Download Review for Exam 3 Chem 1721/1821

Review for Exam 3
Chem 1721/1821
The following are the major concepts with which you should be well acquainted from Chapters 15, 16, 17, and 18. Although
this exam begins formally with acidic and basic nature of salt solutions (chapter 16, section 4), the fundamental knowledge of
aqueous solutions, pH, Ka, Kb, and acid and base ionization equilibria in solution will still be used, and therefore I have sitll
included those points here.
Chapters 15 and 16.1 - 16.4: Acids and Bases
Fundamentals
• Bronsted-Lowry definitions of acids and bases as H+ donors and acceptors
• Acid and base ionization (dissociation) equations including identification of conjugate pairs
• Ka and Kb expressions
• relationship of acid strength to: Ka value, [H3O+], [HA]eq, equilibrium position, strength of conjugate base
• relationship of base strength to: Kb value, [OH–], [B]eq, equilibrium position, strength of conjugate acid
• relative [H3O+] and [OH–] in acidic, basic, and neutral solutions
• identification of strong acids and bases (only a small number so KNOW THEM!!)
Water
•
•
•
Water is amphoteric (or amphoprotic). What does that mean?
Auto-ionization of water, and water auto-ionization constant Kw
Kw = 1.0 x 10–14 at 25°C; Kw = [H3O+][OH–]
pH (and pOH, pKa, pKb) scale
• logarithmic scale
• pH = –log[H3O+] and [H3O+] = 10–pH; pOH = –log[OH–] and [OH–] = 10–pOH
• pKa = –logKa and Ka = 10–pKa; pKb = –logKb and Kb = 10–pKb
• pH + pOH = pKw
• Where do acidic, basic, and neutral solutions fall on the pH scale?
Polyprotic acids
• dissociate in a stepwise manner
• there is a dissociation equation and Ka value corresponding to each step
• acids are progressively weaker (smaller Ka) in each successive dissociation step
Salts
•
•
•
•
ionic compounds dissociate into constituent ions
ions can produce acidic, basic or neutral soltuions
ionization equations corresponding to acid or base behavior of cations or anions
relationship between Ka and Kb for a conjugate acid/base pair
Structure/Strength relationships for acids
Lewis acid/base definitions
Calculations based on these concepts include:
• pH of strong acid and strong base solutions
• pH of weak acid and weak base solutions: this includes identifying the major species in solution, identifying the
major source of H3O+ or OH— in solution and setting up an equilibrium calculation
• pH of mixtures of weak acids or weak bases
• pH of polyprotic acid solutions
• pH of salt solutions
• Calculation of % dissociation or % ionization or weak acids and weak bases
• Determination of Ka or Kb using pH or % ionization values
Chapters 16 (16.5 - 16.7) and 17: Applications of Aqueous Equilibria; Solubility Equilibria
Common Ion Effect and Buffer Solutions
• understand how the presence of a common ion effects the solution with respect to equilibrium position, % ionization
and pH
• buffer solutions are common ion solutions – usually composed of a conjugate acid/base pair
• buffer solutions resist changes in pH when small amounts of strong acid (H+) or strong base (OH–) are added
• upon addition of strong acid or strong base a neutralization reaction occurs resulting in changing [HA] and [A–]
• Henderson-Hasselbalch equation for calculation of pH of buffer solutions
• effective buffering occurs when pH of solution = pKa + 1
• other terms to understand: buffer capacity, buffer failure
Titrations
• 3 classes of titrations: strong acid + strong base, weak acid + strong base, and weak base + strong acid
• calculation of pH before (including half-way point), at, and beyond stoichiometric point
• characteristic profiles of pH or titration curves
• indicators – chosen appropriately based on relationship of pKa of indicator to pH at stoichiometric point
Solubility Equilibria
• ionic solid in equilibrium with constituent ions in solution (saturated solution); equilibrium constant is the solubility
product constant, Ksp
• solubility quantified in mol/L or g/L
• relative solubilities of ionic compounds – when can you compare Ksp’s directly and when is a calculation required?
