1. Redox
&
Electrochemistry

2. You can’t have one… without the other!
REDuction Oxidation=
Redox
Reactions that involve the transfer of electrons
Reduction & Oxidation must occur simultaneously
The total # of e- lost must be = to the total e- gained. • • • • +2 • • -2 • • • Ca + • O • Ca • O • • • •
You can’t have one… without the other!

3. Oxidation Oxidation = lose e-
Any chemical reaction in which an element loses electrons & the oxidation # increases (more positive)
Oxidation # increases -4 -3 -2 -1 +1 +2 +3 +4
Oxidation = lose e-

4. Reduction Reduction = gain e-
Any chemical change in which an element gains electrons & the oxidation # decreases (more negative)
Oxidation # decreases -4 -3 -2 -1 +1 +2 +3 +4
Reduction = gain e-

5. Ger ! LEO says GER : Lose Electrons = Oxidation Sodium is oxidized
Gain Electrons = Reduction
Chlorine is reduced

6. Rules for Assigning Oxidation #’s:
1) Uncombined elements (not bonded with any other element) have an oxidation # of ZERO.
Ex.
H2 + Cl2 = 2HCl
H2 = 0
Cl2 = 0

7. Rules for Assigning Oxidation #’s:
2) Sum of the oxidation #’s in a compound must equal ZERO. +2 -1
Ex.
CaCl2
+2(1) - 1(2)
= 0
Be sure to multiply the oxidation # by the # of atoms indicated by the subscript.

8. Rules for Assigning Oxidation #’s:
3) All metals in Group 1( ) have and oxidation # of +1.
Alkali metals
4) All metals in Group 2( ) have and oxidation # of +2.
Alkali earth
metals

9. Rules for Assigning Oxidation #’s:
5) Fluorine always has an oxidation # of –1. 6) Hydrogen has an oxidation # of +1 in all compounds except:
in metal hydrides hydrogen is -1
(metal and hydrogen)
Ex. +1 -1 +2 -1
LiH
CaH2
+1(1) - 1(1)
+2(1) - 1(2)
= 0
= 0

10. Rules for Assigning Oxidation #’s:
7) In ionic compounds, monoatomic ions have oxidation #’s equal to their ionic charge. +1 -1 +1 -1
Ex.
Ex.
LiBr
KBr

11. Rules for Assigning Oxidation #’s:
8) Oxygen has an oxidation # of –2 in all compounds except: +1 -1
in peroxides (H2O2) oxygen is –1.
+1(2) - 1(2)
= 0 +2 -1
with Fluorine (OF2) oxygen is +2.
+2(1) - 1(2)
= 0

12. Rules for Assigning Oxidation #’s:
9)The sum of the oxidation #’s of all the atoms must equal the charge of the ion. +6 -2 2-
Ex.
SO4
? - 2(4) = -2
? - 8 = -2
6 - 8 = -2
- 2 = -2
So sulfur is a +6.

13. Identifying Redox Reactions:
1) Inspect oxidation numbers from reactants to products
2) Oxidation numbers are located in the upper right hand corner & track the movement of electrons

14. Identifying Redox Reactions:
3) Single replacement reactions are always redox reactions ___ Zn + ___ HCl 
4) Double replacement reactions are NEVER redox reactions
___ NaOH + ___ HCl 

15. Identifying Redox Reactions:
5) Changes in oxidation numbers
oxidation
Ex. +1 -1
2Na + Cl2 = 2NaCl
reduction
Na went from to : Na was +1
oxidized
Cl went from to : Cl was -1
reduced

16. Half-Reactions: 1) Show the EXCHANGE OF ELECTRONS in a redox reaction.
2) Follow the LAW OF CONSERVATION OF MASS. This means that there must be the SAME NUMBER OF ATOMS on both sides of the reaction arrow.
3) Follow the LAW OF CONSERVATION OF CHARGE. In half reactions, the NET CHARGE must be the same on both sides of the reaction arrow.

