1. ELECTROCHEMISTRY
To play the movies and simulations included, view the presentation in Slide Show Mode.

2. Electrochemistry Background
Many of the things people deal with in real life are related to electrochemical reactions.
Batteries - flashlights, watches, car batteries, calculators, cell phones, garage door openers.
Aluminum cans – aluminum is extracted by an electrochemical process.
Chrome – found on cars or motorcycle parts is electroplated on the item.
Therefore, this field of chemistry is often called ELECTROCHEMISTRY

3. Electron Transfer Reactions
Redox reactions – reactions in which there are a simultaneous transfer of electrons from one chemical species to another.
Composed of two different reactions.
1) Oxidation reaction - loss of electrons
2) Reduction reaction – gain of electrons
These reactions are coupled, as the electrons that are lost in the oxidation reaction are the same electrons gained in the reduction reaction. Redox reaction.

4. You can’t have one… without the other!
Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons.

5. How to remember the terminology
LEO the lion says GER!

6. Oxidation – Losing Electrons
Oxidation has three definitions:
The loss of electrons
The gain of oxygen atoms
The loss of hydrogen atoms
In electrochemistry we will deal primarily with the definition that describes the loss of electrons.

7. The loss of electrons Na Na+ + e-
One way to define oxidation is where a chemical substance loses electrons going from reactant to product during a reaction.
For example, when sodium metal reacts with chlorine gas to form sodium chloride (NaCl), the sodium metal loses electrons, which the chlorine gains.
Na Na+ + e-
The sodium metal has been oxidized

8. Reduction – Gaining Electrons
Reduction has three definitions:
The gain of electrons
The loss of oxygen atoms
The gain of hydrogen atoms
In electrochemistry we will deal primarily with the definition that describes the gain of electrons.

9. The gain of electrons Ag+ + e- Ag (metal)
One way to define reduction is where a chemical substance gains electrons going from reactant to product during a reaction.
In the process of electroplating silver onto a teapot, the silver cation is reduced to silver metal by the gain of an electron.
Ag+ + e Ag (metal)
The silver cation has been reduced

10. Examples of Redox Reactions
Consider the example of the reaction where copper metal reacts with a silver nitrate solution:
Cu(s) + 2 Ag Cu Ag(s)
This overall reaction is really composed of two half-reactions:
Cu(s) Cu e- (oxidation)
2 Ag e Ag(s) (reduction)

11. Examples of Redox Reactions
Consider the example of the reaction where zinc metal reacts with a copper(II) sulfate solution:
Zn(s) + Cu Zn Cu(s)
This overall reaction is really composed of two half-reactions:
Zn(s) Zn e- (oxidation)
Cu e Cu(s) (reduction)

12. Terminology for Redox Reactions
OXIDATION—loss of electron(s) by a species
REDUCTION—gain of electron(s)
OXIDIZING AGENT—electron acceptor; substance that is reduced. Copper cation in the last slide allows oxidation of zinc.
REDUCING AGENT—electron donor; substance that is oxidized. Zinc in the last slide allows reductions of copper.
Both the oxidizing and reducing agents are on the left (reactant) side of the redox equation

13. Why Study Electrochemistry?
Batteries
Corrosion
Industrial production of chemicals such as Cl2, NaOH, F2 and Al
Biological redox reactions
The heme group

14. OXIDATION-REDUCTION REACTIONS
Direct Redox Reaction
Oxidizing and reducing agents in direct contact.
Cu(s) Ag+(aq)
Cu2+(aq) Ag(s)

15. OXIDATION-REDUCTION REACTIONS
Indirect Redox Reaction
galvanic or voltaic cell
A battery functions by transferring electrons through an external wire from the reducing agent to the oxidizing agent.

16. Galvanic Cells
An apparatus that allows a redox reaction to occur by transferring electrons through an external connector (wire).
voltaic or galvanic cell: Chemical reaction produces an electric current
electrolytic cell: Electric current used to cause chemical change.
Batteries are voltaic cells

17. Basic Concepts of Galvanic Cells
Anode (-)
Cathode (+)

18. CHEMICAL CHANGE ELECTRIC CURRENT
With time Cu metal plates out and the Zn strip “disappears.”
Zn is oxidized: Zn(s) Zn2+(aq) + 2e-
Cu2+ is reduced: Cu2+(aq) + 2e Cu(s)

19. CHEMICAL CHANGE ELECTRIC CURRENT
To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire.
This is accomplished in a GALVANIC or VOLTAIC cell.
A group of such cells is called a battery.

20. Zn --> Zn2+ + 2e-
Cu2+ + 2e- --> Cu
Oxidation
Anode
Negative
Reduction
Cathode
Positive
<--Anions
Cations-->
RED CAT
•Electrons travel thru external wire.
Salt bridge allows anions and cations to move between electrode compartments.

21. Terms Used for Voltaic Cells

22. CELL POTENTIAL, E E˚cell > 0
A quantitative measure of the amount of electricity (volts) that the voltaic cell can produce.
E˚cell = E˚cathode + E˚anode
E˚cell > 0

23. CELL POTENTIAL, E
For Zn/Cu cell, potential is V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M.
This is the STANDARD CELL POTENTIAL, Eo
—a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C.

25. Calculating Cell Voltage
Balanced half-reactions can be added together to get overall, balanced equation.
Zn(s) ---> Zn2+(aq) + 2e- (oxidation)
Cu2+(aq) + 2e- ---> Cu(s) (reduction)
Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s)
If we know Eo for each half-reaction, we could get Eo for net reaction.

