Unit 5 – Part 2: Redox Reactions and Electrochemistry

1. Unit 5 – Part 2: Redox Reactions and Electrochemistry • Oxidation Numbers • Oxidizing and Reducing Agents • Balancing Redox Reactions • Acidic solutions • Basic solutions • Galvanic Cells • Nernst Equation

2. Redox Reactions • Oxidation-Reduction Reactions (Redox Reactions) • reactions that involve the transfer of electrons between two reactants • an element in one reactant is oxidized while an element in another reactant is reduced

3. Redox Reactions • Many practical or everyday examples of redox reactions: • Corrosion of iron (rust formation) • Forest fire • Charcoal grill • Natural gas burning • Batteries • Production of Al metal from Al2O3 • Metabolic processes combustion

4. Redox Reactions • Many different types of redox reactions: • Oxidation by Molecular Oxygen 4 Fe (s) + 3 O2(g) 2 Fe2O3(s) • Oxidation of Metals by Acids or Water Zn (s) + 2 H+(aq) Zn2+(aq) + H2(g) 2 Na (s) + 2 H2O (l) 2 NaOH (aq) + H2(g) • Metal Displacement: Fe (s) + Ni2+(aq) Fe2+(aq) + Ni (s)

5. Redox Reactions • Oxidation: • the loss of electrons • oxidation number increases • the gain of oxygen • the loss of hydrogen • Reduction: • the gain of electrons • oxidation number decreases • the gain of hydrogen • the loss of oxygen

6. Redox Reactions • LEO: Lose Electrons Oxidation • GER: Gain Electrons Reduction GER LEO LEO says GER

7. Redox Reactions • Electrons are not explicitly shown in chemical equations. • Oxidation Numbers are used to keep track of electrons gained and lost during redox reactions. • Oxidation number • a hypothetical number assigned to an individual atom present in a compound using a set of rules. • May be positive, negative, or zero

8. Rules for Oxidation Numbers • Oxidation numbers are always reported for individual atoms or ions notgroups of atoms or ions!!!!!!!!!!! • For an atom in its elemental form, the oxidation number is always zero. • H2: oxidation # = 0 for each H atom • Cu: oxidation number = 0 • Cl2: oxidation # = 0 for each Cl atom

9. Rules for Oxidation Numbers • For any monoatomic ion, oxidation # = charge of the ion • K+ oxidation # = +1 • Cl- oxidation # = -1 • S2- oxidation # = -2

10. Rules for Oxidation Numbers • Group 1A Metal Cations: • Always +1 • Group 2A Metal Cations: • Always +2 • Hydrogen (H) • +1 when bonded to nonmetals • -1 when bonded to metals

11. Rules for Oxidation Numbers • Oxygen (O) • -1 in peroxides (O22-) • -2 in all other compounds • Fluorine (F) • always -1

12. Rules for Oxidation Numbers • The sum of the oxidation numbers of all atoms in any chemical species (ion or neutral compound) is equal to the charge on that chemical species • H2O: 1 + 1 + -2 = 0 • MgCl2: 2 + -1 + -1 = 0 • MnO4-: 7 + -2 + -2 + -2 + -2 = -1

13. Oxidation Numbers • For many compounds, you will be able to directly apply the rules to determine the oxidation number of all atoms except for one. • Use the last rule to determine the oxidation number of that last element.

14. Oxidation Numbers Example: Determine the oxidation state of all elements in SO3.

15. Oxidation Numbers Example: Determine the oxidation state of all elements in Cr2O72-.

16. Oxidation Numbers Example: Determine the oxidation state of all elements in Cu(NO3)2

17. Redox Reactions • Changes in oxidation number are used to determine if a redox reaction has occurred and to identify the elements oxidized and reduced. • Reduction (GER) • gain of electrons • oxidation number decreases (is reduced) • Oxidation (LEO) • loss of electrons • oxidation number increases

18. Redox Reactions Example: Is the following a redox reaction? If so, which element is oxidized? Which element is reduced

19. Redox Reactions • Oxidizing Agent (oxidant): • the reactant that causes another reactant to be oxidized • the reactant that contains the element that is reduced • Reducing Agent (reductant): • the reactant that causes another substance to be reduced • the reactant that contains the element that is oxidized

20. Redox Reactions Example: Did a redox reaction occur in the following reaction? If so, what is the oxidizing agent? What is the reducing agent? 2 Na (s) + 2 H2O (l)  2 NaOH (s) + H2 (g)

21. Redox Reactions • When writing the equation for a redox reaction, you must • balance the atoms on both sides • balance the loss and gain of electrons • For “simple” redox reactions, the loss and gain of electrons is “automatically” balanced when you balance the atoms Zn (s) + 2 H+ (aq) Zn2+ (aq) + H2 (g)

22. Redox Reactions • Most redox reactions are more complex to balance. Sn2+(aq) + Fe3+(aq) Sn4+(aq) + Fe2+(aq) Sn2+(aq) + 2 Fe3+(aq) Sn4+(aq) + 2 Fe2+(aq)

23. Redox Reactions • Redox reactions can be broken up into 2 half-reactions: • a reaction that shows either oxidation or reduction alone • Overall reaction: • Oxidation half reaction: • Reduction half reaction:

24. Redox Reactions • Notice that electrons lost = electrons gained in a balanced redox reaction:

25. Balancing Redox Reactions • Procedure for AcidicSolutions: • Divide the equation into 2 incomplete half reactions • one for oxidation • one for reduction • Balance each half-reaction: • balance all elements except H and O • balance O atoms by adding H2O • balance H atoms by adding H+ • balance charge by adding e- to side with more positive overall charge

26. Balancing Redox Equations • Multiply each half reaction by an integer so that • # e- lost = # e- gained • Add the half reactions together. • Simplify where possible by canceling species appearing on both sides of equation • Check the equation • # of atoms • total charge on each side

27. Balancing Redox Equations Example: Balance the following redox reaction: Cr2O72- + Fe2+ Cr3+ + Fe3+ (acidic soln)

28. Balancing Redox Equations

29. Balancing Redox Equations

30. Balancing Redox Reactions

31. Balancing Redox Reactions • Procedure for BasicSolutions: • Divide the equation into 2 incomplete half reactions • one for oxidation • one for reduction

32. Balancing Redox Reactions • Balance each half-reaction: • balance all elements except H and O • balance O atoms by adding H2O • balance H atoms by adding H+ • add 1 OH- to both sides for every H+ added • combine H+ and OH- on same side to make H2O • cancel the same # of H2O from each side • balance charge by adding e- to side with the more positive overall charge different

33. Balancing Redox Equations • Multiply each half reaction by an integer so that • # e- lost = # e- gained • Add the half reactions together. • Simplify where possible by canceling species appearing on both sides of equation • Check the equation • # of atoms • total charge on each side

34. Balancing Redox Reactions Example: Balance the following redox reaction. NH3 + ClO- Cl2 + N2H4 (basic soln)

35. Balancing Redox Reactions

36. Balancing Redox Reactions

37. Balancing Redox Reactions

38. Balancing Redox Reactions Example: Balance the following redox reaction which takes place in acidic solution. NO2- + Cr2O72- NO3- + Cr3+

39. Balancing Redox Reactions Example:Balance this redox reaction which occurs under basic conditions. Pb(OH)42- + ClO- PbO2 + Cl-