1. RedOx
Chapter 18

2. Oxidation- Reduction Reactions
Redox or oxidation-reduction reactions are reactions that involve a transfer of electrons.
Oxidation is the loss of electrons.
Reduction is the gain of electrons.
(think of the charge, OIL RIG)
So in the reaction
4 K + O2 → 4 K O2-
Potassium get oxidized, oxygen get reduced

3. Of course… We would normally write this expression
4 K + O2 → 4 K O2-
as…
4 K + O2 → 2 K2O
It doesn’t change anything
Potassium still gets oxidized, oxygen still gets reduced.

4. Oxidation States CH4 +2 O2 → CO2 + 2 H2O
The above reaction also involves a transfer of electrons, so it is a redox reaction.
To see this we need to assign oxidation numbers to each atom present.
Oxidation states (numbers)- hypothetical charge an atom would have if all bonds were ionic.
It is used as a method to keep track of electrons in oxidation reduction reactions.

5. Rules for assigning Oxidation States
1.The oxidation state of an uncombined atom is 0.
2. The oxidation state of a monoatomic ion is the same as its charge.
3. The sum of the oxidation states of a neutral compound must be 0.
4. The sum of the oxidation states of a polyatomic ion must equal its charge.
5. In binary compounds, the element with the greater electronegativity is assigned a negative oxidation state equal to its charge as an ion.

6. Special Elements
6. Diatomic elements are assigned an oxidation state of zero if they are not bonded to anything else.
7. Alkali metals are almost always +1
8. Alkaline Earth metals are almost always +2
9. Halogens are normally -1 unless bonded to a more electronegative halogen
10.Except for the above rules, Oxygen is assigned an oxidation state of -2 unless it in a compound containing peroxide (O22-), then it gets a -1.
11. Hydrogen gets a charge of +1when bonded to nonmetals, it if is bonded to a metal it is -1.

7. Determining Oxidation StatesC F2
AlF3
CO2
H2O2
K3PO4
SO42-
CaCl2
NH3
NaH
CaSiO3

8. Using oxidation states
In the reaction…
2 Na +2 H2O →2 NaOH + H2
Note the changes
Sodium went from 0 to 1
2 of the hydrogen atoms went from +1 to 0 (the other two were unchanged)

9. So… Sodium must have lost 2 electrons 2 Na → 2Na+ + 2 e-
And Hydrogen gained two electrons
2 H2O +2 e-→ 2 OH- + H2
Sodium is oxidized, hydrogen is reduced in this reaction
Oxidation is an increase in oxidation state
Reduction is a decrease in oxidation state

10. Balancing Redox Equations by Half Reactions Method

11. Balancing Equations
Redox reactions don’t follow normal rules for balancing equations because we also have to pay attention of the electron transfer.
For example
Ce4+ + Sn2+ → Ce3+ + Sn4+
Is it balanced?
No, look at the charges

12. Half reactions
Half reactions are exactly what the sound like, half of the reaction.
Half reactions also include electrons, e-, as reactants or products.
So for
Ce4+ + Sn2+ → Ce3+ + Sn4+
we have
Ce4+ + e- → Ce3+
Sn2+ → 2e- + Sn4+

13. Determining # of electrons
Just look at the oxidation numbers.
For
Ce4+ → Ce3+
I have to decrease my charge by one so I must add an electron.
Ce4+ + e- → Ce3+
Sn2+ → 2e- + Sn4+
My charge increases by two, so electrons were lost.

14. Simple Rule Electrons lost must equal electrons gained!
So to make this work
2 Ce4+ +2 e- →2 Ce3+
Sn2+ → 2e- + Sn4+
Put our two half reactions back together to make a “whole” reaction again.
2 Ce4+ + Sn2+ → 2 Ce3+ + Sn4+

15. Balancing Redox reactions in an acid or a base

16. Redox reactions in acidic solutions
I will tell you if it is in an acidic solution.
These have special rules.
Balance all elements except hydrogen and oxygen.
Balance oxygen by adding H2O (which is always prevalent in an acidic solution)
Balance hydrogen by adding H+
Then balance the charge adding electrons and proceed as normal.