26. TABLE OF STANDARD REDUCTION POTENTIALS
oxidizing
ability of ion E o
(V) Cu 2+
+ 2e Cu
+0.34 Zn
+ 2e Zn
-0.76
reducing ability
of element
To determine an oxidation from a reduction table, just take the opposite sign of the reduction!

27. Zn/Cu Electrochemical Cell+
Anode, negative
Cathode, positive
Zn(s) ---> Zn2+(aq) + 2e- Eo = V
Cu2+(aq) + 2e- ---> Cu(s) Eo = V
Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s)
Eo = V

28. All ingredients are present. Which way does reaction proceed?
Eo for a Voltaic Cell
Cd --> Cd2+ + 2e- or
Cd2+ + 2e- --> Cd
Fe --> Fe2+ + 2e- or
Fe2+ + 2e- --> Fe
All ingredients are present. Which way does reaction proceed?

29. Eo for a Voltaic Cell From the table, you see
• Fe is lower on the list than Cd (oxidized)
• Cd is higher on the list (reduced)
Since Fe is being oxidized the half- reaction listed in the table as well as the cell potential listed needs to be reversed.
The table lists reduction half-reactions

30. Calculating Cell Voltage
Balanced half-reactions can be added together to get overall, balanced equation.
Fe(s) ---> Fe2+(aq) + 2e Eo = V
Cd2+(aq) + 2e- ---> Cd(s) Eo = V
Cd2+(aq) + Fe(s) ---> Fe2+(aq) + Cd(s)
If we know Eo for each half-reaction, we
could get Eo for net reaction.
Eo = V

31. More About Calculating Cell Voltage
Assume I- ion can reduce water.
2 H2O e- ---> H OH Cathode
2 I > I e Anode
2 I H2O --> I OH H2
Assuming reaction occurs as written,
E˚ = E˚cat+ E˚an= ( V) - ( V) = V
Minus E˚ means rxn. occurs in opposite direction
(the connection is backwards or you are recharging the battery)

32. Charging a Battery
When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal.
In your car, the battery charger is called an alternator. If you have a dead battery, it could be the battery needs to be replaced OR the alternator is not charging the battery properly.

33. Dry Cell Battery Anode (-) Zn ---> Zn2+ + 2e- Cathode (+)
2 NH e- ---> NH3 + H2

34. Alkaline Battery Anode (-): Zn + 2 OH- ---> ZnO + H2O + 2e-
Nearly same reactions as in common dry cell, but under basic conditions.
Anode (-): Zn + 2 OH- ---> ZnO + H2O + 2e-
Cathode (+): 2 MnO2 + H2O + 2e- --->
Mn2O3 + 2 OH-

35. Mercury Battery Common type of battery in watches and pacemakers
Anode:
Zn is reducing agent under basic conditions
Cathode:
HgO + H2O + 2e- ---> Hg OH-

36. Lead Storage Battery Anode (-) Eo = +0.36 V
Pb + HSO > PbSO4 + H+ + 2e-
Cathode (+) Eo = V
PbO2 + HSO H+ + 2e > PbSO4 + 2 H2O

37. Ni-Cad Battery Anode (-) Cd + 2 OH- ---> Cd(OH)2 + 2e- Cathode (+)
NiO(OH) + H2O + e- ---> Ni(OH)2 + OH-

38. H2 as a Fuel Cars can use electricity generated by H2/O2 fuel cells.
H2 carried in tanks or generated from hydrocarbons

39. Stop

40. Balancing Equations for Redox Reactions
Some redox reactions have equations that must be balanced by special techniques.
MnO Fe H Mn Fe H2O
Mn = +7
Fe = +2
Mn = +2
Fe = +3

41. Balancing Equations Cu + Ag+ --give--> Cu2+ + Ag
Consider the reduction of Ag+ ions with copper metal.
Cu + Ag give--> Cu Ag

42. Balancing Equations
Step 1: Divide the reaction into half-reactions, one for oxidation and the other for reduction.
Ox Cu ---> Cu2+
Red Ag+ ---> Ag
Step 2: Balance each element for mass. Already done in this case.
Step 3: Balance each half-reaction for charge by adding electrons.
Ox Cu ---> Cu e-
Red Ag+ + e- ---> Ag

43. Balancing Equations
Step 4: Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires.
Reducing agent Cu ---> Cu e-
Oxidizing agent 2 Ag e- ---> 2 Ag
Step 5: Add half-reactions to give the overall equation.
Cu Ag > Cu Ag
The equation is now balanced for both charge and mass.

44. Add H2O on O-deficient side and add H+ on other side for H-balance.
Balancing Equations
Balance the following in acid solution—
VO Zn ---> VO Zn2+
Step 1: Write the half-reactions
Ox Zn ---> Zn2+
Red VO > VO2+
Step 2: Balance each half-reaction for mass.
Red
2 H+ +
VO > VO H2O
Add H2O on O-deficient side and add H+ on other side for H-balance.

45. Balancing Equations Step 3: Balance half-reactions for charge.
Ox Zn ---> Zn e-
Red e H+ + VO > VO H2O
Step 4: Multiply by an appropriate factor.
Red 2e H VO > 2 VO H2O
Step 5: Add balanced half-reactions
Zn H VO > Zn VO H2O

46. Tips on Balancing Equations
Never add O2, O atoms, or O2- to balance oxygen.
Never add H2 or H atoms to balance hydrogen.
Be sure to write the correct charges on all the ions.
Check your work at the end to make sure mass and charge are balanced.
PRACTICE!