₃^-2 + H₂O ⇄ HCO₃^- + OH^-

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

Mg(OH)₂ is strong base, however, acts like a weak base due to its low solubility, and releases OH^- slowly,

whereas NaOH is a strong base which releases OH^- in high concentrations which is corrosive.

Mg(OH)_2(s) ⇄ Mg⁺² + 2OH^-

NaOH → Na⁺ + OH^-

WS # 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.

Data

1. Formula of compound Fe(NO₃)₃

pH 2

Classification acid salt

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

Fe(H₂O)₆³⁺ ⇄ Fe(H₂O)₅(OH)²⁺ + H⁺

2. Formula of compound NaCH₃COO

pH 10

Classification basic salt

Reaction or reactions NaCH₃COO → Na⁺ + CH₃COO^-

CH₃COO^- + H₂O ⇄ CH₃COOH + OH^-

3. Formula of compound K₂HPO₄

pH 10

Classification basic salt

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

HPO₄^-2 + H₂O ⇄ H₂PO₄^- + OH^-

4. Formula of compound HCl

pH 0

Classification strong acid

Reaction or reactions HCl → H⁺ + Cl^-

5. Formula of compound Al₂(SO₄)₃

pH 3

Classification acid salt

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

Al(H₂O)₆³⁺ ⇄ Al(H₂O)₅(OH)²⁺ + H⁺

6. Formula of compound Na₂CO₃

pH 12

Classification basic salt

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

CO₃^-2 + H₂O ⇄ HCO₃^- + OH^-

7. Formula of compound P₂O₅

pH 2

Classification acid anhydride

Reaction or reactions P₂O₅ + H₂O → H₂P₂O₆

H₂P₂O₆ ⇄ H⁺ + HP₂O₆^-

8. Formula of compound Cu(NO₃)₂

pH 4

Classification acid salt

Reaction or reactions Cu(NO₃)₂ → Cu²⁺ + 2NO₃^-

                     Cu(H₂O)₆²⁺ ⇄ Cu(H₂O)₅(OH)⁺ + H⁺

9. Formula of compound Fe₂(SO₄)₃

pH 3

Classification acid salt

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

Fe(H₂O)₆³⁺ ⇄ Fe(H₂O)₅(OH)²⁺ + H⁺

10. Formula of compound N₂O₅

pH 0

Classification acid anhydride

Reaction or reactions N₂O₅ + H₂O → H₂N₂O₆ → 2HNO₃

HNO₃ → H⁺ + NO₃^-

11. Formula of compound Zn(OH)₂

pH 12

Classification weak base

Reaction or reactions Zn(OH)₂ ⇄ Zn⁺² + 2OH^-

12. Formula of compound KHSO₄

pH 2

Classification acid salt

Reaction or reactions KHSO₄ → K⁺ + HSO₄^-

HSO₄^- ⇄ H⁺ + SO₄^2-

13. Formula of compound NaHCO₃

pH 10

Classification basic salt

Reaction or reactions NaHCO₃ → Na⁺ + HCO₃^-

HCO₃^- + H₂O ⇄ H₂CO₃ + OH^-

14. Formula of compound CaCO₃

pH 10

Classification basic salt

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

CO₃^-2 + H₂O <⇄ HCO₃^- + OH^-

15. Formula of compound CaO

pH 12

Classification basic anhydride

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

Ca(OH)₂ → Ca⁺² + 2OH^-

16. Formula of compound Al₂(SO₄)₃

pH 3

Classification acidic salt

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

Al(H₂O)₆³⁺ ⇄ Al(H₂O)₅(OH)²⁺ + H⁺

17. Formula of compound NaCl

pH 7

Classification neutral salt

Reaction or reactions NaCl → Na⁺ + Cl^-

WS # 8 - pH and pOH Calculations

Complete the chart:

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

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

Ba(OH)₂ ⇄ Ba⁺² + 2OH^-

0.20M 0.20M 0.40M

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

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

HCl → H⁺ + Cl^-

0.030M 0.030M

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

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

NaOH → Na⁺ + OH^-

0.20M 0.20M 0.20M

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

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

HCl → H⁺ + Cl^-

300.0 x 0.20 M = 0.075 M 0.075 M pH = -Log[H⁺] = 1.12

800.0

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

HCl → H⁺ + Cl^-

200.0 x 0.020 M = 0.0080 M 0.0080 M pH = -Log[H⁺] = 2.10

500.0

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

Ba(OH)₂ → Ba²⁺ + 2OH^-

150.0 x 0.40 M = 0.12 M 0.12 M 0.24 M

500.0

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

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

Sr(OH)₂ → Sr²⁺ + 2OH^-

250.0 x 0.20 M = 0.083 M 0.083 M 0.1667 M

600.0

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

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

Ba(OH)₂ D Ba²⁺ + 2OH^-

(25)0.40 M 0.20 M 0.40 M

(50)

