1. ELECTROCHEMISTRY Oxidation-reduction or REDOX reactions result in the
To play the movies and simulations included, view the presentation in Slide Show Mode.
Oxidation-reduction or REDOX reactions result in the
generation of an electric current (electricity) or recharging
of a battery.

2. Fireflies
Glow to attract mates
Anglerfish
Emit light to attract prey
Squid, jellyfish, bacteria, and shrimp
Luminous
How?
Redox reactions
Transfer of electrons in a redox reaction provides energy

3. Electrochemical Processes
Chemical reactions either release or absorb energy
Energy very often in the form of heat, light and sound but sometimes it is in the form of electricity
Electron transfer reactions or redox reactions
result in the generation of an electric current spontaneously or
are caused by imposing an electric current nonspontaneously

4. Spontaneous Redox Reactions
How do we know if a reaction is spontaneous?
Use the Activity series chart (Reference TABLE J)
Any metal or nonmetal in the series will displace a metal BELOW it from a compound or solution of that compound.
Consequently, Metals higher are easier to react and therefore LOSE their electrons (thus OXIDIZED) and the metal that is displaced from compound is REDUCED
Displacement of a metal or nonmetal by a metal/ nonmetal higher in the series corresponds to a SPONTANEOUS reaction
For example:
Would zinc spontaneously displace copper from a copper II sulfate (CuSO4) solution?
Yes, Zinc is higher on the list than copper and would SPONTANEOUSLY displace copper
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
Zinc is oxidized while the copper is reduced (a redox reaction)

5. Which are spontaneous & if spontaneous then Identify Reaction Products ?
Yes
Li + AlCl3 
Cs + CuCl2 
I2 + NaCl 
Cl2 + KBr 
Fe + CaBr2 
Mg + Sr(NO3)2 
F2 + MgCl2 
Al + LiCl
Yes
Cu + CsCl
NonSpontaneous
Yes
Br + KCl
NonSpontaneous
NonSpontaneous
Yes
Cl2 + MgF2

6. Predicting Spontaneous Redox Reactions
In the 2 beakers a strip of Cu was placed in a solution AgNO3.
Cu + AgNO3 
Is the reaction Spontaneous?
Yes, Ag is below Cu, therefore a Spontaneous reaction would occur
Cu + AgNO3 Cu(NO3)2 + Ag
(Note: the fuzzy material is the silver)
So Who cares!!!!!!! Why is this useful?
Well, this reaction is a redox reaction
Cu + AgNO3 Cu(NO3)2 + Ag
Since copper loses electrons to the silver, if we separate the copper and silver from each other by a wire, we can create BATTERIES
Oxidation Cu ---> Cu e-
Reduction Ag+ + e- ---> Ag

7. Electrochemical Cell: Batteries
A battery is a voltaic cell where a chemical or redox reaction allows flow of electrons (electricity) SPONTANEOUSLY through an external wire
Voltaic (AKA Electrochemical)
Chemical energy changed into electrical energy
SPONTANEOUS – happens naturally or on its own
A chemical reaction PRODUCES an electric current

8. Oxidation Cu ---> Cu2+ + 2e- Reduction Ag+ + e- ---> Ag --- Ag
A strip of copper is in a copper II nitrate solution on one side and silver strip is in a silver nitrate solution on the other side with the silver and copper connected by an external wire. According to our equation before, Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s) Remember Copper is oxidized (loses e-) while the silver is reduced (gains e-) Cu + Ag > Cu Ag Ag
Ag+
NO3-
---
Oxidation
Cu ---> Cu e-
Reduction
Ag+ + e- ---> Ag

9. Two Parts to a Voltaic or Battery Cell
Oxidation on one side and Reduction at the other side
- Consists of two containers (each half cells) – each side with a piece of metal (AKA an electrode) that is placed in aqueous solution of a salt of that metal.
Wire or External Conductor – connects electrodes.
Salt bridge (or U-tube) that connects the aqueous solutions. Ag
Ag+
NO3-
---
Wire: connects the electrodes, carries electrons (electric current)
Salt bridge: Allows positive and negative ions to pass from one cell to another but prevents solutions from mixing completely (maintain neutrality)

10. An Ox Ate a Red Cat Anode – Oxidation
The anode = location for the oxidation (loss of electrons) half-reaction.
Cu ---> Cu e-
Reduction – Cathode
The cathode = location for the reduction (gain of electrons) half-reaction.
Ag+ + e- ---> Ag
Surface at which oxidation or reduction half-reaction occurs.
Anode loses electrons (losses mass or gets smaller) Oxidation: Cu0  Cu2+ + 2e-
while Cathode or place where e- gained (gains mass). Reduction: Ag+1 +1e-  Ag0 Ultimately a battery dies when anode is completely gone

11. Salt Bridge:
Pathway for flow of positive & negative IONS between half-cells.
Necessary to maintain electrical neutrality.
Reaction will not proceed without salt bridge. (Positive ions toward electrode where are being gained and negative ions vice versa)
The metals (Ag & Cu in this case) or ELECTRODES are called:
CATHODE (+)(reduction) & ANODE (-) (oxidation) : Electrodes may be connected by a wire and symbol shown as