17. Example In an acidic solution Cr2O7 2- + Cl- → Cr3+ + Cl2
Half reactions
Cr2O7 2- → Cr3+
Cl- → Cl2

18. Here we go Cr2O7 2- → Cr3+ Cr2O7 2- → 2 Cr3+ Cr2O7 2- → 2 Cr3+ + 7 H2O
Cr2O H+→ 2 Cr H2O
Cr2O H++ 6 e- → 2 Cr H2O

19. Other side Cl- → Cl2 2 Cl- → Cl2 2 Cl- → Cl2 + 2 e-
I have to equal 6 e- so multiply by 3
6 Cl- → 3 Cl2 + 6 e-

20. Combine my half reactions
Cr2O H++ 6 e- → 2 Cr H2O
6 Cl- → 3 Cl2 + 6 e-
And you get
Cr2O H++ 6 Cl- → 2 Cr Cl2 +7 H2O
The electrons cancel out .

21. Example
In an acidic solution
MnO4- + H2O2 → Mn2+ + O2

22. Top Equation

23. Bottom Equation

24. Add them together

25. Balancing Redox Equations in a basic solution
Follow all rules for an acidic solution.
After you have completed the acidic reaction add OH- to each side to neutralize any H+.
Combine OH- and H+ to make H2O.
Cancel out any extra waters from both sides of the equation.

26. Example We will use the same equation as before In a basic solution
MnO4- + H2O2 → Mn2+ + O2
2 MnO H++ 5 H2O2
→ 2 Mn O2 + 8 H2O

27. Basic solution Since this is a basic solution we can’t have excess H+.
We will add OH- to each side to neutralize all H+
2 MnO H++ 5 H2O2 + 6OH-
→ 2 Mn O2 + 8 H2O + 6OH-
We added 6 OH- because there were 6H+

28. Cont.
H+ + OH- → H2O
Combine the hydroxide and hydrogen on the reactant side to make water
2 MnO H2O + 5 H2O2
→ 2 Mn O2 + 8 H2O + 6OH-
Cancel out waters on both sides
2 MnO H2O2
→ 2 Mn O2 + 2 H2O + 6OH-

29. Another example In a basic solution MnO4 − + SO32-→MnO4 2− + SO42-
Half reactions
MnO4 − → MnO4 2−
SO32-→ SO42-

30. Electrochemistry

31. Terminology
You may have noticed oxygen never gets oxidized, it always gets reduced.
The reason for this is because oxygen is an oxidizing agent.
An oxidizing agent is something that causes something else to be oxidized.
An oxidizing agent readily accepts (or takes) electrons from something else.
In the process, the oxidizing agent gets reduced.
A reducing agent is something that causes something else to be reduced.

32. Electrochemistry
~The study interactions of chemical and electrical energy.
Electrochemistry deals with 2 types of processes
1. The production of an electric current from an oxidation reduction reaction
2. The use of an electric current to produce a chemical reaction.

33. Production of Current
Oxidation Reactions involve a transfer of electrons.
Electric current is a movement of electrons.
In order to produce a usable current, the electrons must be forced across a set path (circuit).
In order to accomplish this, an oxidizing agent and something to oxidize must be separated from a reducing agent with something to reduce.

34. Pictures
An Oxidation Reduction reaction in the same container will have electrons transferring, but we can’t harness them.
Separating the oxidation from the reduction, but connecting them by a wire would allow only electrons to flow.
Oxidizing agent
Reducing Agent
Oxidation
Reduction

35. Closer look
X → X+ + e-
X+ + e- → X
Oxidation
Reduction
We now have excess electrons being formed in the oxidizing solution and a need for electrons in the reducing solution with a path for them to flow through.
However, if electrons did flow through the wire it would cause a negative and positive solution to form.

