Exam 2 Review Guide Chapter 15 Typical Math Problems 1) What is the pH of a 0.750M solution of HCl? 2) a) What is the pH of a 0.750 M solution of a weak monoprotic acid with a pKa =6.15? b) What is the %-­‐ionization of this weak acid? c) What is the corresponding pKb for the conjugate base of this weak acid? 3) a) What is the pKb for a weak base if the pH of a 0.125M solution of this weak base was found to be 8.16? 4) When a weak acid is added to water to make a 0.250M solution, the acid was found to be 1.5% ionized. What is the Ka and pKa of this weak acid? 5) At 60°C, Kw = 9.6x10-­‐14 (Kw increases as the temperature increases). a) For pure water at 60°C, what is the [H+] and pH. b) If a sample of an aqueous solution was found to have a pH of 7.00 at 60°C, is this solution acidic, basic, or neutral given your answer to a)? Typical Concept Problems 1) Which of the following acids is stronger and why? HCl or HBr 2) Which of the following bases is stronger and why? OH-­‐ or NH2-­‐ 3) Which of the following acids is stronger and why? H3AsO4 or H3AsO3 4) Which of the following acids is stronger and why? Lauric Acid, pKa=4.95 or Naproxen, pKa=4.50 5) Which of the following 0.10M solutions would have the lowest pH? 0.10 M HF; 0.10M HBr; 0.10M HCl 6) Label the following salt solutions as neutral, basic, or acidic: NaCl, NaF, NaCO3, NH4Cl, NaOH 7) Write the conjugate base formulas for the following: HClO4, H3AsO3, CH3CO2H, C6H5CO2H, HS-­‐ Chapter 16 Typical Math Problems 1) What is the pH of a buffer solution which is 0.12M in benzoic acid, pKa = 4.19, and is 0.11M in sodium benzoate? 2) Find the %-­‐diss of a buffer if the pKa = 7.10 and the pH of the buffer is 7.50. 3) A weak monoprotic acid has a pKa = 6.15. 50.00 mL of an 0.1250M aqueous solution of this weak acid is titrated with 0.1000M NaOH. a) What is the equivalence point volume and ½ equivalence point volume for this titration? Find the pH b) before the titration begins; c) after 20.00 mL of the NaOH has been added, d) after 62.50 mL of the NaOH has been added; and e) after 85.00 mL of the NaOH has been added. 4) A 0.2500 g sample of a weak acid was titrated with 0.1000M NaOH. After 9.15 mL was added, the pH was measured to be 5.67; and 20.00 mL was required to reach the equivalence point. a) What is the molecular weight of this acid; and b) find the pKa of this acid. 5) Determine the molar solubility of barium sulfate, Ksp = 1.10x10-­‐10 6) If the solubility of ferrous hydroxide is 0.00143g/L at 25°C, a) calculate the molar solubility; and b) calculate the Ksp for ferrous hydroxide. 7) The Ksp for AgCl is 1.56x10-­‐10. If HCl is added to a solution of AgNO3, precipitating out AgCl, what will be the [Ag+] when the chloride concentration is 0.10M? 8) The equations and Ka values for 3 acids are as follows: Hydrogen carbonate ion, Ka=4.2x10-­‐7 Phosphoric acid, Ka=6.2x10-­‐8 Hydrogen sulfate ion, Ka=1.3x10-­‐2 Which of the above acids would be best for preparing a buffer with a pH of 7.2 Typical Concept Problems 1) For a weak acid, or a buffer, why is the half-­‐equivalence point important? 2) What is true at the equivalence point of the titration of a weak monoprotic acid? Why is the pH above 7? 3) Why should dilution of a buffer with pure water not change the pH? 4) Why would a mixture of NH3 and NH4Cl constitute a buffer; why would a mixture of HCl and NH3 constitute a buffer; why would a mixture of NaOH and NH3 not constitute a buffer? 5) Will a mixture of HCl and NaOH make a buffer? Why or why not? 6) Why is CaCO3 more soluble in HCl than in pure water? Chapter 17 Problems 1. Answer the following questions in terms of thermodynamic principles.
a. Consider the reaction represented below, which is spontaneous at 298 K.
CO2(g) + 2 NH3(g) —> CO(NH2)2(s) + H2O(l); ΔH° = -134 kJ
i. For the reaction, indicate whether the standard entropy change, ΔS°, is positive, or
negative, or zero. Justify your answer.
ii. Which factor, the change in enthalpy, ΔH°, or the change in entropy, ΔS°, provides
the principal driving force for the reaction at 298 K? Explain.
iii. For the reaction, how is the value of the standard free energy change, ΔG°, affected
by an increase in temperature? Explain.
b. Some reactions that are predicted by their sign of ΔG° to be spontaneous at room
temperature do not proceed at a measurable rate at room temperature.
i. Account for this apparent contradiction.
ii. A suitable catalyst increases the rate of such a reaction. What effect does the catalyst have on
ΔG° for the reaction? Explain.
2. Use the following rxn to answer the questions.
CO(g) + 2 H2(g) ! CH3OH(l)
For this reaction, ΔH° = -128.1 kilojoules
ΔHf° (kJ/mol)
ΔGf° (kJ/mol)
S° (J/mol•K)
CO(g)
-110.5
-137.3
197.9
CH3OH(l)
-238.6
-166.2
126.8
The data in the table above were determined at 25 °C.
a) Calculate ΔG° for the reaction above at 25 °C.
b) Calculate Keq for the reaction above at 25 °C.
c) Calculate ΔS° for the reaction above at 25 °C.
d) In the table above, there is no data for H2. What are the values of ΔHf°, ΔGf°, and of the
absolute entropy, S°, for H2 at 25 °C?
3. Use the following table and rxn to answer the questions.
Substance Standard Heat of Formation, ΔHf°, in kJ mol¯ˉ1 Absolute Entropy S°, in J mol¯ˉ1 K¯ˉ1 C(s) 0.00 5.69 CO2(g) -­‐393.5 213.6 H2(g) 0.00 130.6 H2O(l) -­‐285.85 69.96 O2(g) 0.00 205.0 C3H7COOH(l) ? 226.3 The enthalpy change for the combustion of butyric acid at 25 °C, ΔH°comb, is -2,183.5
kilojoules per mole.
C3H7COOH(l) + 5 O2(g) ! 4 CO2(g) + 4 H2O(l)
a) From the data above, calculate the standard heat of formation, ΔHf°, for butyric acid at 25 °C.
b) Write a correctly balanced equation for the formation of butyric acid from its elements.
c) Calculate the standard entropy change, ΔSf°, for the formation of butyric acid at 25 °C. The
entropy change, ΔS, for the combustion reaction above is -117.1 J K¯1 at 25 °C.
d) Calculate the standard free energy of formation, ΔGf°, for butyric acid at 25 °C.
4. Use the following formation rxn to answer the questions.
1/2 Cl2(g) + 3/2 F2(g) → ClF3(g) ClF3 can be prepared by the reaction represented by the equation above. For ClF3 the
standard enthalpy of formation, ΔH°f, is -163.2 kilojoules/mole and the standard free energy
of formation, ΔG°f, is -123.0 kilojoules/mole.
a) Calculate the value of the equilibrium constant for the reaction at 298 K.
b) Calculate the standard entropy change, ΔSf°, for the above reaction at 298 K.
c) If ClF3 were produced as a liquid rather than as a gas, how would the sign and magnitude of
ΔS for the reaction be affected? Explain.
5. What is true about ΔG and Q at equilibrium? What is true about Q and K at Standard State?