AP Chapter 15 & 16: Acid-Base Equilibria 15 and16.pdf · AP Chapter 15 & 16: Acid-Base Equilibria 3 •Warm-ups and problems will be collected before you take the test. •Read Chapter

AP Chapter 15 & 16: Acid-Base Equilibria Name_______________________

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AP Chapter 15 & 16: Acid-Base Equilibria 2

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•Warm-ups and problems will be collected before you take the test.

•Read Chapter 15.1-15.11: Acids and Bases, and read Chapter 16.1-16.5 & 16.9: Acid-Base Equilibria

Answer the following problems in the space provided. For problems involving an equation, carry out the

following steps: 1. Write the equation. 2. Substitute numbers and units. 3. Show the final answer with units.

There is no credit without showing work.

Bronsted Acids and Bases

1. Write an equation for the reaction of the following compounds with water, and identify the acid-base

conjugate pairs.

(a) hydrocyanic acid

(b) methanamine

2. Write the formula for the conjugate acid of each of the following bases:

a. HS- b. HCO3- c. CO3

-2

d. H2PO4- e. HPO4

2- f. PO43-

pH: A Measure of Acidity

3. Write an equation for the ion-product constant for water, and state its value at 25oC.

4. The ion-product constant for water at 25oC is 1.0E-14 and at 40oC is 3.8E-14.

(a) Is the autoionization of water exothermic or endothermic? Explain.

(b) Calculate the pH of water at 40oC.

5. Calculate the pH of each of the following solutions:

a. 2.8 x 10-4 M Ba(OH)2 b. 5.2 x 10-4 M HNO3.

Strength of Acids and Bases

6. Classify each of the following species as a strong or weak acid or strong or weak base:

a. LiOH b. HCN c. H2O d. HClO4

e. CH3CH2NH2 f. HBr g. HF h. HCOOH


7. Which of the following statements are true regarding a 1.0 M solution of a strong acid HA?

a. [A-] > [H+]

b. The pH is 0.00.

c. [H+] = l.0 M

d. [HA] = 1.0 M

8. Which of the following statements are true regarding a 1.0 M solution of a weak acid HA?

a. [A-] > [HA]

b. The pH is 0.00.

c. [HA] > [H+]

d. [H+] = [A-]

9. Predict whether the following reaction will proceed from left to right to a significant extent. Explain.

F-(aq) + H2O(l) HF(aq) + OH-(aq)

10. Determine whether the equilibrium constant for the following reaction is greater or less than 1.0. Explain.

CH3COOH(aq) + NO2-(aq) CH3COO-(aq) + HNO2(aq)

Weak Acids and Acid Ionization Constants

11. A 0.0560 g quantity of acetic acid is dissolved in enough water to make 50.0 mL of solution. Calculate the

concentrations of H+, CH3COO-, and CH3COOH at equilibrium. (Ka for acetic acid = 1.8 x 10-5.)

12. What is the original molarity of a solution of formic acid (HCOOH) whose pH is 3.26 at equilibrium?


13. Calculate the percent ionization of hydrofluoric acid at the following concentrations. Explain the trend.

a. 0.60 M b. 0.046 M

14. A solution of acetic acid with a pH of 2.73 is compared to a solution of hydrochloric acid with a pH of 2.73.

a. Predict which solution is more concentrated. Explain.

b. Calculate the molar concentration of each solution.

c. Which solution has a greater concentration of undissociated acid? Explain.

15. A solution of 0.13 M acetic acid is compared to a solution of 0.13 M hydrochloric acid.

a. Predict which solution is more acidic (lower pH). Explain.

b. Calculate the pH of each solution.

Weak Bases and Base Ionization Constants

16. The pH of a 0.30 M solution of a weak base is 10.66. What is the Kb of the base?

17. What is the original molarity of a solution of ammonia whose pH is 11.22?


18. A solution of methylamine (CH3NH2) has a pH of 10.64. How many grams of methylamine are there in 100.0

mL of the solution?

Acid-Base Conjugate Pairs

19. Write the equation that relates Ka and Kb for a conjugate pair. The Ka of carbonic acid is 4.2E-7. Write the

formula of its conjugate base and calculate Kb for this base.