36. That’s not possible Or at least it would require a lot of energy.
A negative solution would theoretically be formed by adding electrons, and a positive one by removing electrons.
The negative solution would then repel the electrons and stop them from flowing in, and a positive solution would attract the electrons pulling them back where they came from.
Making it so the charged solutions wouldn’t form.
In order for this to work, I would need a way for ions to flow back and forth but keeping the solutions mostly separated.

37. Salt Bridge
Salt Bridge- a connector for two solutions previously discussed that allows ions to pass back and forth.
This can be accomplished by a tube filled with an electrolyte (positive and negative ions) or a porous disc connecting the two solutions.

38. e-
Closer look e- e-
Salt Bridge
X → X+ + e-
X+ + e- → X
Oxidation
Reduction
Now electrons can flow across the wire from the oxidation reaction to the reduction reaction.
As the oxidation reaction becomes positive, it removes negative ions and adds positive ions to the salt bridge.
The reduction reaction does the reverse.

39. Closer look Zooming in on the oxidizing side
Salt Bridge
- ion
+ ion
Oxidation Side e-
+ ion
- ion
- ion
+ ion
+ ion
- ion
Zooming in on the oxidizing side
This would make the salt bridge positive…

40. Closer look (Zooming in on the reducing side)
Salt Bridge
- ion
+ ion
- ion e-
+ ion
- ion
+ ion
- ion
Reduction side
+ ion
(Zooming in on the reducing side)
if the reverse wasn’t happening on this side.

41. Close up of salt bridge
+ ion
+ ion
+ ion
- ion
- ion
- ion
- ion
- ion
+ ion
+ ion
+ ion
- ion
The ions keep flowing in the salt bridge to keep everything neutral.
Electrons do also travel across the salt bridge.
This decreases the cell’s effectiveness.

42. Electrochemical battery
This is the basic unit of a battery.
It is also called a galvanic cell, most commercial batteries have several galvanic cells linked together.
Batteries always have two terminals.
The terminal where oxidation occurs is called the anode.
The terminal where reduction occurs is called the cathode.

43. Cell Potential (Ecell)
Cell potential (electromotive force, emf) is the driving force in a galvanic cell that pulls electrons from the reducing agent in one compartment to the oxidizing agent in the other.
The volt (V) is the unit of electrical potential.
Electrical charge is measured in coulombs (C).
 A volt is 1 joule of work per coulomb of charge transferred: 1 V = 1 J/C.
A voltmeter is a device which measures cell potential.

44. How much voltage? Voltage of a cell depends on the half reactions.
You will have a chart of several half reactions reduction potentials for the test.
Obviously you cannot have two reductions. One will need to be turned into an oxidation.
To do that flip the half reaction and flip the sign of the half reaction.

45. The Chart

46. Which reaction to flip Eocell = Eoreduction + Eooxidation
The Ecell (voltage of the cell) will always be positive. If is negative the cell won’t happen on it’s own.
So if you have a reaction of Zn/Zn2+ and Cu/Cu2+.
Zn e-  Zn E = -.76 V
Cu e-  Cu E = .34 V

47. Zinc will need to be flipped to an oxidation to make the cell positive
Zn Zn e- E = .76 V
Cu e-  Cu E = .34 V
Ecell = 1.10 V
The overall reaction of the cell is
Zn + Cu2+  Cu + Zn2+

48. Write the equation for and figure out the electric potential of a cell based on…
Sn4+/Sn2+ & Pb2+/Pb
Zn2+/Zn & Cr3+/Cr
Li+/Li & Co3+/C2+

49. Corrosion

50. Corrosion
Corrosion-An oxidization of a metal, and the oxide flaking off.
Oxidized metal is commonly called rust
Most commonly oxygen will oxidize a metal.
Either by
[Metal] + O2 → [Metal]O Or
[Metal] + H2O → [Metal]O + H2

51. Resisting corrosion
Most metals resist corrosion by an oxide layer forming on the outside that protects the metal inside.
It protects the inside metal by preventing the oxygen (or other oxidizing agent) from being able to reach it.

