Oxidation-Reduction Reactions

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 = NO3- 6- + x = -1 • 2O2- + Nx = NO2- 4- + x = -1 N = +5 N = +3

What are the oxidation states of H2S and SO42-? • 2H+ + Sx = H2S 2+ + x = 0 • 4O2- + Sx = SO42- 8- + x = -2 S = -2 S = +6

Oxidation Reactions • Oxidation can be thought of as involving molecular oxygen • 3Fe2O3 2Fe3O4 + 1/2O2 (hematite) (magnetite) All as Fe3+ One as Fe2+ + two as Fe3+ High O/Fe ratioLower O/Fe ratio Oxidized 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- 2Fe3O4+ H2O • LEO – lose electron oxidation – the Fe3+ is oxidized • GER – gain electron reduction – the Fe2+ is reduced • OIL – oxidation is loss • RIG – reduction is gain

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: • Here Zn solid releases electron, which is consumed by dissolved Cu2+. Zn(s) + Cu2+(aq) Cu(s) + Zn2+(aq)

Physical model of process Ammeter e- Salt bridge – keeps charge balance in solution. e- cations anions Dissolves Precipitates Increases Decreases

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: Half reactions Zn = Zn2+(aq) + 2e- 2e- + Cu2+(aq) = Cu Zn + Cu2+(aq) = Zn2+(aq) + Cu

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

Next – balance elements (oxidation half reaction: • And charge: 2I- I2 2I- I2 + 2e-

Balance reduction half reaction • First balance oxygen, then add H+ to balance hydrogen, then add electrons for electrical neutrality: H2O2 H2O H2O22H2O 2H++ H2O2 2H2O 2e- + 2H+ + H2O2 2H2O

Combine two half reactions to get net reactions: 2I- I2 + 2e- 2e- + 2H+ + H2O2 2H2O 2H+ + 2I- + H2O2 2H2O + I2 Flow of electrons – Oxygen is electron acceptor, reduced; I- is electron donor, oxidized

Common reaction in natural waters is reduction of Fe3+ by organic carbon • With half reactions: 4Fe3+ + C + 2H2O 4Fe2+ + CO2 + 4H+ 4Fe3+ + 4e- 4Fe2+ C + 2H2O 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 can’t be measured • Difference in electrical potential between metal electrode and solution is zero • DGfº of H+ = 0 • DGfº of e- = 0

SHE Half reaction in solution: H+ + e- = 1/2H2(g) By definition, aH+ = 1 Allows electrons to flow but chemically inert

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

In cell A, platinum wire is inert – transfers electrons to or from solution only. • If wire has no source of electrons • Pt wire develop an electrical potential – “tendency” for electrons to enter or leave solution • Define the potential as “activity of electrons” = ae- • 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: • If reaction goes to left, Fe2+ gives up e- • If reaction goes to right, Fe3+ acquires e- • If no source or sink of e-, (switch open), volt meter measures the potential (tendency) Fe3+ + e- = Fe2+

Since we have a reaction • can write an equilibrium constant Fe3+ + e- = Fe2+ aFe2+ Keq = aFe3+ ae-

aFe2+ • Rearranged: • 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 ae-= Keq-1 aFe3+

Consider half cell B: • Direction of reaction depends on tendency for wire to gain or lose electrons • Equilibrium constant H+ + e- = 1/2H2(g) 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: • Direction of reaction depends on which half cell has highest activity of electrons Flow of electrons Fe3+ + 1/2H2(g) =Fe2+ + H+

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 > or < that of SHE

Convention • Eh > 0 if ae- of the half cell < SHE • I.e. if electrons flow from the SHE to the fluid • For thermodynamics: • Is equivalent to: Fe3+ + 1/2H2(g) =Fe2+ + H+ Fe3+ + e-=Fe2+

Expressions for activities of electrons: • Eh or pe • pe = [F/(2.303RT)]*Eh • @ 25ºC, pe = 16.9 Eh; Eh = 0.059pe • F = Faraday’s constant = 96,485 coul/mol • Coulomb = charge /electron = quantity of electricity transferred by 1 Amp in 1 second.

Oxidation-Reduction Reactions