1. Electrochemistry Review

2. Oxidation Reduction Reactions
Oxidation - A process in which chemical entities lose e-
Cu was oxidized to Cu2+

3. Oxidation Reduction Reactions
Reduction – A process in which chemical entities gain e-
Each Ag+ was reduced to Ag

4. Oxidation Reduction Reactions
reaction in which one reactant is oxidized and the other is reduced is called a redox reaction.
Therefore, in any redox reaction, the number of e- lost must equal the number of e- gained. It must balance.

5. Cu(s) + 2 Ag+ (aq)  2 Ag(s) + Cu2+ (aq)
If we look at the transfer of electrons in this reaction we will see that Cu looses e- and Ag gains e-
Cu(s) Ag+ (aq)  2 Ag(s) + Cu2+ (aq)
Cu becomes Cu2+, a loss of 2 e-
Each Ag+ becomes Ag, a gain of 2 e

6. TIP:
To remember which is oxidation and which is reaction you can use of these memory aids.
LEO says GER
Lose Electrons Oxidation
Gain Electrons Reduction OR
OIL RIG
Oxidation Is Loss – Reduction Is Gain

7. Oxidizing and Reducing Agents
A substance that removes electrons from another substance is known as a oxidizer or a Oxidizing agent
A substance that gives electrons to another substance is known as an reducer or a reducing agent.

8. Steps to Identifying Oxidation and Reduction
Step 1) Identify repeating entities
Write the total ionic equation
Eliminate ions common on both sides
Step 2) Label charges
Uncombined have a charge of 0
Step 3) Identify loss and gain of electrons

9. Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
Example
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
Step 1)
Zn + Cu2+(aq) + SO42-(aq)  Zn2+(aq) + SO42-(aq) + Cu(s)
Step 2)
Zn0(s) + Cu2+(aq)  Zn2+(aq) + Cu0(s)

10. Step 3)
Zn becomes Zn2+, loss of 2 e- (oxidation)
Cu2+ becomes Cu, gain of 2e- (reduction)

11. Oxidation Numbers
The sum of all of the oxidation numbers in a molecule must be zero because the molecule does not have an overall charge.
If it was a polyatomic ion it would have a charge but the oxidation numbers can still be calculated.

12. Rules for Assigning Oxidation #’s
I) The oxidation number of a atom in an uncombined element is always Zero.
Example: Na, K, O2, H2, Cl2, S8, Li,

13. Rules for Assigning Oxidation #’s
II) The oxidation number of a simple ion is the charge of the ion.
Example: Ca2+ = + 2
Na+ = + 1
Cl- = - 1

14. Rules for Assigning Oxidation #’s
III) The oxidation number of Hydrogen in most compounds is +1.
Example: H­2O, H2SO4, NH3, H = +1
Except in metal hydrides: NaH, H = -1

15. Rules for Assigning Oxidation #’s
IV) The oxidation number of oxygen in most compounds is -2
Example : MgO, HNO3, OH- = -2
Except in peroxides: H2O2, = -1

16. Rules for Assigning Oxidation #’s
V) The oxidation number of group 1 elements is +1
Example: Na, Li, = + 1
The oxidation number of group 2 elements is +2
Example: Mg, Ca = + 2

17. Rules for Assigning Oxidation #’s
VI) The sum of oxidation numbers in a compound must = 0
Example: H2O = 2 (+1) + (-2) = 0

18. Rules for Assigning Oxidation #’s
The sum of oxidation numbers in a polyatomic ion must equal the charge of the ion.
Example: OH- = (-2) + (+1) = -1

19. Identifying Redox Reactions Using Oxidation Numbers
Not all reactions are redox
By using oxidation numbers, you can tell if a reaction is redox. A redox reaction will have a change in oxidation numbers where a reaction that is not redox will not.
Example: AgNO3 + HCl  HNO3 + AgCl
Ag+ + NO3- + H+ + Cl-  H+ + NO3- + Ag+ + Cl-

20. Metal Activity Series
The Activity series is an arrangement of metals in order of their tendency to react (become oxidized).
The most reactive metals (most easily oxidized) are at the top of the list.
The least reactive are at the bottom.

21. Metal Activity Series
This allows us to predict whether a single displacement reaction will occur not. For it to occur, the metal must be higher on the list than the metal in the compound that it is trying to displace.

22. Activity series Fluorine Chlorine Bromine Iodine Metals
Decreasing Activity
Halogens
Lithium
Fluorine
Potassium
Chlorine
Calcium
Bromine
Sodium
Iodine
Aluminum
Zinc
Chromium
Iron
Nickel
Tin
Lead
Hydrogen *
Copper
Mercury
Silver
Platinum
Gold

23. The Activity series is used to predict the products of single displacement reactions.
A + BC  AC + B
In general, an element that is higher on the activity series will displace an element that is lower. The lower element is, thus left as a pure metal.

24. Galvanic cell This apparatus is called a galvanic cell.
A device that converts chemical energy from redox reactions into electrical energy.

