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Title: electrochemistry
Description: well ..defined...notes..on electochemistry
Description: well ..defined...notes..on electochemistry
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Electrochemistry
a Chem1 Supplement Text
Stephen K
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Potential differences at interfaces
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Cell description conventions
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5
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8
8
3 Standard half-cell potentials
Reference electrodes
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Cell potentials and the electromotive
Cell potentials and free energy
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Latimer diagrams
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4 The Nernst equation
Concentration cells
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Determination of solubility products
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Measurement of pH
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5 Batteries and fuel cells
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The fuel cell
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34
7 Electrolytic cells
Electrolysis involving water
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Industrial electrolytic processes
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Electrolytic refining of aluminum
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38
Chemistry and electricity
The connection between chemistry and electricity is a very old one, going back
to Allesandro Volta’s discovery, in 1793, that electricity could be produced by
placing two dissimilar metals on opposite sides of a moistened paper
...
This was
surely one of the most significant experiments in the history of chemistry, for it
implied that the atoms of hydrogen and oxygen were associated with positive
and negative electric charges, which must be the source of the bonding forces
between them
...
In electrolysis, the applied voltage was thought to overpower the
attraction between these opposite charges, pulling the electrified atoms apart in
the form of ions (named by Berzelius from the Greek for “travellers”)
...
N
...
1
CHEMISTRY AND ELECTRICITY
-
Zn
3
dissolution of Zn as Zn2+ causes
electric charges to build up in the
two phases which inhibits further
dissolution
e-
eZn2+(aq)
Zn2+(aq)
Zn in metal
Figure 1: Oxidation of metallic zinc in contact with water
Meanwhile, the use of electricity as a means of bringing about chemical
change continued to play a central role in the development of chemistry
...
It was left to Davey’s former assistant, Michael Faraday, to show that
there is a quantitative relation between the amount of electric charge and the
quantity of electrolysis product
...
Electroneutrality
Nature seems to very strongly discourage any process that would lead to an
excess of positive or negative charge in matter
...
A small number of zinc atoms go
into solution as Zn2+ ions, leaving their electrons behind in the metal:
Zn(s) −→ Zn2+ + 2 e−
(1)
As this process goes on, the electrons which remain in the zinc cause a negative
charge to build up which makes it increasingly difficult for additional positive
ions to leave the metallic phase
...
Very soon, therefore, the process comes
to a halt, resulting in a solution in which the concentration of Zn2+ is so low
(around 10−10 M ) that the water can still be said to be almost “pure”
...
One way to arrange this is to drain off the excess
electrons through an external circuit that forms part of a complete electrochemical
cell; this we will describe later
...
A
1
CHEMISTRY AND ELECTRICITY
4
suitable electron acceptor would be hydrogen ions; this is why acids attack many
metals
...
The degree of charge unbalance that is allowed produces differences in electric potential of no more than a few volts, and corresponds to concentration unbalances of oppositely charged particles that are not even detectable by ordinary
chemical means
...
The additional work raises the free energy ∆G of the process, making
it less spontaneous
...
A simple way to accomplish this would be immerse the zinc in a solution of
copper sulfate instead of pure water
...
The reaction is
a simple oxidation-reduction process, a transfer of two electrons from the zinc
to the copper:
Zn(s) −→ Zn2+ + 2 e−
Cu2+ + 2 e− −→ Cu(s)
The dissolution of the zinc is no longer inhibited by a buildup of negative charge
in the metal, because the excess electrons are removed from the zinc by copper
ions that come into contact with it
...
The net reaction
Zn(s) + Cu2+ −→ Zn2+ + Cu(s)
quickly goes to completion
...
A process of this
kind is known generally as an electrode process
...
The result is an interfacial potential difference which, as we saw above,
can materially affect the rate and direction of the reaction
...
In particular,
2
ELECTROCHEMICAL CELLS
Ð
Zn
direction of electron flow
in external circuit
+
porous fritted
glass barrier
5
Cu
direction of conventional current flow
Cu2+
Zn 2+
NO3Ð
Zn ê Zn2+ ,NO3Ð êê Cu2+, NO 3Ð ê Cu
Figure 2: A simple electrochemical cell
manipulation of the interfacial potential difference affords an important way of
exerting external control on an electrode reaction
...
This may not seem like very much,
but it is important to understand that what is important is the distance over
which this potential difference exists
...
a only a few atomic diameters
...
In many forms of matter, they are the result
of adsorption or ordered alignment of molecules caused by non-uniform forces in the
interfacial region
...
