Your SlideShare is downloading. ×
Oxidation and reduction
Upcoming SlideShare
Loading in...5

Thanks for flagging this SlideShare!

Oops! An error has occurred.


Introducing the official SlideShare app

Stunning, full-screen experience for iPhone and Android

Text the download link to your phone

Standard text messaging rates apply

Oxidation and reduction


Published on

Published in: Education, Technology, Business

No Downloads
Total Views
On Slideshare
From Embeds
Number of Embeds
Embeds 0
No embeds

Report content
Flagged as inappropriate Flag as inappropriate
Flag as inappropriate

Select your reason for flagging this presentation as inappropriate.

No notes for slide


  • 1. Oxidation and reduction (Redox reactions) By Azfar Ali Bajwa 1
  • 2. Oxidation and Reduction (Redox)Early chemists saw “oxidation” reactionsonly as the combination of a materialwith oxygen to produce an oxide. • For example, when methane burns in air, it oxidizes and forms oxides of carbon and hydrogen CH4 + O2 → CO + H2 + H2O 2
  • 3. Oxidation and Reduction (Redox)But, not all oxidation processes that useoxygen involve burning: • Elemental iron slowly oxidizes to compounds such as iron (III) oxide, commonly called “rust”• Bleaching stains in fabrics• Hydrogen peroxide also releases oxygen when it decomposes 3
  • 4. Oxidation and Reduction (Redox)• A process called “reduction” is the opposite of oxidation, and originally meant the loss of oxygen from a compound• Oxidation and reduction always occur simultaneously• The substance gaining oxygen (or losing electrons) is oxidized, while the substance losing oxygen (or gaining electrons) is reduced. 4
  • 5. Oxidation and Reduction (Redox)• Today, many of these reactions may not even involve oxygen• Redox currently says that electrons are transferred between reactants Mg + S→ Mg2+ + S2- (MgS)• The magnesium atom (which has zero charge) changes to a magnesium ion by losing 2 electrons, and is oxidized to Mg2+• The sulfur atom (which has no charge) is changed to a sulfide ion by gaining 2 electrons, and is reduced to S2- 5
  • 6. Oxidation and Reduction (Redox) 0 0 1 1 2 Na Cl 2 2 Na ClEach sodium atom loses one electron: 0 1 Na Na eEach chlorine atom gains one electron: 0 1 Cl e Cl 6
  • 7. LEO says GER : Lose Electrons = Oxidation 0 1Na Na e Sodium is oxidized Gain Electrons = Reduction0 1Cl e Cl Chlorine is reduced 7
  • 8. LEO says GER : - Losing electrons is oxidation, and the substance that loses the electrons is called the reducing agent. - Gaining electrons is reduction, and the substance that gains the electrons is called the oxidizing agent. Mg is the Mg is oxidized: loses e-, becomes a Mg2+ ion reducing agent Mg(s) + S(s) → MgS(s)S is the oxidizing agent S is reduced: gains e- = S2- ion 8
  • 9. Not All Reactions are Redox Reactions - Reactions in which there has been no change in oxidation number are NOT redox reactions. Examples: 1 5 2 1 1 1 1 1 5 2Ag N O3 (aq) Na Cl (aq) Ag Cl ( s) Na N O3 (aq) 1 2 1 1 6 2 1 6 2 1 22 Na O H (aq) H 2 S O 4 (aq) Na 2 S O 4 (aq) H 2 O(l ) 9
  • 10. Corrosion•Damage done to metal is costly to preventand repair•Iron, a common construction metal oftenused in forming steel alloys, corrodes bybeing oxidized to ions of iron by oxygen. • This corrosion is even faster in the presence of salts and acids, because these materials make electrically conductive solutions that make electron transfer easy 10
  • 11. Corrosion•Luckily, not all metals corrode easily •Gold and platinum are called noble metals because they are resistant to losing their electrons by corrosion •Other metals may lose their electrons easily, but are protected from corrosion by the oxide coating on their surface, such as aluminum – •Iron has an oxide coating, but it is not tightly packed, so water and air can penetrate it easily 11
  • 12. Corrosion•Serious problems can result if bridges, storagetanks, or hulls of ships corrode •Can be prevented by a coating of oil, paint, plastic, or another metal •If this surface is scratched or worn away, the protection is lost•Other methods of prevention involve the“sacrifice” of one metal to save the second •Magnesium, chromium, or even zinc (called galvanized) coatings can be applied 12
  • 13. Assigning Oxidation NumbersOxidation: Loss of electrons.Reduction: Gain of electrons.Reductant: Species that loses electrons.Oxidant: Species that gains electrons.Valence: the electrical charge an atom wouldacquire if it formed ions in aqueous solution.• An “oxidation number” is a positive or negative number assigned to an atom to indicate its degree of oxidation or reduction. 13
  • 14. Rules for Assigning Oxidation Numbers1) The oxidation number of any uncombined element is zero.2) The oxidation number of a monatomic ion equals its charge. 0 0 1 1 2 Na Cl 2 2 Na Cl 14
  • 15. Rules for Assigning Oxidation Numbers3) The oxidation number of oxygen in compounds is -2, except in peroxides, such as H2O2 where it is -1.4) The oxidation number of hydrogen in compounds is +1, except in metal hydrides, like NaH, where it is -1. 1 2 H2O 15
