Transcript Chapter 20 electrochemistry powerpoint

ELECTROCHEMISTRY INVOLVES TWO
MAIN TYPES OF PROCESSES
• A. Voltaic(galvanic) cells – which are
spontaneous chemical reactions
(battery)
• B. Electrolytic cells – which are nonspontaneous and require external e−
source (DC power source)
• C. BOTH of these fit into the category
entitled Electrochemical cells
Electrochemistry
Oxidation Numbers
In order to keep
track of what loses
electrons and what
gains them, we
assign oxidation
numbers.
Electrochemistry
Oxidation and Reduction
• A species is oxidized when it loses electrons.
 Here, zinc loses two electrons to go from neutral
zinc metal to the Zn2+ ion.
Electrochemistry
Oxidation and Reduction
• A species is reduced when it gains electrons.
 Here, each of the H+ gains an electron and they
combine to form H2.
Electrochemistry
Oxidation and Reduction
• What is reduced is the oxidizing agent.
 H+ oxidizes Zn by taking electrons from it.
• What is oxidized is the reducing agent.
 Zn reduces H+ by giving it electrons.
Electrochemistry
Assigning Oxidation Numbers
1. Elements in their elemental form have
an oxidation number of 0.
2. The oxidation number of a monatomic
ion is the same as its charge.
Electrochemistry
Voltaic Cells
In spontaneous
oxidation-reduction
(redox) reactions,
electrons are
transferred and
energy is released.
Electrochemistry
Voltaic Cells
• We can use that
energy to do work if
we make the
electrons flow
through an external
device.
• We call such a
setup a voltaic cell.
Electrochemistry
Voltaic Cells
• A typical cell looks
like this.
• The oxidation
occurs at the
anode.
• The reduction
occurs at the
cathode.
Electrochemistry
Voltaic Cells
The flow of
electrons is always
from the anode to
the cathode
through the wire
Electrochemistry
Voltaic Cells
Once even one
electron flows from
the anode to the
cathode, the
charges in each
beaker would not be
balanced and the
flow of electrons
would stop.
Electrochemistry
Voltaic Cells
• Therefore, we use a
salt bridge, usually a
U-shaped tube that
contains a salt
solution, to keep the
charges balanced.
 Cations move toward
the cathode.
 Anions move toward
the anode.
Electrochemistry
Voltaic Cells
• In the cell, then,
electrons leave the
anode and flow
through the wire to
the cathode.
• As the electrons
leave the anode, the
cations formed
dissolve into the
solution in the
anode compartment.
Electrochemistry
Voltaic Cells
• As the electrons
reach the cathode,
cations in the
cathode are
attracted to the now
negative cathode.
• The electrons are
taken by the cation,
and the neutral
metal is deposited
on the cathode.
Electrochemistry
Voltaic Cells Animation
Electrochemistry
Electromotive Force (emf)
• Water only
spontaneously flows
one way in a
waterfall.
• Likewise, electrons
only spontaneously
flow one way in a
redox reaction—from
higher to lower
potential energy.
Electrochemistry
Cell Potential
(Electromotive Force (emf))
• The potential difference between the
anode and cathode in a cell is called the
electromotive force (emf).
• It is also called the cell potential, and is
designated Ecell.
Electrochemistry
Cell Potential
(Electromotive Force (emf))
• Each potential is measured against a
standard , which is the standard
hydrogen electrode [consists of a piece
• of inert Platinum that is bathed by
hydrogen gas at 1 atm]. The hydrogen
electrode is assigned a value
• of ZERO volts.
Electrochemistry
Cell Potential
(Electromotive Force (emf))
• standard conditions—1 atm for gases, 1.0M for solutions
and 25°C for all (298 K)
• • naught (°)--we use the naught to symbolize standard
conditions [Experiencing a thermo flashback?]
• That means
• Ecell, Emf, or εcell become
• Ecello , Emfo , or εcello when measurements are
taken at standard conditions. You’ll soon learn how
these change when the conditions are nonstandard!
