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Oxidation-Reduction Reactions (Redox)

Oxidation-Reduction Reactions (Redox). What is the difference between acid/base reactions and redox reactions?. Acid/base reactions proton transfer (p + ) Redox reactions electron transfer (e - ). Flow of electrons.

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Oxidation-Reduction Reactions (Redox)

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  1. Oxidation-Reduction Reactions (Redox)

  2. What is the difference between acid/base reactions and redox reactions? • Acid/base reactions • proton transfer (p+) • Redox reactions • electron transfer (e-)

  3. Flow of electrons • Electrons respond to differences in potential by moving from the region of high potential to the region of low potential. High Ep Low Ep e- - +

  4. Flow of electrons low electronegativity high electronegativity e- Cl Li Lithium loses the e- tug-of-war with chloride.

  5. Terminology • Cations: • positively charged ions • generally metals • NH4+ is the exception • Anions: • negatively charged ions • non-metals • complex ions

  6. Oxidation: • When a substances loses e-. • Reduction: • When a substance gains e-.

  7. oxidized reduced

  8. Ca(s) + 2H+(aq)  Ca2+(aq) + H2(g) • Ca(s) has lost two e- to 2 H+(aq) to become Ca2+(aq). Ca(s) has been oxidized to Ca2+(aq) • At the same time 2 electrons are gained by 2 H+(aq) to form H2(g) . We say H+(aq) is reduced to H2(g) .

  9. Half-reactions • Ca(s)→ Ca2+(aq) + 2e- • Oxidation half reaction • 2H+(aq) + 2e- → H2(g) • reduction half reaction

  10. Half-reactions add together Ca(s)→ Ca2+(aq) + 2e- 2H+(aq) + 2e- → H2(g) Ca(s) +2H+ + 2e-Ca2+ + 2e- +H2(g) Ca(s) +2H+(aq)Ca2+(aq)+H2(g) +

  11. Half-reactions add together Cu(s)→ Cu2+(aq) + 2e- Ag+(aq) + e- → Ag(s) Cu(s) +2Ag+(aq) + 2e-Cu2+(aq) + 2e- +2Ag(s) Cu(s) +2Ag+(aq) Cu2+(aq)+2Ag(s) + ( ) x 2

  12. Electron Transfer and Terminology • Lose Electrons: Oxidation • Gain Electrons: Reduction. OIL - OXIDATION IS LOSS OF ELECTRONS RIG - REDUCTION IS GAIN OF ELECTRONS

  13. Iron • Iron comes from iron ore which is taken out of the ground by mining. • The pure iron is obtained by heating the ore at very high temperatures in a furnace with limestone to remove impurities. • The molten iron is taken out of the bottom of the furnace. It is further processed depending on how it is to be used.

  14. Why is gaining electrons called reduction? • Reduction originally meant the loss of oxygen from a compound. • 2Fe2O3(s) + C(s)→ 4Fe(s) + 3CO2(g) • Iron ore is reduced to metallic iron. The size of the pile gets smaller, hence the word reduction.

  15. Why is losing electrons called oxidation? • Oxidation originally meant the combination of an element with oxygen. • 4Fe(s) + 3O2(g)→ 2Fe2O3(g) C(s) + O2(g) → CO2(g)

  16. It Takes Two: Oxidation-Reduction • In all reduction-oxidation (redox) reactions, one species is reduced at the same time as another is oxidized.

  17. It Takes Two: Oxidation-Reduction Oxidizing Agent: • the species which causes oxidation is called the oxidizing agent. • substances that gains electrons • the oxidizing agent is always reduced

  18. It Takes Two: Oxidation-Reduction • Reducing Agent: • the species which causes reduction is called the reducing agent. • the reducing agent is always oxidized. • substances that give up electrons

  19. Example Cu(s) + 2 Ag+(aq)→ Cu2+(aq) + Ag(s) oxidated reduced R.A. O.A.

  20. Summary: Redox Theory • A redox reaction is a chemical reaction in which electrons are transferred. • Number of electrons lost by one species equals number of electrons gained by the other species. • Reduction is a process in which e- are gained. • Oxidation is a process in which e- are lost • A reducing agent donates e- and is oxidized. • A oxidizing agent gains e- and is reduced. WS 15-1

  21. Spontaneous Reactions

  22. Only one of these two reactions is possible. Which one? Cu(s) + 2 Ag+(aq)→ Cu2+(aq) + 2 Ag(s) Cu2+(aq) + 2 Ag(s) → Cu(s) + 2 Ag+(aq) Data table values EO, page 7 of your data books. 1) Cu(s) -- >> Cu 2+(aq) + 2 e- -0.34 EO ( R) 2) Ag +(aq) + e- -- >> Ag(s) +0.80 EO 1) Cu 2+(aq) + 2e- -- >> Cu (s) + 0.34 EO 2) Ag (s) -- >> Ag +(aq) + e-- 0.80 EO (R)

  23. Electric potential (V), Eo • the electric potential under standard conditions of a half-reaction in which reduction is occurring. • Standard conditions: • 25oC with all ions at 1 M concentrations and all gases at 1 atm pressure

  24. Standard Reduction Potentials • We cannot measure the potential of an individual half-cell! • We assign a particular cell as being our reference cell and then assign values to other electrodes on that basis.

  25. The Standard Hydrogen electrode • Eo (H+/H2) half-cell = 0.000 V e- p{H2(g)} = 1.00 atm H2 (g) [H+] = 1.00 Pt gauze

  26. Electric potential (V), Eo • If the net potential is a positive number then the reaction is spontaneous. • If the net potential is a negative number then the reaction is non-spontaneous. • Half cell potentials are not doubled or tripled as per balancing. We are only comparing potentials.

