1. Chapter 21: Electrochemistry
Table of Contents
Chapter 21: Electrochemistry
21.1 – Voltaic Cells
21.2 – Types of Batteries
21.3 – Electrolysis

2. Chapter 21: Electrochemistry
Table of Contents
Chapter 21: Electrochemistry
21.1 – Voltaic Cells

3. Try it out! Zn + NiSO4 → Ni + ZnSO4
Basic Assessment Questions
Try it out!
Identify what is oxidized, what is reduced, the oxidizing agent & the reducing agent.
Zn + NiSO4 → Ni + ZnSO4
Ox: Zn
Red: Ni2+
Ox ag: Ni2+
Red ag: Zn
Define
Redox reactions
Oxidation
Reduction
Half-reactions

4. Electrochemistry is using chemistry to create electricity
Electrochemistry: Basic Concepts
Electrochemistry is using chemistry to create electricity
Electrical current: The flow of electrons in a particular direction
Redox reactions can be used to produce an electrical current.
This is what occurs in a battery—one form of an electrochemical cell

5. Electrochemistry: Basic Concepts
Definitions
An electrochemical cell is a device that uses redox reactions to create electricity OR electric energy to cause chemical reactions
A voltaic or galvanic cell converts chemical energy to electrical energy by a spontaneous redox reaction. It is a type of electrochemical cell.

6. Balanced: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
Total ionic:
Zn(s) + Cu2+(aq) + SO42-(aq) → Zn2+(aq) + SO42-(aq) + Cu(s)
Net ionic: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) ox
red
Half Reactions:
Zn(s) → Zn2+(aq) + 2e-
Half Reactions:
Zn(s) → Zn2+(aq) + 2e- Cu2+(aq) + 2e-→ Cu(s)

7. Electrochemical Cell
Zinc strip
Copper strip
Cannot transfer e-
SO42-
SO42-
1M Zn2+
1M Cu2+
Zn(s) → Zn2+(aq) + 2e- Cu2+(aq) + 2e-→ Cu(s)

8. Electrochemical Cell
Allows e- transfer
Copper wire
(+) charge builds up in one solution, (-) charge builds up in the other

9. Allows ions to pass from one side to the other
Electrochemical Cell
Allows ions to pass from one side to the other
Ions pass through plugs in the bridge, but solutions do not mix
Salt bridge
KCl
Circles = plug, algae gel

10. Electrochemical Cell Zn → Zn2+ + 2e- Cu2+ + 2e- → Cu e- flow K+ Cl-
e- e-
Cl- K+
e- e- Zn
Cl- K+ Cu
Circles = plug, algae gel
Zn2+
Cu2+
Zn → Zn2+ + 2e-
Cu2+ + 2e- → Cu

11. What is this called? What does it do?
Why can redox reactions create electricity?
e- flow
KCl Zn Cu
Circles = plug, algae gel
1M Zn2+
1M Cu2+
Zn → Zn2+ + 2e-
Cu2+ + 2e- → Cu

12. Voltaic/Galvanic Cell
e- flow
KCl
electrodes
electrodes
half-cells Zn Cu
Circles = plug, algae gel
1M Zn2+
1M Cu2+
electrolytes

13. Electrodes are immersed in electrolytes: usually a solution of ions
Electrochemistry: Basic Concepts
Definitions
Half-cells: Where oxidation and reduction reactions separately take place (2 parts)
Electrode: an object in the half-cell that conducts electrons to or from another substance
Electrodes are immersed in electrolytes: usually a solution of ions

14. Voltaic/Galvanic Cell Voltaic Cell
e- flow
KCl Zn Cu
oxidation
ANODE
Circles = plug, algae gel
reduction
CATHODE
1M Zn2+
1M Cu2+
Zn → Zn2+ + 2e-
Cu2+ + 2e- → Cu

15. Anode: The electrode where oxidation takes place
Electrochemistry: Basic Concepts
Definitions
Anode: The electrode where oxidation takes place
Cathode: The electrode where reduction takes place

16. In the next slide, identify… Salt bridge What is oxidized
Electrochemistry: Basic Concepts
Practice
In the next slide, identify…
Salt bridge
What is oxidized
What is reduced
Cathode
Anode
Direction of e- flow through wire