Electrochemistry
Reduction Potential
Chart
• elements that have the most positive reduction
potentials are easily reduced (in general, non-metals)
• elements that have the least positive reduction
potentials are easily oxidized (in general, metals)
• The table can also be used to tell the strength of
various oxidizing and reducing agents.
• Can also be used as an activity series. Metals having
less positive reduction potentials are more active
and will replace metals with more positive potentials.
Electrochemistry
Standard Reduction Potentials
Reduction
potentials for
many
electrodes
have been
measured and
tabulated.
Electrochemistry
Writing a Cell Diagram-Standard Cell notation
“ion sandwich in alphabetical order”
• The reduced species is written on the right hand
side—this electrode is the cathode
• The oxidized species is placed on the left hand side
and this is the anode.
• For example the cell diagram of a zinc and copper
galvanic cell:
• A vertical line (I ) represents the different phases
present in each half cell-phase boundary. (if in same
phase use comma instead)
• A double vertical line II represents the salt bridge
connecting the two half cells
•
Anode I solution II cathode solution I cathode
Electrochemistry
 Zn(s)
I Zn +2 II Cu+2 I Cu (s)
• Zinc and Lead(s) (plumbous)
•
• Copper and iron+3 and +2 (use copper
II)
Electrochemistry
What do you do when one of your electrodes lacks a SOLID
METAL??
You will need an inert conductor (usually platinum ) put in
parentheses (Pt)
Different species of the same phase are separated by a comma
• Fe+3 (aq)
+
Cu(s)  Cu+2 + Fe2+
Electrochemistry
Calculating Standard Cell Potential
(Ecello , Emfo , or εcello )
• Decide which element is oxidized or reduced using the standard
reduction potential chart. Element with more positive reduction
potential gets reduced!
• Write both equations as is from the chart with their voltage
• Reverse the equation that will be oxidized and change the sign
of the voltage
• Balance the two half reactions BUT DO NOT MULTIPLY
VOLTAGE VALUES
• Add the two half reactions and their voltages together
• Ecello = Eooxidation + Eoreduction where o means standard
conditions 1atm , 1M and 25oC
Electrochemistry
Calculate Cell Potential E cello
•
Zinc and lead (s) Plumbous
• Copper and iron+3 and +2 (use copper
II)
1. H2 (g) + I2 (g)  2H+ (aq) + 2 I – (aq)
Electrochemistry
Calculate Cell Potential E cello
Draw a diagram of the galvanic cell for the reaction and
label completely
• Fe+3
(aq)
+
Cu(s)  Cu+2 + Fe2+
Electrochemistry
Cell Potential , Electrical Work,
And Free energy
Cell potential is measured in volts (V).
It is the work that can be accomplished
when electrons are transferred through
a wire.
Work(J)
1V=1
Charge ( C)
Electrochemistry
Cell Potential , Electrical Work,
And Free energy
1 joule of work is required or
produced when one coloumb of
charge is transferred between two
points
Work(J)
1V=1
Charge ( C)
Electrochemistry
Cell Potential , Electrical Work, And
Free energy
• Using the table of standard reduction
potentials, calculate DG o for the
reaction:
• Cu+2(aq) + Fe(S)  Cu (s) + Fe+2 (aq)
•
• IS it spontaneous?
Electrochemistry
Cell Potential , Electrical Work, And
Free energy
•
•
•
•
•
•
DG o = - nFEo
G= Gibb’s Free energy
n= numbers of moles of electrons
F= Faraday constant 96500 J/V . Mol
The redox reaction will be
spontaneous when DG o = _________
The redox reaction will be
spontaneous when Eo = __________
Electrochemistry
Cell Potential , Electrical Work, And
Free energy
DG
o
= - nFEo
+
DG o = - RT ln K
=
• You can derive the equation: for Ecell at
standard conditions at equilibrium
Eo = RT ln K
nF
Electrochemistry
o
E
K
Greater
than 1
=1
Less than
one
= RT ln K
nF
Eo
Positive
DG o
negative
0
Negative
0
positive
Conclusion
Electrochemistry
•Nernst Equation
E = E −
RT
nF
ln Q
•Why?