17. Writing Half Rxns: Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s)
oxidation
Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s)
reduction
1) Assign oxidation numbers
2) Draw brackets to identify oxidation & reduction
3) Write :
Ox:
Red:

18. Writing Half Rxns: Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s)
oxidation
Given:
Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s)
reduction
4) Draw :
Ox:
Mg(s)
Mg2+(aq)
+ 2e-
5) Copy what is on each side of arrow:

19. Writing Half Rxns: Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s)
oxidation
Given:
Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s)
reduction
4) Draw :
Red:
2Ag+(aq)
+ 2e-
2Ag(s)
5) Copy what is on each side of arrow:

20. Writing Half Rxns: Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s)
oxidation
Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s)
reduction
7) Figure out how many e- you have.
8) Balance charge. (Oxid. e- are lost and go on right side; Red. e- are gained and go on the left side)
**Mass & charge are balanced.
Lost e-
leo
(Oxid # )
Ox:
Mg0(s)  Mg+2(aq) + 2e-
Red:
2Ag+(aq) + 2e-  2Ag(s)
(Oxid # )
ger
gain e-

21. Example 1: Leo Ger Ox: Al0  Al+3 + 3e- Red: Fe+2 + 2e-  Fe0
oxidation
__Al + __Fe __Al __Fe
reduction
1) Write ½ rxns.
Lost e-
Ox:
Al0  Al e-
Red:
Fe e-  Fe0
gain e-

22. __Al + __Fe+2 __Al+3 + __Fe
oxidation
__Al + __Fe __Al __Fe
reduction
2) The e- lost MUST = the e- gained.
Lost e- 2
[Al0  Al e-]
Ox:
Red: 3
[Fe e-  Fe0]
gain e-

23. __Al + __Fe+2 __Al+3 + __Fe
oxidation
__Al + __Fe __Al __Fe
reduction
3) Distribute the coefficient. 2
[Al0  Al e-]
Ox:
Red: 3
[Fe e-  Fe0]
[2Al0  2Al e-]
[3Fe e-  3Fe0]

24. __Al + __Fe+2 __Al+3 + __Fe
leo 2 3 2 3
__Al + __Fe __Al __Fe
ger
4) Enter the coefficients into original equation.
[2Al0  2Al e-]
[3Fe e-  3Fe0]

25. Table J and Spontaneous Reactions
Shows how active each metal and nonmetal is.
Elements at the top on Table J are more reactive than the elements below them and replace elements below them in compounds.

26. Yes Li K Li + KCl Cl + Li K - Is Li above K in Table J ?
Example 1:
Li + KCl Cl + Li K
- Let’s look at the metals.
- Is Li above K in Table J ?
Yes Li K
- is more active than .
The Li will replace K in cmpds.
This rxn is Spontaneous !
No energy needed.

27. Yes F2 Cl F2 + 2NaCl 2Na + F Cl2 - Is F2 is above Cl in Table J ?
Example 2:
F NaCl Na + F
Cl2
- Let’s look at the nonmetals.
- Is F2 is above Cl in Table J ?
Yes F2 Cl
- is more active than
The F2 will replace Cl in cmpds.
This rxn is Spontaneous !
No energy needed.

28. Examples 3 & 4: Ca + MgCO3 CO3 + Ca Mg F Cl2 + 2NaF + 2Na Cl2
This rxn is Spontaneous F
Cl NaF Na
Cl2
This rxn is Not Spontaneous

29. Trends in Oxidation & Reduction
Active metals:
  Lose electrons easily
  Are easily oxidized
Active nonmetals:
  Gain electrons easily
  Are easily reduced

30. Electrochemistry
There are 2 types of electrochemical cells, voltaic (galvanic) and electrolytic.

31. Voltaic Cells (Galvanic)
Cells that spontaneous convert chemical energy into electrical energy or electrical current
Examples:
wet cell batteries-lead storage
dry Cell Batteries-remotes, radios

32. Voltaic Cells Electrodes: Cathode REDuction = CAThode
Less active of the 2 metals (Table J)
Spontaneously attracts electrons to it
Positive electrode in a voltaic cell
Reduction takes place at the cathode
REDuction = CAThode