17. Voltaic/Galvanic Cell
oxidized
reduced

18. Energy in a cell We need to be able to…
find out how much energy we can get out of a cell
write the balanced equations of cells from half-reactions (tomorrow)

19. More (+) means it’s more likely to take electrons
Electrochemistry: Basic Concepts
Equations of Cells
Standard Reduction Potentials: lists which are more likely to take electrons
reduced
Cu2+ + 2e- → Cu
Cu2+ | Cu
More (+) means it’s more likely to take electrons

20. Copper-Hydrogen cell under standard conditions
Electrochemistry: Basic Concepts
Copper-Hydrogen cell under standard conditions
H2 (g) + Cu2+(aq) → 2H+(aq) + Cu(s)
(net ionic equation)
Cell notation
H2 | H+ || Cu2+ | Cu
reduction half-cell
CATHODE
oxidation half-cell
ANODE
salt bridge and wire

21. Eºcell = Eºreduction - Eºoxidation
Electrochemistry: Basic Concepts
Cell Potential
To find a voltaic cell’s standard potential (how many volts it can use):
Eºcell = Eºreduction - Eºoxidation
HINT: Write down everything. Show all work. When you do it in your head, you get confused.

22. Examples: Find the cell potential of a copper-zinc cell.
Electrochemistry: Basic Concepts
Examples: Find the cell potential of a copper-zinc cell.
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
Step 1: Cell notation: Zn | Zn2+ || Cu2+ | Cu
Step 2: ID ox and red ox
red
Step 3: Find reduction potentials
Zn2+|Zn V Cu2+ | Cu V
Step 4: Equation
Eºcell = Eºred - Eºox

23. Examples: Find the cell potential of a copper-zinc cell.
Electrochemistry: Basic Concepts
Examples: Find the cell potential of a copper-zinc cell.
Step 4: Equation
Eºcell = Eºred - Eºox
= EºCu2+|Cu - EºZn2+|Zn
= V – ( V)
= V
The standard reduction potential of the copper-zinc cell is V

24. Try it out! Calculate the cell potential.
Additional Assessment Questions
Try it out!
Calculate the cell potential.
2Ag+(aq) + Co(s) → Co2+(aq) + 2Ag(s)
+ 1.08V

25. In a Electrochemical cell… when Eº = (+), it is spontaneous
Electrochemistry: Basic Concepts
Potential Difference
A spontaneous reaction: proceeds naturally. Voltaic cells are always spontaneous.
A non- spontaneous reaction: requires an outside influence to proceed. (eg. forcing the electrons in the opposite direction)
In a Electrochemical cell…
when Eº = (+), it is spontaneous
when Eº = (-), it is not spontaneous

26. Spontaneity in a Cell Eºcell = + 1.322 V spontaneous
Electrochemistry: Basic Concepts
Spontaneity in a Cell
Eºcell = V
Eºcell = V
Eºcell = V
Eºcell = V
Eºcell = V
spontaneous
not spontaneous

27. Additional Assessment Questions
Question 2b
Calculate the cell potential to determine if each of these redox reactions is spontaneous.
Mn2+ + 2Br- → Br2 + Mn
Fe2+ + Sn2+ → Fe3+ + Sn
Ni2+ + Mg → Mg2+ + Ni
Pb2+ + Cu+ → Pb + Cu2+
V, nonspontaneous
V, nonspontaneous
V, spontaneous
V, nonspontaneous

28. Practice problems Part 1 Oxidation half: Cu → Cu2+ + 2e-
Voltaic Cells handout
Due tomorrow at the end of class
Finished? Read ch 21 – it really will help!
I will not be covering 21.2 or 21.3 in class, but there are WebAssign questions on it - you’re going to have to read it!
Part 1
Oxidation half: Cu → Cu2+ + 2e-
Reduction half: Ag+ + e- → Ag

30. Activity
At your tables, use whatever you have in your backpacks to make model of an electrochemical cell
Be able to explain to another group what each object represents in the “electrochemical cell” and what it does in the “cell”
You have 3 minutes

31. We will go through steps to follow and examples
Electrochemistry: Basic Concepts
Sometimes, you must find a balanced reaction of cells from only reduction half-reactions.
We will go through steps to follow and examples

32. Electrochemistry: Basic Concepts
Equations of Cells
Standard Reduction Potentials: lists which are more likely to take electrons
reduced
Cu2+ + 2e- → Cu
Cu2+ | Cu
When comparing standard reduction potentials in a cell, the ½ reaction that is…
more (+) is reduction
more (-) is oxidation

