Transcript ELECTROCHEMISTRY AND ITS APPLICATIONS GENERAL …

ELECTROCHEMISTRY
Chapter 18
Introduction
* Electrochemistry: the study of the
relationship between electron flow and
redox reactions.
Significance and Applications?
What is a redox reaction? An oxidation –
reduction reaction in which electrons are
transferred (same # e- lost and gained).
Zn(s) + Cu2+(aq)
→ Zn2+(aq) + Cu(s)
I. Oxidation and Reduction Reactions
A. Introduction
1. Reactions involving the transfer of
electrons.
2. One species has a desire to lose electrons
an another species has a desire to gain
electrons. With the correct conditions,
energy in the form of current flow is
produced.
3. Driving force is tendency for electron
flow.
4. Examples: batteries, combustion, rusting of
iron, metabolism of food in the body.
B. Definitions
1. Oxidation - The loss of electron(s).
Oxidation Reaction
M  M n+ + n e While M is oxidized, it acts as an
reducing agent because it causes
something else to be
reduced (i.e.,it gives electrons to
something else).
Fe  Fe2+ + 2 e-
2. Reduction - The gain of electron(s).
Reduction Reaction
X + n e -  X nWhile X is reduced, it acts as an oxidizing
agent because it causes something else to
be oxidized (i.e., it takes electrons from
something else).
Cl2 + 2 e-  2 ClAg+ + e-  Ag
L
E
O
loss of
electrons (products)
oxidation
the lion says
G
gain of
E
electrons (reactants)
R!!! reduction
C. Properties of Oxidation-Reduction Rxns
1. “Redox reaction” always consists of the two
processes (oxidation and reduction) which
occurs simultaneously. One species must
lose electrons so that another species can
gain those same # of electrons.
2. A redox reaction may be recognized if a
metal is being oxidized or reduced.
Metal gains or loses electrons as it goes
from reactants to products.
How can we recognize this in a reaction?
Example Problems
For the redox reaction given, identify
which reactant is being oxidized, which
reactant is being reduced, which is the
oxidizing agent, and which is the
reducing agent?
Cu2+(aq) + Mg(s)  Mg2+(aq) + Cu(s)
MnO4-(aq) + Fe2+(aq)  Mn2+(aq) + Fe3+(aq)
D. Related Terminology For Redox Rxns
* Recall complete redox reaction always
includes both oxidation and reduction.
* Half-reaction: One of two parts of an
oxidation - reduction reaction, one part of
which involves a loss of electrons and the
other a gain of electrons.
For the reaction: Fe + Cu2+  Fe2+ + Cu
1. Write the reduction half reaction.
2. Write the oxidation half reaction.
For the reaction:
Zn(s) + Cu2+(aq)
→ Zn2+(aq) + Cu(s)
d. What is the reducing agent?
Reducing agent: The reactant that is oxidized.
It loses electrons so that some other reactant
can be reduced.
e. What is the oxidizing agent?
Oxidizing Agent: The reactant that is reduced.
It gains electrons so that some other reactant can
be oxidized.
E. Balancing Redox Reactions
1. Introduction
Oxidation Number Method
Arbitrary book-keeping system for
electrons which can be used to identify
what is being oxidized or reduced.
Half-Reaction Method
*** You will be responsible for a simplified
version of this.
2. Simplified Rules for Balancing Redox
A. Split the reaction into half-reactions. (by inspection)
B. Balance number of atoms of each element on both
sides of each equation by changing coefficients.
C. Balance charges on both sides of each equation
by adding the appropriate electrons. (Electrons
added to more positive side.)
D. Same number of electrons must be transferred in
redox reaction. Multiply one or both half
reactions by a whole number to get same number
of electrons transferred.
E. Add half-reactions together for balanced redox
reaction. Recheck to see that electrons cancel out
and atoms of each element balance.
Example Problems
1. Balance the following redox reactions:
Sn2+(aq) + Fe3+(aq) → Sn4+(aq) + Fe2+(aq)
Sb3+(aq) + Sr(s) → Sb(s) + Sr2+(aq)
2. How many mol es of electrons are being
transferred in the following balanced
redox reaction?
2 Al(s) + 3 Cl2(g) → 2 AlCl3(aq)
II. Voltaic (Galvanic) Cells
A. Introduction
1. An electrochemical cell in which a
product-favored (spontaneous) redox
reaction generates an electric current.
2. The reaction produces an electron flow
through an outside conductor (wire).
3. Examples: batteries
B. Example of a Voltaic Cell
1. Reaction
Zn(s)  Zn2+(aq) + 2 e__Cu2+(aq) + 2 e-  Cu (s)________
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
If you just ran the redox reaction in a test tube,
the reaction would occur, but no current (energy)
would be captured. See slide #15
However in a galvanic (voltaic) cell, energy is
captured in the form of current flow. Slide # 16.
