1 / 15

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 )

yolanda-henderson
Download Presentation

REDOX

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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)

More Related