[OH^-] = 0.40

pOH = 0.40

pH = 13.60

WS # 9 pH Calculations for Weak Acids

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

HCN D H⁺ + CN^-

I 0.20 M 0 0

C x x x

E 0.20 - x x x

x²

= 4.9 x 10^-10

0.20 - x

x = 9.9 x 10^-6 M

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

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

[H⁺] = 2.8 x 10^-2 M [OH^-] = 3.6 x 10^-13 M pH = 1.56 pOH = 12.44

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

[H⁺] = 3.8 x 10^-3 M [OH^-] = 2.6 x 10^-12 M pH = 2.42 pOH = 11.58

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

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

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

Mg(OH)₂ D Mg²⁺ + 2OH^-

x x 2x

Ksp = [Mg²⁺][OH^-]²

5.6 x 10^-12 = 4x³

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

pH = 10.35

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

H₂X D H⁺ + HX^- Note- only lose one proton for any weak acid!!

I 0.200 M 0 0

C x x x

E 0.20 - x x x

Small Ka approximation x = 0

x²

= 1.8 x 10^-6

0.20

x = 6.0 x 10^-4 M

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

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.

Sr(OH)₂ → Sr²⁺ + 2OH^-

350.0 x 0.20 M = 0.0875 M 0.0875 M 0.175 M

800.0

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

WS # 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.

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

HCN D H⁺ + CN^-

I 0.20 M 0 0

C 0.0000100 M 0.0000100 M 0.0000100 M

E 0.19999 0.0000100 M 0.0000100 M

Ka = (0.0000100)² = 5.0 x 10^-10

0.19999

Ka = 5.0 x 10^-10

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

Ka = 3.5 x 10^-4

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

Ka = 1.8 x 10^-5

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

Ka = 7.3 x 10^-10

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

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

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

I 0.10 M 0 0

C 0.0002075 M 0.0002075 M 0.0002075 M

E 0.09979 0.0002075 M 0.0002075 M

Ka = (0.0002075)² = 4.3 x 10^-7

0.09979

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

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

pOH = 14.00 - pH = 2.773

[OH^-] = 10^-pOH = 0.001686 M

NH₃ + H₂O ⇄ NH₄⁺ + OH^-

I 0.20 M 0 0

C 0.001686 M 0.001686 M 0.001686 M

E 0.1983 M 0.001686 M 0.001686 M

Kb= (0.001686)² = 1.4 x 10^-5

0.1983

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

pOH = 14.00 - pH = 2.544

[OH^-] = 10^-pOH = 0.002858 M

CN^- + H₂O ⇄ HCN + OH^-

I 0.40 M 0 0

C 0.002858 M 0.002858 M 0.002858 M

E 0.3971 M 0.002858 M 0.002858 M

Kb= (0.002858)² = 2.1 x 10^-5

0.3971

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

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

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

I 0.10 M 0 0

C 8.55 x 10^-6 M 8.55 x 10^-6 M 8.55 x 10^-6 M

E 0.10 M 8.55 x 10^-6 M 8.55 x 10^-6 M

Ka = (8.55 x 10^-6)² = 7.3 x 10^-10

0.10

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

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

[H⁺] = 10^-2.45 = 0.003548 M

CH₃COOH ⇄ H⁺ + CH₃COO^-

I x 0 0

C 0.003548 M 0.003548 M 0.003548 M

E x - 0.003548 M 0.003548 M 0.003548 M

Keq = [H⁺][CH₃COO^-]

[CH₃COOH]

1.8 x 10^-5 = (0.003548)(0.003548)

[CH₃COOH]

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

1 L 1 mole

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

WS # 11 K_b For Weak Bases

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

1. 1. NaNO₂ (the basic ion is NO₂^-)

2.

Kb(NO₂^-) = Kw = 1.0 x 10^-14

Ka(HNO₂) 4.6 x 10^-4

3.