The resulting net
electric charge prevents the particles from coming together and coalescing, which they
would otherwise tend to do under the influence of ordinary van der Waals attractions
...
However, if we have two such metal-solution interfaces, we can easily
measure a potential diference between them
...
A typical cell might consists of two pieces of metal, one zinc and
the other copper, each immersed each in a solution containing a dissolved salt
of the corresponding metal (see Fig
...
The two solutions are connected by a
tube containing a porous barrier that prevents them from rapidly mixing but
allows ions to diffuse through
...
However, if we connect
the zinc and copper by means of a metallic conductor, the excess electrons that
remain when Zn2+ ions go into solution in the left cell would be able to flow
through the external circuit and into the right electrode, where they could be
delivered to the Cu2+ ions that are converted into Cu atoms at the surface of
the copper electrode
...
If the external circuit is broken, the reaction stops
...
By connecting a battery or other source of
current to the two electrodes, we can even force the reaction to proceed in its
non-spontaneous, or reverse direction
...
Electric charge q is measured in coulombs
...
Careful
experiments have determined that
1 F = 96467 c
For most purposes, you can simply use 96,500 c as the value of the faraday
...
A current of one ampere corresponds to the flow of one coulumb per second
...
2, how much mass would the zinc electrode lose if a
current of 0
...
5 hours?
2
ELECTROCHEMICAL CELLS
7
Solution
...
15 amp) × (5400 sec) = 810 c
or
(810 c)/(96500 c F −1 ) =
...
0042, corresponding to a weight loss of
(
...
37 g M −1 ) =
...
2 to operate, not only must there be an external
electrical circuit between the two electrodes, but the two electrolytes (the solutions) must be in contact
...
Positive charge
(in the form of Zn2+ is added to the electrolyte in the left compartment, and
removed (as Cu2+ ) from the right side
...
Put in a slightly different way, the charge carried by the electrons through
the external circuit must be accompanied by a compensating transport of ions
between the two cells
...
This ionic transport involves not only the
electroactive species Cu2+ and Zn2+ , but also the counterions, which in this
example are NO−
...
More detailed studies reveal that both processes
occur, and that the relative amounts of charge carried through the solution by
positive and negative ions depends on their relative mobilities, which express the
velocity with which the ions are able to make their way through the solution
...
In the simplest cells, the barrier between the two solutions can be a porous membrane, but for precise measurements, a more complicated arrangement, known as a
salt bridge, is used
...
4) filled with
a concentrated solution of KCl and fitted with porous barriers at each end
...
This potential difference would combine with the two half-cell potentials so as introduce a degree of uncertainty into any
measurement of the cell potential
...
Cell description conventions
In order to make it easier to describe a given electrochemical cell, a special
symbolic notation has been adopted
...
2 would
be
Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s)
In this notation, the vertical bars indicate phase boundaries; the double vertical
bar in the middle denotes the phase boundary between the two solutions
...
Note
carefully that this is entirely independent of the physical location of the two
electrodes in the actual cell in Fig
...
There are several other conventions relating to cell notation and nomenclature that you are expected to know:
• The anode is where oxidation occurs, and the cathode is the site of reduction
...
• If electrons flow from the left electrode to the right electrode when the cell
operates in its spontaneous direction, the potential of the right electrode
will be higher than that of the left, and the cell potential will be positive
...
This means that if the electrons are
flowing from the left electrode to the right, a galvanometer placed in the
external circuit would indicate a current flow from right to left
...
The substance that loses or receives the electron is called the
electroactive species
...
There are a number of other kinds of electrodes which are widely
encountered in electrochemistry and analytical chemistry
...
If neither of the electroactive species is a metal, some other
metal must serve as a conduit for the supply or removal of electrons from the
system
...
Such a half cell would be represented as
Pt(s) | Fe2+ (aq) || · · ·
and the half-cell reaction would be
Fe2+ (aq) + e− −→ Fe3+ (aq)
The reaction occurs at the surface of the electrode (Fig
...
The electroactive
ion diffuses to the electrode surface and adsorbs (attaches) to it by van der Waals
and coulombic forces
...
This process always costs energy;
if a lot of work is required, then only a small fraction of the ions will attach to
the surface long enough to undergo electron transfer, and the reaction will be
slow
...
Gas electrodes Some electrode reactions involve a gaseous species such as H2 ,
O2 , or Cl2
...
A typical reaction of considerable commercial
importance is
Cl− (aq) −→ 1 Cl2 + e−
2
3
STANDARD HALF-CELL POTENTIALS
10
Similar reactions involving the oxidation of Br2 (l) or I2 (s) also take place at
platinum surfaces
...
