1. Electrochemistry

2. Electrochemistry Electrochemistry
the branch of chemistry that studies the electricity-related application of oxidation-reduction reactions
Redox reactions involve the transfer of electrons from the oxidized substance to the reduced substance
If you place a zinc into
CuSO4, the Cu+2 will come out
of solution (become reduced) and copper metal (the black substance) is deposited onto the zinc metal. The zinc metal is oxidized and goes into solution.

3. Electrochemistry
If you take the temperature of the CuSO4/Zinc solution during the reaction, you would have find that the temperature increased.
However, if we separate the two metals by a porous barrier and connect them by a wire, we can change that heat energy into electrical energy.
This is called a “Daniel” cell after its inventor, John Frederic Daniell, a British chemist

4. Electrochemistry
Electrode: metal conductor used to establish electrical contact with the non-metallic (solution) part of the circuit.
Electrolyte: a substance that dissolves in water to give a solution that will conduct electricity (positive and negative ions – can be an ionic salt, an acid or a base)
Half-cell: a single electrode in a solution of its own ions. One half-cell contains the oxidation reaction and one contains the reduction reaction.

5. Electrochemistry Anode: the electrode where oxidation takes place
Cathode: the electrode where reduction takes place
RED CAT and AN OX

6. Electrochemistry Electrochemical Cell
A system of electrodes and electrolytes in which either chemical reaction produces electrical energy by the flow of electrons
In the cell on the right:
What are the electrodes made of?__________________________________________________________________________________________
What are the,electrolytes?_______________________________________________________________________________

7. Electrochemistry There are two kinds of electrochemical cells:
1. Voltaic
2. Electrolytic

8. Voltaic Cells
Uses redox reactions that are spontaneous (they happen by themselves without any outside energy input)
The electrolyte solutions used are Copper Sulfate and Zinc Sulfate

9. Voltaic Cells
How do we know if the reactions will be spontaneous? We check the metal activity series to see if one metal can displace the other.
According to the activity table, zinc can displace copper in the cell to the right.
Does this make the reaction spontaneous?

10. Voltaic Cells Remember: RED CAT and AN OX LEO the lion goes GER
In this cell, cations (+) in solution are reduced (gain electrons) at the surface of the cathode (RED CAT) to become metal ions.
In this cell, the reduction half reaction is:
Cu+2(aq) e- → Cu(s)

11. Voltaic Cell
RED CAT and AN OX / LEO the lion goes GER
Electrons given up by the zinc at the anode travel through the wire to the cathode.
However, without the salt bridge, the circuit would not be complete and would not work.

12. Voltaic Cell
Anions (negatively charged ions) move toward the anode to replace the electrons that are flowing through the wire.
Cations move toward the cathode since the positive charges have been lost during reduction.

13. Voltaic Cell
This occurs as the sulfate ions (SO4-) move through the salt bridge to the anode and the Na+2 ions move through to the cathode.
The solution in the salt bridge is sodium sulfate

14. Voltaic Cells Now let’s look at it all together:
Zinc loses two electrons at the anode and forms Zn+2 (which goes into solution)
The electrons flow through the wire to the copper cathode where it gains two electrons to form solid copper metal.
Ions flow through the salt bridge to replenish the lost charges.
The flow of electrons through the wire generates electricity.

15. Electrode Potentials (9:19)
Reduction potential: the measurement of the tendency for a half reaction to occur as a reduction half reaction
Standard Reduction Potential – a half-cell potential measured relative to a potential of zero for a standard hydrogen electrode
The hydrogen electrode is used as a reference to determine the reduction potential of other substances.
Called a SHE (standard hydrogen electrode)

16. Calculating Cell Voltage
To determine how much voltage (E0) a voltaic cell will produce, we used the standard reduction potentials (found in a table).
Calculation:
E0cell = E0cathode - E0anode

17. Calculating Cell Voltage
How much voltage is produced in the voltaic cell of copper and zinc?
1. Find the cells in the Table of Reduction Potentials
*******THE MORE POSITIVE E0 is the CATHODE*********
Cu e- → Cu E0 = V
Zn + 2e- → Zn E0 = V
2. Calculate:
E0cell = E0cathode - E0anode
E0cell = (0.34) - (-0.76) = V
WHEN EVER E0 IS POSITIVE, THE REACTION IS
SPOTANEOUS!