  • 16. Rules for Assigning Oxidation Numbers5) The sum of the oxidation numbers of the atoms in the compound must equal 0. 1 2 2 2 1 H2O Ca(O H ) 2 2(+1) + (-2) = 0 (+2) + 2(-2) + 2(+1) = 0 H O Ca O H 16
  • 17. Rules for Assigning Oxidation Numbers6) The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge. ? 2 ? 2 2 N O3 S O4 X + 3(-2) = -1 X + 4(-2) = -2 N O S O thus X = +5 thus X = +6 17
  • 18. Reducing Agents and Oxidizing Agents•An increase in oxidation number = oxidation• A decrease in oxidation number = reduction 0 1 Na Na eSodium is oxidized – it is the reducing agent 0 1 Cl e ClChlorine is reduced – it is the oxidizing agent 18
  • 19. Identifying Redox Equations• In general, all chemical reactions can be assigned to one of two classes: 1) oxidation-reduction, in which electrons are transferred: • Single- replacement, combination, decompositio n, and combustion 2) this second class has no electron transfer, and includes all others: • Double-replacement and acid-base 19
  • 20. Identifying Redox Equations•In an electrical storm, nitrogen and oxygenreact to form nitrogen monoxide: N2(g) + O2(g) → 2NO(g)•Is this a redox reaction? YES! •If the oxidation number of an element in a reacting species changes, then that element has undergone either oxidation or reduction; therefore, the reaction as a whole must be a redox 20
  • 21. BALANCING OVERALL REDOX REACTIONSExample - balance the redox reaction below: Fe + Cl2 Fe3+ + Cl-Step 1: Assign valences, Fe0 + Cl20 Fe3+ + Cl-Step 2: Determine number of electrons lost or gained by reactants. Fe0 + Cl20 Fe3+ + Cl- 3e- 2e-Step 3: Cross multiply. 2Fe + 3Cl20 2Fe3+ + 6Cl- 21
  • 22. HALF-CELL REACTIONSThe overall reaction: 2Fe + 3Cl20 2Fe3+ + 6Cl-may be written as the sum of two half-cell reactions: 2Fe 2Fe3+ + 6e- (oxidation) 3Cl20 + 6e- 6Cl- (reduction)All overall redox reactions can be expressed as the sum of two half-cell reactions, one a reduction and one an oxidation. 22
  • 23. Another example - balance the redox reaction: FeS2 + O2 Fe(OH)3 + SO42- Fe+2S20 + O20 Fe+3(OH)3 + S+6O42- 15e- 4e- 4FeS2 + 15O2 Fe(OH)3 + SO42- 4FeS2 + 15O2 4Fe(OH)3 + 8SO42- 4FeS2 + 15O2 +14H2O 4Fe(OH)3 + 8SO42- + 16H+This reaction is the main cause of acid generation in drainage from sulfide ore deposits. 23
  • 24. Final example: C2H6 + NO3- HCO3- + NH4+ 24
  • 25. C-32H6 + N+5O3- 2HC+4O3- + N-3H4+ 14e- 8e- 8C2H6 + 14NO3- 2HCO3- + NH4+ 8C2H6 + 14NO3- 16HCO3- + 14NH4+8C2H6 + 14NO3- + 12H+ + 6H2O 16HCO3- + 14NH4+ 25
  • 26. STRENGTH OF REDUCING AND OXIDIZING AGENTS• Reduced form of those substances with highly negative reduction potentials are good reducing agents (are easily oxidized) (GH2)• The oxidized form of those substances with highly positive reduction potentials are good oxidizing agents (are easily reduced ) (A) 26
  • 27. ELECTROMOTIVE FORCEElectromotive force (EMF): The electrical potential generated by thehalf reactions of an electrochemical cell.Consider: Zn + Cu2+ Zn2+ + CuAt equilibrium the electrochemicalcell has to obey [ Zn 2 ] K 1037.24 [Cu 2 ] 27
  • 28. Reduction PotentialsThe numerical values for E’oreflect the reduction potentialsrelative tohalf reaction which is taken as-0.414 volt at pH 7 . 28
  • 29. Reduction potentials 29
  • 30. Standard hydrogen electrode- The standard hydrogen electrode (abbreviated SHE), is a redox electrode which forms the basis of the thermodynamic scale of oxidation- reduction potentials- potentials of any other electrodes are compared with that of the standard hydrogen electrode at the same temperature.- Hydrogen electrode is based on the redox half cell : 2H+(aq) + 2e- → H2(g) 30
  • 31. Deriving the NERNST EQUATION• The relative tendency of the overall oxidation reduction reaction to go may be calculated from the difference between the reduction potentials of the component half reactions :• As they are independent of the number of electrons in the half reaction so G of the reaction can be calculated by G n 31
  • 32. Deriving the NERNST EQUATION G n • Where n is the number of electrons transferred per mole (i.e the number of equivalents per mole )• is the Faraday’s constant• must be positive to yeild a ‘spontaneous’ reaction (i.e., negative G ) 32
  • 33. Deriving the NERNST EQUATION• Relating now the Keq with G 33
  • 34. Deriving the NERNST EQUATION 34
  • 35. Effect of PH on E’o• Equation 35 can be used to predict the E’o of a half reaction at one pH if E’o at another pH is known .... For example consider• The non standard-state E value is related to E’o by :• If [A] and [AH2] are each one molar then 35
  • 36. Effect of PH on E’o• For each increase of one pH unit ,E’o becomes more negative by 0.059 volt• And 0.0295 volt for NAD+/NADH and NADP+/NADPH half reactions ...this value will be used next in numerical 36
  • 37. Numerical Problems 37
  • 38. Numerical Problems 38
  • 39. 39
  • 40. Numerical Problems 40
  • 41. 41
  • 42. Numerical Problems 42
  • 43. 43
  • 44. 44
  • 45. 45
  • 46. Refrences• Biochemical calculations by Irvin H seagel 46