1. Unit 13: Electrochemistry
(Redox)

2. Oxidation Number (State)
identifies how many electrons are either lost or gained by an atom or ion during a chemical reaction

3. Review from Unit 5: Rules for assigning oxidation numbers
An element by itself has an oxidation number of zero.
Ions have an oxidation number equal to the charge.
Group 1 metals in compounds always have a +1 oxidation number, and Group 2 metals in compounds are always +2.
Fluorine always has an oxidation number of -1 in compounds.

4. Rules for assigning oxidation numbers:
Hydrogen is +1 in compounds unless it is combined with a metal, when it is -1.
Oxygen is usually -2 in compounds, but when combined with F it is +2. Oxygen is -1 in the peroxide ion.
The sum of the oxidation numbers in a neutral compound must be zero.
The sum of the oxidation numbers in a polyatomic ion is equal to the overall charge.

5. How to Recognize a Redox Rxn
Ex) Na + Cl2  NaCl REDOX Ex) Mg + Ni(NO3)3  Mg(NO3)2 + Ni REDOX Ex) H2SO4 + NaOH  Na2SO4 + H2O NOT REDOX

6. How to Recognize a Redox Rxn
Redox occurs when electrons are transferred between reactants
Because of electron transfer, the oxidation numbers of certain atoms or ions change
Helpful Hints:
Must be redox if an atom is by itself on one side but in a cmpd on the other
Single Replacement and Synthesis: always redox
Double Replacement: never redox

7. LEO says GER!! Oxidation Reduction the gain of e- the loss of e-
GER: Gain Electrons Reduction
If a particle is reduced, oxidation number decreases
The particle undergoing reduction is called the oxidizing agent
the loss of e-
LEO: Lose Electrons Oxidation
If a particle is oxidized, oxidation number increases
The particle undergoing oxidation is called the reducing agent
LEO says GER!!

8. Redox Half-Reactions
Used to show how many electrons are lost (oxidation) or gained (reduction)
Half-reactions must obey BOTH the Laws of Conservation of Mass and Charge
# of electrons lost = # of electrons gained
Oxidation: e- written on the right
Reduction: e- written on the left

9. Writing Half-Reactions
Ex) Ca + CuCl2  Cu + CaCl2
Ex) Mg + Ni(NO3)3  Mg(NO3)2 + Ni
Ex) Na + Cl2  NaCl

10. Spontaneous Reactions (Table J)
METALS: Strong tendency to lose electrons
More reactive metal gets OXIDIZED (loses e-)
Less reactive metal gets REDUCED (gains e-)
Element higher up starts as the atom (neutral) and becomes the ion (positive charge)
NONMETALS: Strong tendency to gain electrons
More reactive nonmetal gets REDUCED
Less reactive nonmetal gets OXIDIZED

11. What is an electrochemical cell?
Electrochemical Cell: any device that converts chemical energy into electrical energy or electrical energy into chemical energy
Voltaic (Galvanic): chemical  electrical
Electrolytic:
electrical  chemical

12. Voltaic Cells Chemical energy  electrical energy
Electrical energy is produced by a spontaneous redox reaction
Electrons flow on their own

13. Voltaic Cells AN OX: oxidation at the anode
Anode has a negative charge
Anode decreases in mass as oxidation occurs
RED CAT: reduction at the cathode
Cathode has a positive charge
Cathode increases in mass as reduction occurs
Electrons always flow from anode  cathode

14. Electrolytic Cells Anode: Ag (s)  Ag+ (aq) + 1e-
Uses electrical energy to drive a non-spontaneous redox rxn
Electrons are pushed by an outside power source (battery)
Electrons still flow from anode  cathode, but anode is (+) and cathode is (-)
Anode: Ag (s)  Ag+ (aq) + 1e-
Cathode: Ag+ (aq) + 1e-  Ag(s)
Electroplating

15. Electrolysis of NaCl 2NaCl  2Na + Cl2 Anode: 2Cl-  Cl2 + 2e-
Cathode: 2Na+ + 2e-  2Na