The electrode reaction consists in the oxidation and reduction of the silver:
AgCl(s) + e− −→ Ag(s) + Cl− (aq)
The half cell would be represented as
· · · || Cl− (aq) | AgCl(s) | Ag(s)
Although the usefulness of such an electrode may not be immediately apparent,
this kind of electrode finds very wide application in electrochemical measurements, as we shall see later
...
This voltage, which we usually refer to as the cell potential, is the
potential difference between the electrodes, and is the difference between the
half-cell potentials of the right and left sides:
Ecell = ∆V = Vright − Vleft
(2)
Each of the half-cell potentials is in turn a potential difference between the
electrode and the solution, so for our example cell the above relation can be
expanded to
Ecell = VCu − Vsoln + Vsoln − VZn
(3)
It is important to understand that individual half-cell potentials are not directly
measurable; there is no way you can determine the potential difference between a
piece of metal and a solution
...
The fact that individual half-cell potentials are not directly measurable does
not prevent us from defining and working with them
...
In particular, if we adopt a reference halfcell whose potential is arbitrarily defined as zero, and measure the potentials
of various other electrode systems against this reference cell, we are in effect
3
STANDARD HALF-CELL POTENTIALS
11
potential difference
H2
H2
Platinum sheet coated
with colloidal Pt
standard hydrogen
reference electrode
test cell
salt bridge
containing
saturated KCl
Figure 4: Cell for measurement of standard potentials
measuring the half-cell potentials on a scale that is relative to the potential of
the reference cell
...
When this electrode is set up under standardized conditions, it becomes
the standard hydrogen electrode, sometimes abbreviated SHE
...
71 v
Zn(s)
−
...
44
Cd(s)
−
...
126
Pb2+ + 2e− →
H2 (g)
0
...
222
AgCl(s)+ e− →
2Cl− (g)+ Hg(l) +
...
337
2I− (s)
+
...
771
Fe3+ + e− →
Ag(s)
+
...
23
O2 (g)+ 4H+ + 4e− → 2 H2 O(l)
2Cl− (g)
+1
...
The conventional
way of doing this, as shown in Table 3, is to write the half-cell reactions as reductions, and to place them in the order of increasing (more positive) potentials
...
Reference electrodes
In most electrochemical experiments our interested is concentrated on only one
of the electrode reactions
...
The major requirements of a reference
electrode is that it be easy to prepare and maintain, and that its potential be
3
STANDARD HALF-CELL POTENTIALS
saturated calomel
electrode
saturated KCl
solution
Hg - Hg2Cl2 mixture
glass
frit
KCl
crystals
fiber wick for contact
with external solution
13
silver-silver chloride
electrode
KCl
solution
Ag wire coated with
AgCl
fiber wick for contact
with external solution
Figure 5: Reference electrodes: the silver-silver chloride and calomel electrodes
stable
...
The most common way of accomplishing this is to use an electrode reaction
involving a saturated solution of an insoluble salt of the ion
...
The coating is done by making the silver the anode in an electrolytic cell containing HCl; the Ag+ ions combine with Cl− ions as fast as they are formed at
the silver surface
...
Hg | Hg2 Cl2 | KCl || · · ·
Hg + Cl− −→
1
2 Hg2 Cl2
+ e−
(7)
The potentials of both of these electrodes have been very accurately determined
against the hydrogen electrode
...
Also, there is need for a
supply of hydrogen gas which makes it somewhat cumbersome and hazardous
...
3
STANDARD HALF-CELL POTENTIALS
14
Problem Example 2
Find the standard potential of the cell
Cu(s) | Cu2+ || Cl− | AgCl(s) | Ag(s)
and predict the direction of electron flow when the two electrodes are
connected
...
The net reaction corresponding to this cell will be
2 Ag(s) + 2 Cl− (aq) + Cu2+ (aq) −→ AgCl(s) + Cu(s)
Since this involves the reverse of the AgCl reduction, we must reverse the
corresponding half-cell potential:
Ecell = (
...
222) v =
...
Note carefully
that in combining these half-cell potentials, we did not multiply the E ◦
for the Cu2+ /Cu couple by two
...
Cell potentials and the electromotive series
The remarkable thing about Table 3 is that similar tables, containing the same
sequence of reactions, were in common use long before electrochemical cells were
studied and half-cell potentials had been measured
...
This sequence is known as the activity series of the metals, and expresses
the decreasing tendency of the species listed in this column to lose electrons–
that is, to undergo oxidation
...
Consider, for example, the oxidation of Cu2+ by metallic
zinc that we have mentioned previously
...
