redox

REDOX and Electrochemistry

REDOX

The term REDOX comes from the combination of the terms Reduction and Oxidation. This process involves the movement of electrons from one particle to another. Since we can't have one without the other we combine the terms. LEO the lion says GER is a device to remember Losing Electrons is Oxidation and Gaining Electrons is Reduction

A. Oxidation States (a.k.a. oxidation number)

The charge that an atom has, or appears to have when involved in a compound.
The sum of these numbers is zero for a molecule.

Rules for assigning Oxidation States:

If only one is listed in the reference tables, use it.
	Sodium has a +1 charge.

If more than one is listed, use the one that gives the molecule a net charge of zero (or an ion, the charge listed)

In FeCl3 Cl is a -1 therefore Fe must be +3
in order to make the molecule equal zero.

Polyatomic ions (table E) must be broken-up
and charges assigned to each element.

In NaNO3 we need to break down the polyatomic ion (net charge +1)
further to find out the charge on N and O individually

Na= +1

N= +5

O3= (-2x3) -6

Hydrogen will always be a +1 charge unless it is
combined with a metal to form a hydride.

In the compound
Hydrochloric Acid:
HCL
H = +1

In the compound
Sodium Hydride:
NaH
H = -1

Oxygen is a -2 except in peroxides
or when combined with fluorine.

In peroxides
oxygen is a -1
H2O2
H = +1 & O = -1

With Fluorine
Oxygen is a +2
OF2
O = +2 & F = -1

B. Oxidation: loss of electrons (increase oxidation number)

Example:
2 Mg	+ O2	2 MgO
0...	0	+2	-2
2 Mg	+ O2	2 Mg	+ 2 O
Oxidation ½ reaction:
0...		+2
Mg		Mg	+ 2 e-

lose electron

C. Reduction: gain of electrons (decrease oxidation number)

Example:
2 Mg	+ O2	2 MgO
0...	0	+2	-2
2 Mg	+ O2	2 Mg	+ 2 O
Reduction ½ reaction:
0...			-2
O2	+ 4e-		2 O

NH3 + O2 NO + H2O
1. Assign the Oxidation States
-3	..+1		0		+2	.-2		+1	.-2
N.	H3	+	O2		N.	O	+	.H2	O
2. Write the balanced ½ reactions for the oxidation and the reduction
				-3		+2
Oxidation ½:				N		N	+		5 e-
				0					-2
Reduction ½:				O2	+	4 e-			2O
3. Balance the e- in the two ½ reactions
	-3		+2			-3		+2
4 x	(N		N+	5 e-)	=	4N		4N +	20 e-
	0			-2		0			-2
5 x	(O2+	4 e-		2O)	=	5O2 +	20 e-		10 O
4. Place the coefficients into the original equation
.-3	+1	0		+2	-2			+1	-2
4 N	H3	+ 5 O2		4 N	O	+	_?_	H2	0
5. Drop the Oxidation States
4 NH3 + 5 O2 4 NO + ___ H20
6. Balance anything left by inspection
4 NH3 + 5 O2 4 NO + _6_ H20

A.Voltaic Cell (a.k.a. Galvanic)
Makes a Battery

1. The Positive electrode is called the Cathode and is where Reduction takes place. It will grow Fatter as the reaction progresses.

2. The Negative electrode is called the Anode and is where Oxidation takes place. It will grow Smaller as the reaction progresses.

3. The Salt Bridge/Porous Barrier allows for the flow of ions between solutions.

Anode
Oxidation

Reduction
Cathode

Voltaic | Electrolytic | Electrolplating | Electrolysis | Top

B. Electrolytic Cell Uses a Power Supply
The electrolytic cell can be used for Electroplating or Electrolysis
1. The Negative electrode is called the Cathode and is where Reduction takes place. It will grow Fatter as the reaction progresses.
2. The Positive electrode is called the Anode and is where Oxidation takes place. It will grow Smaller as the reaction progresses.
Voltaic \| Electrolytic \| Electrolplating \| Electrolysis \| Top
Electroplating: Using electricity to coat a cheap metal with a more expensive one
A N O X		R E D C A T
Anode Oxidation	Reduction Cathode
Voltaic \| Electrolytic \| Electrolplating \| Electrolysis \| Top
Electrolysis: Using electricity to split fused salts
A N O X		R E D C A T
Anode Oxidation	Reduction Cathode

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Last Revised July 28, 2005