Kb = 2.2 x 10^-11

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

3. 3. NaHCO₃ Kb = 2.3 x 10^-8

4. NH₃ Kb = 1.8 x 10^-5

5. NaCN Kb = 2.0 x 10^-5

6. Li₂HPO₄ Kb = 1.6 x 10^-7

7. KH₂PO₄ Kb = 1.3 x 10^-12

8. K₂CO₃ Kb = 1.8 x 10^-4

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

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

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

Ka = 4.4 x 10^-7

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

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

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

[H⁺] = 10^-2.87 = 0.001349 M

CH₃COOH ⇄ H⁺ + CH₃COO^-

I 0.10 M 0 0

C 0.001349 M 0.001349 M 0.001349 M

E 0.09865 M 0.001349 M 0.001349 M

Ka = [H⁺][CH₃COO^-]

[CH₃COOH]

Ka = (0.001349)( 0.001349)

(0.09865)

Ka = 1.8 x 10^-5

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

H₂SO₄ + 2NaOH ® Na₂SO₄ + 2HOH

0.200 L x 0.120 mol = 0.0240 mol 0.400 L x 0.140 mol = 0.0560 mol

L L

I 0.0240 mole 0.0560 mole

C 0.0240 mole 0.0480 mole

E 0 0.0080 mole

[OH^-] = 0.0080 mole = 0.013 M

0.6000 L

pOH = 1.88

pH = 12.12

WS # 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^-

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

Ka(HCN) 4.9 x 10^-10

CN^- + H₂O D HCN + OH^-

I 0.20 0 0

C x x x

E 0.20 - x x x

x² = 2.0408 x 10^-5

0.20 - x

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

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

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

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

3. 0.067 M KCH₃COO

Kb = 5.55 x 10^-10 [OH^-] = 6.1 x 10^-6 M pOH = 5.21 pH = 8.79 [H⁺] = 1.6 x 10^-9 M

4. 0.40 M KHCO₃

Kb = 2.3 x 10^-8 [OH^-] = 9.6 x 10^-5 M pOH = 4.02 pH = 9.98 [H⁺] = 1.0 x 10^-10 M

5. 0.60 M NH₃

Kb = 1.786 x 10^-5 [OH^-] = 3.3 x 10^-3 M pOH = 2.49 pH = 11.51 [H⁺] = 3.1 x 10^-12 M

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

H₂X ⇄ H⁺ HX^-

I 0.100 M 0 0

C - 0.0002075 0.0002075 0.0002075

E 0.09979 0.0002075 0.0002075

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

(0.09979)

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

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

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

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

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.

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

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

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

X^- + H₂O D HX + OH^-

I 0.65 M 0 0

C 0.02884 M 0.02884 M 0.02884 M

E 0.6212 M 0.02884 M 0.02884 M

(0.02884)²

Kb =

(0.6212)

Kb = 1.3 x 10^-3

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

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

Trial Volume of HCl added

#1 37.80 mL Reject

#2 35.49 mL

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

35.50 mL Average

2HCl + Ba(OH)₂ → BaCl₂ + 2H₂O

0.03550 L 3.16 g

3.16 g Ba(OH)₂ x 1 mole x 2 moles HCl

Molarity =

171.3g 1 mole Ba(OH)₂

0.03550 L

[HCl] = 1.04M

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

0.00625 L

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

No excess pH = 7.000

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

pH = 12.456

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

pH = 12.323

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

A standard solution is one of known molarity. If you make the solution from a weighed amount of solid and dilute it to a final volume in a volumetric flask it is a standard solution. If you titrate a solution to determine its concentration it is a standard solution.

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

[H+] = 6.3 x 10^-13 M

WS # 13 Review

1. List the properties of acids/bases.

Acids- conduct electricity, taste sour, change the color of indicators, neutralize bases, react with active metals like Mg to produce H₂ gas.

Bases- conduct electricity, taste bitter, change the color of indicators, neutralize acids, feel slippery.

2. Define the following:

Arhenius strong acid- completely ionizes to form H⁺

Arhenius weak base- partially ionizes to form OH^-

Bronsted strong acid- completely donates a proton to a base

Bronsted weak base- partially accepts a proton to an acid

Conjugate pair – an acid base pair that differs by one proton

Amphiprotic- a chemical species that can be an acid or base

Standard solution- a solution of known molarity

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

a) HCO₃^- Base Ka = 5.6 x 10^-11 Kb = 2.3 x 10^-8

b) H₂PO₄^- Acid Ka = 6.2 x 10^-8 Kb = 1.3 x 10^-12

c) HPO₄^2- Base Ka = 2.2 x 10^-13 Kb = 1.6 x 10^-7

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

Base OH- NaOH → Na⁺ + OH^-

Acid H+ HCl → H⁺ + Cl^-

5. Match each equation:

Acid/base complete HCl + NaOH →NaCl + HOH

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

Solubility product H⁺ + OH^- → HOH

Hydrolysis AgCl_(s) → Ag⁺ + Cl^-

Acid/Base formula H₂0 → H⁺ + OH^-

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

6. HCl and HF. Describe each acid as:

a) strong/weak b) high/low ionization c) large or small Ka

d) good/poor conductor e) strong or weak electrolyte

7. 0.2M HCl and 1.0M HF. Which is the most concentrated? Which is the strongest acid?

