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Development Team Paper No: 16 Environmental Chemistry Module: 23 Aquatic Redox Chemistry Development Team Prof. R.K. Kohli Principal Investigator & Prof. V.K. Garg & Prof. Ashok Dhawan Co- Principal Investigator Central University of Punjab, Bathinda Prof. K.S. Gupta Paper Coordinator University of Rajasthan, Jaipur Dr. Alka Sharma Content Writer University of Rajasthan, Jaipur Content Reviewer Prof. K.S. Gupta University of Rajasthan, Jaipur Anchor Institute Central University of Punjab 1 Environmental Chemistry Environmental Sciences Aquatic Redox Chemistry Description of Module Subject Name Environmental Sciences Paper Name Environmental Chemistry Module Name/Title Aquatic Redox Chemistry Module Id EVS/EC-XVI/23 Pre-requisites A basic knowledge of redox chemistry 1. To define redox (e- transfer) reactions, half reactions 2. To define and understand aquatic redox chemistry 3. To define redox potential, cell reactions and cell potential 4. To define standard electrode potential E0 and Nernst equation Objectives 5. To understand pE-scale, significance and measurement of pE values 6. To understand redox ladder 7. To understand redox potential in natural water, its interpretation and significance 8. To define and understand applications of redox potentials in aquatic systems Redox reactions, redox process, redox potential, pE-scale, redox ladder, aquatic redox Keywords system 2 Environmental Chemistry Environmental Sciences Aquatic Redox Chemistry Module 23: Aquatic Redox Chemistry Contents 1. Introduction 2. Half Reactions 3. Reduction Potential 4. Standard Reduction Potential 5. Interpretation and Significance 6. Cell Reaction and Cell Potential 7. Hydrogen Electrode and determination of Electrode Potential 8. pE-Scale 9. Significance of pE Values 10. Measurement of pE-Values 11. Solved Problem 12. pE and pH Relationship 13. Pourbaix diagram 14. Interpretation and Significance 15. Solved Problem 16. The Limits of pE and pH in Natural Waters 17. Redox Potentials in Natural Systems 18. Redox Ladder 19. Effect of redox on Metal Pollution 20. Suggesting Reading 3 Environmental Chemistry Environmental Sciences Aquatic Redox Chemistry Introduction In aquatic environmental systems, including soils, sediments, aquifers, rivers, lakes, and water treatment systems, the most important and interesting chemical reactions occurring are the oxidation-reduction (redox) reactions. These reactions are central to major element cycling, to many sorption processes, to trace element mobility and toxicity, to most remediation schemes, and to life itself. Why study Redox Reactions? They fuel and constrain more or less all life processes. They are a major determinant of chemical species present in natural environments. Redox reactions are core to many emerging domains of the aquatic sciences research, which includes all aspects of the aquatic sciences: involving the hydrosphere, aquatic (i.e., aqueous) aspects of environmental processes in the atmosphere, lithosphere, biosphere, etc. The Aquatic Redox Chemistry has multidisciplinary roots (straddling mineralogy to microbiology) and interdisciplinary applications (e.g., in removal of contaminants from sediment, soil or water). To characterize the oxidation-reduction status of surface environments, the geochemists, soil scientists and limnologists have used redox potential (Ered) measurements. The redox potential of aquatic, marine and soil systems is a measure of electrochemical potential or electron availability within these systems. The redox potential (Ered) is determined from the concentration of oxidants and reductants in the environment. Oxygen, nitrate, nitrite, manganese, iron, sulphate, and CO2 are some of the prominent inorganic oxidants; while organic substrates and reduced inorganic compounds are the well-known reductants. The redox-potential is the evaluation of the equilibrium potential, i.e. reduction/oxidation potential, built at the interface between an electrode (a noble metal) and the solution consisting of electroactive redox species and is measured under standard state conditions (at 25 0C, 1 atmospheric pressure and one unit activity for all species) with respect to the standard hydrogen electrode. The term ‘redox’ is the occurrence of both the chemical changes: oxidation and reduction, in a chemical reaction. The redox reactions i.e. oxidation-reduction reactions 4 Environmental Chemistry Environmental Sciences Aquatic Redox Chemistry entail the changes of oxidation states of reactants in a reaction and it is the sum up of two half reactions. The two processes: oxidation (loss of electrons) and reduction (gain of electrons) takes place together during the same reaction, i.e. oxidation/reduction never takes place in isolation; hence, these reactions are called oxidation-reduction