AP Chapter 15 & 16: Acid-Base Equilibria Name

AP Chapter 15 & 16: Acid-Base Equilibria Name Warm-Ups (Show your work for credit) Date 1. Date 2. Date 3. Date 4. Date 5. Date 6. Date 7. Date 8.

AP Chapter 15 & 16: Acid-Base Equilibria 2 Warm-Ups (Show your work for credit) Date 1. Date 2. Date 3. Date 4. Date 5. Date 6. Date 7. Date 8.

AP Chapter 15 & 16: Acid-Base Equilibria 3 Warm-ups and problems will be collected before you take the test. Read Chapter 15.1-15.11: Acids and Bases, and read Chapter 16.1-16.5 & 16.9: Acid-Base Equilibria Answer the following problems in the space provided. For problems involving an equation, carry out the following steps: 1. Write the equation. 2. Substitute numbers and units. 3. Show the final answer with units. There is no credit without showing work. Bronsted Acids and Bases 1. Write an equation for the reaction of the following compounds with water, and identify the acid-base conjugate pairs. (a) hydrocyanic acid (b) methanamine 2. Write the formula for the conjugate acid of each of the following bases: a. HS - b. HCO 3 - c. CO 3-2 d. H 2PO 4 - e. HPO 4 2- f. PO 4 3- ph: A Measure of Acidity 3. Write an equation for the ion-product constant for water, and state its value at 25 o C. 4. The ion-product constant for water at 25 o C is 1.0E-14 and at 40 o C is 3.8E-14. (a) Is the autoionization of water exothermic or endothermic? Explain. (b) Calculate the ph of water at 40 o C. 5. Calculate the ph of each of the following solutions: a. 2.8 x 10-4 M Ba(OH) 2 b. 5.2 x 10-4 M HNO 3. Strength of Acids and Bases 6. Classify each of the following species as a strong or weak acid or strong or weak base: a. LiOH b. HCN c. H 2O d. HClO 4 e. CH 3CH 2NH 2 f. HBr g. HF h. HCOOH

AP Chapter 15 & 16: Acid-Base Equilibria 4 7. Which of the following statements are true regarding a 1.0 M solution of a strong acid HA? a. [A - ] > [H + ] b. The ph is 0.00. c. [H + ] = l.0 M d. [HA] = 1.0 M 8. Which of the following statements are true regarding a 1.0 M solution of a weak acid HA? a. [A - ] > [HA] b. The ph is 0.00. c. [HA] > [H + ] d. [H + ] = [A - ] 9. Predict whether the following reaction will proceed from left to right to a significant extent. Explain. F - (aq) + H 2O(l) HF(aq) + OH - (aq) 10. Determine whether the equilibrium constant for the following reaction is greater or less than 1.0. Explain. CH 3COOH(aq) + NO 2- (aq) CH 3COO - (aq) + HNO 2(aq) Weak Acids and Acid Ionization Constants 11. A 0.0560 g quantity of acetic acid is dissolved in enough water to make 50.0 ml of solution. Calculate the concentrations of H +, CH 3COO -, and CH 3COOH at equilibrium. (K a for acetic acid = 1.8 x 10-5.) 12. What is the original molarity of a solution of formic acid (HCOOH) whose ph is 3.26 at equilibrium?

AP Chapter 15 & 16: Acid-Base Equilibria 5 13. Calculate the percent ionization of hydrofluoric acid at the following concentrations. Explain the trend. a. 0.60 M b. 0.046 M 14. A solution of acetic acid with a ph of 2.73 is compared to a solution of hydrochloric acid with a ph of 2.73. a. Predict which solution is more concentrated. Explain. b. Calculate the molar concentration of each solution. c. Which solution has a greater concentration of undissociated acid? Explain. 15. A solution of 0.13 M acetic acid is compared to a solution of 0.13 M hydrochloric acid. a. Predict which solution is more acidic (lower ph). Explain. b. Calculate the ph of each solution. Weak Bases and Base Ionization Constants 16. The ph of a 0.30 M solution of a weak base is 10.66. What is the K b of the base? 17. What is the original molarity of a solution of ammonia whose ph is 11.22?

