CHM 152-Exam 3 Review Jeopardy Template

Acids and Bases

Salts

Buffers

Titrations

100

Give the formula for the conjugate base of each of the following Bronsted-Lowry acids:

a. HSO₄^-

b. H₂SO₃

c. H₂PO₄^-

d. NH₄⁺

a. SO₄^2-

b. HSO₃^-

c. HPO₄^2-

d. NH₃

100

Using the Kb values given, calculate the values of Ka for each of the following ions:

a. Propylammonium ion (C₃H₇NH₃⁺): Kb = 5.1*10^-4

b. Hydroxylammonium ion (NH₃OH⁺): Kb =9.1*10^-9

c. Anilinium ion (C₆H₅NH₃⁺): Kb = 4.3*10^-10

d. Pyridinium ion (C₅H₅NH⁺): Kb = 1.8*10^-9

a. 1.96*10^-11

b. 1.1*10^-6

c. 2.3*10^-5

d. 5.6*10^-6

100

Using the Kb values given, calculate [OH-] and the pH for each of the following solutions:

a. 0.24 M methylamine(CH₃NH₂); Kb = 3.7*10^-4

b. 0.040 M pyridine(C₅H₅N); Kb = 1.8*10^-9

c. 0.075 M hydroxylamine(H₃NO); Kb = 9.1*10^-9

a. 11.97 pH; [OH-]= 9.24*10^-3

b. 8.93 pH; [OH-]= 8.48*10^-6

c. 9.42 pH; [OH-]= 2.61*10^-5

100

Calculate the pH of a solution that is 0.25 M in HF and 0.10 M in NaF.

pH = 3.06

100

A 60.0 mL sample of a monoprotic acid is titrated with 0.150 M NaOH. If 20.0 mL of base is required to reach the equivalence point, what is the concentration of the acid?

0.0500 M

200

For each of the following reactions, identify the Bronsted-Lowry acids and bases and the conjugate acid base pairs:

a. CH₃CO₂H(aq) + NH₃(aq) ⇄ NH₄+(aq) + CH₃CO₂^-(aq)

b. CO₃^2-(aq) + H₃O⁺(aq) ⇄ H₂O(l) + HCO₃^-(aq)

a. acid, base, conj. acid, conj. base

b. base, acid, conj. base, conj. acid

200

Classify each of the following ions according to whether they react with water to give a neutral, acidic, or basic solution:

a. F^-

b. Br^-

c. NH₄⁺

d. K(H₂O)₆⁺

e. SO₃^2-

f. Cr(H₂O)₆³⁺

a. basic

b. neutral

c. acidic

d. neutral

e. basic

f. acidic

200

Calculate the [H3O⁺] and pH of each polyprotic acid solution.

a. 0.125 M H₂CO₃; Ka₁ = 4.3*10^-7, Ka₂ = 5.6*10^-11

b. 0.125 M H₃C₆H₅O₇; Ka₁= 7.4*10^-4, Ka₂ = 1.7*10^-5, Ka₃ = 4.0*10^-7

a. [H3O⁺]_eq= 0.125 M; pH=0.903

b. [H3O⁺]_eq = 0.125 M; pH=0.903

200

Calculate the pH of 0.500 L of a 0.200 M NH₄Cl-0.200 M NH₃ buffer.

pH = 9.25

200

Consider the titration of 60.0 mL of 0.150 M HNO₃ with 0.450 M NaOH.

a. How many moles of HNO₃ are present at the start of the titration?

b. How many mL of NaOH are required to reach the equivalence point?

c. What is the pH at the equivalence point?

a. 9*10^-3 mol

b. 20.0 mL

c. pH = 7.00

300

For each of the following reactions, identify the Bronsted-Lowry acids and bases and the conjugate acid base pairs:

a. HSO₃^-(aq) + H₂O(l) ⇄ H₃O⁺(aq) + SO₃^2-(aq)

b. HSO₃-(aq) + H₂O(l) ⇄ H₂SO₃(aq) + OH^-(aq)

a. acid, base, conj. acid, conj. base

b, base, acid, conj. base, conj. acid

300

Calculate the pH in a solution that is 0.10 M in acetic acid (CH₂CO₂H; Ka = 1.8*10^-5) and 0.10 M benzoic acid (C₆H₅CO₂H; Ka = 6.5*10^-5).

