1. Redox Reactions and Electrochemistry
Chapter 17

2. Unit Objectives Identify redox reactions that occur in daily life.
Identify what is being oxidized (reducing agent) and reduced (oxidizing agent) in a redox reaction.
Write the half and overall balanced equations for redox reactions.
Explain how an electrochemical cell works.
Explain the difference between a galvanic and electrolytic cell.

3. Unit Objectives Write the cell notation for a galvanic cell.
Calculate the standard reduction potential for a galvanic cell.
Relate Gibbs Free Energy, the equilibrium constant, and standard reduction potential.
Determine if a reaction will occur spontaneously using the Nernst Equation.

4. Redox Reactions in Everyday Life
Reduction-Oxidation Reactions (Redox) occur all around us
Burning of fuels
Converting food to energy
Photosynthesis
Batteries
Occur Together

5. Redox Reactions in Everyday Life
Reduced items, such as food and fuels, are high in energy
Oxidized items – carbon dioxide, water (byproducts) are low in energy
The energy released during redox reactions are what power our homes, cars, and bodies.

6. Redox Reactions Involve the reactions of metals with non-metals
There are three ways to view a redox reaction
If something is oxidized, something else must be reduced

7. Redox Reactions - Electrons
An increase in the oxidation number means loss of electrons and oxidation has occurred.
A decrease in the oxidation number means electrons have been gained and reduction has occurred.

8. Redox Reactions - Electrons
LEO the lion goes GER
Lose Electrons – Oxidation
Gain Electrons – Reduction
H2 + F2 → 2HF
Oxidation Reaction:
H2 → 2H+ + 2e-
Reduction Reaction:
F2 + 2e- → 2F-

9. Identify what is being oxidized and reduced in the following reactions:
H2 + Ag+  Ag + H+
Fe + CuSO4 → FeSO4 + Cu
H2 + O2  H2O
Cu + AgCl  CuCl2 + Ag

10. Biological Redox Reactions
Cellular Respiration
C6H12O6 + 6 O2 → 6 CO2 + 6 H2O + Energy
Photosynthesis
6 CO2 + 6 H2O + Light Energy→ C6H12O O2
Fermentation
C6H12O6 → 2 C2H5OH + 2 CO2

11. Rust
Iron is oxidized to produce Iron (II) hydroxide, then Iron (III) Hydroxide, which is typically written as Fe2O3 x H2O
Salt water can act as an electrolyte to facilitate this reaction.

12. Oxidizing and Reducing Agents

13. Oxidizing and Reducing Agents

14. Oxygen as an Oxidizing Agent
One of the most common oxidizing agents (undergoes reduction).
Oxygen occupies about 50% by mass of the accessible portion of the Earth and almost two-thirds of your body.
Found in carbohydrates, fats, sugars, proteins contained in food.
Used in combustion of fuels to power our industries, schools, and homes.
Also causes corrosion, food spoilage and food decay.

15. Oxidants Another name for oxidizing agents
Used to destroy microorganisms
Cleaners such as bleach
Antioxidants (such as Vitamin C) can prevent oxidation to living tissue

16. Hydrogen as a Reducing Agent
Most abundant element in the universe, but highly flammable
Often used to release metals from their ores after mining
WO3 + 3H2  W + 3H2O

17. Identify the oxidizing and reducing agents in the following reactions:
H2 + O2  H2O
Al + 3O2  Al2O3
Cu + AgNO3  Cu(NO3)2 + Ag

18. Redox Reactions

19. What are your questions?

20. Writing Half Reactions
Cr3+ + Zn  Cr + Zn2+ Step 1: Split reaction into half-reactions (reduction and oxidation) and balance the matter Zn  Zn2+ (oxidation) Cr3+  Cr (reduction) Step 2: Balance the charge or oxidation number with electrons Zn  Zn2+ + 2e (oxidation) 3e + Cr3+  Cr (reduction) Step 3: Check atom balance and charge balance on both sides of the equations.

