Practice Problems: Applications of Aqueous Equilibria

Transcription

1 Practice Problems: Applications of Aqueous Equilibria CHEM 1B 1. Ammonia (NH3) is a weak base with a Kb = 1.8 x 1 5. a) Write the balanced chemical equation for the reaction of ammonia with water. Using the I.C.E. method, calculate the ph and % ionization of a 1.75 M NH3 solution in 2.5 M NH4Cl. b) How would the % ionization and ph have been different without the NH4Cl? What is this effect called? The % ionization would be larger / smaller (circle one), and the ph would be higher / lower (circle one) without the. c) Is this solution a buffer? Explain. d) Other than with the ICE method, how else could you have solved this problem? Method:

2 2. A buffer is made from acetic acid (HC2H3O2) and sodium acetate (NaC2H3O2). 1.5 moles of acetic acid and 2.5 moles of sodium acetate are added to enough water to make 1.5 liter of solution. The ionization constant (Ka) for acetic acid is 1.8 x 1 5. a) Write the equation for the ionization of acetic acid in water, and using a stress-shiftequilibrium arrow diagram, show how the buffer would minimize the effect of adding a base to the solution. Stress: Shift: Equilibrium: b) Using the Henderson-Hasselbalch equation, find the ph of this buffer solution. c) Calculate the percent ionization of the acetic acid in this solution. d) At what ph would this solution have the largest buffering capacity in both directions?

3 3. A buffer is needed to maintain ph = a) There are four reagents to choose from to make the buffer: 1. M H3AsO4 (aq) (Ka1 = 6. x 1 3, Ka2 = 1.1 x 1 7, Ka3 = 3. x 1 12 ), NaH2AsO4 H2O (s), Na2HAsO4 7H2O (s), and Na3AsO4 12H2O (s) What conjugates should be chosen and why? Conj. Acid Conj. Base b) It is decided that the concentration of the conjugate acid should be.1 M. Use the Henderson-Hasselbalch equation to calculate the concentration of the conjugate base required to achieve the desired ph. c) Use stoichiometry to calculate the amount of the solid salt of the conjugate base needed to make 1. ml of buffer. d) Use the dilution equation to calculate the amount of stock aqueous solution needed to make 1. ml of buffer that is.1 M conjugate acid. e) Other than the amounts of solid and aqueous solution calculated in Parts c) and d) above, what makes up the rest of the volume of the buffer? f) Would this buffer have a better buffering capacity against an acid or a base (circle one)?

4 4. Based on the titration curve to the right and the tables of Ka values below, determine the following: a) The most appropriate indicator for the titration. b) Whether the unknown acid is a strong acid or a weak acid (circle one). How do you know? ph Titration of 1. ml of an Unknown Acid c) The most likely identity of the unknown acid. d) The original concentration of the 1. ml sample of unknown monoprotic acid ml of.5 M NaOH added Indicator Ka Metacresol Purple 3.1 x1 2 Thymol Blue 2.2 x 1 2 Methyl Orange 3.5 x 1 4 Bromophenol Blue 7.9 x 1 5 Bromocresol Green 1.3 x 1 5 Methyl Red 1. x 1 5 Chlorophenol Red 5.6 x 1 7 Bromocresol Purple 4. x 1 7 Bromothymol Blue 5. x 1 8 Cresol Red 3.5 x 1 9 Phenolphthalein 3 x 1 1 Acid Ka chlorous acid 1.2 x 1 2 hydrofluoric acid 7.2 x 1 4 nitrous acid 4. x 1 4 lactic acid 1.38 x 1 4 benzoic acid 6.4 x 1 5 acetic acid 1.8 x 1 5 hypochlorous acid 3.5 x 1 8 ammonium ion 5.6 x 1 1

5 14 Titration of 25. ml of an Unknown Diprotic Acid ph 12 F 1 E 8 D A B C ml of.1 M NaOH added 5. Based on the titration curve for a diprotic acid above, answer the following: a) Identify in the table to the right which species exist in solution at each point on the curve (using H2A to represent a generic diprotic acid). Point A B C D Species from the unknown Species from the titrant b) Given the table of ionization constants, identify the unknown acid. E F c) Calculate the concentration of the original diprotic acid solution. Acid Ka1 Ka2 oxalic acid 6.5 x x 1 5 sulfurous acid 1.5 x x 1 7 ascorbic acid 7.9 x x 1 12 carbonic acid 4.3 x x 1 11

6 6. Based on the titration curve to the right and the tables of Kb values below, determine the following: a) The most appropriate indicator for the titration. b) Whether the unknown base is a strong base or a weak base (circle one). How do you know? ph Titration of 15. ml of an Unknown Base c) The most likely identity of the unknown base. d) The original concentration of the 15. ml sample of unknown base ml of.25 M HCl added Indicator Ka Metacresol Purple 3.1 x1 2 Thymol Blue 2.2 x 1 2 Methyl Orange 3.5 x 1 4 Bromophenol Blue 7.9 x 1 5 Bromocresol Green 1.3 x 1 5 Methyl Red 1. x 1 5 Chlorophenol Red 5.6 x 1 7 Bromocresol Purple 4. x 1 7 Bromothymol Blue 5. x 1 8 Cresol Red 3.5 x 1 9 Phenolphthalein 3 x 1 1 Base Kb ethylamine 6.4 x 1 4 dimethylamine 5.4 x 1 4 methylamine 3.7 x 1 4 trimethylamine 6.5 x 1 5 ammonia 1.8 x 1 5 hydroxylamine 9.1 x 1 9 phenylamine 4.3 x 1 1

7 7. The Ksp of silver chromate (Ag2CrO4) is 1.9 x a) Write the balanced complete ionic equation for the dissociation of Ag2CrO4. b) Using the I.C.E. method and the equilibrium expression, find the molar solubility of Ag2CrO4. (Hint: you are solving for X) 8. The molar solubility of chromium (III) hydroxide in water at 25 C is 1.26 x 1 8 M Cr(OH)3. a) Write the balanced complete ionic equation for the dissociation of Cr(OH)3. b) Using the I.C.E. method and the equilibrium expression, find the Ksp of Cr(OH)3. 9. The molar solubilities of two ionic compounds can be directly compared, but not always the Ksp values. Circle the two compounds for which the Ksp values could be compared and indicate which one of the two is more soluble in water. PbCrO4 Mg(OH)2 Cu3(PO4)2 Ag2SO4 Al(OH)3 K sp = 2.8 x 1 13 K sp = 5.6 x 1 12 K sp = 1.4 x 1 37 K sp = 1.2 x 1 5 K sp = 1.9 x 1 33