18. Additional Information on the Table of Reduction Potentials
- The substances on top of the Table of Reduction Potentials are the most easily reduced (gain electrons)
- This goes along with electronegativity.
Fluorine has a high ability to attract electrons to itself, so its E0 is V.

19. Additional Information on the Table of Reduction Potentials
Oxidation
- The bottom of the table shows which substances are most easily oxidized (loss of electrons).
- Here you will find the Group 1 and Group 2 metals, such as Lithium, which have a low electronegativity and will give up their electrons easily.

20. Additional Information on the Table of Reduction Potentials
Example:
According to the Table of Reduction Potentials, which of the following ions is most easily oxidized?
a. F
b. Cl-
c. Br
d. I-
Answer: I- because it is lower on the list than the other atoms.

21. Additional Information on the Table of Reduction Potentials
Spontaneous Reactions
Substances to the left and those below them on the right react spontaneously with each other
Half Reaction
E0 Volts
F2 + 2e- → 2F-
+2.87
Au+3 + 3e- → Au
+1.50
Cl2 + 2e-→ 2Cl-
+1.36
F2 reacts spontaneously with Cl- and Cl- reacts spontaneously with F2.

22. Additional Information on the Table of Reduction Potentials
Example:
According to the Table of Reduction Potentials, which metal will react spontaneously with Ag+ ions, but not with Zn+2 ions?
Half Reaction
E0 (volts)
Ag+ + e- → Ag
+0.80
Cu e- → Cu
+0.34
Zn e- → Zn
-0.76
Remember: the arrow can only go from the upper left to the lower right or from the lower right to the upper left!

23. This is the end of information for Honors unless your instructor tells you that you must know about electrolytic cells.

24. Electrolytic Cells
If a reaction is NOT spontaneous, we can force the reaction to happen by applying a power source, such as a battery, in the wire between half cells.

25. Electrolytic Cells
The process in which an electric current is used to produce a redox reaction is called electrolysis.
In electrolysis, the cathode and the anode SWITCH.

26. Electrolytic Cells In this cell:
Electrons from the negative end of the battery flow through the wire to the electrode, which is NOW NEGATIVE and BECOMES THE CATHODE.
The electrode attached to the positive end of the battery loses electrons to it and it is NOW POSITIVE and BECOMES THE ANODE.

27. Electrolytic Cells Let’s put it all together:
Electrons move from the battery to the cathode.
Positive ions (potassium ions) move to the cathode to gain electrons and become reduced (GER).
Negative ions (chlorine ions) move toward the anode and give up electrons to replace the ones lost to the battery. They become oxidized (LEO).
The electrons go through the wire back to the battery (positive end).
K e-→ K
RED CAT
at Cathode
2Cl- → Cl2 + 2e-
AN OX
at Anode
K+ Cl-

28. Electroplating An electrolytic cell is used for electroplating.
- an electrolytic process by which a metal ion is reduced and solid metal is deposited on a surface

29. Electroplating
To plate a silver spoon: make the spoon the cathode (GER). The silver ions from anode will gain an electron at the cathode and silver metal will deposit on the spoon.
The electrolyte is Silver Nitrate which provides more silver ions (Ag+).
AgNO3
solution +

30. Comparison of Cells Both have: Differences: Voltaic Cell
Anode – where oxidation takes place (AN OX)
Cathode – where reduction takes place (RED CAT)
Differences:
Voltaic Cell
Electrolytic Cell
Chemical reactions produce electricity
Electricity produces chemical reactions
Anode is negative
Cathode is positive
Anode is positive
Cathode is negative