Oxidation reduction reactions BY Muhammad Fahad Ansari 12IEEM14


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Oxidation reduction reactions BY Muhammad Fahad Ansari 12IEEM14

  1. 1. Electrochemistry• MuhaMMad Fahad ansari • 12iEEM14
  2. 2. Electrochemistry• One of the major classes of chemical reactions involves an electron exchange between the reactants.• These reactions are called oxidation-reduction reactions or electrochemical, reactions. they can be further subdivided into two types or classes: Reactions those that produce electrical energy and those that are produced by electrical energy. One-type of reaction is spontaneous and may be used to produce electricity; The-second must be forced to occur and requires electricity to make it take place
  3. 3. Spontaneous Chemical Reactions Produce Electrical-EnergyFor example:• The electrical energy used to start a car is produced by an oxidation-reduction reaction in the car’s battery.• Similarly, the electricity that causes the light in a flash light comes from electrochemical reactions in the dry cell of the flashlight.
  4. 4. Non spontaneous Chemical Reaction Produced by Electrical-EnergyThe type of electrochemical reactions requireelectrical energy, or the use of electricalenergy to bring about chemical reactions,which is called electrolysis.Electroplating is another importantapplication in which electricity is used to bringabout a chemical reaction.
  5. 5. Electrolysis Many of the active elements (those with strong tendency to lose or gain electrons) are available only as the result of electrolysis.For example: Na and Cl found in nature only in the form of their ions (Na+, Cl-) in compounds.• Electricity is used to take an electron away from the Cl- ion and their by re-form the atom, (Cl).• Similarly, electricity is required to force the Na+ to gain an electron to become a sodium (Na) again.
  6. 6. Electroplating• Electroplating involves electrolysis that is carried out so that the ions of metal are converted into atoms on the surface to be plated.• Cu (s)  Cu (aq) + 2e- (half equation/reaction)• Ag+ (aq) + e-  Ag (s) (half equation/reaction)• Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)• Cu(s)RA + 2Ag+(aq)OA <===> Cu2+(aq)OA + 2Ag(s)RA
  7. 7. Redox Reactions• Some common redox reactions are those that occur in batteries, when metals rust, when metals are plated from solutions, and• Combustion of organic molecules such as hydrocarbons (like methane and gasoline) and carbohydrates (like wood).
  8. 8. Basic Oxidation-Reduction Reactions ConceptsChemical reactions involves an electronexchange between the reactants, in which onesubstance loses electrons and anothersubstance gains electrons.
  9. 9. • Oxidation: is the loss of electrons by atoms, ions, or molecules.• Reduction: is the gain of electrons by atoms, ions, or molecules.• Oxidizing agent: Substance that causes another substance to lose electrons• Reducing agent: Substance that loses electrons in a chemical reaction; it reduces another substance by losing electrons to it.
  10. 10. Reactions Reactants Product 2Na + Cl2 → 2 Na+ Cl-Loses GainsOxidized ReducedReducing agent Oxidizing agent
  11. 11. Electron transfer or oxidation/reduction reactionsOr transfer of charge - an electron - from one species to another• 2Mg(s) + O2(g)  2MgO(s)• Fe(s) + O2(g)  Fe oxides (s)• C(s) + O2(g)  CO2(g)
  12. 12. CH4 + O2  H2O + CO2 Since H has an oxidation # of 1+, the oxidation # of C in CH4 is 4-, while in CO2 it is 4+.Clearly C has been oxidized by the oxidizing agent O2.O2 has been reduced to form both products.
  13. 13. Step-wise oxidations of the carbon. CH4 + O2  H2O + CO2the oxidation numbers for the central C in eachmolecule are 4-, 2-, 0, 2+, and 4+ as you proceedfrom left to right
  14. 14. Oxidation Number or Oxidation StateSimple and Arbitrary Rules:• The oxidation number of an element in its free or uncombined form is zero. For example: Na0, Mg0, S0, O20, and so on.• The oxidation number of mono-atomic ions equals the number of electrons it has lost or gained. For example: Na1+, Mg2+, Al3+, Cl1-, S2-, and so on.
  15. 15. • The oxidation number of oxygen in compounds is usually -2. The exceptions are the peroxides, such as H2O2, and compounds of oxygen and fluorine, such as OF2.• The oxidation number of hydrogen in compounds is usually +1. The hydrides, such as NaH, are exceptions.• In the formula for a compound, the sum of the positive oxidation numbers must equal the sum of the negative oxidation numbers. For example: Mg2+S2-, Na1+Cl1-, K1+Mn7+O48-, H22+S6+O48-, and so on.• In complex ions such as SO42-, PO43-, ClO31-, the algebraic sum of the oxidation numbers of the individual atoms in the ion equals the charge on the ion.
  16. 16. ExamplesThe oxidation number of S in SO42- , Cl in ClO31- and P in PO43- are• SO42- x +4(-2) = -2 (net charge on ion) x -8 = -2 x = 8 -2 = 6• ClO31- x + 3(-2) = -1 (net charge on ion) X-6=-1 X=6–1=5• PO43- x + 4(-2) = -3 (net charge on ion) X-8=-3 X= 8 – 3 = 5
  17. 17. Table-Oxidation States of Some Chemical Elements Element Oxidation state Species Formula -2 Hydrogen sulfide H2S 0 Elemental sulfur S Sulfur +4 Sulfur dioxide SO2 +6 Sulfate ion SO42- -4 Methane CH4 0 Soot, graphite C Carbon +2 Carbon monoxide CO +4 Carbon dioxide CO2 -3 Ammonia NH3 0 Nitrogen gas N2 +2 Nitric oxide NONitrogen +3 Nitrite ion NO2- +4 Nitrogen dioxide NO2 +5 Nitrate ion NO3- -2 Almost all compounds -Oxygen -1 Hydrogen peroxide H2O2 0 Oxygen gas O2 0 Hydrogen gas H2Hydrogen +1 Hydrogen ion H+ -1 Chlorine ion Cl- 0 Chlorine gas Cl2Chlorine +1 Hypochlorous acid HOCl +7 Perchloric acid HClO4
  18. 18. Microbial Redox Process• Important redox reactions that are carried out by microorganisms are summarized here: The notation [CH2O] is used to denote a fragment of an arbitrary carbohydrate.
