Bronsted Lowry Acid Conjugate Base Problems

Direction:  For exercises that require problem solving, you need to show the calculation set up leading to the answer in order to receive credit. Make sure your answers contain the correct units. YOU WILL NEED TO HAVE A TABLE OF IONIZATION CONSTANTS FOR THIS HOMEWORK.

  1. In each reaction, identify the Bronsted-Lowry acid and its conjugate base, and the Bronsted-Lowry base     and its conjugate acid.
  1. C5H5N(aq) +  H2O(l)    C5H5NH+(aq)  +  OH(aq)
  1. H2O(l) +  CO32‒(aq)    HCO3(aq)  +  OH(aq)
  1. NH3(aq) +  H2O(l)    NH4+(aq)  +  OH(aq)
  1. CH3CO2H(aq) +  H2O(l)    CH3CO2(aq)  +  H3O+(aq)
  1. Classify each substance is a weak acid or weak base. Then (a) write a balanced equation for its reaction in      water (include all physical labels) and (b) an equilibrium expression (Ka or Kb) for the reaction.
  1. HCN
  1. CH3-NH2 (methylamine)
  1. NH4+
  1. For each solution, determine the [H3O+], [OH], and pH. Show problem set ups and set the answers with          correct sig figs. Provide correct units where applicable. Assume Kw = 1.00 x 10‒14
  1. 35 M HNO3

Answers: [H3O+] = 0.35 M; [OH] = 2.9 x 10‒14 M; pH = 0.46

  1. A solution that is 0.045 M HClO4 and 0.068 M HCl

Answers: [H3O+] = 0.113 M; [OH] = 8.85 x 10‒14 M; pH = 0.947

  1. A solution that is 0.655% HNO3 by mass (i.e. in a 100-g solution, HNO3 makes up 0.655 g). The density             of the solution is 1.01 g/mL. [Hint: Consider a solution that weighs exactly 100. g (3 sig figs)]

Answers: [H3O+] = 0.105 M; [OH] = 9.52 x 10‒14; pH = 0.979

  1. Determine the [H3O+] and pH of a 0.100 M benzoic acid. Benzoic acid is a monoprotic acid with a formula      HC7H5O2 (acidic proton is indicated by bold H). Benzoic acid has a Ka = 6.5 x 10‒5 at 25 

Condensed structure

of benzoic acid

Answers:  [H3O+] = 2.5 x 10‒3 M; pH = 2.59

  1. Determine the [H3O+] and percent ionization of a 0.125 M HCN solution at 25  At this temperature,            hydrocyanic acid has Ka = 4.9 x 10‒10.

Answers:  [H3O+] = 7.8 x 10‒6 M; % ionization = 0.0063%

  1. A 0.148 M solution of a generic weak monoprotic acid (HA) has a percent ionization of 1.55% at certain          Determine the ionization constant (Ka) for the acid.

AnswerKa = 3.61 x 10‒5

  1. Find the pH of each mixture of acids. Show problem set up and explain briefly your strategy for determining     the solution [H3O+] and its pH.
  1. 150 M HNO2 and 0.085 M HNO3

Answer:  pH = 1.07

  1. 050 M CH3CO2H and 0.050 M HCN

Answer:  pH = 3.02

  1. Consider a solution of 4.8 x 10‒4 M Sr(OH)2. Calculate the [OH], [H3O+], pH, and pOH.

Answers:  [OH] = 9.6 x 10‒4 M; [H3O+] = 1.0 x 10‒11 M; pOH = 3.02; pH = 10.98

  1. What volume, in mL, of 0.855 M KOH solution is required to make 3.55 L of a solution with pH = 12.4 ?

Answer: 104 mL

  1. Caffeine (C8H10N4O2) contains nitrogens and acts a weak base with a pKb = 10.4. The molar mass of caffeine is 194.19 g/mol. Calculate the pH of a solution containing a caffeine concentration of 455 mg/L.

                                              C8H10N4O2(aq)  +  H2O(l)    HC8H10N4O2+(aq)  +  OH(aq)

 

 

 

 

 

             

Answer: pH = 7.48

  1. Determine whether each ion acts as a weak acid, weak base, or pH-neutral substance in aqueous solution.     For each ion, explain its acid or base behavior and, in the case of a weak acid or weak base, write a balanced      equation that shows how the ion acts as an acid or a base in water.
  1. ClO
  1. Co3+
  1. NO2
  1. C2O42‒
  1. Mg2+

(Question #11 continued)

  1. HSO4
  1. SO42‒
  1. BrO2
  1. Cr3+
  1. g) NH4+
  1. Determine whether each of the following substances, when dissolved in water, will produce an acidic, basic,    or pH-neutral For each case, explain its acid or base behavior. In case where the substance produces       an acidic or basic solution, write a balanced equation that shows how it acts as an acid or a base. For problems       (c) and (d), you may need to provide Ka and Kb values in your explanation.
  1. K2CO3
  1. Na2SO3
  1. CaHPO4
  1. NH4CN
  1. Calculate the equilibrium concentrations of H2SO3, HSO3, SO3‒2, and the pH of a 0.500 M H2SO3