21. Combining Half Reactions
Step 4: Multiply each reaction so the electrons gained the reduction half-reaction = electrons lost in oxidation half-reaction.
2(Cr3+ + 3e  Cr) 2Cr3+ + 6e  2Cr
3(Zn  Zn2+ + 2e) 3Zn  3Zn2+ + 6e
Step 5: Combine the reactions, canceling the electrons.
2Cr3+ + 6e  2Cr
3Zn  3Zn2+ + 6e
2Cr3+ 3Zn  2Cr + 3Zn2+

22. Balance the following redox reaction
Balance the following redox reaction. Identify what is being oxidized and what is being reduced. Fe + S8  FeS

23. Balance the following redox reaction
Balance the following redox reaction. Identify what is being oxidized and what is being reduced.
MgCl2 + Fe  Mg + FeCl3

24. Balance the following redox reaction
Balance the following redox reaction. Identify what is being oxidized and what is being reduced. Mg + O2  MgO

25. Balance the following redox reaction
Balance the following redox reaction. Identify what is being oxidized and what is being reduced. Fe3+ + Sn2+  Fe2+ + Sn4+

26. Balancing Redox Reactions in Acidic Solution
Identify what is being oxidized and reduced and write the half reactions
Complete and balance each half reaction
Balance everything except O and H
Balance the O by adding water to one side of the reaction
Balance the H by adding H+ to one side of the reaction.
Balance the charges by adding electrons on the more positive side.
Combine the half reactions together to create the final overall reaction.

27. Cr2O72-(aq) + 6Fe2+(aq)  2Cr3+(aq) + 6Fe3+(aq)

28. 5Fe2+(aq) + MnO4-(aq)  5Fe3+(aq) + Mn2+(aq)

29. In a concentrated solution, zinc metal reduces nitrate ion to ammonium ion, and zinc is oxidized to Zn2+. Write the balanced net ionic equation for this reaction.

30. I2(s) + NO3-(aq)  IO3-(aq) + NO2(g)
Iodic Acid, HIO3, can be prepared by reacting I2 with concentrated nitric acid. Write the balanced equation for this if the skeleton reaction is
I2(s) + NO3-(aq)  IO3-(aq) + NO2(g)

31. Balance the following redox reaction that occurs in acidic solution:
H2S + NO3-  S8 + NO2

32. Balancing Redox Reactions in Basic Solution
Begin by balancing the reaction as if it was in an acid solution
Add the same number of OH- ions to each side of the reaction as you have H+ ions
Simplify the H+ and OH- ions to be written as water molecules.
Cancel any H2O molecules that appear on both sides of the reaction.
Combine the half reactions together to create the final overall reaction.

33. MnO4-(aq) + SO32-(aq)  MnO2(s) + SO42-
Permanganate ion oxidizes sulfite ion in basic solution according to the following skeletal question. Write the balanced equation for this redox reaction.
MnO4-(aq) + SO32-(aq)  MnO2(s) + SO42-

34. Balance the following redox reaction that occurs in basic solution:
Mn2+ + ClO3- MnO2 + ClO2

35. Balance the following redox reaction that occurs in basic solution:
H2O2 + ClO2  ClO2- + O2

36. What are your questions?

37. Electrochemistry
Field of chemistry involving the study of chemical reactions that are driven by an electrical current.
The reactions that occur are redox reactions.
Occur through the use of an electrochemical cell.

