Download Dr. Arrington Exam 3

Chemistry 124
Honor Pledge: ______________________________________
Dr. Arrington
Exam 3 (100 points), Thermodynamics and Acid Base Equilibria
Thursday, March 24, 2011
Show all work on numeric problems in Section II to receive full or partial credit. Give all answers with correct units.
Your exam should contain 6 pages of questions.
(2 points) State the 2nd Law of Thermodynamics:
___________________________________________________________________________________________________________
___________________________________________________________________________________________________________
Multiple Choice: (2 points each) Place answer in the blank next to the problem. Only answers in this blank will be graded.
_____ 1. Which of the following reactions will entropy decrease?
1. PCl5(g) → PCl3(g) + Cl2(g)
2. 3C2H2(g) → C6H6(g)
3. H2O(l) → H2O(s)
(A) 1 only
(B) 2 only
(C) 3 only
(D) 1 and 2
(E) 1 and 3
(F) 2 and 3
_____ 2. The pH of 0.1 M nitric acid (HNO3) is 1.0 and the pH of 0.1 M phosphoric acid (H3PO4) is 1.6. The pH of the
phosphoric acid solution is higher because
(A) HNO3 dissociates more fully than H3PO4.
(C) HNO3 is a weaker acid than H3PO4.
(B) H3PO4 has more hydrogen atoms.
(D) H3PO4 produces some OH− in water.
_____ 3. A reasonable assumption(s) for the concentration of the species in a weak diprotic acid, H2A, in water is(are):
1.
[H2A] >> [H+]
(A) 1 only
(B) 2 only
2. [HA−] = [H+]
(C) 3 only
3. [A2−] = Ka2
(D) 1 and 2
(E) 2 and 3
(F) 1, 2 and 3
_____ 4. For the two pH titration curves of acid A and B with NaOH one can conclude:
1.
2.
3.
acid A is a weaker acid than acid B
acid B is a strong acid
moles acid A > moles acid B
(A) 1 only
(D) 1 and 2 only
(B) 2 only
(E) 1 and 3 only
(C) 3 only
(F) 1, 2 and 3
_____ 5. A reasonable estimate for the Ka of acid A is
(A) 1 x 10−3
(B) 1 x 10−4
(C) 1 x 10−5
(D) 1 x 10−6
(E) 1 x 10−8
(F) 1 x 10−10
(E) 1 x 10−7 M
enough information is
(F) Not
provided to determine [HB]
_____ 6 The initial concentration of acid B in the above titration is?
(A) 0.2 M
(B) 0.1 M
(C) 2.0 M
(D) 0.01 M
0
_____ 7. Which of the following reaction(s) is(are) enthalpy driven at 298K?
(A) (i) only
(i)
PCl5(g) qe PCl3(g) + Cl2(g)
(ii)
HF(aq) qe F (aq) + H+(aq)
(iii)
2O3(g) qe 3O2(g)
−
(B) (ii) only
(C) (iii) only
∆H° = +87.84
, ∆S° = 843.6
∆H° = −12.55
, ∆S° = −102.5
∆H° = −285.42
, ∆S° = 137.56
(D) i and iii
(E) ii and iii
(F) all are enthalpy driven
_____ 8. The reason that enthalpy is resoponsible for making a chemical reaction spontaneous is that
(A)
(B)
(C)
(D)
(E)
(F)
exothermic reactions decrease the entropy of the surrounding and make the entropy of the universe decrease.
exothermic reactions increase the entropy of the surrounding and make the entropy of the universe increase.
exothermic reactions increase the entropy of the surrounding and make the entropy of the universe decrease.
endothermic reactions decrease the entropy of the surrounding and make the entropy of the universe decrease.
endothermic reactions increase the entropy of the surrounding and make the entropy of the universe increase.
endothermic reactions increase the entropy of the surrounding and make the entropy of the universe decrease.
_____ 9. Benzene has a normal melting point of 278 K. At 298 K what is ∆G° for the melting process described by the
equation C6H6(s) –––> C6H6(l) ; ∆H° = 9.9 kJ/mol?
(A) 35.6 J/mol
(B) −35.6 J/mol (C) −712 J/mol
(D) 712 J/mol
(E) 178 J/mol
(F) −178 J/mol
_____ 10. For the reaction described in Problem 9 one can conclude that for the reaction C6H6(s) sd C6H6(l) at 298K?
(A) ∆Suniv > ∆Ssys > ∆Ssurr
(D) ∆Ssys > ∆Suniv > ∆Ssurr
(B) ∆Suniv > ∆Ssurr > ∆Ssys
(E) ∆Ssurr > ∆Ssys > ∆Suniv
(C) ∆Ssys > ∆Ssurr > ∆Suniv
(F) ∆Ssurr > ∆Suniv > ∆Ssys
_____ 11. Which condition must be true for all reactions which have a negative value of ∆G?
1.) ∆Go > 0
2.) Q < K
3.) Q > K
4.) Q = K
(A) 1only
(B) 2 only
(C) 3 only
(d) 1 and 2
(e) 1 and 3
(F) 1 and 4
______ 12. Which plot below shows Ka for hydrofluoric acid, HF(aq), as a function of temperature (see data in Problem 7)?
_____ 13. After 0.448 grams of a monoprotic acid are dissolved in 25 mL of water it requires 38.2 mL of 0.085 M NaOH
to reach the end-point in a titration. The molar mass of the acid is
(A) 526 g/mol
(B) 345 g/mol
(C) 201 g/mol
(D) 138 g/mol
(E) 39 g/mol
(F) 7.2 g/mol
_____ 14. Which of the following aqueous solutions will form a buffer when mixed in equal amounts?
1.
2.
