1. Unit 11- Redox and Electrochemistry
Anode
Cathode
Electrochemical cell
Electrode
Electrolysis
Electrolyte
Electrolytic cell
Half-reaction
Oxidation
Oxidation number
Redox
Reduction
Salt bridge
Voltaic cell

2. What’s the point ? REDOX reactions are important in …
30/09/99
What’s the point ?
REDOX reactions are important in …
C3H8O + CrO3 + H2SO4  Cr2(SO4)3 + C3H6O + H2O
Purifying metals (e.g. Al, Na, Li)
Producing gases (e.g. Cl2, O2, H2)
Electroplating metals
Electrical production (batteries, fuel cells)
Protecting metals from corrosion
Balancing complex chemical equations
Sensors and machines (e.g. pH meter)

3. What is redox? Oxidation- loss of electrons by an atom or ion
Reduction- gain of electrons by an atom or ion
**since one can’t occur without the other
Combine terms to Redox
Mnemonic: LEO the lion says GER
Lose Electrons Oxidation
Gain Electrons Reduction

4. Oxidation numbers On periodic table
Determines what is oxidized and reduced in a reaction
If they change it’s a redox reaction
What type of
 reaction is this (besides redox)???

5. Assigning Oxidation numbers
Identify the formula
If element is free (uncombined) its ox # is 0
Monotomic ions- ox # is same as ion charge
Metals in Groups 1,2 and 3 have ox #’s of +1, +2 and +3 respectively
Fluorine is always -1 in a compound
Hydrogen is always +1 unless it’s combined with a metal then it’s -1
Oxygen is usually -2, except when combined with a more electronegative element then it’s +2
*sum of oxidation #’s in a compound must be 0
*sum of oxidation #’s in a polyatomic ion must equal its charge

6. Try these:
HNO3
CO2
K2PtCl6
PCl5
H2SO4

7. Redox reactions
Once you determine oxidation numbers you can see what element was oxidized and what was reduced
Oxidizing agent- substance that was reduced (gained electrons)
Reducing agent substance that was oxidized (lost electrons)
Oxidizing agent (element reduced) because it gains electrons from oxidized element so it allowed for oxidation to take place because it was a place for the electrons to go

8. Half-reactions Show oxidation or reduction of redox rx Ex:
Shows conservation of mass and charge
Charge does not have to be 0

9. Balancing redox rx’s
Assign oxidation numbers to determine what is oxidized and what is reduced.
Write the oxidation and reduction half-reactions.
Balance each half-reaction.
Balance charge by adding electrons.
Multiply the half-reactions by integers so that the electrons gained and lost are the same

10. Example: Cu + AgNO3  Cu(NO3)2 + Ag
Add the half-reactions, subtracting things that appear on both sides.
Make sure the equation is balanced according to mass.
Make sure the equation is balanced according to charge.
Example:
Cu + AgNO3  Cu(NO3)2 + Ag

11. Practical use for redox reactions
Electrochemical cells
Involves a chemical reaction and flow of electrons
2 types:
Voltaic- spontaneous
Electrolytic- requires electric current (nonspontaneous)
Each have 2 electrodes- site of oxidation and reduction
Oxidation occurs at the anode
Reduction occurs at the cathode
An Ox Red Cat
Anode- oxidation, reduction-cathode

12. Voltaic cells
Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop

13. Voltaic cells
Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced. (completes the circuit)
Cations move toward the cathode.
Anions move toward the anode.

14. Voltaic Cells
In the cell, then, electrons leave the anode and flow through the wire to the cathode.
As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.

15. Voltaic Cells
As the electrons reach the cathode, cations in the cathode solution are attracted to the now negative cathode.
The electrons are taken by the cation, and the neutral metal is deposited on the cathode.

16. Activity series helps identify anode and cathode
Metal higher on chart- oxized (anode)
Metal lower on chart- site of reduction (cathode)

17. Determining electric potential
Voltmeter is used
Voltage is compared to the reduction of H which is 0 volts
The more “+” the reading; reduction is more likely
E is positive- happens naturally
E is negaitve- not spontaneous
Standard electrode potentials-Eo= ability for an electrode to gain electrons

18. Reduction potentials for many electrodes has already been measured

19. Cell potentials
At standard conditions can be determined using this equation:
The strongest oxidizers have the most positive reduction potentials.
The strongest reducers have the most negative reduction potentials.
Ecell 
= Ered (cathode) − Ered (anode)

20. Cell Potentials Ered = −0.76 V  Ered = +0.34 V 
For the oxidation in this cell,
For the reduction,
Ered = −0.76 V 
Ered = V 

21. Cell Potentials Ecell  = Ered (cathode) − (anode)
= V − (−0.76 V)
= V

22. Examples of Voltaic Cells: Dry Cells
Dry cells use two electrodes and a “paste” as an electrolyte.
Some pastes are acidic and others are alkaline.
Carbon is generally used as the cathode and zinc as the anode.

23. Lead-Acid Batteries
Lead-Acid batteries usually contain six cells.(2 V each)
The battery contains lead plates, lead oxide plates, dividers, and a sulfuric acid electrolyte.
The lead plate is the anode and the lead oxide plate is the cathode.
Each cell is connected to form one cathode and one anode on the top or side of the battery.

24. Fuel Cells
Fuel cells bring in the oxidizing and reducing agents as gases
Graphite is typically the anode and cathode for the reaction which produced electricity.
Fuel cells are clean and efficient.

25. Corrosion Corrosion is defined as the disintegration of metals.
Corrosion is typically caused by oxygen (O2).
A familiar example of corrosion is iron rusting.
Corrosion is a result of a redox reaction involving a metal.
Iron oxide (rust)

26. Corrosion con’t…

27. Corrosion Prevention
Typical corrosion protection involves plating the iron with another metal.
The production of steel (iron and carbon) reduces the rate of corrosion of the iron.
Aluminum, zinc, titanium are some metals which corrode slowly, or have different properties used to protect iron.

28. …Corrosion Prevention

29. Electrolytic Cells Electricity is used to force a chemical reaction
Electrolysis
Used to obtain metals from molten salts
Starting/keeping a car running
Plating metals
When starting a car- spont rx occurs in battery to provide electricity to start the car
To keep the car running-alternator (non-spont rx) recharges the battery
Also making aluminum
Electrolysis of water into h and o gas

30. Electroplating Item to be plated is cathode
Metal that will plate is anode
Put in solution containing ions- electrolyte

31. Electroplating con’t Benefits Drawbacks Resists corrosion
Improves appearance
Cheaper
Drawbacks
Plating isn’t always even
Can wear off
Solutions are toxic