  19. 19. Photosynthetic Production of Biomass• Photosynthetic microorganisms (algae and some bacteria) carryout photosynthesis reactions, in these reactions, energy-rich carbohydrate molecules are produced by combining carbon dioxide and water, using energy derived from sunlight.• From a Redox Perspective: CO2 + H2O  [CH2O] + O2 Carbon is reduced from oxidation state +4 to 0, and oxygen is oxidized from -2 to 0
  20. 20. Aerobic Respiration:In the presence of oxygen, microorganisms degradebiomass to form carbon dioxide and water.Chemical energy that is released can be used by theorganisms. [CH2O] + O2  CO2 + H2OThis process is the reverse of photosynthesis, carbonis oxidized and oxygen is reduced.
  21. 21. Sources of Nitrogen• Atmosphere N2 gas• Water NH3, NH4, NO2, NO3 ions• Soil NH4+ , NO2, NO3 ions• Plants Proteins contains (-NH2) Amino group (O=C-OH) Carboxyl group• Animals Proteins, Urea, and Amines Urea contains (NH2-C=O-NH2)
  22. 22. Nitrogen Fixation• In the atmosphere, nitrogen is almost entirely in the form of N2 and is in oxidation state 0.• The nitrogen in biological system is mostly in the form of an amine –NH2, Which is very closely related to ammonia (NH3) and ammonium ion (NH4+) here nitrogen is in oxidation state -3.• Nitrogen in water and soil is in form of nitrate (NO3-) in which nitrogen is in oxidation state +5.
  23. 23. Microorganisms play an essential role in the movement ofnitrogen among these oxidation states.Compounds such as ammonia and nitrate contain asingle nitrogen atom as fixed nitrogen species.Certain groups of bacteria are capable of converting gaseousnitrogen to fixed nitrogen, in the form of the ammonium ion.Energy from the oxidation of biomass to CO2 is used to reducethe nitrogen in N2 to ammonium.3[CH2O] + 2N2 + 3H2O + 4H+  3CO2 + 4NH4+
  24. 24. NitrificationIn the nitrification, nitrogen in the ammonium ion isoxidized from -3 to +5, with oxygen as oxidizer. NH4+ + 2O2  NO3- + 2H+ + H2OPlants absorb nitrogen more efficiently in the form ofnitrate than an ammonium, so redox reaction canenhance the effectiveness of ammonia-basedagricultural fertilizers.
  25. 25. Nitrate Reduction or DenitrificationWhen oxygen is not available as the oxidizer todegrade biomass, microorganisms can use nitrate asthe oxidizer (electron acceptor).Nitrate Reduction is used in some wastewatertreatment systems to convert fixed nitrogen to N2 gas,which can then be safely released to the atmosphere.This process is called denitrification, since nitrogen isremoved from the aqueous system.Nitrogen in municipal wastewater begins in a reducedstate (-3), the overall process involved two steps:
  26. 26. Nitrification in an aerobic reactor, followed bydenitrification in an anaerobic reactor, four nitrogen atoms,being reduced from +5 to 0, can fully oxidize five carbonatoms from 0 to +4 NH4+ + 2O2  NO3- + 2H+ + H2O (aerobic reaction)5[CH2O] + 4NO3- + 4H+  5CO2 + 7H2O + 2N2 (anaerobic reaction)
  27. 27. Sulfate Reduction• Some environments that contain biodegradable materials lake both oxygen and nitrate to serve as the oxidizing agent, in such cases, sulfate may serve that role.• The conversion of one sulfur atom from +6 in sulfate to -2, in hydrogen sulfide oxidizes two carbon atoms from 0 to +4 oxidation states. 2[CH2O] + 2H+ + SO42-  2CO2 + 2H2O + H2S• This reaction can occur in stagnant anaerobic marine sediments that are supplied with decaying biomass, algae or seaweed accumulation.
  28. 28. Methane Formation (Methanogenesis)• In the absence of oxygen, nitrate, and sulfate, biomass can still be converted to carbon dioxide as: 2[CH2O]  CO2 + CH4• This is an interesting redox reaction, since the two carbon atoms begins in oxidation state zero (0). One carbon atom is oxidized to +4, and the other is reduced to -4.• Methane generation process is exploited in seawater treatment to convert excess microbiological material to gases, which are more easily handled for disposal.
  29. 29. Sequence Of Redox Reactions (CH2O = Unidentified Organic Matter With Zero-valent Carbon)Aerobic respiration CH2O + O2 ↔ CO2 + H2ODenitrification 5CH2O + 4NO3– ↔ 2N2 + 4HCO3– + CO2 + 3H2OManganese (IV) CH2O + 2MnO2 + 3CO2 + H2O ↔ 2Mn2+ + 4HCO3–reductionIron (III) reduction CH2O + 4Fe(OH)3 + 8H+ ↔ 4Fe2+ + 8HCO3– + 3H2OSulphate reduction 2CH2O + SO42– + H+ ↔ H2S + 2HCO3–Methane fermentation 2CH2O ↔ CH4 + CO2
  30. 30. Thank you,well come for questions