REDOX vObjective vTo understand the concept of Oxidation-Reduction (Redox), Oxidation Numbers, half reactions in chemical reactions, and know the main examples of Redox reactions which are important to Environmental Engineering. vReferences (additional background to Mannahan; Sawyer et al)  Holum J.R. Fundamentals of General, Organic and Biological Chemistry  Dickson T.R. Introduction to Chemistry

Atoms, Electrons and Bonds vAtoms have Protons, Neutrons and Electrons. vElectrons are in orbitals or levels.  These become full with 2, 8, 8, 18 ……electrons  Partly filled orbitals are energetically unfavourable.  Whenever possible, Electrons are gained or lost to achieve the above configurations. electron Proton neutron

Atoms, Electrons and Bonds vThe Configuration of atoms and the electron numbers make certain atoms behave similarly. GROUPElementElectrons vAlkaline metalsLi, Na, K, +1 vAlkaline earthsBe, Mg, Ca, Sr+2 vTransition metalsFe, Mn, Cr, Momid way vNon-metalsN, P, Smid way vHalogensF, Cl, Br, I-1 vNoble GasesHe, Ne, Ar0

Atoms, Electrons and Bonds vBasis of these properties is the requirement to satisfy a full complement of electrons in the outer shell. vTendancy to either: 1.want more electrons (Electronegativity) 2.want to lose electrons Electronegativity generally increases L to R and bottom to top in the periodic table.

Oxidation vCombination of an element or molecule with Oxygen.  H 2 + 1/2 O 2 = H 2 O vExtended to include reactions involving the loss of an Electron.  Ag Ag + + e -

Oxidation Number vDefinition Oxidation number is the charge an atom would have in a compound if the electrons in each bond belonged to the more Electronegative atom. Example HF H F F + +1 H

Oxidation Number Rules 1.Elemental forms have oxidation number of zero.  e.g. H 2, Cl 2, N 2, Fe (metal) 2.The oxidation number of monatomic ions equals their charge.  e.g. Na +, K + are +1; Ca 2+, Cu 2+ are +2; Cl - is In their compounds the oxidation number of any atom of: Group IA is +1 (Na +, K + etc.); Group IIA is +2 (Ca 2+ Mg 2+, etc)

Oxidation Number Rules 4.The oxidation number of any non-metal in its binary compounds with metals, equals the charge of the monatomic ion.  e.g. in Cr Br 3, Br has oxidation number -1, (like Br - ). 5.In compounds the oxidation number of: Oxygen is almost always -2 Hydrogen is almost always +1 F is always -1 6.Sum of oxidation numbers in an ion equals the charge of the ion.  e.g. in NO 3 -, N is +5, O is -2 (-2 x 3 = -6), sum = -1

Oxidation and Reduction vOxidation is the increase in oxidation number during a reaction. Cu 2+ + Fe Cu + Fe Iron has been oxidized Copper has been reduced In this Reaction Cu 2+ is an Oxidizing Agent, it causes the Iron to be Oxidized (lose e - ). Iron is a Reducing Agent, it causes the Cu 2+ to be Reduced (gain e - ).

Oxidising and Reducing Agents ReactionProductsReducing AgentOxidizing Agent 2 Na + Cl 2 2 NaClNa Cl 2 2 K + H 2 2 KHK H 2 4 Li + O 2 2 Li 2 OLi O 2 2 Na + O 2 Na 2 O 2 NaO 2 2 Na + 2 H 2 O 2 Na OH - + H 2 Na H 2 O 2 Mg + O 2 2 MgOMgO 2 3 Mg + N 2 Mg 3 N 2 MgN 2 Ca + 2 H 2 O Ca OH - + H 2 CaH 2 O 2 Al + 3 Br 2 Al 2 Br 6 Al Br 2 Mg + 2 H + Mg 2+ + H 2 MgH + Mg + H 2 OMgO + H 2 MgH 2 O

Reactivity Series (metals) vCu 2+ and Fe will react.  Cu 2+ + FeCu + Fe 2+  Cu 2+ SO FeCu + Fe 2+ SO 4 2- vWill Fe 2+ and Cu react ? No. Why not  Need to consider the half Reactions.  Iron’s tendancy to lose electrons is greater than Copper’s. So Iron wins.  These properties can be found from tables of Standard Electrode Potentials (E o ) sometimes called Standard Reduction (Redox) Potentials.

the Electrochemical Cell vCouples of reactive ions can be made to release some of the electron energy for useful work.  Cu/Cu 2+ =  Zn/Zn 2+ =  Cell = (-0.76) = 1.1V mV Cu 2+ Zn 2+ ZnCu Salt Bridge

Electrochemical Iron Oxidation vIron corrosion Fe + O 2 + H + Fe 2+ + H 2 O vSacrificial Protection (Zn plate, Galvanized)  Zn + Fe 2+ Zn 2+ + Fe  Because Fe e - Fe has the more positive E o, it will go as a reduction reaction and Zn e - Zn will go in reverse (oxidation).

Nernst Equation vA measured Electrode Potential will take account of the concentrations of the half-reaction species. vEnvironmental Redox Levels Can be measured by a Platinum electrode against a reference half- reaction. vEnvironmental concentrations are small, so the value will drift as the reading is taken.

Electron Activity pE vthe concept of pE is analagous to pH. vIt is a reflection of the electron activity. pE = - log (a e ) pE = 16.9 E(at 25C) vIn practice environmental pE ranges range from: > 10 (Oxidising conditions, aerobic) to< -5 (Reducing conditions, anaerobic) in other words(E = +0.8V to - 0.4V)