1. Electrochemistry

2. Oxidation-Reduction 2Na (s) + Cl2 (g)  2NaCl (s)
What happened to each of the reactants in this situation?
Neutral atoms became ions by the transferring of electrons, this is a Redox Reaction (Oxidation-Reduction)

3. Oxidation-Reduction
The concept of Oxidation states provides a way to keep track of electrons during Redox reactions
Oxidation states: in a covalent compound it is defined as the imaginary charges the atoms would have if the shared electrons were divided equally between identical atoms bonded to each other, or for different atoms, electrons are assigned to the atom in each bond that has the greater attraction for electrons
Example: H2O

4. Oxidation-Reduction The oxidation state of… Summary Examples
Is it in its elemental form?
Element:0
Na (s) , O2 (g)
A monatomic ion is the same as its charge
Monatomic ion: charge of ion
Na+ , Cl-
Fluorine is -1 in its compounds
Fluorine: -1
HF, PF3
Oxygen is usually -2 in its compounds (exception of peroxides H2O2, and ozoneO3 Both are -1)
Oxygen: -2
H2O, CO2
Hydrogen is +1 in its covalent compounds
Hydrogen: +1
H2O, HCl, NH3

5. Oxidation and Reduction
GER
Gain of Electrons is Reduction
LEO
Loss of Electrons is Oxidation

6. Oxidation-Reduction CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (g)-4 +1 +4 +1 -2 -2
Oxidation is an increase in oxidation state (a loss of electrons) (Oil)
C is oxidized and CH4 is considered a reducing agent
Reduction is a decrease in oxidation state (a gain of electrons) (Rig)
O is reduced and O2 is considered an oxidizing agent

7. Balancing Redox Reactions
It is often difficult to balance Redox reactions by inspection alone , therefore a special technique called Half Reaction Method is used.
A half reaction can be written for both the oxidized and reduced substances
Ce4+ (aq) + Sn2+ (aq)  Ce3+ (aq) + Sn4+ (aq)
Reduction Half Reaction: Ce4+ (aq)  Ce3+ (aq)
Oxidation Half Reaction: Sn2+ (aq)  Sn4+ (aq)

8. Balancing Redox Reactions
In Acidic Solutions:
1. Write separate equations for the oxidation and reduction half reactions
2. For each half reaction:
a. balance all elements except H and O
b. balance O using H2O
c. balance H using H+
d. balance charge using e-
3. If necessary, multiply one or both balanced half reactions to equalize the number of e- transferred
4. Add half reactions, cancel identical species
5. Check that elements and charges are balanced

9. Balancing Redox Reactions
In Basic Solutions:
1. Use the half reaction method for acidic solutions to obtain the final balanced reaction as if H+ ions were present
2. To both sides add OH- ions equal to the number of H+ ions
3. Form H2O on each side, cancel the equal number of H2O molecules that form
4. Check that the elements and charges are balanced

10. Voltaic cells 2Ag1+ + Cu(s) 2Ag(s) + Cu 2+ Electrons are
transferred from
one half-cell to
the other using
an external metal
conductor. Ag Cu
Ag+
Cu2+