Polyprotic Acids

20. Calculate and compare the pH of a 1.0 M HC1 solution with that of a 1.0 M H2SO4 solution.

21. Calculate the pH of a 0.100 M Na2CO3 solution.

Molecular Structure and Acid Strength

22. Why is HF a weak acid and HCl a strong acid, even though F is more electronegative?

23. List the strength of the following acids from least to greatest. Explain.

a. H2SO4, H2SeO4, and H2TeO4 b. H3PO4, H3AsO4, and H3PO3


24. Which of the following is the stronger base: NF3 or NH3? Hint: draw Lewis structures to determine the more

polar lone pair.

Acid-Base Properties of Salt Solutions

25. Define “salt hydrolysis.” Write an equation for the hydrolysis of sodium hypochlorite.

26. State whether the following salts are acidic, basic, or neutral. For those that are not neutral write a balance

equation showing why they are acidic or basic.

a. NaCN

b. CH3NH3Cl

c. NaHCO3

d. NaHSO4

e. NH4ClO3

f. NaBr

27. Calculate the pH of a 0.42 M NH4C1 solution.

28. Predict whether a solution containing the salt K2HPO4 will be acidic, neutral, or basic.

29. Calculate the pH of a 0.20 M ammonium acetate (CH3COONH4) solution. Hint: determine Kb and Ka of the

cation and anion respectively, and think!

30. Use appropriate Ka and/or Kb values to calculate the equilibrium constant for the following reaction. Hint: find

chemical equations that “add up” to this net equation.

CH3COOH(aq) + NO2-(aq) CH3COO-(aq) + HNO2(aq)


Acidic and Basic Oxides

31. Arrange the oxides in each of the following groups in order of increasing basicity. Explain.

a. K2O, A12O3, BaO b. CrO3, CrO, Cr2O3.

32. Write chemical equations for reactions of the following compounds.

(a) Calcium oxide solid and water.

(b) Sulfur dioxide gas and water.

(c) Calcium oxide solid and sulfur dioxide gas.

pH and Solubility

33. The pH of a saturated solution of a metal hydroxide, MOH, is 9.68. Calculate the Ksp for the compound.

34. Mixed Fe3+ and Zn2+ ions in aqueous solution can be separated by selectively precipitating their hydroxides.

Find the approximate pH range suitable for the separation of Fe3+ and Zn2+ by precipitation of Fe(OH)3 from a

solution that is initially 0.010 M in both Fe3+ and Zn2+.

35. Which of the following insoluble salts are more soluble in acid solution (H+) than in pure water? For the salts

whose solubility is affected by acid, write a chemical equation showing why its solubility is increased.

a. CuI c. Zn(OH)2

b. BaC2O4 d. Ca3(PO4)2


The Common Ion Effect

36. Write the net ionic equation for the hydrolysis of sodium nitrite.

Use Le Chatelier's principle to predict the effect of the following changes on the extent of hydrolysis of sodium

nitrite solution. Explain each by also writing a chemical equation(s).

a. HC1 is added

b. NaOH is added

c. NaCl is added

d. the solution is diluted.

37. Describe the effect on pH (increase, decrease, or no change) that results from each of the following:

a. adding potassium acetate to an acetic acid solution

b. adding ammonium nitrate to an ammonia solution

c. adding sodium formate (HCOONa) to a formic acid (HCOOH) solution

d. adding potassium chloride to a hydrochloric acid solution

e. adding barium iodide to a hydroiodic acid solution

38. Write both the regular and log forms of the Henderson-Hasselbalch equation, and state its assumptions.

39. Determine the pH of

a. a 0.20 M NH3 solution b. a solution that is 0.20 M NH3 and 0.30 M NH4C1


40. Determine the pH of

a. a 0.150 M propanoic acid (CH3CH2COOH) solution (Ka = 1.3E-5).

b. a solution that is 0.150 M propanoic acid and 0.30 M sodium propanoate.

Buffers

41. State whether each system is an acidic buffer, basic buffer, or not a buffer.

a. KC1/HC1

b. NH3/NH4NO3

c. Na2HPO4/NaH2PO4

d. KNO2/HNO2

e. KHSO4/H2SO4

f. HCOOK/HCOOH

42. Calculate the pH of the following two buffer solutions. Which is the more effective buffer, that is, has a

greater buffer capacity? Why?

a. 2.0 M CH3COONa/2.0 M CH3COOH

b. 0.20 M CH3COONa/0.20 M CH3COOH

c. Write chemical equations showing how this buffer system neutralizes additions of H+ and OH-.

