{"id":514,"date":"2019-03-10T12:16:12","date_gmt":"2019-03-10T12:16:12","guid":{"rendered":"http:\/\/esp16.epgpbooks.inflibnet.ac.in\/?post_type=chapter&p=514"},"modified":"2019-04-25T07:36:41","modified_gmt":"2019-04-25T07:36:41","slug":"aquatic-redox-chemistry","status":"publish","type":"chapter","link":"https:\/\/ebooks.inflibnet.ac.in\/esp16\/chapter\/aquatic-redox-chemistry\/","title":{"rendered":"Aquatic Redox Chemistry"},"content":{"raw":"
\"epgp<\/a>\r\n<\/span><\/div>\r\nContents<\/strong>\r\n
    \r\n \t
  1. Introduction<\/li>\r\n \t
  2. Half Reactions<\/li>\r\n \t
  3. Reduction Potential<\/li>\r\n \t
  4. Standard Reduction Potential<\/li>\r\n \t
  5. Interpretation and Significance<\/li>\r\n \t
  6. Cell Reaction and Cell Potential<\/li>\r\n \t
  7. Hydrogen Electrode and determination of Electrode Potential<\/li>\r\n \t
  8. pE-Scale<\/li>\r\n \t
  9. Significance of pE Values<\/li>\r\n \t
  10. Measurement of pE-Values<\/li>\r\n \t
  11. Solved Problem<\/li>\r\n \t
  12. pE and pH Relationship<\/li>\r\n \t
  13. Pourbaix diagram<\/li>\r\n \t
  14. Interpretation and Significance<\/li>\r\n \t
  15. Solved Problem<\/li>\r\n \t
  16. The Limits of pE and pH in Natural Waters<\/li>\r\n \t
  17. Redox Potentials in Natural Systems<\/li>\r\n \t
  18. Redox Ladder<\/li>\r\n \t
  19. Effect of redox on Metal Pollution<\/li>\r\n \t
  20. Suggesting Reading<\/li>\r\n<\/ol>\r\nIntroduction<\/strong>\r\n

    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)<\/strong> reactions<\/strong>. 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.<\/p>\r\n \r\n\r\nWhy study Redox Reactions?<\/em>\r\n\r\nThey fuel and constrain more or less all life processes.\r\n\r\nThey are a major determinant of chemical species present in natural environments.\r\n\r\n \r\n

    Redox reactions <\/strong>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<\/em><\/strong> has multidisciplinary roots (straddling mineralogy to microbiology) and interdisciplinary applications (e.g., in removal of contaminants from sediment, soil or water).<\/p>\r\n

    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.<\/p>\r\n

    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 \u2018redox\u2019<\/strong> is the occurrence of both the chemical changes: oxidation and reduction, in a chemical reaction. The redox reactions i.e. oxidation-reduction reactions<\/p>\r\n

    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 <\/strong>or redox reactions. <\/strong>The potential of the overall reaction at the standard state<\/p>\r\n

    \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0+ E0red can be depicted as E0\u00a0 = E0 ox\u00a0 \u00a0 (where E0 ox \u00a0and E0red\u00a0 are the potentials of oxidation half-reaction and reduction half-reaction respectively)<\/p>\r\n\"\"\r\n

    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<\/strong> or reductants<\/strong>. Similarly, those which are oxygen providers are called oxidizing agents<\/strong> or oxidants.<\/strong><\/p>\r\nThe modern electronic concept of oxidant<\/strong> and reductant<\/strong> is: one which accepts electrons is oxidant<\/strong> and one which donates electron is reductant.<\/strong>\r\n\r\nFor example:\u00a0 \u00a0 \u00a0 \u00a0Cl2\u00a0 + 2 I\u00a0 \u00a0------>\u00a0 \u00a02 Cl\u00a0\u00a0\u00a0 + I2\r\n\r\nHere, Cl2 is an oxidant and I\u00a0 is a reductant; as I\u00a0 is oxidized to I2 by Cl2 and Cl2 is reduced to Cl\u00a0 \u00a0by I ).\r\n\r\n\"\"\r\n