52. Examples Aluminum very readily loses electrons.
You would expect it to “rust” easily.
However, aluminum is a very useful metal because it doesn’t corrode like other metals can.
An aluminum oxide layer forms on the outside, stopping further oxidation from occurring.
This oxide gives aluminum a dull color.

53. Steel
Steel corrodes very readily because iron oxide doesn’t stick to the surface.
It instead falls off exposing new metal to be oxidized.
This makes iron less useful and explains why ancient people would prefer other metals.
However, the abundance and other properties of iron have made it useful.

54. Preventing oxidation
Iron can be protected by painting the surface or coating it with a different material to prevent the corrosion.
Galvanized steel is steel coated with zinc to prevent oxidation.
Zinc actually oxidizes more readily than iron.

55. Galvanic corrosion
Two different metals placed next to each other with an electrolytic solution connecting will cause an oxidation reduction reaction to occur.
Just like the galvanic cell.
Electrons will flow from a more active metal to a less active metal.
One metal will end up oxidizing the other, but in the process will itself become reduced.
This rapidly oxidized or rusts the one metal but prevents the less active metal from oxidizing (rusting)

56. Galvanic corrosion

57. Galvanic corrosion You can also see galvanic corrosion on a battery.
Batteries that are hooked up to a circuit for an extended period of time tend to become rusted.

58. High temperature corrosion
An oxidation reaction like any other reaction occurs faster when heated.
Metals that are constantly heated tend to rust more quickly.

59. Noble metals There are certain metals that don’t form an oxide.
Gold and silver are noble metals.
Silver will oxidize with sulfur, but not with oxygen.
Gold does not readily oxidize in nature.

60. Electrolysis

61. Electrolysis
Electrolysis-Forcing a current through to produce a chemical reaction.
Water can be electrolyisized
H2O → H2 + O2
This reaction is very important for fuel cell cars.
It uses electricity to create a combustible fuel for an internal combustion engine.

62. Refining metals
Metals are found as metal oxides (ores) in nature commonly.
An electrolysis reaction is commonly used to produce metals from these ores.
Sodium metal can be produced by melting sodium chloride and passing an electric current through the melt.

63. Hall-Heroult Process Before 1886 aluminum was a very expensive metal.
Even though it is very abundant on the Earth’s surface, it is only found as bauxite, an oxide.
Since aluminum is so reactive no reducing agent could easily turn the ore into a metal.
It was so valuable the Napoleon served his honored guests aluminum silverware and gave the others gold or silver.

64. Charles Hall
A student in a chemistry course at Oberlin College in Ohio was told by his professor, that if anyone could a cheap method to manufacture aluminum from bauxite they could make a fortune.
Using crude galvanic cells Charles Hall was able to achieve this using an electrolysis reaction.
Yes, he did make a fortune with it.

65. Batteries

66. Batteries history
Battery- combination on 2 or more electrochemical cells that convert chemical energy into electrical energy.
Luigi Galvini and Allesandro Volta are credited with the invention of the first batteries.
Galvini came up with the galvanic cell. Volta connected them together in a series.
The name battery was coined by Benjamin Franklin, because the batteries at the time were a series of connected jars which reminded him of a battery of cannons.

67. Types of batteries Two major types are:
Wet Cell batteries- use a liquid electrolyte to allow the ions to freely exchange during the redox reaction.
Car batteries or batteries with a liquid inside.
Dry Cell battery- use a paste that immobilizes the electrolyte.
AA, AAA, C, D, 9V etc.

68. The electrolyte This is the salt bridge discussed earlier.
It allows ions to flow freely while the electrons travel across our load, the thing you are trying to power.
The electrolyte normally needs to be acidic or basic to make the redox reaction occur.
Sulfuric acid is commonly used, it is commonly called battery acid.

