Electrochemistry
o REDOX (reduction-oxidation) reactions involve the transfer of electrons.
This transfer can release as heat or energy.
o Moving electrons = current
o There can be no reduction of a substance without an oxidation
OIL: Oxidation is loss of electrons and thus an increase in oxidation state
RIG: Reduction is the gain of electrons and thus a decrease in oxidation state
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An Oxidizing Agent is the reactant that is getting reduced, the substance that
is causing the oxidation process to occur
A Reducing Agents is the reactant that is getting oxidized, the substance that
is causing the reduction process to occur
Calculating Oxidation states
o all atoms in their standard state at room temperature have an oxidation
number of O
o Oxygen will usually have an oxidation number of -2 unless working with the
peroxide ion in which case the oxidation number is -1
o Fluorine always has the oxidation number of -1
o Hydrogen will have an oxidation number of +1 when it is with other
nonmetals and a -1 when its combined with a metal
o In a neutral molecule, all oxidation numbers will total to a net charge of zero
o In a polyatomic ion, all oxidation numbers will total to a net charge of the
polyatomic ion
Practice: State the oxidation numbers for the underlined atom
BeCl2
AgBr
NO
BaCrO4
NH4Cl
To balance a REDOX
o Not only must you balance the reaction to satisfy the law of conservation of
mass BUT you must also make sure the charges are balanced as well
o In simple REDOX reactions, the balancing of atoms with coefficients balances
the charge as a result
____ Fe + ____ CuSO4  ____ Fe2(SO4)3 + ____ Cu
2 Fe + 3 Cu+2  2Fe+3 + 3Cu
o In complex, REDOX reactions, the half reaction method is the best approach.
Complex REDOX reactions
o You must identify the oxidation states of each atom in the reactions
o Identify the atom that undergoes oxidation and the one that undergoes
reduction
o Create a half reaction for each process.
o Balance all atoms other than H and O
o In an acidic reaction, add water to balance the oxygen atoms
o Add hydrogen ions to balance the hydrogen
o Balance the number of electrons in each reaction
o Combine the reactions and simplify
o Check to make sure all atoms and charges are equal on both sides
o If basic solution, do all of the above steps except after you combine the
reactions, Add hydroxide ions to both sides to neutralize the hydrogen ions
Let’s use the same simple redox to practice the half reaction method
___ Fe + ____ CuSO4  ____ Cu + ____ Fe2(SO4)3
Oxidation Half Reaction
Fe  Fe+3
Fe  Fe+3 + 3e
2Fe  2Fe+3 + 6e
Reduction Half Reaction
Cu+2  Cu
Cu+2 + 2e  Cu
3 Cu+2 + 6e  3Cu
Add the reactions together and simplify!
Examples
Balance the following REDOX reaction in acidic solution
MnO4-1 + SO2  Mn+2 + SO4-2
Balance the following REDOX reaction in basic solution
Br2  Br -1 + BrO3-1
Voltaic Cells
o Also known as galvanic cell or a battery
o Change chemical energy into electrical energy if reactants are separated
o Electrodes are called the cathode and anode
o A cathode is where reduction occurs (RED CAT)
o An anode is where oxidation occurs (AN OX)
o Electrodes are placed in separate cells
o Electrons flow from the anode to the cathode through a wire
o Need a salt bridge or porous disk to connect the 2 cells (beakers). This
allows for exchange of ions so build up of charges in each half cell does not
happen
Calculating Cell Potential or Ecell
o The potential difference between the electrodes provides the driving force
that pushes the electrons through the wire
o Measured in volts(v) by a volt meter
1V = 1J/C of charge
o Cell potential must be positive for a spontaneous reaction
o Standard E cell is calculated at standard conditions (25 C, 1 atm of gas, or
1M solution)
o The standard reduction potentials are measured indirectly relative to a
standard hydrogen electrode. They are given to you on the reference sheet.
Notice they are REDUCTION values.
o The more positive the reduction potential, the more apt to being
reduced(better oxidizing agent.
o Standard reduction potentials are intensive properties. They do not depend
on the amount present. Coefficients are not important !
To calculate Cell Potential or Ecell
o Write reduction half reactions for each element
o Look at the chart to determine the voltage for each.
o Identify the most positive value. This is the element that is most easily
reduced and it will be considered the cathode
o Reverse the other atoms reaction and reverse the sign. This is the element
that will be oxidized and will be considered to anode.
o Add the 2 values together.
