Transcript Chap 21 Electrochemistry

Electrochemistry
Chapter 21
Electrochemistry and Redox
Oxidation-reduction:
“Redox”
Electrochemistry:
study of the interchange between chemical
change and electrical work
Electrochemical cells:
systems utilizing a redox reaction to produce
or use electrical energy
Redox Review
Redox reactions:
electron transfer processes
Oxidation:
Reduction:
loss of 1 or more egain of 1 or more e-
Oxidation numbers: imaginary charges
(Balancing redox reactions)
Oxidation Numbers (O.N.)
1.Pure element
O.N. is zero
2.Monatomic ion
O.N. is charge
3.Neutral compound: sum of O.N. is zero
Polyatomic ion:
sum of O.N. is ion’s charge
*Negative O.N. generally assigned to more
electronegative element
Oxidation Numbers (O.N.)
4.Hydrogen
assigned +1
(metal hydrides, -1)
5.Oxygen
assigned -2
(peroxides, -1; OF2, +2)
6.Fluorine
always -1
Oxidation-reduction
Oxidation is loss of eO.N. increases (more positive)
Reduction is gain of eO.N. decreases (more negative)
Oxidation involves loss
Reduction involves gain
OIL
RIG
Redox
Oxidation is loss of ecauses reduction
“reducing agent”
Reduction is gain of ecauses oxidation
“oxidizing agent”
Balancing Redox Reactions
1.
Write separate equations (half-reactions) for
oxidation and reduction
2.
For each half-reaction
a. Balance elements involved in e- transfer
b. Balance number e- lost and gained
3.To balance emultiply each half-reaction by whole numbers
Balancing Redox Reactions: Acidic
4.Add half-reactions/cancel like terms (e-)
5. Acidic conditions:
Balance oxygen using H2O
Balance hydrogen using H+
Basic conditions:
Balance oxygen using OHBalance hydrogen using H2O
6.Check that all atoms and charges balance
Examples
Acidic conditions:
MnO 4(aq)

2
Fe(aq)
2

 Mn (aq)
acid

3
Fe(aq)
Basic conditions:

Ag (s)  CN (aq)

 O 2(g) 
 Ag(CN) 2(aq)
base
Types of cells
Voltaic (galvanic) cells:
a spontaneous reaction generates electrical energy
Electrolytic cells:
absorb free energy from an electrical source to
drive a nonspontaneous reaction
Common Components
Electrodes:
conduct electricity between cell and
surroundings
Electrolyte:
mixture of ions involved in reaction or
carrying charge
Salt bridge:
completes circuit (provides charge balance)
Electrodes
Anode:
Oxidation occurs at the anode
Cathode:
Reduction occurs at the cathode
Active electrodes: participate in redox
Inactive:
sites of ox. and red.
Voltaic (Galvanic) Cells
A device in which chemical energy
is changed to electrical energy.
Uses a spontaneous reaction.
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e–
e–
e–
e–
Porous disk
Oxidation
e–
(a)
Reducing
agent
Anode
Reduction
Oxidizing
agent
(b)
e–
Cathode
Zn2+(aq) + Cu(s)  Cu2+(aq) + Zn(s)
Zn
gives up electrons to Cu
— “pushes harder” on e— greater potential energy
— greater “electrical potential”
Spontaneous reaction due to
— relative difference in metals’ abilities to give e— ability of e- to flow
Cell Potential
Cell Potential / Electromotive Force (EMF):
The “pull” or driving force on electrons
Measured voltage (potential difference)
E cell
work or electrical potential energy
J

 V
unit of charge moved
C
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e–
e–
Zn(s)
e–
Ecell = +1.10 V
+
Zn 2
–
SO 4 2
1.0 M Zn 2
solution
Anode
e–
+
Cu 2
–
SO 4 2
+
+
1.0 M Cu 2
solution
Cathode
Cu(s)
Cell Potential, E0cell
E0cell
cell potential under standard conditions
elements in standard states (298 K)
solutions:
1M
gases:
1 atm
Standard Reduction Potentials
E0 values for reduction half-reactions with
solutes at 1M and gases at 1 atm
Cu2+ + 2e  Cu
E0 = 0.34 V vs. SHE
SO42 + 4H+ + 2e  H2SO3 + H2O
E0 = 0.20 V vs. SHE
E0cell and DG0
E0cell > 0
DG0 < 0
Spontaneous
E0cell < 0
DG0 > 0
Not
E0cell = 0
DG0 = 0
Equilibrium
Calculating E0cell
E0cell = E0cathode - E0anode
Br2(aq)+2V3+ +2H2O(l)  2VO2+(aq)+ 4H+(aq)+ 2Br-(aq)
Given: E0cell = +1.39 V
E0Br2 = +1.07 V
What is E0V3+ and is the reaction spontaneous?
E0 values
More positive:
Stronger oxidizing agent
More readily accepts e-
More negative:
Stronger reducing agent
More readily gives eStronger R.A. + O.A.  Weaker R.A. + O.A.
Free Energy and Cell Potential
wmax  DG  nFE
0
n:
F:
number of moles of eFaraday’s constant
96485 C
mol of e-
0
DG0, E0, and K
DG  RTlnK  nFE
RT
0
so E 
lnK
nF
0
At equilibrium:
At 298 K:
DG0 = 0 and K = Q
0.0592
E 
logK
n
0
0
Nernst Equation
Under nonstandard conditions
DG  DG  RTlnQ
0
 nFE   nFE  RTlnQ
0
E cell
RT
E 
lnQ
nF
298K
E cell
0.0592
E 
lnQ
n
0
0
Concentration Cells
. . . a cell in which both compartments
have the same components but at
different concentrations
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e–
e–
e–
Porous
disk
e–
Ag
Ag
0.1 M Ag +
–
0.1 M NO 3
Anode
1 M Ag +
–
1 M NO 3
Cathode
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Reference solution of
dilute hydrochloric acid
Silver wire coated with
silver chloride
Thin-walled membrane
Batteries
A battery is a galvanic cell or, more
commonly, a group of galvanic
cells connected in series.
Fuel Cells
Galvanic cells
Reactants are continuously supplied.
2H2(g) + O2(g)  2H2O(l)
anode:
cathode:
2H2 + 4OH  4H2O + 4e
4e + O2 + 2H2O  4OH
Corrosion
Some metals, such as copper, gold, silver and
platinum, are relatively difficult to oxidize.
These are often called noble metals.
Electrolysis
Forcing a current through a cell to
produce a chemical change for
which the cell potential is negative.
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Electrodes of
graphite rods
To external
power source
Carbon
dioxide
formed at
the anodes
Carbon-lined iron tank
Molten Al 2O3/Na 3AlF 6
mixture
Molten
aluminum
Plug
Stoichiometry
How much chemical change occurs with the flow
of a given current for a specified time?
current and time  quantity of charge 
moles of electrons  moles of analyte 
grams of analyte