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

|________________________|_________________________|__

pH 0 7 14

SA WA WB SB

9. Which ions are amphiprotic?

HPO₄^-2 HCl F^- HS^- H₂S H₂O

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

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

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

13. H₂SO₃ + HS^- <====> H₂S + HSO₃^-

a) Are the reactants or products favoured?

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

14. 0.20M HCl pH = 0.70

15. 0.20M Ba(OH)₂ pH = 13.60

16. 0.20M H₂CO₃ pH = 3.53

17. 0.40M KHCO₃ pH = 9.98

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

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

20. a) For distilled water : pH = 7.00 pOH =7.00 [H+] = 1.0 x 10^-7 M [OH-] = 1.0 x 10^-7 M

b) For 1M HCl: pH = 0.0 pOH =14.0 [H+] = 1 M [OH-] = 1.0 x 10^-14 M

c) For 1M NaOH pH = 14.0 pOH =0.0 [H+] = 1.0 x 10^-14 M [OH-] = 1 M

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

Kb = 5.9 x 10^-3

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

Ka = 5 x 10^-9

24. 100 mL of 0.200M NaOH and 100 mL of 0.100 M KOH is mixed with 100.0 mLof 0.100M HCl.

Calculate the pH of the resulting solution.

pH = 12.52

25. How many grams of NaHCO₃ are required to make 100mL of .200M solution?

1.68 g

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

0.0375 L

27. In a titration 25.0mL of 0.200M H₂SO₄ is required to neutralize 10.0mL NaOH. Calculate the concentration of the base.

1.00 M

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

3.42 M

29. 2SO_2(g) + O_2(g) ⇄ 2SO_3(g)

4.00 moles of SO₂ and 5.00 moles O₂ are placed in a 2.00 L container at 200ºC and allowed to reach equilibrium. If the equilibrium concentration of O₂ is 2.00M, calculate the Keq.

Keq = 0.500

Ws # 14 Buffers

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

2.

Acid Conjugate Base Salt

HCN CN^- NaCN

H₂CO₃ HCO₃^- KHCO₃

NH₄⁺ NH₃ NH₄Cl

HF F^- NaF

CH₃COOH CH₃COO^- NaCH₃COO

H₂C₂O₄ HC₂O₄^- Na HC₂O₄

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

HCN ⇄ H⁺ + CN^-

H₂CO₃ ⇄ H⁺ + HCO₃^-

NH₃ + H₂O ⇄ NH₄⁺ + OH^-

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

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

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

Final pH

a) 2 drops of 0.10M HCl are added 8.98

b) 1 drop of 0.10M NaOH is added 9.01

c) 10 mL of 1.0 M HCl are added 2.00

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

Equation: H₂CO₃ ⇄ H⁺ + HCO₃^-

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

Indicators

1. Definition (Indicator) A weak acid whose conjugate base is a different color.

2. Equilibrium equation HInd ⇄ H⁺ + Ind^-

3. Colors for methyl orange HInd red Ind^- yellow

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

Color Relationship

pH = 2.0 red [Hind] > [Ind^-]

pH = 3.7 orange [Hind] = [Ind^-]

pH = 5.0 yellow [Hind] < [Ind^-]

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

[H⁺] increases

[HInd] increases

[Ind^-] decreases

Color Change yellow to red

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

[H⁺] decreases

[HInd] decreases

[Ind^-] increases

Color Change red to yellow

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

[Hind] = [Ind^-] Ka = [H⁺]

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

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

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

pH = 10

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

pH = 8

Ws # 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 -------> 7 bromothymol blue

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

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

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

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

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

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

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

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

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

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

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

Draw a titration curve for each of the following.

12. Adding 100 ml 1.0 M NaOH to 50 mL 1.0 M HCl 13. Adding 100 ml 1.0 M NaOH to 50 mL 1.0 M HCN

pH

Volume of base added Volume of base added

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

pH pH

Volume of base added Volume of base added

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