1. Aim Redox 1 – Why is redox so important in your life?

2. Moving Electrons and You
Cell phones require electricity
Electricity – the movement of electrons
Batteries in cell phones are composed of metals and other electrolytes that allow electrons to flow through the cell phone

3. Oxidation States
Remember – electron movements are also why we have bonding in many compounds
If you lose an electron,
you lose negative charge
You become positive or more positive (less negative)
Metals tend to do this
Example: if sodium loses an electron:
Na0  Na e-
If you gain an electron,
you gain negative charge
You become negative or more negative (less positive)
Nonmetals tend to do this
Example: if chlorine gains an electron:
Cl0 + 1 e-  Cl-1

4. Oxidation States A Review of the Rules of Oxidation States or Numbers
Free elements are always zero: Fe, Cl2 , Ca
Ions are the charge assigned to them: Fe 3+ Ca 2+
polyatomic ions (Table E): NO3- SO42-
Some elements have only one oxidation state: Group 1 and 2
Some elements usually have a particular oxidation state with exceptions:
oxygen has a -2 oxidation state
Hydrogen has a +1 oxidation state
Compounds’ oxidation numbers always add up to zero

5. In chemical reactions, each element in each compound has an oxidation state
Examples
3H2 + N  2NH3
H = 0 N = N = -3, H = +1
Who loses electrons? H0  H e-
Who gains electrons? N0 + 3 e-  N-3 Cu(NO3)2 + Mg  Mg(NO3)2 + Cu
Cu = Mg = Mg = Cu = 0
N = N = +4
0 = O = -2
Who loses electrons? Mg0  Mg e-
Who gains electrons? Cu e-  Cu0

6. Oxidation-Reduction Reactions (Redox)
In a redox reaction:
Some substances are lose electrons
They are OXIDIZED
Oxidation: the loss of electrons, become more positive
LEO – Lose Electrons: Oxidation
Some substances are gain of electrons
They are REDUCED
Reduction: the gain of electrons, become more negative
GER – Gain Electrons: Reduction

7. Identifying Redox Rxns
Redox - another type of reaction
Elements in the compound change oxidation states
Example: N2 + 3H2  2NH3
(-3)(+1)
Nitrogen gains e-, becomes more negative
Hydrogen loses e- becomes more positive
Example: 2H2O  O2 + 2H2
(+1)(-2)
Hydrogen gains e-, becomes more negative
Oxygen loses e- becomes more positive

8. Half Reactions in Redox Reactions
Half reaction - reactions that show either a gain or loss of electrons
Example 1: What are the half reactions?
2Mg(s) + O2(g)  2MgO(s)
Oxidation half reaction:
LEO – someone loses electrons
2Mg0  2Mg2+ + 4e-
Reduction half reaction:
GER – someone gains electrons
O20 + 4e-  2O2-

9. Half Reactions in Redox Reactions
Example 2:
3Cu(s) + 2Al(NO3)3(aq)  3Cu(NO3)2(aq) + 2Al(s)
Oxidation half reaction:
LEO – someone loses electrons
Cu0  2Cu e-
Reduction half reaction:
GER – someone gains electrons
Al e-  Al0

10. What are the half reactions in each?
N2 + 3 H2  2 NH3
Reduction: N0 + 3 e-  N -3
Oxidation: H0  H+1 + e-
Mg + 2 HCl  H2 + MgCl2
Oxidation: Mg0  Mg e-
Reduction: H+1 + e-  H0
2 H2O  2 H2 + O2
Oxidation: O-2  O0 + 2 e-
Reduction: H+1 + e-  H0
CH4 + 2O2  CO2 + 2 H2O
Reduction: O0 + 2 e-  O -2
Oxidation: C-4  C e-

11. Aim Redox 2 – What metals make a better battery?

12. Metal Reactivity
Table J – shows the activity of metals and nonmetals in terms of moving electrons
Active metals can give up electrons to less active metals
Related to the position on Table J
Higher metals oxidize (lose electrons) more easily
to lower ones (which gain the electrons and are reduced)

13. Zn oxidized and Cu reduced K oxidized and Mg reduced
Metal Reactivity
Examples: which metal is oxidized/reduced in each of the following pairs:
Zn oxidized and Cu reduced
K oxidized and Mg reduced
Cu reduced and Fe oxidized
Acid reactivity can also be
determined from Table J
Metals above hydrogen will react with acids to produce H2 gas
Which metals don’t react with acids?

