1. OXIDATION AND REDUCTION
UNIT 13
OXIDATION AND REDUCTION

2. AIM# 1 : How can we assign oxidation numbers?
In chemical reactions, the goal is for all atoms to reach stability. Remember, in chemistry this is done by achieving an octet.
Example: K2S
AIM# 1 : How can we assign oxidation numbers?

3. Oxidation is defined as the loss of electrons (the charge is increased).
Reduction is defined as the gain of electrons (the charge is reduced).
Helpful hint
LEO says GER
LEO – Loss of Electrons is Oxidation
GER – Gain of Electrons is Reduction

4. Oxidation Numbers: - all atoms are assigned oxidation numbers (states)
** remember these are the charges assigned to an atom of an element
- can be positive
(metal in a compound)
-negative
(nonmetal in a compound)
or neutral (single atom)

5. Rules for assigning oxidation numbers:
Any uncombined element has an oxidation state of zero
Na + Cl2  NaCl
2. Group 1 metals = +1 in compounds
Group 2 metals = +2 in compounds
3. Fluorine is -1 in compounds

6. Hydrogen is +1 in compounds but is -1 when combined with a metal
Halogens (Group 17) are -1 in compounds if they’re the most electronegative element
Hydrogen is +1 in compounds but is -1 when combined with a metal
a. Hydrogen is +1 on HCl but -1 in LiH

7. 6. Oxygen is -2 in compounds
Oxygen is +2 with fluorine (OF2) since fluorine is more electronegative
Oxygen is -1 in H2O2 (hydrogen peroxide)
7. The sum of all oxidation states in compounds is zero
8. The sum of all oxidation states in polyatomic ions is the charge of the polyatomic ion on the reference table (Table E)

8. DETERMINE THE OXIDATION NUMBERS FOR EACH OF THE FOLLOWING:CO
FeO NO
MgO K
Fe2O3
CO2
NO2
N = 0
C = +2, O = -2
Fe = +2, O = -2
N= +2, O= -2
Mg= +2, O= -2
K= 0
Fe= +3, O= -2
C= +4, O= -2
CHALLENGE
Na= +1, C= +4, O= -2
Ca= +2, S= +6, O= -2
O= +2, F= -1

9. CHALLENGE
Na2CO3
CaSO4
OF2
MnO4-

10. AIM# 2: How can we recognize a RedOx reactions?
Redox reactions are used for electrochemistry and are driven by a change in charge
Oxidation- losing electrons, any chemical change in which the oxidation number is increased
2Mg O2  2MgO
Example:
Rusting: occurs when metals react with oxygen. Certain metals corrode more than others
Synthesis (Ex. Can be used to create electricity in fuel cell battery, electrons lost by H2 go through the wire into the device to be powered then into O2 where they are gained – recharge it by adding more H2 to cell)
Decomposition (Ex. Cell carried out by adding electricity to liquid form of NaCl. Na+1 forced to gain electron and reduce to Na0, Cl-1 forced to lose electrons to form Cl0 , electrolytic decomposition
Single Replacement (Ex. Make voltaic cells (more than one connected together make a battery) source of portable electricity called zinc/copper cell, generates 1.10 volts of electricity

11. Reduction- gaining electrons, any chemical change in which an oxidation number decreases 2Fe2O3 + 3C  4Fe + 3CO2

12. 1. A + B  AB what type of reaction?
Redox Reaction types:
1. A + B  AB what type of reaction?
Ex. 2H2 + O2  2H2O
2. AB  A + B what type of reaction?
Ex. 2NaCl  2Na + Cl2
3. A + BC  AC + B what type of reaction?
Ex. Zn + Cu(NO3)2  Zn(NO3)2 + Cu

13. Rules for recognizing RedOx reactions:
Not all reactions are redox reactions (DR are not redox reactions). Assign oxidation numbers to all elements and see if any have changed from reactant to product.
If an element is alone on one side and in a compound on the other side, its redox.
3. If the oxidation number increased, it lost electrons, therefore it was oxidized
4. If the oxidation number decreased, it gained electrons, therefore it was reduced

14. Rules for Identifying Which Species is Oxidized and Which Species is Reduced:
 EX) Cu + 2AgNO3  Cu(NO3)2 + 2Ag
- Determine oxidation numbers for each
- If there is a polyatomic ion that remains constant on both sides just look up the charge of the polyatomic ion
Agents: Species that is oxidized  reducing agent
Species that is reduced  oxidizing agent
Spectator ion: Ion that does not change its charge