AP Chapter 15 & 16: Acid-Base Equilibria 6 18. A solution of methylamine (CH 3NH 2) has a ph of 10.64. How many grams of methylamine are there in 100.0 ml of the solution? Acid-Base Conjugate Pairs 19. Write the equation that relates K a and K b for a conjugate pair. The K a of carbonic acid is 4.2E-7. Write the formula of its conjugate base and calculate K b for this base. Polyprotic Acids 20. Calculate and compare the ph of a 1.0 M HC1 solution with that of a 1.0 M H 2SO 4 solution. 21. Calculate the ph of a 0.100 M Na 2CO 3 solution. Molecular Structure and Acid Strength 22. Why is HF a weak acid and HCl a strong acid, even though F is more electronegative? 23. List the strength of the following acids from least to greatest. Explain. a. H 2SO 4, H 2SeO 4, and H 2TeO 4 b. H 3PO 4, H 3AsO 4, and H 3PO 3

AP Chapter 15 & 16: Acid-Base Equilibria 7 24. Which of the following is the stronger base: NF 3 or NH 3? Hint: draw Lewis structures to determine the more polar lone pair. Acid-Base Properties of Salt Solutions 25. Define salt hydrolysis. Write an equation for the hydrolysis of sodium hypochlorite. 26. State whether the following salts are acidic, basic, or neutral. For those that are not neutral write a balance equation showing why they are acidic or basic. a. NaCN b. CH 3NH 3Cl c. NaHCO 3 d. NaHSO 4 e. NH 4ClO 3 f. NaBr 27. Calculate the ph of a 0.42 M NH 4C1 solution. 28. Predict whether a solution containing the salt K 2HPO 4 will be acidic, neutral, or basic. 29. Calculate the ph of a 0.20 M ammonium acetate (CH 3COONH 4) solution. Hint: determine K b and K a of the cation and anion respectively, and think! 30. Use appropriate K a and/or K b values to calculate the equilibrium constant for the following reaction. Hint: find chemical equations that add up to this net equation. CH 3COOH(aq) + NO 2- (aq) CH 3COO - (aq) + HNO 2(aq)

AP Chapter 15 & 16: Acid-Base Equilibria 8 Acidic and Basic Oxides 31. Arrange the oxides in each of the following groups in order of increasing basicity. Explain. a. K 2O, A1 2O 3, BaO b. CrO 3, CrO, Cr 2O 3. 32. Write chemical equations for reactions of the following compounds. (a) Calcium oxide solid and water. (b) Sulfur dioxide gas and water. (c) Calcium oxide solid and sulfur dioxide gas. ph and Solubility 33. The ph of a saturated solution of a metal hydroxide, MOH, is 9.68. Calculate the K sp for the compound. 34. Mixed Fe 3+ and Zn 2+ ions in aqueous solution can be separated by selectively precipitating their hydroxides. Find the approximate ph range suitable for the separation of Fe 3+ and Zn 2+ by precipitation of Fe(OH) 3 from a solution that is initially 0.010 M in both Fe 3+ and Zn 2+. 35. Which of the following insoluble salts are more soluble in acid solution (H + ) than in pure water? For the salts whose solubility is affected by acid, write a chemical equation showing why its solubility is increased. a. CuI c. Zn(OH) 2 b. BaC 2O 4 d. Ca 3(PO 4) 2

AP Chapter 15 & 16: Acid-Base Equilibria 9 The Common Ion Effect 36. Write the net ionic equation for the hydrolysis of sodium nitrite. Use Le Chatelier's principle to predict the effect of the following changes on the extent of hydrolysis of sodium nitrite solution. Explain each by also writing a chemical equation(s). a. HC1 is added b. NaOH is added c. NaCl is added d. the solution is diluted. 37. Describe the effect on ph (increase, decrease, or no change) that results from each of the following: a. adding potassium acetate to an acetic acid solution b. adding ammonium nitrate to an ammonia solution c. adding sodium formate (HCOONa) to a formic acid (HCOOH) solution d. adding potassium chloride to a hydrochloric acid solution e. adding barium iodide to a hydroiodic acid solution 38. Write both the regular and log forms of the Henderson-Hasselbalch equation, and state its assumptions. 39. Determine the ph of a. a 0.20 M NH 3 solution b. a solution that is 0.20 M NH 3 and 0.30 M NH 4C1