• presence of a common ion, pH of solution, and complex ion formation may effect a compound’s solubility
• precipitation of an ionic compound occurs from a supersaturated solution with Q (ion product) > Ksp
• dissolution of solid occurs in an unsaturated solution with Q < Ksp
Calculations based on these concepts include:
• pH of common ion and buffer solutions
• ΔpH of buffer solution after the addition of strong acid or strong base
• pH of solution at any given point during an acid + base titration
• pH corresponding to visible indicator color change
• calculation of Ksp given solubility or ion concentration data
• calculation of molar solubility or g/L solubility of ionic compounds in water or a common ion solution
• calculation of ion product (Q) for prediction of precipitation
• calculation of mass or concentration of ion required to initiate precipitation
Chapter 18: Thermodynamics and Equilibrium
* you should review the fundamental concepts of thermochemistry that you learned in Chapter 6 *
1st Law of Thermodynamics: the energy of the universe is constant
Entropy and the 2nd Law of Thermodynamics
• the entropy of the universe is increasing; ΔSuniv = ΔSsys + ΔSsurrr
• entropy is the thermodynamic quantity related to randomness and disorder
• qualitative prediction of the sign of ΔS for physical or chemical change
• relationship between spontaneity and temperature given by: ΔS = q/T or ΔS = –ΔH/T
• ΔS can be calculated using standard molar entropy data: ΔS° = Σn•S°prod – Σn•S°react
Gibbs Free Energy
• thermodynamic property of the system that is related to spontaneity
• for a spontaneous process ΔG< 0
• calculation of ΔG: ΔG = ΔH – TΔS; ΔG° = Σn•ΔGf°prod – Σn•ΔGf°react; Hess’s Law approach
• ΔG is dependent on reaction composition: ΔG = ΔG° + RTlnQ where Q = reaction quotient
• ΔG is related to the equilibrium position of a process: at equilibrium ΔG = 0
• so . . . ΔG is related to the equilibrium constant: ΔG° = –RTlnK
A process will proceed spontaneously in the direction that maximizes the entropy of the universe and minimizes the free
energy of the system; this is where the process reaches equilibrium.
Some Review Problems, Chapters 15, 16, 17, and 18 (See App. 4 & 5 as necessary for Ka, Kb, Ksp, ΔH° f, ΔS°, ΔG° f data)
1.
Write the balanced equation for the reaction of CN— with water. Identify all species as acid, base, conjugate acid, or
conjugate base.
2.
For each of the following pairs, give the stronger acid: H2SO3 or H2SeO3, HIO4 or HIO3, H2S or H2O
3.
Phthalic acid, H2C8H4O4 is a diprotic acid used in the synthesis of phenolphthalein indicator. Ka1 = 1.2 x 10—3 and Ka2 =
3.9 x 10—6. Calculate [H3O+] and [C8H4O42–] in a 0.015 M H2C8H4O4 (aq).
4.
For each of the following salts indicate whether the aqueous solution will be acidic, basic or neutral:
a. Fe(NO3)3,
b. Na2CO3
c. Ca(CN)2
d. NH4ClO4
5.
a.
b.
6.
Calculate the concentration of pyridine (C5H5N) in and the pH of 0.15 M pyridinium bromide (C5H5NHBr).
7.
Calculate the percent ionization of each of the following solutions:
a.
0.80 M HF
b.
a mixture of 0.80 M HF and 0.10 M HCl.
8.
Calculate the pH of a sol’n that is 0.10 M CH3NH2 (methylamine) and 0.15 M CH3NH3Cl (methylammonium chloride).
9.
Calculate the pH of a buffer solution that is 0.10 M NH3 and 0.10 M NH4+. Calculate the change in pH that occurs when
12.0 mL of 0.20 M HCl is added to 125.0 mL of this buffer.
10.
Calculate the pH of a buffer solution that is 0.10 M monochloroacetic acid and 0.15 M sodium chloroacetate.