E cell = E red + E ox
Line Notation
o Anode components listed on left
o Cathode components listed on right
o Double line represents salt bridge or porous disk
Cu(s)| Cu+2 (aq) ||Zn+2 (aq)
| Zn(s)
Examples
1. Which is the better oxidizing agent? Zn+2 or Pb+2
2. Which is the better reducing agent? Ag(s) or Au(s)
3. A voltaic cell has 2 half cells. The first half cell has a copper electrode in a 1.0
M copper(II) nitrate solution. The second half cell contains the tin electrode
in a 1.0 M tin(II) nitrate solution. The salt bridge between the 2 half cells
completes the circuit. Assume tin is at the anode. Draw this out and
calculate the cell voltage.
4. In a voltaic cell, a zinc electrode is placed in a 1.0 M solution of zinc nitrate
while a copper electrode is placed in a 1.0 M solution of copper(II) nitrate. A
salt bridge is in place. Calculate the cell potential.
Spontaneity of a REDOX reaction
o You can decide if a REDOX reaction will occur spontaneously by finding the
cell potential
o If the cell potential (Ecell)a positive value, it indicates a spontaneous reaction
o If the cell potential is a negative value, it indicated a non spontaneous
reaction
Relationship between EMF (Ecell )and Free Energy( G)
o When delta G is negative, a reaction is considered spontaneous
o When delta G is positive, a reaction is considered nonspontaneous
o So when delta G is negative, E cell is positive
G = - n F E
n= moles of electrons
F= faraday’s constant = 96500 C/mol or J/molV
E = standard cell potential (v)
Example5 : Calculate the standard free energy change for the following reaction
Zn(s) + 2 Ag+1 (aq)  Zn+2(aq) + 2 Ag(s)
Nernst Equation
E = E – (RT) lnQ
nF
At 25 C
E = E – (.0592) log Q
n
n= number of electrons transferred during the reaction
F= Faraday’s constant
Q= reaction quotient
Q = [P] / [R]
E = cell potential at nonstandard conditions
o Used to calculate the EMF (Ecell )under nonstandard conditions
o All concentrations are 1M in standard potentials
o Most important idea to this equation: As the concentration of the products of
a REDOX reaction increase, the cell voltage decreases; As the concentration of
the reactants in a REDOX reaction increase, the voltage increases
o As cells discharge, concentration changes and so does Ecell
o Cells spontaneously discharge until they achieve equilibrium, the cell is
“dead”
Example 6: calculate the cell potential if the concentration of the zinc solution is
1.0M and the concentration of the copper solution is .10M
Zn(s) + Cu+2 (aq)  Zn+2(aq) + Cu(s)
Electrolysis
o Electrolytic cells require a voltage source to allow a nonspontaneous
reactions to occur
o Forcing a current through a cell to produce a chemical change for which the
cell potential is negative
Voltaic Cell
Source of electrons are from the
oxidation at the anode to the cathode
spontaneously
Cathode is positive; Anode is negative
Electrodes are placed in 2 cells
Electrolytic Cell
Source of electrons are from a dc power
supply which forces the electrons from
anode to cathode
Cathode is negative; Anode is positive
Electrodes are placed in 1 cell with
molten or a salt solution
o In both types, oxidation takes place at the anode and reduction takes place at
the cathode
o In order for electrolysis to start, a minimum voltage is required.
Diagram of an Electrolytic Cell
Electroplating
o
o
o
o
Thin coat of metal plates out on the surface of another.
The metal to be plated is at the cathode
Current: amount of charge per time. It is measured in Amperes(amps)
1 Amp equals 1C/Js
I = q/t
I = current(amps)
q = charge (coulombs)
t= time (seconds)
Stoichiometry of the electroplating process: How much metal can plate out?
1. Convert current and time to quantity of charge in coulombs.
(amps x time) = coulombs
2. Convert amount of charge in coulombs to moles of electrons
Coulombs x
1 mole e= mole e96,500 coulombs
3. Convert moles of electrons to moles of substance
Mole e- x mol substance
= mol substance
mole e4. Convert moles of substance to grams of substance
Moles of substance x molar mass = mass of substance
Examples:
7. 10.0 amps of current is passed through an electrolytic cell filled with
molten LiCl for 500. Seconds. What is the mass of lithium collected at
the anode?
8. 2.65 g Ag are collected at the cathode in an electrolytic cell filled with
aqueous silver nitrate. How long does it take to collect this much if .25
amps are applied?
Electrolysis of Water
o Requires the presence of a soluble salt or dilute acid to serve as an
electrolyte
Some Commercial Electrolytic Processes of Importance are
a.
b.
c.
d.
production of aluminum from the bauxite ore
electro refining of metals
metal plating like chrome
electrolysis of sodium chloride