By the same token, the tendency of Zn2+ to accept electrons is relatively small
...
We would therefore expect the
reaction
Zn(s) + Cu2+ −→ Zn2+ + Cu(s)
to proceed in the direction indicated, rather than in the reverse direction
...
The presence of the half-cell potentials in Table 3 allows us to attach numbers
to our predictions
...
Cell potentials and free energy
From the above, it should be apparent that the potential difference between the
electrodes of a cell is a measure of the tendency for the cell reaction to take place:
the more positive the cell potential, the greater the tendency for the reaction
to proceed to the right
...
Thus E ◦ and ∆G◦ measure the
same thing, and are related in a simple way:
∆G◦ = −nFE ◦
(8)
A few remarks are in order about this very fundamental and important relation:
• The negative sign on the right indicates that a positive cell potential (according to the sign convention discussed previously) implies a negative free
energy change, and thus that the cell reaction will proceed to the right
...
The right side of Eq 8 refers to the movement of n moles of
charge across the cell potential E ◦ , and thus has the dimensions of work
...
“Useful” here means work other than P-V work that is
simply a consequence of volume change, which cannot be channelled to
some practical use
...
The more rapidly the cell
operates, the less electrical work it can supply
...
To relate these units to electrical units, recall that the coulomb
is one amp-sec, and that power, which is the rate at which work is done,
is measured in watts, which is the product of amps and volts
...
From Table 3, E ◦ for the cell
Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s)
is 1
...
0 v
...
0 volt = −212000 watt-sec
which is also the free energy change under these conditions
...
If Eq 8 is solved for E ◦ , we have
E◦ = −
∆G◦
nF
This explains why we do not have to multiply the E ◦ s of half reactions by
stoichiometric factors when we are finding the E ◦ of a complete cell; since n is
in the denominator, we can think of cell potentials as free energy changes per
mole of charge transferred
...
222 v
−(+
...
115 v
−nFE ◦ = ∆G◦
−42800 J
+65000 J
+22200 J
Note, however, that if we are combining two half reactions to obtain a third
half reaction, the E ◦ values are not additive, since we are not eliminating electrons
...
Problem Example 4
Calculate E ◦ for the electrode Fe3+ /Fe from E ◦ values for Fe3+ /Fe2+ and
Fe2+ /Fe
...
Tabulate the E ◦ values and calculate the ∆G◦ s as follows:
Fe3+ + e− −→ Fe2+
Fe2+ + 2 e− −→ Fe(s)
Fe3+ + 3 e− −→ Fe(s)
Eo1 =
...
440 v
E◦3 = ?
∆G◦ 1 = −
...
880F
...
109nF, so E ◦ 3 = −
...
036 v
...
The greater the value
of E ◦ , the greater the tendency of the substance on the left to acquire electrons,
and thus the stronger this substance is as an oxidizing agent
...
Both involve the transfer of a species from a source, the donor, to a
sink, the acceptor
...
From the relation
∆G◦ = −RT ln Ka , you can see that the acid dissociation constant is a measure
of the fall in free energy of the proton when it is transfered from a donor HA to
the solvent H2 O, which represents the reference (zero) free energy level of the
proton in aqueous solution
...
By virtue of the defined
value of E ◦ = 0 for H+ /H2 , the latter level can also be regarded as the zero free
energy level of the electron
...
6 shows a number of redox couples on an electron free-energy scale
...
• Thus, in Fig
...
Similarly, if Fe3+ and I− are both present, one would expect a
3
higher reductant to reduce Fe3+ before I− , as long the two reactions take
3
place at similar rates
...
In the latter role, water
can serve as an electron sink to any metal listed above it
...
A spectacular
example of this is the action of water on metallic sodium
...
For example, an aqueous solution of Cl2 will slowly decompose into hypochlorous acid (HOCl)
with the evolution of oxygen
...
3
STANDARD HALF-CELL POTENTIALS
18
electron
sinks
electron
sources
Li
300
Na
H2
This diagram contains the same
information as a tab le of standard
p otentials, b ut its different
arrangement can p rovide more
insight into the meaning of E¡ values
...
0
HÐ
Ð2
...
There is, however, one very
imp ortant difference: p rotonexchange reactions are very fast,
b ut ox idation -reduction reactions
can b e extremely slow, the order in
which a comb ination ox idants
b ecome reduced cannot b e
p redicted from the thermodynamic
information given here
...
5
100
Ð1
...
5
Fe 2+
Fe
0
...
5
I2
Fe 3+
Ag +
Ni
Pb
H2
Cu
I
Fe 2+
Ag
1
...
5
2
...