It is used to calculate the voltage generated by the combination of two
half-cells when the conditions are not standard!!!!!!!
• n= numbers of electrons transferred between species
•
F= Faraday constant 96500 J/V . mol
•
R=ideal gas constant 8.314 /K mol
•
T= Kelvin Temperature
•
Eo= the voltage generated if the conditions
» were standard
•
Q= reactions quotient [products]x
•
[reactants]y
Electrochemistry
If the reaction below is carried out using solutions that are 5.0 M Zn +2
and 0.3 M Cu+2 at 298 K, what is the actual cell voltage?
Cu(s) + Zn+2
Zn (s) + Cu+2 
• Step 1: Work out Eo cell assuming standard conditions:
•
•
•
Step 2: Calculate Q (don’t forget _____ and _______ are
omitted from Q)
•
•
•
•
Step 3: Find n
•
•
Step 4: Plug everything in!
Electrochemistry
(Answer 1.06 V)
Assigning Oxidation Numbers
3. Nonmetals tend to have negative
oxidation numbers, although some are
positive in certain compounds or ions.
 Oxygen has an oxidation number of −2,
except in the peroxide ion in which it has
an oxidation number of −1.
 Hydrogen is −1 when bonded to a metal,
+1 when bonded to a nonmetal.
Electrochemistry
Assigning Oxidation Numbers
3. Nonmetals tend to have negative
oxidation numbers, although some are
positive in certain compounds or ions.
 Fluorine always has an oxidation number
of −1.
 The other halogens have an oxidation
number of −1 when they are negative;
they can have positive oxidation
numbers, however, most notably in
oxyanions.
Electrochemistry
Assigning Oxidation Numbers
4. The sum of the oxidation numbers in a
neutral compound is 0.
5. The sum of the oxidation numbers in a
polyatomic ion is the charge on the
ion.
Electrochemistry
Balancing Oxidation-Reduction
Equations
Perhaps the easiest way to balance the
equation of an oxidation-reduction
reaction is via the half-reaction method.
Electrochemistry
Balancing Oxidation-Reduction
Equations
This involves treating (on paper only) the
oxidation and reduction as two separate
processes, balancing these half reactions,
and then combining them to attain the
balanced equation for the overall reaction.
Electrochemistry
Half-Reaction Method
1. Assign oxidation numbers to
determine what is oxidized and what is
reduced.
2. Write the oxidation and reduction halfreactions.
Electrochemistry
Half-Reaction Method
3. Balance each half-reaction.
a.
b.
c.
d.
Balance elements other than H and O.
Balance O by adding H2O.
Balance H by adding H+.
Balance charge by adding electrons.
4. Multiply the half-reactions by integers
so that the electrons gained and lost
are the same.
Electrochemistry
Half-Reaction Method
5. Add the half-reactions, subtracting
things that appear on both sides.
6. Make sure the equation is balanced
according to mass.
7. Make sure the equation is balanced
according to charge.
Electrochemistry
Half-Reaction Method
Consider the reaction between MnO4− and C2O42− :
MnO4−(aq) + C2O42−(aq)  Mn2+(aq) + CO2(aq)
Electrochemistry
Half-Reaction Method
First, we assign oxidation numbers.
+7
+3
+2
+4
MnO4− + C2O42-  Mn2+ + CO2
Since the manganese goes from +7 to +2, it is reduced.
Since the carbon goes from +3 to +4, it is oxidized.
Electrochemistry
Oxidation Half-Reaction
C2O42−  CO2
To balance the carbon, we add a
coefficient of 2:
C2O42−  2 CO2
Electrochemistry
Oxidation Half-Reaction
C2O42−  2 CO2
The oxygen is now balanced as well.
To balance the charge, we must add 2
electrons to the right side.
C2O42−  2 CO2 + 2 e−
Electrochemistry
Reduction Half-Reaction
MnO4−  Mn2+
The manganese is balanced; to balance
the oxygen, we must add 4 waters to
the right side.