69. Why not HCl
HCl would be a very poor choice because of the redox reaction
2 HCl → H2 + Cl2
Hydrogen typically gets reduced
2 H+ +2e- → H2
But chlorine getting oxidized is very dangerous
2 Cl- → 2e- + Cl2
Because of the poisonous gas produced.

70. Wet Cell Batteries Car batteries are wet cell batteries.
The obvious problem with these batteries is the need to be keep them upright or the electrolyte, sulfuric acid, will leak out.
However the power they produce is quite substantial.

71. Lead-Acid
The standard battery used in a car was invented in 1859 by Gaston Planté.
It uses a Lead plate and a Lead Dioxide plate in a sulfuric acid solution.
Here is the unbalanced redox reaction
Pb + PbO2 + H2SO4 ⇌ PbSO4
Reduction half
PbO2 + H2SO4 ⇌ PbSO4
Oxidation half
Pb +H2SO4 ⇌ PbSO4

72. Rechargeable
The nice thing about this battery is it is easily rechargeable.
PbSO4 will readily form Pb and PbO2 if electric current is added back to the cell.
This happened completely by chance since there was no practical way to recharge the battery when it was invented.
Later the generator would be invented and from that a car’s alternator and easily recharge the battery while you drive.

73. Alkaline Batteries Normal AA AAA C and D batteries are alkaline.
These are dry cell batteries
The reaction is
Zn + MnO2 →ZnO + Mn2O3
This occurs in a paste of KOH.
This reaction is not reversible!

74. These may leak if you try to recharge them.

75. Strangely enough
A single AA, AAA, C or D “battery” is not a battery by definition.
They are all single cells.
They are not a battery until you connect them together, like you have to in most devices.
A 9 V battery is a battery because it has 6 cells linked together in the rectangular case.
Car batteries also have 6 cells linked together.

76. Lithium Ion Batteries
Commonly used in cell phones, laptops and other portable electronic devices.
Not to be confused with Lithium single use batteries (like energizer e2).
These batteries are rechargeable.
There use a lithium compound as the cathode and variety of possibilities for the anode material.

77. Li-Ion

78. Lithium Ion Batteries
These batteries are very light for the power the produce
They can be built to a variety of shapes to fit their device.
Over time, the battery will not be able to hold as much of a charge so it will need to be recharged more often.
It will take less time to recharge when this occurs.

79. Other batteries
Zinc-carbon battery - Also known as a standard carbon battery, zinc-carbon chemistry is used in all inexpensive AA, C and D dry-cell batteries. The electrodes are zinc and carbon, with an acidic paste between them that serves as the electrolyte.
Nickel-cadmium battery (NiCd)- The electrodes are nickel-hydroxide and cadmium, with potassium-hydroxide as the electrolyte (rechargeable).
Nickel-metal hydride battery (NiMh)- This battery is rapidly replacing nickel-cadmium because it does not suffer from the memory effect that nickel-cadmiums do (rechargeable).

80. Other batteries
Lithium-iodide battery - Lithium-iodide chemistry is used in pacemakers and hearing aides because of their long life.
Zinc-air battery - This battery is lightweight and rechargeable.
Zinc-mercury oxide battery - This is often used in hearing-aids.
Silver-zinc battery - This is used in aeronautical applications because the power-to-weight ratio is good.

81. Recycling All batteries break down over time.
Rechargeable batteries normally produce some other compound through an irreversible reaction.
All batteries contain caustic chemicals that are potentially hazardous to the environment.
None should be put into landfills as they will eventually break down and leak over time.
Car batteries are almost all recycled (like 98%). You can recycle them anywhere that sells car batteries (Autozone, Sears etc.).

82. Where to recycle
Power tool batteries (NiCd/NiMH or Li-Ion) can be recycled at Home Depot. As soon as you walk in, to the left there is a bin.
Electronics batteries (Li Ion) can be recycled at Best Buy. The bin is in that area when you first walk in before you get into the actual store.
The e check is also taking cell phone batteries currently.
Regular batteries can be recycled at the hazardous household waste center in Stow.