AP Chapter 15 & 16: Acid-Base Equilibria 10 40. Determine the ph of a. a 0.150 M propanoic acid (CH 3CH 2COOH) solution (K a = 1.3E-5). b. a solution that is 0.150 M propanoic acid and 0.30 M sodium propanoate. Buffers 41. State whether each system is an acidic buffer, basic buffer, or not a buffer. a. KC1/HC1 b. NH 3/NH 4NO 3 c. Na 2HPO 4/NaH 2PO 4 d. KNO 2/HNO 2 e. KHSO 4/H 2SO 4 f. HCOOK/HCOOH 42. Calculate the ph of the following two buffer solutions. Which is the more effective buffer, that is, has a greater buffer capacity? Why? a. 2.0 M CH 3COONa/2.0 M CH 3COOH b. 0.20 M CH 3COONa/0.20 M CH 3COOH c. Write chemical equations showing how this buffer system neutralizes additions of H + and OH -. 43. The ph of blood plasma is 7.40. Assuming the principal buffer system is HCO 3- /H 2CO 3, calculate the ratio [HCO 3- ]/[H 2CO 3]. Is this buffer more effective against an added acid or an added base? Explain

AP Chapter 15 & 16: Acid-Base Equilibria 11 44. A student is asked to prepare a buffer solution at ph = 8.60, using one of the following weak acids and its conjugate: HA (K a = 2.7 x 10-3 ), HB (K A = 4.4 x 10-6 ), HC (K A = 2.6 x 10-9 ). Which acid should she choose? For the acid/conjugate selected, explain how to prepare the buffer. Acid-Base Titrations 45. Given 50 ml of 0.10 M HCl and 50 ml of 0.10 M acetic acid, will the amount of 0.10 M NaOH required to neutralize each solution be the same, more, or less? Explain. 46. Will the ph at the equivalence point of 50 ml 0.10 M HCl be the same, more, or less as the ph at the equivalence point for 50 ml of 0.10 M acetic acid? Explain. 47. A 5.00-g quantity of a diprotic acid was dissolved in water and made up to 250. ml. Calculate the molar mass of the acid if 25.0 ml of this solution required 11.1 ml of 1.00 M KOH for neutralization. Assume that both protons of the acid were titrated. 48. In a titration experiment, 20.4 ml of 0.883 M HCOOH neutralize 19.3 ml of Ba(OH) 2. What is the concentration of the Ba(OH) 2 solution? 49. Calculate the ph for the titration of 0.10 M HCOOH versus 0.10 M NaOH at the following points: a. at the half titration point b. at the equivalence point

AP Chapter 15 & 16: Acid-Base Equilibria 12 50. An 80. ml sample of 0.24 M ethanamine (CH 3CH 2NH 2) is put in a flask and titrated with 0.18 M HCl. (K b of ethanamine is 5.6E-4.) a. What is the original ph in the flask? b. What is the ph in the flask after 25. ml of the acid is added? c. What is the ph at the half-titration point? d. What is the ph at the equivalence point? e. What is the ph after the addition of 127 ml of HCl? f. Sketch the titration curve and indicate each point a-e on the plot.

AP Chapter 15 & 16: Acid-Base Equilibria 13 51. Sketch the titration curve for a weak diprotic acid vs. NaOH, and show how to determine K a1 and K a2 graphically. 52. A 52 ml solution of acetic acid with a ph of 2.73 is compared to a 52 ml solution of hydrochloric acid with a ph of 2.73. a. Predict which solution requires more 0.12 M NaOH to neutralize it. Explain. b. Calculate the volume of the 0.12 M NaOH solution needed to neutralize each. 53. Which chemical species are in significant concentration when: a. A hydrochloric acid solution is titrated to the half equivalence point with a solution of NaOH. b. An acetic acid solution is titrated to the half equivalence point with a solution of NaOH. Acid-Base Indicators 54. Two drops of indicator HIn (K a = 1.0E-9), where HIn is yellow and In - is blue, are placed in100 ml of 0.10 M HCl. a. What is the color initially? b. The solution is titrated with 0.10 M NaOH. At what ph will the color change of the indicator begin? c. What color will the solution be after 200. ml of NaOH solution is added? 55. What indicator dye(s) listed in Table 16.1 of your text is appropriate for acid-base titrations where the equivalence point is at ph 5.0?