11.
How many grams of sodium acetate must be added to 2.0 L of 0.10 M acetic acid to give a solution with a pH of 5.00?
Ignore the volume change due to the addition of sodium acetate.
12.
Calculate the pH of the solution that forms when 20.0 mL of 0.10 M NaOH is added to 25.0 mL of 0.10 M HCl.
13.
A 1.24 g sample of benzoic acid (C6H5COOH) is dissolved in water resulting in 50.0 mL of solution. This solution is
titrated with 0.180 M NaOH. Calculate the pH of the solution (a) after the addition of 35.0 mL of NaOH, (b) at the
stoichometric point, and (c) after the add’n of 60.0 mL of NaOH.
14.
32.0 mL of 0.087 M ethylamine is titrated with 0.15 M HCl. Calculate the pH of the solution (a) after the add’n of 15.0
mL of HCl, (b) at the stoichiometric pt, and (c) after the add’n of 22.5 mL of HCl.
15.
The solubility of silver bromate, AgBrO3, in water is 0.0072 g/L. Calculate Ksp.
16.
Celestite (strontium sulfate) is an important mineral of Sr. Calculate the molar solubility of SrSO4 (Ksp = 2.5 x 10—7).
17.
The solubility of magnesium fluoride, MgF2, in water is 0.0076 g/L. Calculate the solubility (in g/L) of magnesium
fluoride in 0.020 M sodium fluoride.
18.
Lead (II) chromate, PbCrO4, is used as a yellow paint pigment. A solution is prepared that is 5.0 x 10—4 M in lead ion
and 5.0 x 10—5 M in chromate ion. Do you expect precipitation of lead chromate from this solution?
19.
How many moles of calcium chloride can be added to 1.5 L of 0.020 M potassium sulfate before a precipitate is
expected? Assume that the volume of the solution is not changed significantly by the addition of the calcium chloride.
20.
Calculate the pH of the following solutions:
a. 0.75 M C2H5NH3Br
b. 1.50 M Ca(NO2)2,
21.
Calculate the value of Kb for NO2—.
Calculate Ka for C5H5NH+ (pyridinium ion).
c. 0.88 M CoCl2 (for Co(H2O)62+ Ka = 1.3 x 10—9)
What mass (in g) of sodium chloride can be added to 785 mL of 0.0015 M silver nitrate before a precipitate forms?
22.
Sorbic acid (HC6H7O2, pKa = 4.77) is widely used in the food industry as a preservative. For example, its potassium salt
(potassium sorbate) is added to cheese to inhibit the formation of mold. Calculate the pH of 0.37 M KC6H7O2 (aq).
23.
Calculate [H3O+] in the following solutions:
a.
0.035 M HCl and 0.075 M HOCl
b. 0.100 M NaNO2 and 0.0550 M HNO2
24.
You wish to prepare a buffer solution with pH = 9.45. How many grams of (NH4)2SO4 would you add to 425 mL of
0.258 M NH3 to do this?
25.
25.0 mL of 1.25 M HNO3 is titrated with 0.95 M KOH. Calculate the pH of the solution (a) after the add’n of 20.0 mL
of KOH, (b) at the stoichiometric pt., (c) after the add’n of 40.0 mL of KOH.
26.
Write the solubility product expression for each of the following salts:
a, CrF3
b. Cd3(PO4)2
c. SrF2
27.
Chloroform, CHCl3, is a solvent that has been used as an anesthetic. The heat of vaporization of chloroform at its
boiling point (61.2°C) is 29.6 kJ/mol. What is the entropy change when 1 mol CHCl3 vaporizes at its boiling point?
28.
Predict the sign of ΔS° if possible for each of the following reactions:
a.
C2H2 (g) + 2 H2 (g)  C2H6 (g)
b. N2 (g) + O2 (g)  2 NO (g)
c.
2 C2H2 (g) + 3 O2 (g)  4 CO (g) + 2 H2O (g)
d. 2 C (s) + O2 (g)  2 CO (g)
29.