25. Galvanic cell This apparatus is called a galvanic cell.
A device that converts chemical energy from redox reactions into electrical energy.

26. Galvanic cell
A Galvanic cell is a spontaneous reaction
A reaction that proceeds on its own without outside assistance (energy).
The oxidation of zinc and reduction of copper occur in separate beakers called half cells.
One of the two compartments in a galvanic cell
Composed of an electrode and a electrolytic solution.

27. Galvanic cell The metal in each beaker is called a electrode.
A solid electrical conductor where the electron transfer occurs.

28. Galvanic cell
Each electrode has a special name
Anode – The electrode where oxidation occurs (think of anion, a negative ion)
Cathode – The electrode where reduction occurs (think cation, a positive ion)
REDCAT (Reduction Cathode)
Metals and non-reactive conductors such as graphite are often used as conductor electrodes.

29. How does a salt bridge work?
How does a salt bridge work? (do not look at text, they used the wrong metals)
The purpose of the salt bridge is to provide ions to prevent charge from building up.
In a way it is much like simple diffusion.

30. How does a salt bridge work?
Every time a zinc atom is oxidized to an ion it would make the solution more positive, which would then stop the reaction. The nitrate from the salt bridge moves in and balances it.
On the other side, every time a copper ion is reduced it would make the solution more negative and stop the reaction. A sodium on then moves in a negative nitrate ion leaves. This allows the circuit to continue without a build up of charge.

31. Cell Reactions
The chemical equation for this reaction can be broken down into 2 parts, called half cell reactions.
Anode half –reaction = Zn(s)  Zn2+(aq) + 2e- (oxidation)
Cathode half –reaction = Cu2+(aq) + 2e-  Cu(s) (reduction)

32. Cell Reactions
Therefore, as the cell operates the mass of the zinc electrode decreases and the mass of the copper electrode increases.
Anode half –reaction Zn(s)  Zn2+(aq) + 2e-
Cathode half –reaction Cu2+(aq) + 2e-  Cu(s)
Overall cell reaction = Cu2+(aq) + Zn(s)  Cu(s) + Zn2+ (aq)

33. Corrosion
Corrosion – The deterioration of metals as a result of oxidation.

34. A few metals such as copper, zinc and aluminum form protective coatings when they oxidize. This makes them more corrosion resistant than other metals that are lower on the activity series. This is why they are commonly used to coat and protect other metals.

35. Rusting Iron
Rust is produced when iron reacts with oxygen to form rust oxide. This new compound does not stick well to the existing metal and flakes off leaving the metal underneath it to rust. This process continues until all the metal is gone.

36. The Redox Reaction of Rust
A corroding metal is a galvanic cell in which the anode and the cathode are found at different points on the same metal surface.
The metal itself is the conducting material that allows the electrons to flow from the anode to the cathode. The anode is normally starts due to a scratch dent or impurity.

37. The Redox Reaction of Rust

38. The Redox Reaction of Rust
Cathode
O2(aq) + 2H2O(l) + 4e-  4OH-(aq)
Anode
Fe(s)  Fe2+(aq) + 2e-

39. Factors that Affect the Rate of Corrosion
Moisture
Since water takes part in the reaction, it must be present for the reaction to occur. A relative humidity of at least 40% is needed for the reaction to take place.

40. Electrolytes
When salts dissolve in water they become ions which increase the conductivity of water. The chlorine ions also act in a similar manner to a salt bridge in the way they offset the increase of Fe2+ ions at the anode. The sodium ions play a similar role at the cathode as they help to offset the negative charge build up from the hydroxide ions.

41. Contact with Less Reactive Metals
When two different metals come in contact with each other, the more reactive metal becomes oxidized. This is why metal fabricators must use the same type of metal when fabricating materials to avoid corrosion.

42. Mechanical Stress
Bending, shaping, or cutting metal, stresses the structure of the metal which creates weak points. The weak points are then prone to corrosion.

43. Preventing Corrosion Protective Coatings
The simplest method of corrosion resistance is to cover the metal with a protective coating. Once the coating is scratched or exposed the metal will rust, even though the rust may appear to only be at the surface it can actually spread deep into the metal.

44. Galvanizing
The process of coating iron or steel with a thin layer of zinc. This can be done by dipping the metal in a hot vat of molten zinc or by electroplating. When the zinc oxidizes it forms a tough, protective coating.

45. Cathodic Protection
A form of metal corrosion prevention in which the metal being protected is forced to be the cathode of a cell, using either impressed current or a sacrificial anode.

46. Sacrificial Anode
A form of protection where a metal that is more easily oxidized is attached to another metal to protect it. The more reactive metal acts as the anode, thus protecting the other metal by making it the cathode.

47. Sacrificial Anode
This method does not require complete covering of the metal; all it needs is some sort of conductive connection that allows it to pass electrons to the metal that needs protecting. The sacrificial anode will need periodic replacement.

48. Impressed Current
In this method, the metal needing protection is attached to the negative terminal of a power source making it the cathode. Continually pumping electrons into the cathode prevents corrosion