NAME: ENCE 610: Environmental Chemistry FINAL EXAM (Fall 2012) 1. (100 points) Water hardness is caused by the presence of dissolved minerals. Typically, we think of

hardness as a measure of the Ca2+ and Mg2+ concentrations. One way to remove hardness is to precipitate calcium and magnesium as solids. After these solids are precipitated, the “softened” water can be used for various purposes. Consider a water containing 10-4 M Ca2+, 2×10-5 M Mg2+, and 10-3 M TOTCO3. While there are other background electrolytes present, you can assume that they do not significantly affect precipitation of calcium and magnesium solids. Please use the stability constants for formation of complexes and solids provided on the following page (do not use the values from the textbook) – you should consider all of the highlighted species (i.e., the species in the red box) a. (30 points) Draw TOTMe curves for the following metals. Validate (quantitatively) all

decisions.

i. Ca2+

ii. Mg2+

b. (45 points) Draw log C-pH diagrams for the following species. Validate (quantitatively) all decisions. i. All calcium species.

ii. All magnesium species.

c. (20 points) Identify the following key pieces of information: i. The optimal pH for calcium precipitation. What is the corresponding solid?

ii. The optimal pH for magnesium precipitation. What is the corresponding solid?

d. (5 points) Comment on your findings. Specifically, indicate the pH that you would use to design

the precipitation process. Explain this decision.

NAME: ENCE 610: Environmental Chemistry FINAL EXAM (Fall 2012)

NAME: ENCE 610: Environmental Chemistry FINAL EXAM (Fall 2012) 2. (100 points) The use of chlorine as a disinfectant for US drinking water supplies began in the early

1900s in New Jersey. In the 1970s, we discovered that chlorine reacts with natural organic matter to form compounds known as disinfection by-products (DBPs). Many DBPs stemming from chlorine- based disinfection are carcinogens. To combat the formation of these compounds, while still attaining adequate inactivation of microorganisms, environmental engineers have studied several “alternative disinfectants.” One alternative disinfectant is ozone. Unlike chlorine which forms dozens of hazardous DBPs, ozone forms only one major DBP of concern – bromate, a suspected carcinogen. Recently, researchers determined that chlorine can also oxidize bromide to bromate. To better understand these phenomena, your boss has asked you to compare the formation of bromate during chlorine and ozone disinfection.

a. (25 points) Write balanced reactions and identify the equilibrium constants (K) for oxidation of bromide to bromate for the following cases:

i. Hypochlorous acid is reduced to chloride.

ii. Gaseous ozone is reduced to oxygen gas.

b. (25 points) Determine the mass of bromate formed per mass of oxidant for the following:

i. Hypochlorous acid. Express your answer as “mg bromate / mg chlorine (as Cl2).”

ii. Aqueous ozone. Express your answer as “mg bromate / mg O3(aq).”

c. (45 points) Determine whether bromide oxidation is favorable during disinfection of a water containing 2×10-7 M bromide and 3×10-5 M chloride at pH 7.6 for the oxidant conditions described below. Note that the maximum allowable concentration of bromate in drinking water is 10 µg/L. Explain any decisions/assumptions that you make.

i. The chlorine residual is 0.7 mg/L as Cl2.

ii. The residual ozone concentration is 0.2 mg/L as O3(aq).

d. (5 points) Comment on your findings from part c. Which disinfectant would you use for the water described in part c and why? Explain the advantages and disadvantages of your choice.

NAME: ENCE 610: Environmental Chemistry FINAL EXAM (Fall 2012) 3. (100 points) Students in my lab use phosphate-based buffer solutions when running kinetics

experiments. These buffers are useful because phosphate does not interfere with the reactions taking place and because the three pKa values associated with phosphoric acid cover a large portion of the pH scale. In this question, you are asked to investigate the generation of phosphate buffers and to draw titration curves and buffer intensity profiles.

a. (30 points) Indicate the mass of H3PO4, NaH2PO4, Na2HPO4, and Na3PO4 that must be added to 1 L of deionized water to generate the following buffer solutions:

i. 10 mM phosphate buffer (TOTPO4 = 10 mM) at pH 6.2.

ii. 25 mM phosphate buffer (TOTPO4 = 25 mM) at pH 9.1.

b. (35 points) Consider a 100 mL solution of the 25 mM phosphate buffer described above in part a.ii. Starting from pH 9.1, you add 10-1 M NaOH to titrate to pH = 12 and 2×10-1 M HCl to titrate to pH = 2. Draw the titration curve with pH on the y-axis and the volume of titrant (in mL added) on the x-axis. Use negative values to account for the volume of acid added and positive values for volume of base added. Be sure to choose appropriate scales for your plot.

c. (20 points) Plot the buffer intensity for the following solutions:

i. 10 mM phosphate buffer (TOTPO4 = 10 mM)

ii. 25 mM phosphate buffer (TOTPO4 = 25 mM)

d. (15 points) You want to buffer the pH of your solution at pH = 9.1 such that pH does not deviate by more than 0.05 units when 10-4 M of acid/base is added. Is the 10 mM buffer sufficient? What is the minimum buffer strength needed to insure these conditions are met?