15. Directions: 1) Cu + O2  CuO 2) Al + Cl2  AlCl3 3) N2 + H2  NH3
Assign oxidation states to each species in the reaction
Identify the species being oxidized and species being reduced
Identify the reducing agent, oxidizing agent, and spectator ion (if there is one)
Balance
1) Cu + O2  CuO
2) Al + Cl2  AlCl3
3) N2 + H2  NH3
4) H2 + O2  H2O
5) Na + CaCl2  NaCl + Ca 

16. AIM#3: How can we write half reactions?
Redox reactions may be split into 2 half-reactions, one for oxidation and one for reductions
Oxidation
Fe(s) -> Fe3+(aq) + 3e-
In oxidation half reactions, an atom or ion is losing one or more electrons while the oxidation number increases

17. Reduction
Fe3+(aq) + 3e- -> Fe(s)
In reduction half reaction, an atom or ion is gaining one or more electrons while the oxidation number decreases.
Remember, redox reaction must also follow conservation of CHARGE.

18. Assign oxidation numbers to each element on both side of the reaction
Balancing
It is possible to balance charge in a redox reaction by writing half reaction for both the reduction and oxidation components of the reaction. Below are the steps to balancing.
Assign oxidation numbers to each element on both side of the reaction
MgCl Na  2NaCl + Mg
b. Write the half reaction

19. c. Add in the electrons to achieve a balance of charge
c. Add in the electrons to achieve a balance of charge. The net charge on both sides must be equal. Make sure the number of electrons gained is equal to the electrons lost. d. Rewrite the equation to bring both half reactions together.

20. For each of the following:
- Make sure each balanced
- Assign oxidation numbers
- Write the oxidation half rxn/reduction half rxn
Identify the oxidizing agent and reducing agent
Identify the spectator ion (if there is any)
EX1) 2Li +Zn(NO3)2  2LiNO3 +Zn
EX2) 2K + Cl2  KCl
EX3) SnCl2 + 2FeCl3  SnCl4 + 2FeCl2

21. Aim #4: What are electrochemical cells?
ELECTROCHEMISTRY: the branch of chemistry that deals with the relations between electrical and chemical phenomena. These are chemical reactions that involve the flow of electrons.

22. 1. Voltaic Cell
Spontaneous reactions, which produces a flow of electrons. Use chemical energy to convert to electrical energy Has two surfaces in which conduct electricity Electrodes- sites where oxidation and reduction happen Anode: the electrode at which oxidation occurs Cathode: the electrode at which reductions occurs ** Using Table J: The metal that is more active (higher on the table) will be oxidized.

23. These cells use separate metals, connected with a wire to produce a current. The salt bridge allows for the ions to move through. This balances out the charge. This is a battery!

24. Label this cell using the metals: Cu and Ag

25. PARTS OF A VOLTAIC CELL: VAN- “Voltaic, Anode, Negative”
2 half cells (2 beakers)
2 electrodes (anode and cathode, these are the sites where ox and red occur)
Salt Bridge – MIGRATION OF IONS!!! (upside down U-TUBE). This maintains the charge balance.
Wire – flow of ELECTRONS!
Optional (voltmeter, switch)
Anode- negative Cathode- positive
**FLOW OF ELECTRONS IS AKWAYS ANODE TO CATHODE!!!!!***** (RED CAT GETS FAT)

26. RED CAT sat on AN OX REDUCTION occurs at the CATHODE OXIDATION occurs at the ANODE Remember from Table J - the higher metal will be oxidized; anode - the lower metal is the site of reduction; cathode ** FLOW OF ELECTRONS IS ALWAYS FROM ANODE TO CATHODE**** (RED CAT GETS FAT)

28. a. used to plate metals (silverware, jewelry)
ELECTROLYTIC CELL
These reactions are non- spontaneous, there needs to be a power source. They require an electric current. When you need electric, it called electrolysis. This is used to electroplate
This process is
a. used to plate metals (silverware, jewelry)
b. aids to decompose chemical compounds
2NaCl  Cl2 + 2Na

29. APE- “anode, positive, electrolytic”
The cathode is always what is being coated
anode: positive cathode: negative
Anode is a piece of solver
Ag ions produced by oxidation travel through solution to cathode, reduced back to Ag atoms adhere to metal being plated
Pos Ag ions migrate away from anode because like repels like anode positive in the electrolytic cell
Pos Ag ions migrate through solution to cathode (neg)
External power source forces electrons in wire to travel from anode to cathode