pH = 4.74

300

Oxalic acid (H₂C₂O₄) is a diprotic acid that occurs in plants such as rhubarb and spinach. Calculate the pH and the concentration of C₂O₄^2- ions in 0.20 M H₂C₂O₄ (Ka₁ = 5.9*10^-2; Ka₂ = 6.4*10^-5).

[C₂O₄^2-] = 6.4*10^-5 M

pH = 0.699

300

Use the Henderson-Hasselbalch equation to calculate the pH of a buffer solutions that is 0.25 M in formic acid (HCO₂H) and 0.50 M in sodium formate (HCO₂Na).

pH = 4.04

300

A 50.0 mL sample of 0.120 M HBr is titrated with 0.240 M NaOH. Calculate the pH after the addition of the following volumes of base.

a. 0.0 mL

b. 20.0 mL

c. 24.9 mL

d. 25.0 mL

e. 25.1 mL

f. 40.0 mL

a. 0.92

b. 1.77

c. 3.5

d. 7.00

e. 11.5

f. 12.60

400

Calculate the pH of 1.5 M HNO₂ dissociated in water. (Ka = 4.5*10^-4)

pH = 1.59

400

Determine whether each anion is basic or neutral.

a. C₇H₅O^2-

b. I^-

c. NO₃^-

a. basic

b. neutral

c. neutral

400

Amphetamine (C₉H₁₃N) is a weak base with a pKa of 4.2. Calculate the pH of a solution containing an amphetamine concentration of 225 mg/L.

pH = 7.71

400

Use the Henderson-Hasselbalch equation to calculate the ratio of H₂CO₃ to HCO₃^- in blood having a pH of 7.40.

0.093 mol base per 1 mol acid

400

Consider the titration of 40.0 mL of 0.250 M HF with 0.200 M NaOH. How many milliliters of base are required to reach the equivalence point?

50.0 mL

500

Lactic acid (C₃H₅OH), which occurs in sour milk and foods such as sauerkraut, is a weak monoprotic acid. The pH of a 0.10 M solution of lactic acid is 2.43. What are the values of Ka and pKa for lactic acid?

Ka = 1.4 *10^-4and pKa= 3.84

500

Find the pH of each mixture of acids.

a. 0.075 M in HNO₃ and 0.175 M in HC₇H₅O₂

b. 0.020 M in HBr and 0.015 M in HClO₄

c. 0.095 M in HF and 0.225 M in HC₆H₅O

d. 0.100 M in formic acid and 0.050 M in hypochlorous acid

a. 1.42 pH

b. 3.76 pH

c. 1.31 pH

d. 1.30 pH

500

Morphine (C₁₇H₁₉NO₃), a narcotic used in painkillers, is a weak organic base. If the pH of a 7.0*10^-4M solution of morphine is 9.5, what are the values of Kb and pKb?

Kb = 1*10^-6; pKb = 5.8

500

What mass of ammonium chloride should you add to 2.55 L of a 0.155 M NH₃ solution to obtain a buffer with a pH of 9.55? (Assume no volume change.)

1.66 g

500

Consider the titration of 50.0 mL of a 0.100 M solution of the protonated form of the amino acid alanine (H₂A⁺; Ka₁ = 4.6*10^-3, Ka₂ = 2.0*10^-10) with 0.100 M NaOH. Calculate the pH after the addition of each of the following volumes of base:

a. 10.0 mL

b. 25.0 mL

c. 50.0 mL

d. 75.0 mL

e. 100.0 mL

a. pH = 1.74

b. pH = 2.34

c. pH = 6.02

d. pH = 9.70

e. pH = 11.11