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REDOX

REDOX. Objective To 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. References (additional background to Mannahan; Sawyer et al )

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REDOX

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  1. REDOX • Objective • To 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. • References(additional background to Mannahan; Sawyer et al) • Holum J.R. Fundamentals of General, Organic and Biological Chemistry • Dickson T.R. Introduction to Chemistry

  2. Atoms, Electrons and Bonds • Atoms have Protons, Neutrons and Electrons. • Electrons 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

  3. Atoms, Electrons and Bonds • The Configuration of atoms and the electron numbers make certain atoms behave similarly. GROUP Element Electrons • Alkaline metals Li, Na, K, +1 • Alkaline earths Be, Mg, Ca, Sr +2 • Transition metals Fe, Mn, Cr, Mo mid way • Non-metals N, P, S mid way • Halogens F, Cl, Br, I -1 • Noble Gases He, Ne, Ar 0

  4. Atoms, Electrons and Bonds • Basis of these properties is the requirement to satisfy a full complement of electrons in the outer shell. • Tendancy 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.

  5. Oxidation • Combination of an element or molecule with Oxygen. • H2 + 1/2 O2 = H2O • Extended to include reactions involving the loss of an Electron. • Ag Ag+ + e-

  6. H H Oxidation Number • Definition 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 F F + -1 +1

  7. Oxidation Number Rules 1. Elemental forms have oxidation number of zero. • e.g. H2, Cl2, N2, Fe (metal) 2. The oxidation number of monatomic ions equals their charge. • e.g. Na+, K+ are +1; Ca2+, Cu2+are +2; Cl- is -1. 3. In their compounds the oxidation number of any atom of: Group IA is +1 (Na+, K+ etc.); Group IIA is +2 (Ca2+ Mg2+, etc)

  8. 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 Br3, 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 NO3-, N is +5, O is -2 (-2 x 3 = -6), sum = -1

  9. Oxidation and Reduction • Oxidation is the increase in oxidation number during a reaction. Cu2+ + Fe Cu + Fe2+ +2 0 0 +2 Iron has been oxidized Copper has been reduced In this Reaction Cu2+is an Oxidizing Agent, it causes the Iron to be Oxidized (lose e-). Iron is a Reducing Agent, it causes the Cu2+to be Reduced (gain e-).

  10. Oxidising and Reducing Agents Reaction ProductsReducing AgentOxidizing Agent 2 Na + Cl2 2 NaCl Na Cl2 2 K + H2 2 KH K H2 4 Li + O2 2 Li2O Li O2 2 Na + O2 Na2O2 Na O2 2 Na + 2 H2O 2 Na+ + 2 OH- + H2 Na H2O 2 Mg + O2 2 MgO Mg O2 3 Mg + N2 Mg3N2 Mg N2 Ca + 2 H2O Ca2+ + 2 OH- + H2 Ca H2O 2 Al + 3 Br2 Al2Br6 Al Br2 Mg + 2 H+ Mg2++ H2 Mg H+ Mg + H2O MgO + H2 Mg H2O

  11. Reactivity Series (metals) • Cu2+ and Fe will react. • Cu2+ + Fe Cu + Fe2+ • Cu2+ SO42- + Fe Cu + Fe2+ SO42- • Will Fe2+ 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 (Eo) sometimes called Standard Reduction (Redox) Potentials.

  12. the Electrochemical Cell • Couples of reactive ions can be made to release some of the electron energy for useful work. • Cu/Cu2+ = + 0.34 • Zn/Zn2+ = - 0.76 • Cell = 0.34 - (-0.76) = 1.1V mV Salt Bridge Cu Zn Cu2+ Zn2+

  13. Electrochemical Iron Oxidation • Iron corrosion Fe + O2 + H+ Fe2+ + H2O • Sacrificial Protection (Zn plate, Galvanized) • Zn + Fe2+ Zn2+ + Fe • Because Fe2+ + 2e- Fe has the more positive Eo, it will go as a reduction reaction and Zn2+ + 2e- Zn will go in reverse (oxidation).

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

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

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