Hydrogeochemistry - CLAS Users

Hydrogeochemistry - CLAS Users

Oxidation-Reduction Reactions Carbonate reactions are acid-base reactions: Transfer of protons H+ Other acid-base systems are similar: Sulfuric acid - H2SO4

Phosphoric acid - H2PO3 Nitric Acid HNO3 Oxidation-Reduction Reactions Redox reactions are analogous, but are transfer of electrons rather than protons Very important class of reactions

Elements may have variable charges number of electrons (valence state) Valence state controls speciation of elements Examples of primary valence states of some elements C = +4 or -4 S +6 or -2

N +5 or +3, also +4, +2 Fe +3 or +2 Mn +3 or +2, also +7, +6, +4 Minor elements also have various valence states V, Cr, As, Mo, V, Se, Sb, W, Cu All nasty elements Important environmental controls e.g., mining

Valence state very important for mobility, as well as absorption and thus toxicity Fe3+ (oxidized) is highly insoluble Precipitate as Fe-oxide minerals (magnetite, hematite, goethite, lepidocrocite, limonite) Fe2+ (reduced) much more soluble most Fe in solution is +2 valence Common precipitates are Fe-sulfides (pyrite, marcasite)

Assignment of oxidation state Valence state of oxygen is always -2 except for peroxides, where it is -1. E.g., H2O2 and Na2O2 Valence state of hydrogen is +1 in all compounds except when bonded with metals where it is -1.

NaH NaBH4 LiAlH4 Valence state of all other elements are selected to make the compound neutral Certain elements almost always have the same oxidation state

Alkali metals = +1 (left most column) Alkaline earths = +2 (second column from left) Halogens = -1 (2nd column from right) Examples What are the oxidation states of N in NO3- and NO2-? 3O2- + Nx = NO3N = +5 6- + x = -1 2O2- + NNx = +3

NO2- 4- + x = -1 What are the oxidation states of H2S and SO42-? 2H+ + Sx = H2S S = -2 2+ + x = 0 4O2- + SSx = +6 SO42-

8- + x = -2 Oxidation Reactions Oxidation can be thought of as involving molecular oxygen 3Fe2O3 (hematite) All as Fe3+ 2Fe3O4 + 1/2O2 (magnetite) One as Fe2+ + two as Fe3+

High O/Fe ratio Oxidized Lower O/Fe ratio Reduced In this case, the generation of molecular oxygen controls the charge imbalance Also possible to write these reactions in terms of electrons: 3Fe2O3 + 2H+ + 2e-

LEO lose electron oxidation the Fe3+ is oxidized GER gain electron reduction the Fe2+ is reduced OIL oxidation is loss RIG reduction is gain 2Fe3O4+ H2O Generally easiest to consider reactions as transfer of electrons

Many redox reaction do not involve molecular oxygen Problem is that free electrons are not really defined Reactions that consume free electrons represent only half of the reaction A complementary reaction required to

produce a free electron Concept is two half reactions The half reaction simultaneously create and consume electrons, so typically not expressed in reaction Half Reactions Example of redox reaction without oxygen: Zn(s) + Cu2+(aq) Cu(s) + Zn2+(aq) Here Zn solid releases electron, which is consumed by dissolved Cu2+.

Physical model of process Ammeter eecations Dissolves anions Salt bridge keeps charge balance in solution. Precipitates Increase s

Decreas es Ammeter shows flow of electrons from Zn to Cu: Zn rod dissolves Zn2+ increases Cu rod precipitates Cu2+ decreases At the rod, the reactions are: Zn = Zn2+(aq) + 2e2e- + Cu2+(aq) = Cu Zn + Cu2+(aq) = Zn2+(aq) + Cu

Half reaction s Benefits of using half reactions: Half reactions help balance redox reactions Used to create framework to compare strengths of oxidizing and reducing agents Rules for writing and balancing half reactions

1. 2. 3. 4. Identify species being oxidized and reduced Write separate half reactions for oxidation and reduction Balance reactions using (1) atoms and (2) electrical charge by adding e- or H+ Combine half reactions to form net oxidation-reduction reactions