    Thus, in a redox reaction<\/strong> which is brought about by loss and gain of electrons simultaneously, the oxidant is reduced and the reductant oxidized with an exchange of n<\/em> electron (e ) which may be depicted as:<\/p>\r\n\u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0Oxidant\u00a0 + n<\/em> e\u00a0 ------->\u00a0 reductant\r\n\r\nSome examples of Reductant in wetland soil are:\r\n

      \r\n \t
    1. Organic matter & other organic compounds;<\/li>\r\n \t
    2. Reduced inorganic compounds, viz, Mn2+, S2 , CH4, H2, Fe2+, NH4 etc. Oxidants are inorganic compounds, viz, O2, NO3 , FeOOH, SO42 , HCO3 etc.<\/li>\r\n<\/ol>\r\n

      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\u00b0. Since the redox potential depends on the members of the pair in a reaction, hence it\u2019s 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\u00b0 = zero volts). The greater the positive potential, the more anticipated it will be reduced.\u00a0The 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\u2019s free energy G) of redox-reaction at standard states as:\u00a0 G0 (eV) =\u00a0\u00a0\u00a0 nF\u00a0\u00a0 E\u00b0<\/p>\r\n

      where F is the Faraday constant number (96,485 C\/mol or 96,485 J\/mol\/V or \u2248100 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).<\/p>\r\n

      For non-standard redox-reactions, the difference in redox-potential (precisely reduction potential E) is correlated to G as: \u0394G (eV) = -nF E<\/p>\r\nHalf-Reactions<\/strong>\r\n

      A half reaction <\/strong>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.\u00a0<\/strong><\/p>\r\n\"\"\r\n\r\n \r\n\r\nSimilarly, adding the simultaneously occurring two half reactions of metallic zinc atoms reaction with aqueous nickel ions give the net redox-reaction\r\n\r\n\"\"\r\n\r\nThe half-reactions depict the exact oxidation state changes happening in two half-reactions.\r\n\r\nReduction Potential<\/strong>\r\n\r\nDefinition<\/strong>\r\n

      By definition, redox-potential (reduction potential\/electrode potential) is the tendency of a substance to accept electrons. The positive (high) values of electrode potential means that the elements or ions would readily accept electrons; on the other hand, negative (low) values indicate their easy capability of electron donation.<\/p>\r\nFew examples are cited below:\r\n\r\n\"\"\r\n

      The standard reduction potential<\/strong> Eored is measured relative<\/em> to: 2H+(aq) + 2e H2(g) (which has an assigned value Eored= 0.00 V and hence treated as reference)<\/p>\r\n

      Measurement of redox potential, E, allows quantitative appraisal of the force and tendency of the system. It can be measured by the difference between the potential of the hydrogen electrode (or more easily, the calomel electrode) and the potential of a platinum electrode immersed in the medium.<\/p>\r\nInterpretation and Significance<\/strong>\r\n

      Redox (electron-transfer) reactions provide the energetic basis for the life process, and through this, play a decisive role in the geochemical cycle of the elements. Redox reactions are very significant in water-saturated environments, such as sediments, soils, and sludges. Actually, all aquatic organisms obtain their energy for metabolic processes from oxidation-reduction reactions. Photosynthetic organisms catalytically reduce CO2 to reduced organic matter by locking light energy, while non-photosynthetic organisms catalytically decompose the organic products of photosynthesis through energy-yielding redox reactions. One of the major elements in the terrestrial and aquatic environments \u2013 nitrogen \u2013 circulates by many microbially catalyzed redox reactions. In truth, the only non-redox process in the entire nitrogen cycle is NH3 integration with and liberation from N-containing organic matter. The movements of many other elements also involve redox reactions, such as C, Fe, and S. The oxidizing power of anaerobic environments in the biosphere is mainly controlled by five molecules. In decreasing order of energy produced, they are nitrate (NO3\u2013), manganese dioxide (MnO2), ferric hydroxide (Fe(OH)3), sulfate (SO42\u2013) and, under extreme conditions, carbohydrate (CH2O) itself.<\/p>\r\n

