1. What is a redox reaction?
Unit 9: Redox Chemistry
What is a redox reaction?

2. Electrochemistry Electrochemistry:
the study of electricity and how it relates to chemistry.
** There are 2 types of electrochemical cells:
voltaic (aka galvanic, battery): converts chemical energy to electrical energy
electrolytic: converts electrical energy to chemical energy

3. Electricity Electricity: the movement of charges (ions or electrons)
**There are also some chemical reactions that involve a transfer of e-. These reactions are called redox reactions. In redox reactions there are always 2 substances that change ox #’s.
Reduction:
gaining electrons or decreasing in charge
Oxidation:
losing electrons or increasing in charge
Remember “LEO says GER”

4. Practice with Ox #’s (Ex) 4Al(s) + 3O2(g)  2Al2O3(s)
Reducing Agent: causes reduction (or is oxidized)
Oxidizing Agent: causes oxidation (or is reduced)
(Ex) Zn + Ag(NO3)  Ag + Zn(NO3)2
Spectator ions: the ion that does not change charge
(Ex) Ba + Sr(OH)2  Ba(OH)2 + Sr

5. Half Reactions Half Reaction:
show only half of the redox reaction; either the oxidation half or the reduction half.
**Remember for REDUCTION half reactions the electrons go on the reactants side because electrons are being gained.
**For OXIDATION half reactions the electrons go on the products side because electrons are being lost.

6. Practice with half reactions
(Ex) H2 + O2  H2O ox
red
(Ex) Zn + AgNO3  Zn(NO3)2 + Ag
(Ex) HCl + Na  NaCl + H2

7. Balancing redox reactions
You can balance half reactions to show that the number of electrons lost = the number of electrons gained.
(Ex) Cu + Ag(NO3)  Cu(NO3)2 + Ag

8. Practice balancing (Ex) Al + CuSO4  Al2(SO4)3 + Cu
(Ex) H Fe  Fe H2

9. PAI’s
**Sometimes you may have to assign oxidation numbers to the elements in PAI’s. If the PAI only appears on one side of the reaction arrow OR if there is only one other element that changes oxidation numbers, you have to look at the PAI.
(Ex) NH4NO2  N2 + H2O

10. Table J
Table J:
The metal closer to the top of Table J are more reactive and therefore more likely to be oxidized. The nonmetals that are more reactive are more likely to be reduced.
**Any METAL on table J will react spontaneously with a METAL ION that is below it.
**This means that the metal with a charge must be below the neutral metal for there to be a reaction.
**Remember any element that is in a compound is an ION.

11. Practice with Table J
(Ex) According to reference table J, which of the following will react spontaneously with H+?
X + H+  H2 + Xcation
a) Pb c) Fe
b) Sn+2 d) H2
(Ex) Which of the following will replace Ni+2 in the compound Ni(NO3)2?
a) Sn c) Sn
b) Pb+2 d) Cr
(Ex) Which atom/ion pair will Co oxidize spontaneously under standard conditions?
a) Co + Fe c) Co + Cu
b) Co + Pb+2 d) Co + Cr+3

12. Voltaic cell (battery)Zn Cu

13. Voltaic Cell (galvanic cell or battery)
The parts of a voltaic wet cell
Electrode:
Anode:
Cathode:
Salt Bridge:
External Circuit:
Flow of electrons:
Half Cells:
Oxidation Half Cell: Reduction Half Cell:
A piece of metal where ox or red occur
The electrode (metal) where ox occurs ( in size)
The electrode (metal) where red occurs ( in size)
Ions from the salt bridge flow to the solutions completing the circuit
Wire that conducts the electrons
Anode  Cathode
Zn0  Zn+2 + 2e-
Cu+2 + 2e-  Cu0

14. Voltaic cell (battery)
How a voltaic cell works:
Converts chemical energy to electrical energy.
uses a spontaneous redox reaction
electrons move from the metal that is ox to the metal that is red
if we separate the electrodes, the e- travel further and produce more electricity.

15. Voltaic Cell Dry Cell Cathode: Anode: Electrolyte Paste:
Acidic dry cell:
Alkaline dry cell:
Dry Cell is a type of voltaic cell.
**A dry cell is not completely dry, but instead of having a liquid electrolyte, there is an electrolytic paste.
The core
Shell (outside metal)
Acts as the salt bridge and the source of ions
Has an acid mixture as its electrolytic paste
Has an basic mixture as its electrolytic paste

16. Electrolytic Cell
Battery +
Ni0
Ag0

17. Electrolytic Cell Electrolytic Cell Electrode: Electrolyte solution:
Anode:
Cathode:
Battery: Flow of electrons:
External Circuit:
Half Cells:
Reduction Half Rxn: Oxidation Half Rxn:
Allows ions to flow between electrodes.
Must contain ions of the anode for plating
The metal that is used to plate an object
The object that is being plated
Produces energy to force a nonspontaneous rxn to happen
e- start at the anode of the battery and move toward the object being plated
Wire
Ag+ + e-  Ag0
Ag0  Ag+ + e-

18. Electrolytic Cell
Uses a nonspontaneous redox reaction (the rxn will not occur on its own)
Uses electricity to force a chemical reaction to occur
This process is called electrolysis
Used for
Electroplating
Isolation of an element in a compound (electrolysis)
Purification of an element

19. Electrolytic Cell Electroplating: Anode: Cathode: Plating Solution:
**Type of
**Commonly used metals are:
**Factors that must be controlled include:
The process of depositing a thin layer of a metal onto an object
Electrolytic cell
Precious metals like Ag, Cu, Au, Cr, Pt
[ ], pH, electrolyte
Ag0
spoon
Must contain Ag+ for plating the spoon

20. Electrolysis
Electrolysis: chemical decomposition produced by passing an electric current through a liquid or solution containing ions
Battery +
NaCl

21. Similarities and differences
An Ox, Red Cat
e- flow from anode to cathode
e- flow through wire, ions flow through solution
cations ions flow toward the cathode, anions ions flow toward the anode
Anode gets smaller, cathode gets larger
Differences
(V) is a battery, (E) requires a battery
(V) anode is (-) and cathode is (+), (E) anode is (+) and cathode is (-)
(V) produces electricity, (E) uses electricity
(V) is a spontaneous rxn, (E) is a nonspontaneous rxn