Calculate ΔS° for the following reaction using S° data from Appendix 4:
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l)
30.
Using enthalpies of formation calculate ΔH° for the following reaction at 25°C. Also, calculate ΔS° for this reaction
from standard entropies at 25°C. Then use these values to calculate ΔG°.
2 CH3OH (l) + 3 O2 (g)  2 CO2 (g) + 4 H2O (l)
31.
Calculate the standard free energy of the following reactions at 25°C using standard free energies of formation.
a.
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g)
b. CaCO3 (calcite) + 2 H+ (aq)  Ca2+ (aq) + H2O (l) + CO2 (g) (for Ca2+ ΔG°f = —554 kJ)
32.
Consider: H2 (g) + Cl2 (g)  2 HCl (g). Calculate ΔG° and K for this reaction at 25°C.
33.
Calculate ΔG° and K for the following reaction at 25°C: CO (g) + 3 H2 (g)  CH4 (g) + H2O (g).
34.
Calculate ΔG° at 25°C for the equilibrium: CaF2 (s)  Ca2+ (aq) + 2 F— (aq). Calculate Ksp for this reaction at 25oC.
For CaF2 (s) ΔG°f = —1162 kJ/mol; see problem 31 for ΔG°f of Ca2+.
35.
Calculate ΔH°f for CH4 (g) using the following reactions:
CO2 (g) + 2 H2 (g)  CH2O (g) + H2O (g) ΔH° = 35.5 kJ
CO2 (g)  C (s) + O2 (g) ΔH° = 393 kJ
CH2O (g) + 2 H2 (g)  CH4 (g) + H2O (g)
ΔH° = —210 kJ
2 H2 (g) + O2 (g)  2 H2O (g)
ΔH° = —484 kJ
From appropriate standard entropy data calculate ΔS° for the formation of methane from its elements. Then, using ΔH°
and ΔS°, calculate ΔG°f for methane at 25°C.
Answers to Review Problems:
1.
CN– + H2O  HCN + OH–; CN— base, H2O acid, HCN conj. acid, OH— conj. base
2.
H2SO3, HIO4, H2S
3.
a. 3.7 x 10—3 M, b. 3.9 x 10—6 M
4.
a. acidic, b. basic, c. basic, d. acidic
5.
a. 2.5 x 10—11, b. 5.9 x 10—6
6.
9.4 x 10—4 M, 3.03
7.
3.00%, 0.73%
8.
10.47
9.
9.26, ΔpH = 0.17 (pH = 9.09)
10.
3.05
11.
29.87 g
12.
1.96
13.
(a) 4.40, (b) 8.59, (c) 11.74
14.
(a) 10.11, (b) 6.01, (c) 1.96
15.
9.3 x 10—10
16.
5.0 x 10—4
17.
1.0 x 10—3 g/L
18.
ppt'n occurs
19.
4.7 x 10—3 mol
20.
(a) 5.43, (b) 8.94, (c) 4.47
21.
5.1 x 10—6 g
22.
9.18
23.
a. 0.035, b. 2.2 x 10—4
24.
4.66 g
25.
(a) 0.56, (b) 7.00, (c) 13.01
26.
a. Ksp = [Cr3+][F—]3, b. Ksp = [Cd2+]3[PO43—]2, c. Ksp = [Sr2+][F—]2
27. 88.5 J/K•mol
28. a. —; b. can't predict; c. +; d. +
29. —242 J/K
30. ΔH° = —1453 kJ; ΔS° = —161 J/K; ΔG˚ = —1405 kJ
31. a. —801 kJ; b. —56 kJ
32. —190 kJ; 2.0 x 1033
33. —143 kJ; 1.2 x 1025
34. ΔG˚ = 50 kJ, Ksp = 1.7 x 10—9
35. ΔH˚ = —83.5 kJ; ΔS˚ = —82.0 J/K; ΔG˚ = —59.1 kJ