AP Chapter 15 & 16: Acid-Base Equilibria 14 56. A student carried out an acid-base titration by adding NaOH solution from a buret to an Erlenmeyer flask containing HC1 solution and using phenolphthalein as indicator. At the equivalence point, she observed a faint reddish-pink color. However, after a few minutes, the solution gradually turned colorless. What happened? 57. Draw Lewis structures for the following molecules, state whether each is symmetrical or nonsymmetrical, state whether each is polar or nonpolar, state the geometric shape of the molecule, and state the hybridization of the central atom. a. BrCl 3 b. IF 5 58. Explain the following properties of solid potassium and lithium: a. Both solids are malleable. b. Both solids conduct electricity. c. Lithium has a higher melting point than potassium. d. The first ionization energy of lithium is greater than that of potassium. 59. Explain the following properties of solid CsBr and MgBr 2: a. Both solids are brittle. b. MgBr 2 has a higher melting point than CsBr. c. Both solids do not conduct electricity, but when dissolved in water, both solutions do conduct electricity.

AP Chapter 15 & 16: Acid-Base Equilibria 15 Titration Summary Titration problems are difficult unless you think of them in a systematic way. Two problems are the most common (although any combination is possible): weak acid in a flask titrated with strong base in a buret weak base in a flask titrated with strong acid in a buret In both cases, knowing special regions of the titration curve is key to solving problems. As an example, the titration curve for a weak acid in a flask (25 ml of 0.20 M HA) titrated with a strong base (0.35 M NaOH) in a buret is shown in Figure 1. ph d e ml base added Figure 1. Weak acid strong base titration curve. Key points and regions are shown on the curve: a. ph in the flask before addition of strong base. b. buffer region (H-H applies) c. ph at the half-titration point where ph = pk a d. ph at the equivalence point (ph > 7.0 in this case) e. region of excess strong base For regions b-e, do this step first: Determining ph at any point along the titration is a two-step process. First, one must do stoichiometry using the titration equation: Titration Equation: HA + OH - A - + H 2O (one-way arrow) Since OH - is a strong base, this reaction is quantitative, that is, every mole of OH - added subtracts a mole of HA and creates a mole of A -. Thus it is easy to determine the new HA and A - concentrations using simple stoichiometry. Do stoichiometry in moles, and then convert to concentrations using the new total volume. Point a. Before addition of any base (OH - ), the titration equation above hasn t happened yet, so the ph is governed solely by hydrolysis of the weak acid: HA H + + A - Knowing K a, you can set up a table and solve for H + and ph. c a b

AP Chapter 15 & 16: Acid-Base Equilibria 16 Region b: Before the equivalence point (region b), a buffer is created because the OH - has converted part of the HA into A -, so they are both present. Do stoichiometry using the titration equation above to determine the new HA and A - concentrations. Plug these into Henderson-Hasselbalch* (H-H) using K a of HA to determine the H + and the ph. Be sure to keep track of the new volume when determining concentrations. *Remember that H-H isn t so much the form of the equation, (normal vs. log form) as it is the simplifying approximation that hydrolysis of both the weak acid (HA) and conjugate base (A - ) can be ignored if both species are present. In other words, you plug the HA and A - concentrations determined using stoichiometry directly into K a. Point c: Point c is the half-titration point, where exactly half as many moles of strong base has been added as moles of weak acid originally present. Thus the concentration of conjugated base (A - ) created by the titration equation equals the concentration of weak acid (HA) remaining. When plugged into H-H, these two species cancel and ph = pk a. Point d: Point d is the equivalence point where all of the HA has been converted into A - by the titration equation above. Thus the moles of A - at the equivalence point is equal to initial moles of HA. Use M av a = M bv b to determine the volume of OH - solution added. Then add V a + V b to get the new total volume, and then calculate the concentration of A -. The ph is then determined solely by the hydrolysis of A - according to the equation: A - + H 2O HA + OH - Set up a table using this equation and plug into K b for A -. Solve for OH - and then ph. Region e: After the equivalence point (region e), stoichiometry of the titration equation will show an excess amount of OH -. ph is determined by the excess moles of OH - added after the equivalence point. Again keep track of the new volume, determine [OH - ] and then ph. No table is needed. Diprotic Acids Just one last hint, if there is a diprotic acid, refer to Figure 2 to determine K a1 and K a2. Figure 2. ph during titration of a diprotic acid with sodium hydroxide where A = the volume of NaOH needed to react with both of the acid hydrogens B = volume of NaOH needed to react with one of the acid hydrogens C = the volume of NaOH used when all of the first and half of the second hydrogens are neutralized D = the volume of NaOH needed to neutralize half of the first acid hydrogen E = the ph when half of the first hydrogen is neutralized, or pk al F = ph when half of the second hydrogen is neutralized, or pk a2