1. ASPECTS OF AQUATIC REDOX CHEMISTRY

2. PART - I REDOX CONDITIONS IN NATURAL WATERS Redox conditions in natural waters are controlled largely by photosynthesis and bacterial respiration processes

3. Oxidation – Reduction Reactions Oxidation - a process involving loss of electrons. Reduction - a process involving gain of electrons. Reductant - a species that loses electrons. Oxidant - a species that gains electrons. Free electrons do not accumulate in solution. Electrons lost from one species in solution must be immediately gained by another. Ox 1 + Red 2  Red 1 + Ox 2

4. Reduction-Oxidation Potential The potential that is generated between an oxidation or reduction half cell and the standard hydrogen electrode In aqueous solutions, the reduction potential is the tendency of the solution to either gain or lose electrons The potential abundance of electrons or the electron activity

5. PHOTOSYNTHESIS Synthesis of organic matter by photosynthesis The average composition of the organic matter in plankton is approximately: C 106 H 263 O 110 N 16 P 1. Therefore, photosynthesis reaction can be represented by the following and more complex reaction algae Redfield Ratio (Redfield et al., 1963)

6. Redfield Ratio Concept

7. RESPIRATION In general, respiration involves the decomposition of organic matter produced through photosynthesis During respiration, the organic matter is oxidized and an electron acceptor is reduced Example electron acceptors: Respiration can occur under oxygenated (or aerobic) conditions or in the absence of molecular oxygen (anaerobic respiration).

8. RESPIRATION (cont’d) In water containing excessive biomass (e.g. during algal blooms), dead organic matter (OM) is mineralized via microbial respiration in the presence of terminal electron acceptors (TEA) as illustrated in the following general reaction Using CH 2 O as a general formula for OM and different TEA types, one obtains (see next slide)

9. Progressive Microbial Respiration of OM in Natural Waters and Thermodynamics

10. Redox Couples For any half reaction, the oxidized/reduced pair is the redox couple: –Fe 2+  Fe 3+ + e- –Couple: Fe 2+ /Fe 3+ –H 2 S + 4 H 2 O  SO 4 2- + 10 H + + 8 e- –Couple: H 2 S/SO 4 2-

11. Redox Ladder 1 0.5 0 -0.5 E h (V) O2O2 H2OH2O NO 3 - NO 2 - NH 4 + Mn +4 Mn +2 FeOOHFe +2 SO 4 -2 HS - CO 2 CH 4 H+H+ H2H2 HCOO - CH 2 O pE = E h / 0.0591 Oxidized species (TEAs)Reduced species

12. Another Representation of The Redox Ladder H2OH2O H2H2 O2O2 H2OH2O NO 3 - N2N2 MnO 2 Mn 2+ Fe(OH) 3 Fe 2+ SO 4 2- H2SH2S CO 2 CH 4 Oxic Sub-oxic anaerobic Sulfidic Methanic Aerobes Denitrifiers Manganese reducers Sulfate reducers Methanogens Iron reducers The redox-couples are shown on each stair-step, where the most energy is gained at the top step and the least at the bottom step (i.e. the Gibb’s free energy of reaction becomes more positive going down the steps).

13. Half Reactions Often split redox reactions in two: –oxidation half rxn  e- leaves left, goes right Fe 2+  Fe 3+ + e- –Reduction half rxn  e- leaves left, goes right O 2 + 4 e -  2 H 2 O SUM of the half reactions yields the total redox reaction 4 Fe 2+  4 Fe 3+ + 4 e- O 2 + 4 e -  2 H 2 O 4 Fe 2+ + O 2  4 Fe 3+ + 2 H 2 O

14. Steps for Balancing Redox Reactions 1.Indentify principle reactants and products 2.Balance atoms other than Hydrogen and Oxygen 3.Balance oxygen using H 2 O 4.Balance H using H + 5.Balance Charge with electrons

15. 7. Multiply each half cell by an integer so that both half cells contain same number of electrons 8. Add two balanced half cells 9. H+ may be present as product of reaction. If the reaction is known to take place in an alkaline solution, then add the reaction for the dissociation of water to eliminate the H + form the overall redox reaction

16. Examples Write the half reactions corresponding to each of these 2 reactions and show the balanced overall redox reactions Mn (IV) + H 2 S  Mn 2+ + S 0 + H+ H 2 S + O 2  S 8 + H 2 O

17. Example Redox Impact on the Aquatic Cycling of Iron

18. Efficiency of Thermodynamic Predictions Measured E h Vs Calculated E h in Acid Mine Waters ACS, 1979

19. Limitations of Thermodynamic Predictions Measured E h Vs Calculated E h in Groundwaters From Lindberg and Runells, 1984 (Science)