reactions or redox reactions. The potential of the overall reaction at the standard state 0 0 0 0 0 can be depicted as E = E ox + E red (where E ox and E red are the potentials of oxidation half-reaction and reduction half-reaction respectively) For example: In a reaction: Copper (II) oxide + Hydrogen heat Copper + Water Oxidation CuO + H2 Cu + H2O Reduction Consider another reaction: Zinc oxide + carbon Zinc + Carbon monoxide Oxidation ZnO + C Zn + CO Reduction In both these reactions, the hydrogen and carbon removes oxygen from copper(II) oxide and zinc oxide respectively, hence hydrogen and carbon are reducing agents or reductants. Similarly, those which are oxygen providers are called oxidizing agents or oxidants. The modern electronic concept of oxidant and reductant is: one which accepts electrons is oxidant and one which donates electron is reductant. For example: Cl2 + 2 I 2 Cl + I2 Here, Cl2 is an oxidant and I is a reductant; as I is oxidized to I2 by Cl2 and Cl2 is reduced to Cl by I ). 5 Environmental Chemistry Environmental Sciences Aquatic Redox Chemistry Reducing agent (Compound A) Oxidizing agent (Compound B) A B A oxidized (Electrons loss) B reduced (Electrons gain) A B Compound A (Oxidized Form) Compound B (Reduced Form) Thus, in a redox reaction which is brought about by loss and gain of electrons simultaneously, the oxidant is reduced and the reductant oxidized with an exchange of n electron (e) which may be depicted as: Oxidant + n e reductant Some examples of Reductant in wetland soil are: 1. Organic matter & other organic compounds; 2+ 2 2+ 2. Reduced inorganic compounds, viz, Mn , S , CH4, H2, Fe , NH4 etc. 3 2 3 Oxidants are inorganic compounds, viz, O2, NO , FeOOH, SO4 , HCO etc. It may be noted that in the overall redox reaction no free electrons are generated. The movement of electrons is from reductant to oxidant, thereby an electrical potential is developed between the two which is measured in volts and denoted by E°. Since the redox potential depends on the members of the pair in a reaction, hence it’s a relative rather than an absolute value. The standard redox potential values for elements, compounds and ions can be ascertained under standard conditions (unit molar concentration at 1 atm pressure and 25 0C), relative to a standard hydrogen electrode (SHE) potential (which is arbitrarily given a potential, E° = zero volts). The greater the positive potential, the more anticipated it will be reduced. The redox-potentials (also known as electrochemical version of Gibbs free energy G) are used to ascertain the direction and the free energy (the change in the system’s free energy G) of redox-reaction at standard states as: G0 (eV) = nFE° 6 Environmental Chemistry Environmental Sciences Aquatic Redox Chemistry where F is the Faraday constant number (96,485 C/mol or 96,485 J/mol/V or ≈100 kJ/mol/V) and n is the number of electrons involved in the redox-reaction). (The 0 symbol is for the substances involved in the reaction in their standard states). For non-standard redox-reactions, the difference in redox-potential (precisely reduction potential E) is correlated to G as: ΔG (eV) = -nFE Half-Reactions A half reaction is either the reduction or the oxidation reaction component of a redox-reaction. The reactions occurring in an electrochemical cell is frequently described on the basis of the half-reaction concept. A redox reaction is expressed as the difference of two reduction half-reactions. For example: Cu2+(aq) + 2e Cu(s) (reduction of Cu2+) (E0 = + 0.34 V) Zn2+(aq) + 2e Zn(s) (reduction of Zn2+) (E0 = - 0.76 V) Deducing the two half-reactions: Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) Or, in other words, combining the two half-reactions, reduction of Cu2+ and oxidation of Zn(s) would result into a net redox-reaction: Cu2+(aq) + 2e Cu(s) and Zn(s) Zn2+ + 2e Being half components of the redox-reaction, these reactions are termed as Half-reactions. + Similarly, the redox reaction: 2Na + Cl2 Na + 2Cl is made up of two half-reactions: + 2Na 2Na + 2e (oxidation) Cl2 + 2e 2Cl (reduction) Another example of half-reaction is the reduction of MnO4 in which atom transfer accompanies electron transfer; here the oxygen atoms are lost from MnO4 (aq) and H2O(l): + 2+ MnO4 (aq) + 8H (aq) + 5e Mn (aq) + 4H2O(l) 7 Environmental Chemistry Environmental Sciences Aquatic Redox Chemistry Similarly, adding the simultaneously occurring two half reactions of metallic zinc atoms reaction with aqueous nickel ions give the net redox-reaction: Zn(s) Zn2+(aq) + 2e 2 e + Ni2+(aq) Ni(s) -------------------------------------------------- Ni2+(aq) + Zn(s) Ni(s) + Zn2+(aq) The half-reactions depict the exact oxidation state changes happening in two half-reactions. Reduction Potential Definition By definition, redox-potential
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