Environmental Engineering

Question 1: Iron species in solution exist in two major forms with different charges: “ferrous” iron, with a charge of +2, and “ferric” iron with a charge of +3. Ferrous iron typically dominates in anaerobic environments, and ferric in aerobic environments, although the time needed for the transition can be long (so sometimes they are found in environments where they are not stable in the long term). Ferrous iron is immensely more soluble than ferric iron, so when dissolved ferrous iron is converted to ferric iron, it almost all precipitates, generating a rusty color in the water. 

Ferrous iron is fairly common in the leachate from landfills, which it enters because of corrosion of materials made of steel. Carbonate is also common in these leachates, due to the degradation of organic matter (which consumes the available oxygen and therefore stabilizes the ferrous iron).

(a) A leachate at pH 6.8 containing 10^-2 M TOTCO₃ and 5x10^-3 M Fe²⁺. The solution is also in contact with the mineral siderite [FeCO₃(s), with Ksp = 2.57 x 10^-11]. Determine whether there is “too much” or “too little” of the dissolved species present for equilibrium with the solid. (These conditions are referred to as supersaturated and undersaturated, respectively.)

(b) How much Fe²⁺ will be present in the solution after it reaches equilibrium with the solid, assuming that the pH remains at 6.8? Keep in mind that, regardless of whether solid dissolves or precipitates, the dissolved concentrations of both Fe²⁺ and carbonate species change.

(c) In reality, dissolution or precipitation of the solid will cause the pH to change. For the scenario being considered here, in which direction do you expect this change to be?

Question 2: Anaerobic bacteria can generate hydrogen sulfide (H₂S) in sewer pipes when they oxidize organic matter and use sulfate (SO₄^-2 as the electron acceptor. When the sewage enters a wastewater treatment plant, hypochlorous acid (HOCl) is sometimes added to oxidize the H₂S and re-convert it to SO₄^-2, thereby minimizing foul odors and mitigating potential adverse effects of the H₂S on the treatment process. In the process, HOCl is converted to Cl^-

(a) How much HOCl is needed to react with each mg/L of S in the latter reaction?

(b) If the H₂S were not oxidized by HOCl, it would eventually be oxidized by dissolved oxygen during the wastewater treatment process. How much O2 would be consumed per mg/L of S by this reaction?

Question 4: A wastewater containing 60 mg/L of easily degradable organic compounds with an average chemical formula of C₈H₁₈O₅ is discharged at a rate of 0.2 m³/s into a river that has an upstream flow rate of 0.7 m3/s. The water upstream of the discharge point has an ultimate BOD of 4 mg/L and contains 8.0 mg/L DO. The rate constants for BOD utilization and reaeration in the river are 0.14 d^-1 and 0.30 d^-1, respectively. The Henry's law constant for oxygen under the ambient conditions is 1.70x10^-3 mol/L-atm.

(a) Plot the BOD remaining in the water (i.e., L) as a function of travel time downstream from the discharge point, until 95% of the organics have decayed. 

(b) What is the slope of the DO sag curve immediately after the two waters mix?

Question 5: Plot the DO sag curve for a portion of a river where the values of k_d and k_r are 0.25 d^-1 and 0.3 d^-1, respectively. The BOD and DO at the beginning of the reach are 20 mg/L and 6 mg/L, respectively, and the saturation concentration of DO is 10.4 mg/L.

At what point or region in the river is:

(a) the reaeration rate greater than the BOD utilization rate?

(b) The BOD utilization rate greater than or equal to 3 mg/L-d?

(c) The BOD remaining in the solution less than or equal to 12 mg/L?

(d) The reaeration rate greater than or equal to 1.5 mg/L-d?