  27. Cu2+(aq) + 2Ag(s)→ Cu(s) + 2Ag+(aq) Cu2+ + 2e- → CuEo = 0.34 2Ag → 2Ag+ + 2e- Eo = -0.80 Cu2+(aq) + 2Ag(s)→ Cu(s) + 2Ag+(aq) Eo = -0.46 Negative potential, non-spontaneous Compare the two half reactions that make up the reaction. +

  28. Cu(s) + 2Ag+(aq)→ Cu2+(aq) + 2Ag(s) Cu(s)→ Cu2+ + 2e- Eo = -0.34 2Ag+ + 2e- → 2AgEo = 0.80 Cu(s) + 2Ag+(aq)→ Cu2+(aq) + 2Ag(s) Eo = 0.46 Positive potential, spontaneous Compare the two half reactions that make up the reaction.

  29. Problem • Write the oxidation/reduction half reactions and the net ionic equation when zinc is placed in Ni(NO3)2 solution. Identify the O.A. and R.A. and state if the reaction is spontaneous or non-spontaneous.

  30. Problem A piece of zinc is placed in a solution of nickel nitrate Ni(NO3)2 Spectator ion • Ni(NO3)2→ Ni2+(aq) + 2NO3- (aq) Zn(s) + Ni2+(aq) → ? • Oxidation: Zn(s) → Zn2+(aq) + 2e- +0.76 • Reduction: Ni2+(aq) + 2e- → Ni(s) - 0.26 Add half reactions

  31. Problem Zn is Oxidized Zn(s) + Ni2+(aq) → Zn2+(aq) + Ni(s)+0.50 Ni2+ is Reduced R.A. O.A. Positive potential, spontaneous

  32. NOTE*** Spontaneous shortcut • Locate the O.A. on the left and the R.A. on the right of the table. • If the O.A. is higher up on the table than the R.A. then the reaction is spontaneous. O.A. R.A. O.A. R.A. SPONTANEOUS REACTION NON-SPONTANEOUS REACTION

  33. highest attraction for electrons weak attraction for electrons

  34. Problem • Explain what happens when nickel is placed in a zinc nitrate solution. Ni(s) + Zn2+(aq)→ ? + ? REDUCING AGENT OXIDIZING AGENT ARE ON LEFT SIDE O.A. R.A. NICKEL  Ni ZINC NITRATE  Zn2+ and NO3 -

  35.  Ni(s)      Zn2+(aq)  R.A. is above the O.A. On the table NON SPONTANEOUS

  36. Disproportionation redox reactions where the OA and the RA are the same species. ( p 577 – text) Example: Fe2+ (aq) and Fe 2+ (aq) Fe2+ (aq) + 2 e - Fe (s) reduction of Fe2+ 2[ Fe2+ (aq)  Fe3+(aq) + e - ]oxidation of Fe2+ 3 Fe 2+(aq)  Fe(s) + 2 Fe3+(aq) net reaction NON – SPONTANEOUS REACTION

  37. DISPROPORTIONATION TRY THE REACTION WHERE Cu 1+ ACTS AS THE OXIDIZING AND REDUCING AGENTS TRY THE REACTION WHERE Cr2+ ACTS AS THE AS THE OXIDIZING AND REDUCING AGENTS

  38. Predicting redox reactions • List all species present. • Choose the strongest oxidizing and reducing agent. • Write the reduction half reaction, as written in the data book. • Write the oxidation half reaction, reverse the equation in the data book. • Balance number of electrons. • Add the two half reactions together to form the net ionic equation. • Predict if reaction is spontaneous or not.

  39. Problems • A mixture of bromine gas and chlorine gas is added to a solution of copper (II) sulphate and a copper strip. (water) ( CuSO4) (Br2(g)) (Cl2(g) ) ( Cu(s) ) • NOTE( Go down S.O.A. / Go up S.R.A.) Is the reaction spontaneous? Cl2(g) + 2e-→ 2 Cl-(aq) Cu(s)→ Cu2+(aq) + 2e- Br2(g) Cl2(g) H20(l) Cu2+(aq) Cu(s) SOA * Cl2(g) + Cu(s) → 2 Cl-(aq) + Cu2+(aq) SRA * SPONTANEOUS

  40. Problems Is the reaction spontaneous? • Lead is placed in a zinc nitrate solution.(list species) NO3-(aq) H20(l) Zn2+(aq) Pb(s) Non-spontaneous OA is below RA Zn2+(aq) + 2e- Zn (s) SOA Pb(s) Pb 2+(aq) + 2e- SRA R.A. Zn2+(aq) + Pb(s)  Zn(s) + Pb2+ O.A.

  41. Problems • A few drops of Hg(l) are dropped into a solution which is 1.0 M in both sulphuric acid and potasium permanganate. MnO4-(aq) SO42-(aq) H20(l) K+(aq) Hg(l) H+(aq) OA Is the reaction spontaneous? H+ hydrogen ion (From acid) O.A. RA R.A YES

  42. Problems • A few drops of Hg(l) are droped into a solution which is 1.0 M in both sulphuric acid and potasium permanganate. ( ) x2 MnO4-(aq) + 8 H+(aq) + 5e-→ Mn2+(aq) + 4 H2O(l) Hg(l)→ Hg2+(aq) + 2e- ( ) x5 Oxidized 2MnO4-(aq) + 16H+(aq) + 5Hg(l) → 2Mn2+(aq) + 8H2O(l) + 5Hg2+(aq) (Balance electrons) LHS = RHS

  43. Activity Series

  44. General Rules • Metal (+) ions are oxidizing agents. • Nonmetal (-) ions are reducing agents. • Metal elements are reducing agents. • Nonmetal elements are oxidizing agents.

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