Consider reaction H2O2 + I- I2 + H2O First, ID oxidized and reduced species: Iodine is being oxidized from -1 to 0 charge I I2 Oxygen in peroxide

is being reduced to water H2O2 H2O Next balance elements (oxidation half reaction: 2I- I2 And charge: 2I- I2 + 2e-

Balance reduction half reaction First balance oxygen, then add H+ to balance hydrogen, then add electrons for electrical neutrality: H2O2 H2O H2O2 2H2O 2H+ + H2O2 2H2O

2e- + 2H+ + H2O2 2H2O Combine two half reactions to get net reactions: 2I2e- + 2H+ + H2O2 I2 + 2e2H2O 2H+ + 2I- + H2O2 2H2O + I2 Flow of electrons Oxygen is electron acceptor, reduced; I- is electron donor, Common reaction in natural waters is

reduction of Fe3+ by organic carbon 4Fe3+ + C + 2H2O 4Fe2+ + CO2 + 4H+ With half reactions: 4Fe3+ + 4eC + 2H2O 4Fe2+ CO2 + 4H+ + 4e- From thermodynamic conventions, its impossible to consider a single half reaction

There is no thermodynamic data for e- Practically, half reactions are defined relative to a standard The standard is the Standard Hydrogen Electrode (SHE) SHE Platinum electrode in solution containing H2 gas at P = 1 Atm.

Assign arbitrary values to quantities that cant be measured Difference in electrical potential between metal electrode and solution is zero DGf of H+ = 0 DGf of e- = 0 SHE Half reaction in solution H+ + e- = 1/2H2(g) By definition,

aH+ = 1 Allows electrons to flow but Example of how SHE used E= Potential Positive or negative Fe3+ + e- = Fe2+ reaction goes to left, wire removes electrons reaction goes to right, wire adds electrons SHE: H+ + e- = 1/2H2(g)

In cell A, platinum wire is inert transfers electrons to or from solution only. If wire has no source of electrons Define the potential as activity of electrons = ae Pt wire develop an electrical potential tendency for electrons to enter or leave solution

Not a true activity, really a tendency Define pe = -logae-, similar to pH In Cell A solution, Fe is both oxidized and reduced Fe2+ and Fe3+ Reaction is:

Fe3+ + e- = Fe2+ If reaction goes to left, Fe2+ gives up eIf reaction goes to right, Fe3+ acquires eIf no source or sink of e-, (switch open), volt meter measures the potential Since we have a reaction Fe3+ + e- = Fe2+ can write an equilibrium constant Keq = aFe2+ aFe3+ ae-

Rearranged: ae-= Keq-1 aFe2+ aFe3+ ae- is proportional to the ratio of activity of the reduced species to activity of oxidized species ae- is electrical potential (in volts) caused by ratio of reduced to oxidized species

Consider half cell B: H+ + e- = 1/2H2(g) Direction of reaction depends on tendency for wire to gain or lose electrons Equilibrium constant PH21/2 KSHE = aH+ ae -

Switch closed electrons flow from one half cell to the other Electron flow from the side with the highest activity of electrons to side with lowest activities Overall reaction: Flow of electrons Fe3+ + 1/2H2(g) =Fe2+ + H+ Direction of reaction depends on which half cell has highest activity of electrons

Switch open: No longer transfer of electrons Now simply potential (E) generated at Pt wire By convention, potential of SHE (ESHE) = O Potential called Eh, i.e. E (electromotive force) measured relative to SHE (thus the h)

Eh > or < O depends on whether ae- is > Convention Eh > 0 if ae- of the half cell < SHE I.e. if electrons flow from the SHE to the fluid For thermodynamics: Fe3+ + 1/2H =Fe2+ + H+ 2(g)

Is equivalent 3+ -to: Fe + e =Fe2+ Expressions for activities of electrons:

Eh or pe pe = [F/(2.303RT)]*Eh @ 25C, pe = 16.9 Eh; Eh = 0.059pe F = Faradays constant = 96,485 coul/mol Coulomb = charge /electron =

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