The 2 half-reactions occur in separate half-cells.
Redox Reaction With No Current Flow
Redox Reaction With Current Flow
(Voltaic Cell)
C. Requirements of Galvanic Cell
1. Anode: an electrode (conductor such as metal strip
or graphite) where oxidation occurs.
2. Cathode: an electrode (conductor such as metal
strip or graphite) where reduction occurs.
3. Salt Bridge: A tube of an electrolyte (sometimes in
a gel) that is connected to the two half-cells of a
voltaic cell: the salt bridge allows the flow of
ions but prevents the mixing of the different
solutions that would allow direct reaction of
the cell reactants. Charge does not build up in half
cells. Electrical neutrality must be maintained.
Cathode
A
Reduction
Anode
U
T
Oxidation
Cell Diagrams
•
•
•
•
A cell diagram is “shorthand” for an electrochemical cell.
The anode is placed on the left side of the diagram.
The cathode is placed on the right side.
A single vertical line ( | ) represents a boundary between
phases, such as between an electrode and a solution.
• A double vertical line ( || ) represents a salt bridge or
porous barrier separating two half-cells.
D. Galvanic Cells and Electrical Potential
1. Electron flow in galvanic cell can do work /
produce energy.
2. Electrical potential energy measured in volts.
1 volt = (1 joule) / (1 coulomb)
work
unit of charge due to
6.24 x 1018 electrons
3. Coulombs = amperes x seconds
C =Axs
or
A=C/s
E. Standard Cell Voltages
1. Cell voltages can be measured under standard
conditions (1 atm pressure, 250 C, and 1.0 M
concentrations). E0cell.
2. The standard cell potential is determined by the
equation
E0cell = E0red + E0ox
3. If E0cell is positive, the net cell reaction is productfavored (spontaneous).
4. If E0cell is negative, the net cell reaction is reactantfavored (nonspontaneous).
F. Standard Electrode Potentials
1. Standard Electrode Potentials are measured for
half-reactions, relative to a standard hydrogen
electrode potential (which is assigned 0 volts).
2. See Table 18.1 for Standard Electrode Potentials
a. Each half reaction is written as a reduction.
b. Each half reaction could occur in either
direction.
c. The more “+” the standard electrode potential,
the greater the tendency to undergo
reduction, meaning a good oxidizing agent.
d. The more “-” the standard electrode potential,
the greater the tendency to undergo
oxidation, meaning a good reducing agent.
e. If a half-reaction is written in the reverse
direction, the sign of the corresponding
standard electrode potential must change.
(Same magnitude, opposite in sign)
f. If a half-reaction is multiplied by a factor
(coefficient), the standard electrode
potential is not multiplied by that factor.
Selected Electrode (Reduction) Potentials
F2(g) + 2e-  2F- +2.87V
F2/FAu3+ + 3e-  Au(s) +1.50V
Au3+/Au
Hg2+ + 2e-  Hg(l) +0.855V
Hg2+/Hg
2H+ + 2e-  H2(g) 0.0000V
H+/H2
Fe2+ + 2e-  Fe(s) -0.44V
Fe2+/Fe
Li+ + e-  Li(s)
-3.045V
Li+/Li
G. Problem Solving and Voltaic Cells
1. A voltaic (galvanic) cell is constructed using a piece
of Sn in 1.00 M Sn2+ for one half-cell and a piece of
Au in 1.00 M Au3+ for the other half-cell. The
reaction is run under standard conditions of
pressure and temperature.
a.
b.
c.
d.
e.
Determine the net cell potential (voltage).
Determine the net cell reaction.
What is the oxidizing agent for the net reaction?
What is the reaction occurring at the anode?
Sketch the cell information in the following
diagram.
H. Cell Potential and Gibbs Free Energy
A positive E0cell indicates a spontaneous or
product-favored reaction. There must be a
relationship between E0cell and free energy (ΔG0).
ΔG0 = -n F E0cell
n = # moles electrons transferred
F = Faraday constant = 9.65 x 104 C / mol e(also recall that volt = Joules / Coulomb = J / C)
** A positive E0cell would produce a negative ΔG0
1. Calculate the standard free energy change (ΔG0)
for the reaction below under standard conditions.
