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Redox reactions

Redox reactions. Gain oxygen. Gain oxygen. Loss electrons. Oxidation. 2Mg(s) + O 2 (g)  2MgO(s) CH 4 (g) + O 2 (g)  CO 2 (g) + H 2 O( ) Mg(s)  Mg 2+ (aq) + 2e -. Loss oxygen. Gain electrons. Reduction. 2PbO(s) + C(s)  2Pb(s) + CO 2 (g) Pb 2+ (aq) + 2e -  Pb(s).

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Redox reactions

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  1. Redox reactions

  2. Gain oxygen Gain oxygen Loss electrons Oxidation • 2Mg(s) + O2(g)  2MgO(s) • CH4(g) + O2(g)  CO2(g) + H2O() • Mg(s)  Mg2+(aq) + 2e-

  3. Loss oxygen Gain electrons Reduction • 2PbO(s) + C(s)  2Pb(s) + CO2(g) • Pb2+(aq) + 2e-  Pb(s)

  4. Oxidation and Reduction

  5. oxidation Oxidizing agent • Oxidizing agent • A substance that causes oxidation is called an oxidizing agent. • An oxidizing agent is a substance that gains electrons. • 2PbO(s) + C(s)  2Pb(s) + CO2(g) oxidizing agent • Pb2+(aq) + 2e-  Pb(s) oxidizing agent

  6. reduction Reducing agent • Reducing agent • A substance that causes reduction is called a reducing agent. • A reducing agent is a substance that loses electrons. • 2PbO(s) + C(s)  2Pb(s) + CO2(g) reducing agent • Pb(s)  Pb2+(aq) + 2e- reducing agent

  7. Oxidizing agent and Reducing agent

  8. Oxidation numbers • The oxidation number of an element is zero. • oxidations numbers of N in N2, O in O2, S, Na, K, Cu are all zero. • The oxidation number of an atom in ionic form is equal to the charge on the ion. • oxidation number of iron in Fe3+ is +3 • oxidation number of oxygen in O2– is –2

  9. Oxidation numbers • The oxidation numbers of all atoms in a neutral compound added together must be zero. • MgO: Mg = +2, O = -2 • KCl: K = +1, Cl = -1 • ZnS: Zn = +2, S = -2

  10. Oxidation numbers • Some elements have fixed oxidation numbers in compounds • all alkali metals (Group I) in compounds must be +1 • Na in NaCl, K in K2SO4 • hydrogen in most of its compounds +1 • H in H2O, H in HCl, H in NH3 • Exception: H in metal hydride e.g. H in NaH is -1

  11. Oxidation numbers • Some elements have fixed oxidation numbers in compounds • all alkaline earth metals (group II) in compounds must be +2 • Ca in CaCO3, Mg in MgCl2 • fluorine in its compounds must be –1 • F in NaF, F in HF

  12. Oxidation numbers • Some elements have fixed oxidation numbers in compounds • oxygen in most of its compounds –2 • O in H2O, O in MgO • Exception: O in peroxide e.g. O in H2O2 is –1

  13. Oxidation numbers • The sum of oxidation numbers of all atoms in an ion is equal to the charge of the ion. • OH-: O = -2, H = +1 • NO3-: O = -2, N = +5 • SO42-: O = -2, S = +6 • NO2-: O = -2, N = +3 • SO32-: O = -2, S = +4

  14. Oxidation numbers

  15. Oxidation numbers

  16. reduction +2 0 0 +2 Mg(s) + CuSO4(aq)  MgSO4(aq) + Cu(s) Oxidizing agent Reducing agent oxidation Redox reactions • Displacement reaction

  17. reduction 0 +2 +1 0 Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) Oxidizing agent Reducing agent oxidation Redox reactions • Metal with acid

  18. The electrochemical series of metals

  19. The electrochemical series of metals • Cu2+(aq) is lower than H+(aq) ion in the e.c.s., it is a stronger oxidizing agent and thus it discharges (reduces) more readily.

  20. The electrochemical series of non-metals • OH-(aq) is higher than halides ions in the e.c.s., it is a stronger reducing agent and thus it discharges (oxidizes) more readily.

  21. Common oxidizing agents

  22. Common oxidizing agents • Acidified potassium permanganate, KMnO4 • Acidified sodium dichromate, Na2Cr2O7 • Concentrated sulphuric acid, H2SO4 • Dilute or Conc. nitric acid, HNO3 • Halogens: Cl2, Br2, etc. • Iron(III) ion, Fe3+(aq) • Metal ions low in e.c.s. e.g. Ag+(aq), Cu2+(aq) • Oxygen gas, O2(g)

  23. Common reducing agents

  24. Common reducing agents • Metal high in e.c.s. e.g. K(s), Na(s) • Hydrogen gas, H2(g) • Carbon, C(s) and carbon monoxide, CO(g) • Sulphur dioxide, SO2(g) and sulphite, SO32-(aq) • Iodide ion, I-(aq) • Iron(II) ion, Fe2+(aq)

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