33. Ni(s) + 2Ag+(aq) → Ni2+ + 2Ag(s)
Electrochemistry: Basic Concepts
1) Identify which reaction is oxidation and which is reduction, 2) write the balanced equation
Ag+(aq) + e- → Ag(s)
Ni2+ + 2e- → Ni(s)
EºAg+|Ag = V
EºNi2+|Ni = V
reduction
oxidation
Ni(s) → Ni2+ + 2e- 2
Ag+(aq) + e- → Ag(s) 2 2
Ni(s) + 2Ag+(aq) → Ni2+ + 2Ag(s)

34. Magnesium in a solution of Mg2+ Lead in a solution of Pb2+
Electrochemistry: Basic Concepts
1) Identify which reaction is oxidation and which is reduction, 2) write the balanced equation
Magnesium in a solution of Mg2+
Lead in a solution of Pb2+
Mg2+(aq) + 2e- → Mg(s)
Pb2+(aq) + 2e- → Pb(s)

35. Mg(s) + Pb2+ → Mg2+(aq) + Pb(s)
Electrochemistry: Basic Concepts
1) Identify which reaction is oxidation and which is reduction, 2) write the balanced equation
Mg2+(aq) + 2e- → Mg(s)
Pb2+(aq) + 2e- → Pb(s)
EºMg2+|Mg = V
EºPb2+|Pb = V ox
red
Mg(s) → Mg2+(aq) + 2e-
Pb2+(aq) + 2e- → Pb(s)
Mg(s) + Pb2+ → Mg2+(aq) + Pb(s)

36. Chapter 21: Electrochemistry
Table of Contents
Chapter 21: Electrochemistry
21.2 – Types of Batteries
21.3 – Electrolysis
In class we will cover only what you need to know for the test. Please use your book to answer questions for WebAssign.

37. He piled several cells together to make the first battery
Electrochemistry: Basic Concepts
21.2: Batteries
Alessandro Volta:
If one cell generates a current, several cells should make a larger current
He piled several cells together to make the first battery
Batteries are one or more electrochemical cells in a single package that generates electrical current

40. Electrochemistry: Basic Concepts
21.3: Electrolysis
Electrolysis is using electric energy to bring about a chemical reaction.
Electrolytic cell: An electrochemical cell in which electrolysis occurs.
Electrolysis forces a current in the reverse direction (a nonspontaneous reaction) by passing an electric current through it (recharging)

41. An electrolytic cell is just the opposite of a voltaic cell
Electrochemistry: Basic Concepts
Voltaic Cell
Electrolytic Cell
Voltage source e- e- e- e-
An electrolytic cell is just the opposite of a voltaic cell

42. Reduction of silver ions onto cheaper metals forms silverplate.
Electrochemistry: Basic Concepts
Electroplating
Reduction of silver ions onto cheaper metals forms silverplate.
Click box to view movie clip.

43. Practice problems
Finish Ch 21 Practice Problems (due at the end of class)
Begin WebAssign (due Monday, 11 pm)
You will need your book (21.2 & 21.3) for some questions on WebAssign. Now is a great time to look those up.
Review for test is Monday!
Test on Ch 10, 20, 21 is Tuesday!

44. End of Ch 21

45. The Electrolysis Process
Electrochemistry: Basic Concepts
The Electrolysis Process
Click box to view movie clip.

46. Electrochemistry: Basic Concepts
Electrolytic Cell

47. Distinguish between a voltaic cell and an electrolytic cell.
Basic Assessment Questions
Distinguish between a voltaic cell and an electrolytic cell.
spontaneous redox reaction; nonspontaneous redox reaction by electrolysis

48. Reduction of silver ions onto cheaper metals forms silverplate.
Electrochemistry: Basic Concepts
Electroplating
Reduction of silver ions onto cheaper metals forms silverplate.
Click box to view movie clip.

49. Chapter 21: Electrochemistry
Table of Contents
Chapter 21: Electrochemistry
21.2 – Types of Batteries
In class we will cover batteries only briefly. You will not be tested on this. Don’t take notes. Please use your book to answer questions for WebAssign.