H2SO3(aq)  +  H2O(l)    HSO3(aq)  +  H3O+(aq)          Ka1 = 1.6 x 10‒2

HSO3(aq)  +  H2O(l)    SO32‒(aq)  +  H3O+(aq)            Ka2 = 6.4 x 10‒8

Answers:  [H2SO3] =  0.418 M; [HSO3] = 0.082 M; [SO3‒2] = 6.4 x 10‒8 M; pH = 1.09

  1. If each of the salts listed below were dissolved in water to give a 0.10 M solution, which solution would have the highest pH ?  Which would have the lowest pH ?  Explain your choices by providing Ka or Kb values for
  1. Na2S c) NaH2PO4     e) CH3CO2Na
  2. Na3PO4 d) NaF f) AlCl3
  1. For each of the following reactions, predict whether the equilibrium lies predominantly to the left or to the       Explain your predictions using Ka or Kb values.
  1. NH4+(aq) +  Br(aq)    NH3(aq)  +  HBr(aq)
  1. HPO42‒(aq) +  CH3CO2(aq)    PO43‒(aq)  +  CH3CO2H(aq)
  1. [Fe(H2O)6]3+(aq) +  HCO3(aq)    [Fe(H2O)5(OH)]2+(aq)  +  H2CO3(aq)
  1. Equal molar quantities of ammonia and sodium dihydrogen phosphate (NaH2PO4) are mixed. Write a   balanced, net ionic equation for the acid-base reaction that can, in principle, occur. Use Ka and Kb values to           determine if the equilibrium lies to the right or the left.
  1. Identify the Lewis acid and Lewis base from among the reactants in each equation, then briefly explain your
  1. Ag+(aq) +  2 NH3(aq)    [Ag(NH3)2]+(aq)
  1. F(aq) +  BF3(aq)    BF4(aq)

(Question 17 continued)

  1. SO2(aq) +  H2O(l)    H2SO3(aq)
  1. Ni2+(aq) +  6 H2O(l)    [Ni(H2O)6]+2(aq)
  1. Cr(OH)3(s) +  3 H3O+(aq)    Cr3+(aq)  +  6 H2O(l)
  1. Zn(OH)2(s) +  2 OH(aq)    [Zn(OH)4]2–(aq
  1. Of each pair of acids, circle the stronger acid and briefly explain your choice. [Hint: use concepts such as bond            energy, electronegativity, bond polarity, and inductive effect in your explanation.]
  1. HF or  HBr
  1. H2S or  H2Se
  1. HClO or  HIO
  1. H2O or  HF

(Question 18 continued)

  1. HClO3 or  HClO
  1. CH3CO2H or  CF3CO2H
  1. HOCN or  HCN
  1. Calculate the pH of the solution that results from mixing 25.0 mL of 0.14 M formic acid (HCO2H) and 50.0     mL of 0.070 M NaOH. [Note: Formic acid is a weak monoprotic acid that reacts with a base to produce             formate ion as follows:  HCO2H(aq)  +  OH(aq)    HCO2(aq) +  H2O(l); the Ka of HCO2H is 1.8 x 10‒4 (at

25 C)]

Answer: pH = 8.21

  1. A solution of hydrofluoric acid (HF) is prepared by dissolving enough HF in water to produce a solution with pH = 2.30. The Ka of hydrofluoric acid is 7.2 x 10‒4.
  1. Calculate the equilibrium concentrations of HF and F.

Answers: [HF]equil = 0.035 M; [F] = 0.0050 M

  1. Calculate the mass, in g, of HF initially dissolved per liter solution.

Answers:  0.80 g per 1 L

  1. Calculate the pH for a 0.025 M solution of sodium oxalate (Na2C2O4). Oxalic acid (H2C2O4) is a diprotic acid   with the following dissociation constants:

H2C2O4  +  H2O    HC2O4  +  H3O+       Ka1 = 5.9 x 10‒2

HC2O4  +  H2O    C2O42‒  +  H3O+        Ka2 = 6.4 x 10‒5

Answer:  pH = 8.30

  1. The acid form of pyridine, which contains a single acidic proton, can be synthesized with a different     substituent X, where X is an atom such as Cl or a group such as CH3. The table below gives the Ka values for a        variety of substituted pyridines.

Atom X or Group X                              Ka of substituted pyridines

NO2                                                         5.9 x 10‒2

Cl                                                            1.5 x 10‒4

H                                                             6.8 x 10‒6

                                           CH3                                   1.0 x 10‒6

  1. Suppose each substituted pyridine is dissolved in sufficient water to give a 0.050 M solution. Which                  solution would have the highest pH ?  The lowest pH ?
  1. Which of the substituted pyridines is the strongest Bronsted base? Which is the weakest Bronsted base?
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