12. Electron Flow thru wire from Anode to Cathode
Anode to Cathode (3:29 min) e- e- e-
Anode
Cathode

13. Electrolytic cells (rechargeable) or Electrolysis
Use electrical energy to force a chemical reaction to occur (Nonspontaneous).
Electrolytic cells can be identified by the a battery or a power supply.
Apply an external voltage (using a battery) and Reverse the direction of the electrons from a Redox rxn.
Recogonize by seeing that this usually occurs in one container without a salt bridge

14. Electrolysis - Reversing a Voltaic Cell
Voltmeter is here, but a power source (battery) must be here to recharge (reverse) the reaction
Here’s a voltaic cell that must be reversed (recharged) when it no longer works. This time:
Ag is oxidized & Cu is reduced. Ag
AgNO3(aq)

15. Electrolytic cells - Electroplating
How do we plate silver onto a copper spoon if copper is more likely to oxidize than silver?
Electroplating (47 seconds):
Mark Rosengarten Electrolysis (2:56min)
Copper(Cu)
Use a battery or power supply to force electrons to move in the reverse or nonspontaneous direction.
Ag + AgNO3  Nonspontaneous
Website showing Electrolysis:

16. Electroplating Example of ElectrolysisAg
Example of Electrolysis
One metal is plated onto another
Used to protect the surface of the base metal from corrosion
Or to make objects, such as tableware and jewelry more attractive
An electric current is used to produce a chemical reaction
The object to be plated is the cathode (negative)
The metal used for plating is the anode (positive)
Solution contains ions of element to be plated.
AgNO3, or AgClO3, or AgCl etc.
Oxidation:
Ag ---> Ag+ + e-
Silver ions (Ag+) move toward negative electrode (copper key) & effectively stick or plate to the outside of the key.

17. Electroplating: Ag onto Fe
Fe is above Ag in Table J.
This reaction will not happen spontaneously like in a volatic cell
Thus, use an external energy source – a power supply or a battery – to force reaction to occur.
Symbols to look for:

18. An Ox ate a Red Cat
Anode is still electrode at which oxidation occurs.
Cathode is still electrode at which reduction occurs.
Polarity of electrodes is DIFFERENT between Voltaic Cells & Electrolytic Cells!
Anode is positive.
Cathode is negative.
Look at drawings in Regents questions.

19. A POX on Electrolytic Cells
Anode – Positive – Oxidation
In an electrolytic cell, the polarity is determined by the outside power supply.
The anode is hooked up to the positive terminal and the cathode is hooked up to the negative terminal.
Look at drawings in Regents questions.
ANODE
CATHODE

20. Na+ Cl- Electrolysis of Molten Sodium Chloride
Electrolysis of sodium chloride is the only process for the production of sodium metal
What chemical species would be present in a vessel of molten sodium chloride, NaCl (l)?
Na+
Cl-
Let’s examine the electrolytic cell for molten NaCl.

21. At the microscopic level
Molten NaCl
At the microscopic level - +
battery e-
NaCl (l)
cations
migrate
toward
(-)
electrode
anions
migrate
toward
(+)
electrode
Na+
Cl-
Na+ e-
Cl-
(-)
(+)
cathode
anode
Cl-
Na+
2Cl-  Cl e-
Na+ + e-  Na

22. Electrolysis of Molten NaCl
Why does the NaCl have to be molten?
Which electrode will the Na+ ions be attracted to?
Which electrode will the Cl- ions be attracted to?
Ions are mobile
Negative electrode, in an electrolytic cell the cathode
Positive electrode, in an electrolytic cell the anode

23. Electrolysis of Molten NaCl
Cathode half-cell (-)
REDUCTION Na+ + e-  Na
Anode half-cell (+)
OXIDATION 2Cl-  Cl e-
Overall reaction
2Na Cl-  2Na Cl2
X 2
Non-spontaneous reaction!

24. Electrochemistry: http://www. youtube. com/watch
Electrochemistry: (4:29 min)
Cleaning sterling silver jewelery: (5:48=7 min)

25. Construct a Voltaic Cell with Al & Pb
Use Table J to identify anode & cathode.
Draw Cell, put in electrodes & solutions
Use Al(NO3)3 and Pb(NO3)2 for aqueous solutions
Label anode, cathode, direction of electron flow in wire, direction of negative ion flow in salt bridge, positive electrode, negative electrode.
Negative electrode is where electrons originate. Positive electrode attracts electrons.
Write half reactions for the Al electrode and Pb electrode

26.  - Electron flow  wire Pb = Positive ion flow  Al = cathode anode
Salt bridge  -
Pb+2 & NO3-1
Al+3 & NO3-1

27. What are the half-reactions?
Al  Al e-
Pb e-  Pb
Al metal is the electrode – it is dissolving.
Al+3 ions go into the solution.
Pb+2 ions are in the solution. They pick up 2 electrons at the surface of the Pb electrode & plate out.

28. Overall Reaction 2(Al  Al+3 + 3e-) 3(Pb+2 + 2e-  Pb)
OXIDATION:
Half Reaction
REDUCTION:
Half Reaction
_____________________________
2Al + 3Pb+2  2Al+3 + 3Pb

29. 2Al + 3Pb+2  2Al+3 + 3Pb Al Pb Which electrode is losing mass?
Which electrode is gaining mass?
What’s happening to the [Al+3]?
What’s happening to the [Pb+2]? Al Pb
Increasing
Decreasing