NiO2 + 2Cl- + 4 H+  Cl2 + Ni2+ + 2 H2O
E0cell = 0.320 v
I. Equilibrium Constant and Cell Potential
We know
and:
Therefore:
ΔG0 = -RTlnKeq
ΔG0 = -n F E0cell
-n F E0cell = -RTlnKeq
E0cell = RTlnKeq / nF
Rearranging:
ln Keq = n F E0cell / RT
1. Calculate Kc for the following reaction under
standard conditions:
Cu2+(aq) + 2 Ag(s)  Cu(s) + 2 Ag+(aq)
Is the reaction product-favored?
Do you expect a large amount of product formed
at equilibrium?
Thermodynamics, Equilibrium, and
Electrochemistry: A Summary
From any one of the three
quantities Keq, ΔG°, E°cell, we
can determine the others.
J. Concentration Effects on Cell Potential
1. When all the concentrations in a voltaic
cell are 1.0 M, under standard conditions,
the cell potential equals the standard cell
potential.
2. As concentrations of reactants or products
change, the cell potential (voltage changes).
3. When equilibrium is reached, the cell
potential will drop to zero.
Mathematical Relationship
Nernst Equation
Ecell = E°cell
RT
– –––– ln Q
nF
R = gas constant = 8.314 J / mol K
0.0592V
E E

log Q
T = Kelvin temperature
n
n = # moles of electrons transferred
Q = reaction quotient. What is it??
Ecell  E 0
0 cell 
cell
cell
0.0592V
log Q
n
or (easiest to use equation)
Ecell  E
0
cell
0.0592V

log Q
n
Problem:
A galvanic cell contains Ni2+ (aq) in contact
with Ni (s) and Cr3+ (aq) in contact with Cr (s).
When [Ni2+] = 1.0 x 10-4 and [Cr3+] = 2.0 x 10-3,
determine the value of Ecell . Given:
Ni2+(aq) + 2 e-  Ni(s)
E0 = -.25v
Cr3+(aq) + 3 e-  Cr(s)
E0 = -.76v
K. Applications of Voltaic Cells
1. The Lead Storage Battery
(Secondary Battery)
Anode Reaction
Pb(s) + HSO41-(aq)  PbSO4(s) + H+ + 2e-
Cathode Reaction
PbO2(s) + 3H+ + HSO41-(aq) + 2e-  PbSO4(s) + 2H2O (l)
Lead Storage Battery
2. Dry Cell Batteries
3. Fuel Cells
What are they?
Where are they used?
How do they differ from batteries?
Corrosion of an Iron Piling
One way to minimize
rusting is to provide a
different anode reaction.
IV. Electrolytic Cells
A. Electrolytic vs. Voltaic (Galvanic) Cells
Voltaic (Galvanic) Cells
Redox reaction which proceeds spontaneously in a
product favored direction, generating electricity.
Electrolytic Cell
Redox reaction in which an electrical current is
supplied to drive a nonspontaneous, reactionfavored reaction.
B. Properties of Electrolytic Cell
1. Energy requiring (in form of electric current).
2. No physical separation needed for the two
electrode reactions.
3. Usually no salt bridge required.
4. Conducting medium is molten salt or aqueous
solution.
5. For electrolytic redox reaction:
E0cell is negative.
ΔG0 is positive.
Kc is small (<1).
C. Electrolysis of Molten Sodium Chloride
1. Redox Reaction:
2 Na+ + 2 e-  2 Na(l)
cathode (reduction)
2 Cl Cl2(g) + 2 e- anode (oxidation)
2 Na+ + 2 Cl-  2 Na(l) + Cl2(g) net cell rxn
2. Electrolytic Cell:
See next slide
E. Stoichiometry of Electrolysis
1. Electrons Treated as Stoichiometric Factor
a. Given the half-reactions:
Na+ + e- Na(s) and Cu2+ + 2 e-  Cu(s)
How many mol e- required to form 1 mol Na(s)?
How many mol e- required to form 4 mol Cu(s)?
***
***
2 mol e- = 1 mol Cu(s)
b. Chemical change is directly proportional to
electron flow (charge).
1 mol e- = 9.65 x 104 C
2. Coulomb Relationship
***
Charge = current x time C = Amps x seconds
Steps for Electrolysis Calculations
3. Problem Solving
a. How many moles of electrons are required in an
electrolytic cell to deposit 2.00 grams of chromium,
Cr(s), from a solution of CrCl3?
b. What mass of aluminum metal can be produced per
hour in the electrolysis of a molten aluminum salt
by a current of 26 A?
F. Applications of Electrolysis
Electroplating
Refining of Metals
Copper
Aluminum
Electroplating
• Electrolysis can be used to coat
one metal onto another, a
process called electroplating.
• Usually, the object to be
electroplated, such as a spoon, is
cast of an inexpensive metal. It is
then coated with a thin layer of a
more attractive, corrosionresistant, and expensive metal,
such as silver or gold.
Refining of Aluminum