43. The pH of blood plasma is 7.40. Assuming the principal buffer system is HCO3-/H2CO3, calculate the ratio

[HCO3-]/[H2CO3]. Is this buffer more effective against an added acid or an added base? Explain


44. A student is asked to prepare a buffer solution at pH = 8.60, using one of the following weak acids and its

conjugate: HA (Ka = 2.7 x 10-3), HB (KA = 4.4 x 10-6), HC (KA = 2.6 x 10-9). Which acid should she choose?

For the acid/conjugate selected, explain how to prepare the buffer.

Acid-Base Titrations

45. Given 50 mL of 0.10 M HCl and 50 mL of 0.10 M acetic acid, will the amount of 0.10 M NaOH required to

neutralize each solution be the same, more, or less? Explain.

46. Will the pH at the equivalence point of 50 mL 0.10 M HCl be the same, more, or less as the pH at the

equivalence point for 50 mL of 0.10 M acetic acid? Explain.

47. A 5.00-g quantity of a diprotic acid was dissolved in water and made up to 250. mL. Calculate the molar mass

of the acid if 25.0 mL of this solution required 11.1 mL of 1.00 M KOH for neutralization. Assume that both

protons of the acid were titrated.

48. In a titration experiment, 20.4 mL of 0.883 M HCOOH neutralize 19.3 mL of Ba(OH)2. What is the

concentration of the Ba(OH)2 solution?

49. Calculate the pH for the titration of 0.10 M HCOOH versus 0.10 M NaOH at the following points:

a. at the half titration point

b. at the equivalence point


50. An 80. mL sample of 0.24 M ethanamine (CH3CH2NH2) is put in a flask and titrated with 0.18 M HCl. (Kb of

ethanamine is 5.6E-4.)

a. What is the original pH in the flask?

b. What is the pH in the flask after 25. mL of the acid is added?

c. What is the pH at the “half-titration” point?

d. What is the pH at the equivalence point?

e. What is the pH after the addition of 127 mL of HCl?

f. Sketch the titration curve and indicate each point a-e on the plot.


51. Sketch the titration curve for a weak diprotic acid vs. NaOH, and show how to determine Ka1 and Ka2

graphically.

52. A 52 mL solution of acetic acid with a pH of 2.73 is compared to a 52 mL solution of hydrochloric acid with a

pH of 2.73.

a. Predict which solution requires more 0.12 M NaOH to neutralize it. Explain.

b. Calculate the volume of the 0.12 M NaOH solution needed to neutralize each.

53. Which chemical species are in significant concentration when:

a. A hydrochloric acid solution is titrated to the half equivalence point with a solution of NaOH.

b. An acetic acid solution is titrated to the half equivalence point with a solution of NaOH.

Acid-Base Indicators

54. Two drops of indicator HIn (Ka = 1.0E-9), where HIn is yellow and In- is blue, are placed in100 mL of 0.10 M

HCl.

a. What is the color initially?

b. The solution is titrated with 0.10 M NaOH. At what pH will the color change of the indicator begin?

c. What color will the solution be after 200. mL of NaOH solution is added?

55. What indicator dye(s) listed in Table 16.1 of your text is appropriate for acid-base titrations where the

equivalence point is at pH 5.0?


56. A student carried out an acid-base titration by adding NaOH solution from a buret to an Erlenmeyer flask

containing HC1 solution and using phenolphthalein as indicator. At the equivalence point, she observed a faint

reddish-pink color. However, after a few minutes, the solution gradually turned colorless. What happened?

57. Draw Lewis structures for the following molecules, state whether each is symmetrical or nonsymmetrical, state

whether each is polar or nonpolar, state the geometric shape of the molecule, and state the hybridization of the

central atom.

a. BrCl3 b. IF5

58. Explain the following properties of solid potassium and lithium:

a. Both solids are malleable.

b. Both solids conduct electricity.

c. Lithium has a higher melting point than potassium.

d. The first ionization energy of lithium is greater than that of potassium.