      Like Gibb\u2019s free energy (G), the redox potentials (E) are not absolute. The redox potentials are measured with reference to the reduction of hydrogen ions to hydrogen gas at standard state conditions (SHE), i.e., 25 degrees Celsius (\u00b0C), 1 atmospheric pressure, and one unit activity for all species. The potential of this reaction, by convention, is taken as zero. Furthermore, as the redox potentials are comprised of two parts: the oxidation potential, and the reduction potential, the potential of the overall reaction of a cell at the standard state can be\u00a0expressed as:\u00a0 \u00a0E0cel<\/sub>l = E0<\/sup>ox<\/sub> + E0<\/sup><\/span><\/p>\r\n

      A positive Eored indicate that a half-reaction will proceed in the direction indicated (reduction) when paired with the hydrogen half-reaction where as a negative Eored means that a half-reaction will move in the opposite direction indicated (oxidation).<\/p>\r\n

      The numerical values of E\u00b0ox and E\u00b0red of a substance in a half-reaction are the same with opposite in sign (until the conditions are changed). Thus, numerically, the potential for oxidation half-reaction is the negative of the potential of the reduction half-reaction, i.e. E\u00b0ox = E\u00b0red (with opposite sign). The voltage produces by an electrochemical cell is determined by summing up all the potentials in circuit, as:<\/p>\r\n\"\"\r\n\r\n\"\"\r\n

      The value of redox potential (E) (for non-standard states) under concentration conditions can be related to standard potential (E0) by the Nernst equation<\/strong>:<\/p>\r\nE = E0 + RT\/nF * ln {(Ox)\/(Red)} = E0 RT\/nF * ln {(Red)\/(Ox)} = E0 RT\/nF * ln Q\r\n\r\nOr at 298 K, this is expressed as: E = E0 0.0592\/n * log {(Red)\/(Ox)}\r\n

      where, E0 = standard electrode potential, R = ideal gas constant (8.314 J\/mol-K); T = absolute temperature (Kelvin); n = number of electrons involved in the reaction; F = the Faraday constant number and (Red)\/(Ox) is the concentration ration of reduced and oxidized forms of a given reaction pair, i.e. the reaction quotient.<\/p>\r\n

      The important factors upon which the redox potential of a substance depends are: nature of the substance, its affinity for electron, concentration of reductants and oxidants (referred as redox pair) and temperature.<\/p>\r\n \r\n\r\nSolved Problem <\/strong>(based on Nernst equation):<\/strong>\r\n\r\nFor the half reaction 2H+ + 2e\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 = H2. What is E1\/2 at pH 5 and PH2 = 1 atm?\r\n\r\nSolution: <\/strong>pH =\u00a0\u00a0\u00a0 log [H+] = 5, therefore [H+] = 10-5 M.\r\n\r\nIn the above half-reaction, n = 2.\r\n\r\nPutting these values in the Nernst equation: E = E0 0.0592\/2 * log PH2\/(H+)2 E = E0 0.0592\/2 * log (105)2\r\n\r\nSince, for the given half-reaction, the E01\/2 = 0 V, hence, E = 0.000\u00a0 0.0592\/2 * (10)\r\n\r\nTherefore, E =\u00a0\u00a0 0.296 V.\r\n

      Many electrode combinations are possible in electrochemical cells, and it is convenient to specify a standard potential, Eo, for each electrode by referencing it to the SHE (whose standard potential is defined as zero). There are many half-reactions whose electrode potential cannot be measured, because the electron transfer reaction at an electrode is too slow.<\/p>\r\nThese potentials can nevertheless be calculated from the free energy of appropriate redox reactions\r\n\r\nFor example, the formation of NO from N2 and O2 is a redox reaction:\r\n\r\n\"\"\r\n

      From the free energy of the overall reaction 173.4 kJ, we obtain a cell potential of -0.45 V (using equation G <\/em>= \u2013n<\/em>F E). Then knowing that the standard potential of the oxygen electrode is 1.24 V, we can readily evaluate the standard potential for half reaction as 1.69 V, (1.24 V \u2013[\u20100.45 V]), even though it is impossible to measure this potential directly because the electron transfer between the electrode and NO and N2 molecules is too slow to establish a reversible potential.<\/p>\r\nSignificance of Redox Electrodes, Redox Potentials & Redox Reactions in the determination of:<\/strong>\r\n\r\n \r\n