MnO4−  Mn2+ + 4 H2O
Electrochemistry
Reduction Half-Reaction
MnO4−  Mn2+ + 4 H2O
To balance the hydrogen, we add 8 H+
to the left side.
8 H+ + MnO4−  Mn2+ + 4 H2O
Electrochemistry
Reduction Half-Reaction
8 H+ + MnO4−  Mn2+ + 4 H2O
To balance the charge, we add 5 e− to
the left side.
5 e− + 8 H+ + MnO4−  Mn2+ + 4 H2O
Electrochemistry
Combining the Half-Reactions
Now we evaluate the two half-reactions
together:
C2O42−  2 CO2 + 2 e−
5 e− + 8 H+ + MnO4−  Mn2+ + 4 H2O
To attain the same number of electrons
on each side, we will multiply the first
reaction by 5 and the second by 2. Electrochemistry
Combining the Half-Reactions
5 C2O42−  10 CO2 + 10 e−
10 e− + 16 H+ + 2 MnO4−  2 Mn2+ + 8 H2O
When we add these together, we get:
10 e− + 16 H+ + 2 MnO4− + 5 C2O42− 
2 Mn2+ + 8 H2O + 10 CO2 +10 e−
Electrochemistry
Combining the Half-Reactions
10 e− + 16 H+ + 2 MnO4− + 5 C2O42− 
2 Mn2+ + 8 H2O + 10 CO2 +10 e−
The only thing that appears on both sides are the
electrons. Subtracting them, we are left with:
16 H+ + 2 MnO4− + 5 C2O42− 
2 Mn2+ + 8 H2O + 10 CO2
Electrochemistry
Balancing in Basic Solution
• If a reaction occurs in basic solution, one
can balance it as if it occurred in acid.
• Once the equation is balanced, add OH−
to each side to “neutralize” the H+ in the
equation and create water in its place.
• If this produces water on both sides, you
might have to subtract water from each
side.
Electrochemistry
Standard Hydrogen Electrode
• Their values are referenced to a standard
hydrogen electrode (SHE).
• By definition, the reduction potential for
hydrogen is 0 V:
2 H+ (aq, 1M) + 2 e−  H2 (g, 1 atm)
Electrochemistry
Standard Cell Potentials
The cell potential at standard conditions
can be found through this equation:
Ecell
 (cathode) − Ered
 (anode)
 = Ered
Because cell potential is based on
the potential energy per unit of
charge, it is an intensive property.
Electrochemistry
Cell Potentials
• For the oxidation in this cell,
Ered
 = −0.76 V
• For the reduction,
Ered
 = +0.34 V
Electrochemistry
Cell Potentials
Ecell
 = Ered
 (cathode) − Ered
 (anode)
= +0.34 V − (−0.76 V)
= +1.10 V
Electrochemistry
Oxidizing and Reducing Agents
• The strongest
oxidizers have the
most positive
reduction potentials.
• The strongest
reducers have the
most negative
reduction potentials.
Electrochemistry
Oxidizing and Reducing Agents
The greater the
difference between
the two, the greater
the voltage of the
cell.
Electrochemistry
Nernst Equation
• Remember that
DG = DG + RT ln Q
• This means
−nFE = −nFE + RT ln Q
Electrochemistry
Nernst Equation
Dividing both sides by −nF, we get the
Nernst equation:
RT
ln Q
E = E −
nF
or, using base-10 logarithms,
2.303 RT
ln Q
E = E −
nF
Electrochemistry
Concentration Cells
• Notice that the Nernst equation implies that a cell
could be created that has the same substance at
both electrodes.
 would be 0, but Q would not.
• For such a cell, Ecell
• Therefore, as long as the concentrations
are different, E will not be 0.
Electrochemistry
Applications of
Oxidation-Reduction
Reactions
Electrochemistry
Batteries
Electrochemistry
Alkaline Batteries
Electrochemistry
Hydrogen Fuel Cells
Electrochemistry
Corrosion and…
Electrochemistry
…Corrosion Prevention
Electrochemistry