59. Explain the following properties of solid CsBr and MgBr2:

a. Both solids are brittle.

b. MgBr2 has a higher melting point than CsBr.

c. Both solids do not conduct electricity, but when dissolved in water, both solutions do conduct electricity.


Titration Summary

Titration problems are difficult unless you think of them in a systematic way. Two problems are the most common

(although any combination is possible):

weak acid in a flask titrated with strong base in a buret

weak base in a flask titrated with strong acid in a buret

In both cases, knowing special regions of the titration curve is key to solving problems. As an example, the titration

curve for a weak acid in a flask (25 mL of 0.20 M HA) titrated with a strong base (0.35 M NaOH) in a buret is

shown in Figure 1.

pH

mL base added

Figure 1. Weak acid – strong base titration curve.

Key points and regions are shown on the curve:

a. pH in the flask before addition of strong base.

b. buffer region (H-H applies)

c. pH at the “half-titration” point where pH = pKa

d. pH at the equivalence point (pH > 7.0 in this case)

e. region of excess strong base

For regions b-e, do this step first:

Determining pH at any point along the titration is a two-step process. First, one must do stoichiometry using the

titration equation:

Titration Equation: HA + OH- A- + H2O (one-way arrow)

Since OH- is a strong base, this reaction is quantitative, that is, every mole of OH- added subtracts a mole of HA

and creates a mole of A-. Thus it is easy to determine the new HA and A- concentrations using simple

stoichiometry. Do stoichiometry in moles, and then convert to concentrations using the new total volume.

Point a.

Before addition of any base (OH-), the titration equation above hasn’t happened yet, so the pH is governed solely by

hydrolysis of the weak acid:

HA H+ + A-

Knowing Ka, you can set up a table and solve for H+ and pH.

a

b

c

d

e


Region b:

Before the equivalence point (region b), a buffer is created because the OH- has converted part of the HA into A-,

so they are both present. Do stoichiometry using the titration equation above to determine the new HA and A-

concentrations. Plug these into Henderson-Hasselbalch* (H-H) using Ka of HA to determine the H+ and the pH. Be

sure to keep track of the new volume when determining concentrations.

*Remember that H-H isn’t so much the form of the equation, (normal vs. log form) as it is the simplifying

approximation that hydrolysis of both the weak acid (HA) and conjugate base (A-) can be ignored if both species

are present. In other words, you plug the HA and A- concentrations determined using stoichiometry directly into Ka.

Point c:

Point c is the half-titration point, where exactly half as many moles of strong base has been added as moles of weak

acid originally present. Thus the concentration of conjugated base (A-) created by the titration equation equals the

concentration of weak acid (HA) remaining. When plugged into H-H, these two species cancel and pH = pKa.

Point d:

Point d is the equivalence point where all of the HA has been converted into A- by the titration equation above.

Thus the moles of A- at the equivalence point is equal to initial moles of HA. Use MaVa = MbVb to determine the

volume of OH- solution added. Then add Va + Vb to get the new total volume, and then calculate the concentration

of A-. The pH is then determined solely by the hydrolysis of A- according to the equation:

A- + H2O HA + OH-

Set up a table using this equation and plug into Kb for A-. Solve for OH- and then pH.

Region e:

After the equivalence point (region e), stoichiometry of the titration equation will show an excess amount of OH-.

pH is determined by the excess moles of OH- added after the equivalence point. Again keep track of the new

volume, determine [OH-] and then pH. No table is needed.

Diprotic Acids

Just one last hint, if there is a diprotic acid, refer to Figure 2 to determine Ka1 and Ka2.

Figure 2. pH during titration of a diprotic acid with sodium hydroxide

where

A = the volume of NaOH needed to react with both of the acid hydrogens

B = volume of NaOH needed to react with one of the acid hydrogens

C = the volume of NaOH used when all of the first and half of the second hydrogens are neutralized

D = the volume of NaOH needed to neutralize half of the first acid hydrogen

E = the pH when half of the first hydrogen is neutralized, or pKal

F = pH when half of the second hydrogen is neutralized, or pKa2

Documents

AP Chapter 15 & 16: Acid-Base Equilibria 15 and16.pdf · AP Chapter 15 & 16: Acid-Base Equilibria 3 •Warm-ups and problems will be collected before you take the test. •Read Chapter