38. Electrochemistry

39. Electrochemistry Dry Cell Battery Mercury Battery Lead Storage Battery
Lithium-Ion Battery
Fuel Cells

41. Electrochemistry Reaction at Anode (oxidation): Zn(s)  Zn2+(aq)+ 2e-
Reaction at Cathode (reduction):
2MnO2(s) + H2O + 2e-  Mn2O3(s) + 2OH-(aq)
Overall Reaction:
Zn(s) + 2MnO2(s) + H2O  Zn2+ + Mn2O3(s) + 2OH-(aq)

42. Zn(s)∣Zn2+ (1M)∥Cu2+ (1M)∣Cu(s)
Electrochemistry
Cell Voltage
Voltage across the electrodes of a galvanic cell
Also called the cell potential
Measured using a voltmeter
Cell Diagram
Zn(s)∣Zn2+ (1M)∥Cu2+ (1M)∣Cu(s)

43. Electrolysis
The process in which electrical energy is used to carry out a nonspontaneous chemical reaction
Electrolytic Cell
A setup used to carry
out electrolysis

44. Standard Reduction Potentials
Standard Reduction Potentials (E°)
The voltage associated with a reduction reaction at an electrode when all states are 1M and all gases are at 1atm.
Standard emf (E°cell)
E°cell = E°cathode - E°anode

45. Standard Reduction Potentials
E° apply to half-reactions
Change the sign of E° if the reaction is reversed
The more positive E°, the more likely the substance will be reduced.
The value of E° is not affected by the amount of solution present or the number of moles in solution.
Under standard-state conditions, any species on the left of side of a given half reaction will react spontaneously with any species on the right side of a given half reaction as long as that species has a lower E° value.

47. Order the following oxidizing agents by increasing strength under standard-state conditions: Cl2(g) , H2O2(aq) , Fe3+(aq) Order the following reducing agents by increases strength under standard-state conditions: H2(g) , Al(s) , Cu(s)

48. Using table 17.2, calculate the E°cell for the following reaction: Mg(s) + HCl(aq)  MgCl2(aq) + H2(g)

49. Using table 17.2, calculate the E°cell for the following reaction: Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s)

50. What are your questions?

51. Thermodynamics of Redox Reactions
The standard emf can be related to Gibbs Free Energy (∆G) and the equilibrium constant, K.
Relates to standard states
Where R = J/K•mole
F is Faraday’s Constant = 9.647x104 C/mole e-
n is the number of moles of e-
K is the equilibrium constant
∆G° = -nFE°cell
∆G° = -RTlnK
E°cell = log K RT nF

52. Relationship between ∆G°, K and E°cell
Reaction Under Standard State Conditions -
>1 +
Favors the Products =1
At Equilibrium
<1
Favors the Reactants

53. Using the standard reduction potentials provided in table 17
Using the standard reduction potentials provided in table 17.2, calculate the equilibrium constant for the following reaction: Mg(s) + HCl(aq)  MgCl2(aq) + H2(g)

54. Using the standard reduction potentials provided in table 17
Using the standard reduction potentials provided in table 17.2, calculate the equilibrium constant for the following reaction: Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s)

55. Using the standard reduction potentials provided in table 17
Using the standard reduction potentials provided in table 17.2, calculate the Gibbs Free Energy for the following reaction: N2(g) + O2(g)  NH3(g)

56. Using the standard reduction potentials provided in table 17
Using the standard reduction potentials provided in table 17.2, calculate the Gibbs Free Energy for the following reaction: Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s)

57. The Nernst Equation
Relates emf and the concentrations of the reactants in nonstandard states.
Where R = J/K•mole
F is Faraday’s Constant = 9.647x104 C/mole e-
n is the number of moles of e-
Q is the reaction quotient
E = E° InQ RT nF

58. Using the Nernst Equation, determine if the following reaction will proceed spontaneously at 298K if the concentration of Cu2+ is 0.25M and Fe3+ is 0.20M? Cu(s) + Fe3+(aq)  Cu2+(aq) + Fe(s)

59. Using the Nernst Equation, determine if the following reaction will proceed spontaneously at 298K if the concentration of Cu(NO3)2 is 0.015M and AgNO3 is 0.030M? Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s)

60. Using the Nernst Equation, determine if the following reaction will proceed spontaneously at 298K if the concentration of MgCl2 is 0.6M and HCl is 0.55M? Mg(s) + HCl(aq)  MgCl2(aq) + H2(g)

61. What are your questions?