14. Why is Table J important?
1. It describes corrosion
Active metals corrode easily
Corrosion: loss of metallic properties due to action of air, water, and chemicals
Two examples:
Rust: 4Fe + 3O2  2Fe2O3
Aluminum foil: forms a thin coating of aluminum oxide that protects the inner aluminum

15. Why is Table J important?
2. It explains how electrons can flow in electrochemical cells (aka voltaic cells or batteries)
metals will react spontaneously to move electrons between them
More active metals are oxidized
Less active metals are reduced
The movement of electrons between the two metals creates an electric current

16. Why is Table J important?
3. It describes how energy can be added to metals in various electrolytic cells
Electrolysis – the separation of elements in a compound using electricity
Example: 2NaCl  2Na0 + Cl20
Oxidation: 2Na+1 +2e-  2Na0
Reduction: 2Cl-1  Cl20 + 2e-

17. Why is Table J important?
Hydrolysis – the separation of hydrogen and oxygen from water using electricity
Example: 2H2O  2H20 + O20
Oxidation: 2O-2  O20 + 4e-
Reduction: 4H+1 + 4e-  2H20
Electroplating – the coating of one metal on to another through electron movement
Requires a power source
Example: Ag+1 + 1e-  Ag0

18. Spontaneous reactions
Movement of electrons from an active metal to a less active one
Produces electricity and electron flow
Batteries
Non-spontaneous reactions
Movement of electrons from a less active metal to a more active one
Requires electricity and electron flow
Battery chargers, electrolytic cells, hydrolysis

19. Agents of Redox
Reducing agents
lose electrons to other atoms
Are themselves oxidized
Oxidizing agents
gain electrons from other atoms
Are themselves reduced
In the rxn,
CH4 + 2O2  CO2 + 2 H2O
Carbon is being oxidized: C-4  C e-
Therefore, it is the reducing agent (gives up e-)
Oxygen is being reduced: O0 + 2 e-  O -2
Therefore, it is the oxidizing agent (takes up e-)

20. Aim Redox 3c – What is electrochemistry and how can it help us build a better battery?

21. Electrochemical cells
During a single replacement reaction, more active metals transfer electrons to less active metals
the more active metal is oxidized
the less active metal is reduced
If the oxidation and reduction half reactions are physically separated and attached by a wire
electrons will flow through the wire during the reaction

22. Electrochemical cells - Parts
Half cells - separate containers in which each half reaction occurs electrodes
Anode – the metal electrode where oxidation occurs (this metal is oxidized, lose e-) (AN OX)
Cathode – the metal electrode where reduction occurs (this metal is reduced, gains e-) (RED CAT)
U-tube or salt bridge - lets ions travel between half cells to complete the circuit
Electrolyte - Carries the current in its ions
Voltmeter – measures the voltage or flow of electricity from the elechemical cell

23. Electrochemical Cell Parts

24. Electrochemical cells - Function
In each half cell - one half of the redox reaction occurs
Oxidation of the metal at the anode breaks down the metal into ions as it loses electrons
Reduction of the metal at the cathode adds electrons to the metal ions in solution, making them solid metal
Electrons flow from the oxidation side, across the wire, through the electrical device, and to the reduction side
This continues until the flow of electrons stops
This occurs when the cells reach equilibrium

25. Examples of electrochemical cells
Voltaic Cells (Spontaneous Rxns) or Batteries
Definition - a system that uses a chemical reaction to produce electricity
The cathode where reduction occurs is at the positive electrode (CPR)
The anode where oxidation occurs is at the negative electrode (ANO)

26. Examples of electrochemical cells
Automobile battery or lead acid storage battery
The reaction is as follows:
2H2SO4 + Pb + PbO2  2PbSO4 + 2H2O + energy
Pb0 Pb Pb 2+
Half reactions:
Oxidation at anode: Pb0  Pb e-
Reduction at cathode: Pb 4+ +2e-  Pb2+
The reverse reaction is how we recharge the battery:
2PbSO4 + 2H2O + energy  2H2SO4 + Pb + PbO2

27. Electrolytic cells (Nonspontaneous Rxns)
Remember – you cannot find alkali metals and halogens in nature as free elements
They must be separated by electrolysis from compounds
Electrolysis of molten sodium chloride
electricity + 2NaCl  2Na0 + Cl20

28. Examples of Electrochemical Cells
Electrolytic cells
Electroplating
Formation of a metal layer on to a surface
Requires addition of energy
Electricity is the source of electrons
Example - Silverplating
Electricity + Ag+  Ag0