0.1 M NaBr and 0.1 M HBr
0.1 NaCN and 0.1 HCN
(A) 1 only
(B) 2 only
(C) 3 and 4
3. 0.05 NaOH and 0.1 KCN
4. 0.1 HF and 0.05 NaOH
(D) 2 and 4
(E) 2, 3, and 4
(F) All solutions will buffer
_____ 15. In titrating 0.2 M acetic acid (CH3COOH) with 0.2 M NaOH, the solution at the equvalence point can be
characterized as
(A) 0.2 M NaCH3COO
(C) 0.2 M NaCl and H2O
(B) 0.1 M NaCH3COO
(D) 0.1 M CH3COOH and 0.1 M NaOH
_____ 16. Malonic acid (C3H2O4H2) is a diprodic acid with Ka values of: Ka1 = 1.4 × 10−3, Ka2 = 2.0 × 10−6. What is the Kb
of hydrogen malonate (C3H2O4H−) ?
(A) 5.0 x 10−9
(B) 2.0 x 10−10
(C) 1.4x10−11
(D) 4.4 x10−5
(E) 7.1 x 10−12
(F) 5.0 x10−2
_____17. Which of the following pictures best represents a solution of 1.0 M H2SO4 ?
(A)
(B)
(C)
(D)
_____ 18. The indicator alizarin yellow has a Ka of 1.25 x 10−11. The protenated form of alizarin yellow is clear and the
ionized form is yellow. At a pH slightly greater than neutral (pH = 7.4) what is the color of a solution with this
indicator.
(A) clear
(B) pink
(C) yellow
(D) green
_____ 19. Which of the following statements are true regarding a 1.0 M solution of a strong acid HA.
1. [A−] > [H+]
(A) 1 and 2
(B) 2 and 3
2. pH = 0.0
(C) 1 and 4
3. [H+] = 1.0 M
(D) 1, 2 and 3
(E) 1, 3 and 4
4. [HA] = 1.0 M
(F) 2, 3 and 4
_____ 20. Which of the following compounds would act as a base in an aqueous solution?
(A) I only
(B) II only
(C) III only
(D) IV only
(E) I and III
(F) II and IV
Section II Show your work on numeric problems to receive full and partial credit.
21. (16 points) Dr. Hettes needs to write a lab protocol for creating a citrate buffer. Citric acid is a triprotic acid that we can
write as H3Cit. The Ka values for citric acid are: Ka1 = 7.4x10−4, Ka2 = 1.7x10−5, & Ka3 4.0 x10−7. If Dr. Hettes needs a
buffer at pH = 6.7 what species should the buffer solution contain? (Circle an image below.)
Dr. Hettes has begun the protocol as:
1.) Dissolve 3.87 g of Na3Cit (FW 258 g/mol) in 1L of H2O.
2.) To this solution add _____ mL of 0.2 M (NaOH or HCl) to make the pH = 6.7
Help Dr. Hettes complete the protocol by:
1.) Determine if she must add strong acid or strong base to the above sodium citrate solution.
2.) Report the mL of strong 0.2 M solution required to prepare this buffer solution.
__________________________________ mL NaOH or HCl
22. (22 points) Three aqueous titration curves are given (HA = weak acid , B = weak base, and NaA = salt of a weak acid )
(I) HCl is titrated with B
(II) NaA is titrated with HCl
(III) HA is titrated with NaOH
(a) Label each titration curve as (I), (II), or (III) in the little box at the top of each plot.
(b) Make a solution drawing that indicates which species are present in significant quantities (not obtained from an equilibrium
calculation) at the point on the titration curve indicated by the arrow. Do not show spectator ions in your drawing.
___________________________
__________________________
_______________________
(c) pH is determined in three ways during a titration:
(i) pH calculated from the excess concentration of H+ or OH− in solution.
(ii) pH calculated from the Henderson-Hasselbalch equation (buffer equation)
(iii) pH calculated from a weak acid or weak base equilibrium calculation.
Place the calculation type (i, ii or iii) required to determine pH for the solution you have drawn on the line below each diagram.
(d) Titration III is performed starting with 25 mL of 0.15 M HA (Ka = 2.0 x 10−6) and 0.08 M NaOH. How many mL of
NaOH is required to reach the equivalence point of titration?
______________________________________ mL NaOH
(e) Determine the pH of the titration described in part (d) at the equivalence point
pH = ____________________________
22. (20 points) Thermodynamics and equilibria are intimately connected. We will examine this relationship using
the weak acid equilibria of hydrofluoric acid.
(a) Use the thermodynamic data for the weak acid dissociation of HF(aq) (Problem 7) to calculate Ka for the acid at 298 K.
HF(aq) qe F−(aq) + H+(aq)
∆H° = −12.55 , ∆S° = −102.5
Ka = ___________________________
(b) Calculate the % dissociation of a 1M HF(aq) solution at 298 K.
% Dissociation @ 25 °C = ______________________
(c) The great thing about thermodynamics is that we can find K at any temperature we want. Find Ka for
hydrofluoric acid at 65 °C.
Ka65 °C = __________________________
(d) Determine the % dissociation of 1M HF(aq) at 65 °C.
% Dissociation@ 65 °C = _______________________
-Useful Conversions & Scratch Paperh = 6.63 x 10−34 J⋅s
F = 96,500 J/V mole
4.18 joule = 1 calorie
R = 0.0821 L ⋅ atm
K ⋅ mol
1.67 × 10-23 g = 1 amu
1 inch = 2.54 cm
1atm = 760 mmHg
ln K 2 − ln K 1 = −
∆H
R
1 cm3 = 1 mL
1 1
 − 
 T2 T1 
1 pm = 1× 10-12 m
R = 8.314
J
K ⋅ mol
Kw = 1 x 10-14