33. Voltaic Cells Electrodes: Anode ANode = OXidation
More active of the 2 metals (Table J)
Spontaneous loses electrons to cathode
Negative electrode in a voltaic cell
Oxidation takes place at the anode
ANode = OXidation
In a voltaic cell anode =

34. Description of Voltaic Cells
Half Cells:
- 2 half cells
- rxns occur in separate vessels
Half cell 1
Half cell 2

35. Description of Voltaic Cells
Half Cells:
- electrode – a metal strip
- wire connects to electrodes
- allows the flow of electrons
wire
electrode
electrode
Half cell 1
Half cell 2

36. Description of Voltaic Cells
Half Cells:
- voltmeter – measures electric current
- salt bridge
- allows the flow of ions
- prevents mixing of 2 solutions
switch
wire
voltmeter
IONS
electrode
electrode
Salt bridge
Half cell 1
Half cell 2

37. Description of Voltaic Cells
1) 2 half cells
Complete Circuit
2) Wire & electrodes
3) Salt bridge
4) voltmeter
switch
wire
voltmeter
IONS
electrode
electrode
Salt bridge
Half cell 1
Half cell 2

38. Anode to cathode Let’s see what happens in a voltaic cell:
- A rxn between Zn & Cu
- How will e- travel?
Anode to cathode
RED CAT e- e- e- e-
AN OX e- e-
ger
leo Cu Zn
red.
oxid.
CATHODE
ANODE
Zn(NO3)2
Cu(NO3)2

39. According to Table J, which metal is oxidized?Zn
Which metal is reduced? Cu e- e- e- e- e- e- + + +
CATHODE +
ANODE + + + + + + + + + + + +
Zn(NO3)2
Cu(NO3)2

40. Electrolytic Cells
Cells that use ELECTRICAL ENERGY to force a NONSPONTANEOUS CHEMICAL REACTION to occur.
There is NO SALT BRIDGE.
You will always see a POWER SOURCE (battery) hooked up to an electrolytic cell which drives the FORCED REACTION

41. Electrolytic Cells CATHODE electrode where ELECTRONS are SENT
the NEGATIVE electrode (opposite of voltaic cell)
electrode where REDUCTION occurs (RED CAT)
ANODE
electrode where ELECTRONS are DRAWN AWAY FROM
the POSITIVE electrode (opposite of voltaic cell)
electrode where OXIDATION occurs (AN OX)

42. Electrolytic Cells Used for ELECTROLYSIS and ELECTROPLATING
Examples: alternator (keeps the car battery replenished with energy)
Used for ELECTROLYSIS and ELECTROPLATING

43. Electrolytic Cells
Electrolysis: the decomposition of a substance with an electric current. Examples: water

44. Electrolytic Cell
Electroplating-the process of adding a layer (plate) of metal on the surface of another object.
Example: Gold plated jewelry & chrome bumpers

45. Electroplating 1) Oxidation occurs at the anode.
2) The metal (Ag) bar loses e-. e- e- e- e-
3) e- flow from the metal to the object (spoon) e- e-
Cathode is always attached to the neg. terminal of battery. e-
anode
cathode

46. Electroplating
4) The spoon gains e- and becomes negative. e- e- e- e- e-
5) The spoon attracts the positive Ag+ ions in the solution. e- e- e- e-
6) The Ag+ ions stick to the negative spoon – plating it with silver. e- e-
anode
cathode

47. Electroplating
7) The Ag bar anode loses e- and will eventually disappear. e- e- e- e- e- e- e-
At the anode:
Ag     Ag+ + e- e- e- + e-
At the cathode:
Ag+ + e-     Ag + + e-
anode + +
cathode

48. Comparison of Voltaic and Electrolytic Cells
Similarities
Both use redox reactions
The anode is the site of oxidation
The cathode is the site of reduction
The electron flow through the wire is from anode to cathode
Differences
Voltaic cells use spontaneous reactions to produce energy (voltmeter)
Electrolytic use non-spontaneous reactions that requires energy (power source)
Voltaic cells the anode is negative and the cathode is positive
Electrolytic cells the anode is positive and the cathode is negative.