1. a.k.a. Oxidation-Reduction
Redox!
a.k.a. Oxidation-Reduction
a.k.a Electrochemistry

2. Review of Oxidation numbers
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
The oxidation number of oxygen is usually –2. In H2O2 it is –1.
4.4

3. Rules for Assigning Oxidation Numbers
The oxidation number of any uncombined element is zero.
The oxidation number of a monatomic ion equals its charge.

4. Rules for Assigning Oxidation Numbers
The oxidation number of oxygen in compounds is -2, except in peroxides, such as H2O2 where it is -1.
The oxidation number of hydrogen in compounds is +1, except in metal hydrides, like NaH, where it is -1.

5. Rules for Assigning Oxidation Numbers
The sum of the oxidation numbers of the atoms in the compound must equal 0.
2(+1) + (-2) = 0
H O
(+2) + 2(-2) + 2(+1) = 0
Ca O H

6. Rules for Assigning Oxidation Numbers
The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge.
X + 4(-2) = -2
S O
X + 3(-2) = -1
N O
 X = +6
 X = +5

7. Electron Transfer Reactions
Electron transfer reactions are oxidation-reduction or redox reactions.
Results in the generation of an electric current (electricity) or be caused by imposing an electric current.
Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

8. Electrochemical processes are oxidation-reduction reactions in which:
the energy released by a spontaneous reaction is converted to electricity or
electrical energy is used to cause a nonspontaneous reaction to occur 2+ 2-
2Mg (s) + O2 (g) MgO (s)
2Mg Mg2+ + 4e-
Oxidation half-reaction (lose e-)
O2 + 4e O2-
Reduction half-reaction (gain e-)
19.1

9. Terminology for Redox Reactions
OXIDATION—loss of electron(s) by a species; increase in oxidation number; increase in oxygen.
REDUCTION—gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.
OXIDIZING AGENT—electron acceptor; species is reduced. (an agent facilitates something; ex. Travel agents don’t travel, they facilitate travel)
REDUCING AGENT—electron donor; species is oxidized.

10. You can’t have one… without the other!
Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons.
You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation
LEO the lion says GER!
ose
lectrons
xidation
ain
lectrons
eduction
GER!

11. Another way to remember
OIL RIG s s
xidation
ose
eduction
ain

12. Oxidation and Reduction (Redox)
Redox currently says that electrons are transferred between reactants
Mg S → Mg S2-
The magnesium atom changes to a magnesium ion by losing 2 electrons, and is thus oxidized
The sulfur atom is changed to a sulfide ion by gaining 2 electrons, and is thus reduced.

13. Oxidation and Reduction (Redox)
Each sodium atom loses one electron:
Each chlorine atom gains one electron:

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

15. LEO says GER :
- Losing electrons is oxidation, and the substance that loses the electrons is called the reducing agent.
- Gaining electrons is reduction, and the substance that gains the electrons is called the oxidizing agent.
Mg(s) S(s) → MgS(s)
Mg is the reducing agent
Mg is oxidized – loses e-
S is the oxidizing agent
S is reduced – gains e-

16. Not All Reactions are Redox Reactions
- Reactions in which there has been no change in oxidation number are not redox reactions.
Examples:

17. Reducing Agents and Oxidizing Agents
An increase in oxidation number = oxidation
A decrease in oxidation number = reduction
Sodium is oxidized – it is the reducing agent
Chlorine is reduced – it is the oxidizing agent

18. Trends in Oxidation and Reduction
Active metals:
Lose electrons easily
Are easily oxidized
Are strong reducing agents
Active nonmetals:
Gain electrons easily
Are easily reduced
Are strong oxidizing agents

19. Identifying Redox Equations
In general, all chemical reactions can be assigned to one of two classes:
oxidation-reduction, in which electrons are transferred:
Single-replacement, combination, decomposition, and combustion
this second class has no electron transfer, and includes all others:
Double-replacement and acid-base reactions

20. Identifying Redox Equations
In an electrical storm, oxygen and nitrogen react to form nitrogen monoxide:
N2(g) + O2(g) → 2NO(g)
Is this a redox reaction?
If the oxidation number of an element in a reacting species changes, then that element has undergone either oxidation or reduction; therefore, the reaction as a whole must be a redox.
YES!

21. Balancing Redox Equations
It is essential to write a correctly balanced equation that represents what happens in a chemical reaction
Fortunately, two systematic methods are available, and are based on the fact that the total electrons gained in reduction equals the total lost in oxidation. The two methods:
Use oxidation-number changes
Use half-reactions

22. Using half-reactions
A half-reaction is an equation showing just the oxidation or just the reduction that takes place
they are then balanced separately, and finally combined
Step 1: write unbalanced equation in ionic form
Step 2: write separate half-reaction equations for oxidation and reduction
Step 3: balance the atoms in the half-reactions

23. Using half-reactions
continued
Step 4: add enough electrons to one side of each half-reaction to balance the charges
Step 5: multiply each half-reaction by a number to make the electrons equal in both
Step 6: add the balanced half-reactions to show an overall equation
Step 7: add the spectator ions and balance the equation

24. CHEMICAL CHANGE ---> ELECTRIC CURRENT
To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire.
This is accomplished in a GALVANIC or VOLTAIC cell.
A group of such cells is called a battery.

25. - + Galvanic Cells anode oxidation cathode reduction spontaneous
redox reaction
19.2

26. Galvanic Cells
The difference in electrical potential between the anode and cathode is called:
cell voltage
electromotive force (emf)
cell potential
Cell Diagram
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
cathode
19.2

27. Standard Electrode Potentials
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
Anode (oxidation):
Zn (s) Zn2+ (1 M) + 2e-
Cathode (reduction):
2e- + 2H+ (1 M) H2 (1 atm)
Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)
19.3

28. Standard Electrode Potentials
Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.
Reduction Reaction
2e- + 2H+ (1 M) H2 (1 atm)
E0 = 0 V
Standard hydrogen electrode (SHE)
19.3

29. E0 is for the reaction as written
The more positive E0 the greater the tendency for the substance to be reduced
The half-cell reactions are reversible
The sign of E0 changes when the reaction is reversed
Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0
19.3