1 / 37

Oxidation-Reduction Reactions

Oxidation-Reduction Reactions. Carbonate reactions are acid-base reactions: Transfer of protons – H + Other acid-base systems are similar: Sulfuric acid - H 2 SO 4 Phosphoric acid - H 2 PO 3 Nitric Acid HNO 3. Oxidation-Reduction Reactions.

sian
Télécharger la présentation

Oxidation-Reduction Reactions

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. 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

  2. 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

  3. 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

  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

  5. 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)

  6. 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

  7. 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)

  8. 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

  9. 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

  10. 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

  11. 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

  12. Generally easiest to consider reactions as transfer of electrons • Many redox reaction do not involve molecular oxygen

  13. 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

  14. 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)

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

  16. Ammeter shows flow of electrons from Zn to Cu: • Zn rod dissolves – Zn2+ increases • Cu rod precipitates – Cu2+ decreases

  17. At the rod, the reactions are: Half reactions Zn = Zn2+(aq) + 2e- 2e- + Cu2+(aq) = Cu Zn + Cu2+(aq) = Zn2+(aq) + Cu

  18. Benefits of using half reactions: • Half reactions help balance redox reactions • Used to create framework to compare strengths of oxidizing and reducing agents

  19. 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

  20. 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

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

  22. 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

  23. 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

  24. 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-

  25. 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)”

  26. 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

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

  28. 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

  29. 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

  30. 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+

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

  32. 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+

  33. 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-

  34. 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+

  35. 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

  36. 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+

  37. 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.

More Related