50. He piled several cells together to make the first battery
Electrochemistry: Basic Concepts
Batteries
Alessandro Volta:
If one cell generates a current, several cells should make a larger current
He piled several cells together to make the first battery
Batteries are one or more electrochemical cells in a single package that generates electrical current

51. The object to be plated is made the cathode.
Electrochemistry: Basic Concepts
Electroplating
The object to be plated is made the cathode.
At the pure silver anode, oxidation of silver metal to silver ions replaces the silver ions removed from the solution by plating at the cathode.

52. Electrolytic Cleaning
Electrochemistry: Basic Concepts
Electrolytic Cleaning
Electrolysis can be used to clean objects by pulling ionic dirt away from them.
The electrolysis cell for this cleaning process includes a cathode that is the object itself, a stainless steel anode, and an alkaline electrolyte.
When an electric current is run through the cell, the chloride ions are drawn out.
Hydrogen gas forms and bubbles out, helping to loosen corrosion products.

53. Basic Assessment Questions
Electrical current must be applied to the iron fork
To electroplate an iron fork with silver, which of the following must be true?
The silver electrode must have more mass than the fork
The iron fork must act as the anode of the cell
Electrical current must be applied to the iron fork
Iron ions must be present in the cell solution

54. Batteries Perform Work
Electrochemistry: Basic Concepts
Batteries Perform Work

55. Electrochemistry: Basic Concepts
Modern Batteries

56. Other batteries may have a dozen or more cells.
Electrochemistry: Basic Concepts
Modern Batteries
Although the term battery usually refers to a series of galvanic cells connected together, some batteries have only one such cell.
Other batteries may have a dozen or more cells.

57. Electrochemistry: Basic Concepts
Dry Cells
A dry cell is an electrochemical cell in which the electrolyte (electrode) is a moist paste

58. Electrochemistry: Basic Concepts
Carbon-Zinc Dry Cell
A standard D battery is shown both whole and cut in half to reveal the structure of the carbon-zinc dry cell.
Beneath the outside paper cover of the battery is a cylinder casing made of zinc.
The zinc serves as the anode and will be oxidized in the redox reaction. 

59. manganese (IV) oxide (MnO2) and carbon black—acts as a cathode.
Electrochemistry: Basic Concepts
Carbon-Zinc Dry Cell
The carbon rod in the center of the cylinder— surrounded by a moist, black paste of
manganese (IV) oxide (MnO2) and carbon black—acts as a cathode.
Ammonium chloride (NH4Cl) and zinc chloride (ZnCl2) serve as electrolytes.

60. Electrochemistry: Basic Concepts
Carbon-Zinc Dry Cell
Alkaline batteries contain potassium hydroxide (KOH) in place of the ammonium
chloride electrolyte, and they maintain a high voltage for a longer period of time.

61. Nickel-cadmium rechargeable batteries
Electrochemistry: Basic Concepts
Dry Cell
Primary batteries contain redox reactions that are not easily reversed. They must be discarded when the reaction is complete.
Secondary batteries contain redox reactions that are easily reversed. They can be recharged.
Storage batteries
Nickel-cadmium rechargeable batteries

62. Lead Storage Batteries
Electrochemistry: Basic Concepts
Lead Storage Batteries
Anode = grids of porous lead
Cathode = grids of lead (IV) oxide
Produces water and lead (II) sulfate (PbSO4)

63. Lead Storage Batteries
Electrochemistry: Basic Concepts
Lead Storage Batteries
This battery is rechargeable because one can reverse the reaction
Also called lead-acid batteries because the electrolyte is a solution of sulfuric acid
Often used for car batteries

64. Aqueous Lithium Battery
Electrochemistry: Basic Concepts
Aqueous Lithium Battery
Lightweight batteries that store a large amount of energy for their size.
Lithium:
lightest known metal
Lowest SRP
Oxidation of Li

65. Aqueous Lithium Battery
Electrochemistry: Basic Concepts
Aqueous Lithium Battery
Primary & secondary batteries
About 3V
Long lasting
Watches, computers, cameras, etc.

66. Burning fuel is an oxidation-reduction reaction
Electrochemistry: Basic Concepts
Fuel Cell
Burning fuel is an oxidation-reduction reaction
A fuel cell is a voltaic cell in which the oxidation of a fuel is used to produce electric energy.
Burns fuel in a controlled way while converting chemical energy to electric energy, not heat energy
Cell runs as long as there is fuel