5
3
...
FÐ
Ð300
free energy released p er mole of electrons transferred to hydrogen ion
The chemical sp ecies shown on the
left are oxidants (oxidizing agents),
which can b e regarded as
sub stances p ossessing unoccup ied
electron energy levels
...
Ð2
...
The scale on the
left exp resses these values as free
energies p er electron-moleÐ that is,
as conventional E¡ values
...
3
STANDARD HALF-CELL POTENTIALS
19
• A metal that is above the H+ /H2 couple will react with acids, liberating
H2 ; these are sometimes known as the “active” metals
...
3
STANDARD HALF-CELL POTENTIALS
20
Latimer diagrams
Considerable insight into the chemistry of a single element can be had by comparing the standard electrode potentials (and thus the relative free energies)
of the various oxidation states of the element
...
7
...
440
acid solution
Fe
alkaline solution
Fe
acid solution C l
—1
...
9
Fe(O H )
+1
+2
—1
...
0
4
C l“
Ð1
...
56
—1
...
5
3
Cl
Ð
...
887
—1
alkaline solution
+3
+4
—1
...
6
6
C lO
FeO “
+6
—1
...
3
3
Cl “
O
C lO £
—
...
9
8
Figure 7: Latimer diagrams showing relative stabilities of different oxidation
states of an element
oxidation state of the element are written from left to right in order of increasing
oxidation number, and the standard potential for the oxidation of each species
to the next on the right is written in between the formulas
...
) Potentials for reactions involving
hydrogen ions will be pH dependent, so separate diagrams are usually provided
for acidic and alkaline solutions (effective hydrogen ion concentrations of 1 M
and 10−14 M)
...
7)
...
440 v) indicates
that Fe will dissolve in 1 M H+ to form Fe2+
...
771 v), the +2 state will be
the stable oxidatation state of iron under these conditions
...
This indicates that the species
will tend to undergo disproportionation, or self-oxidation and reduction
...
Although the potential for its oxidation is negative, the potential for its reduction to Cl− is positive (+1
...
Thus elemental
chlorine is thermodynamically unstable with respect to disproportionation in
alkaline solution, and the same it true of the oxidation product, HClO−
...
Thus the disproportionion of chlorine mentioned above occurs only
very slowly
...
4
The Nernst equation
The standard cell potentials we have been discussing refer to cells in which all
dissolved substances are at unit activity, which essentially means an “effective
concentration” of 1 M
...
If these
concentrations or pressures have other values, the cell potential will change in
a manner that can be predicted from the principles you already know
...
5 M to a much smaller value:
Zn(s) | Zn2+ ((aq), 0
...
The relation between the actual cell potential E and the standard potential
E ◦ is developed in the following way
...
059
E of Ag+/Ag
half-cell
...
Notice that the cell potential will be the same as E ◦ only if Q is unity
...
05915
n
log Q
(12)
This relation predicts that a cell potential will change by 59 millivolts per
10-fold change in the concentration of a substance involved in a one-electron
oxidation or reduction; for two-electron processes, the variation will be 28 mv
per decade concentration change
...
8, these predictions are
only fulfilled at low concentrations, not only of the electroactive ion, but of all
ionic species
...
Determination of activity coefficients The activity coefficient γ relates the concentration of an ion to its activity in a given solution: a = γc
...
The resulting γs can then be used to convert concentrations into activities for use in
other calculations involving equilibrium constants
...
One way of doing this is by means of a concentration cell such as
Cu(s) | CuNO3 (
...
01 M ) | Cu(s)
(13)
The electrode cell reactions are
cathode:
anode:
net reaction:
Cu2+ (
...
01 M )
Cu2+ (
...
01 M )
which represents the transport of cupric ion from a region of higher concentration to one of lower concentration
...
05915
log Q = 0 −
...
1 = +
...
Analytical applications of the Nernst equation
A very large part of Chemistry is concerned, either directly or indirectly, with
determining the concentrations of ions in solution
...
Cell potentials are fairly easy to measure, and although the
Nernst equation relates them to ionic activities rather than to concentrations,
the difference between them becomes negligible in solutions where the total ionic
concentration is less than about 10−3 M
...
Rather than
measuring the concentration of the relevant ions directly, the more common
procedure is to set up a cell in which one of the electrodes involves the insoluble
salt, and whose net cell reaction is just the dissolution of the salt
...
222 v
−E ◦ =-(+
...
In such cases it is often
possible to determine the ion indirectly by titration with some other ion
...
The titration is carried out in
one side of a cell whose other half is a reference electrode:
Pt(s) | Fe2+ , Fe3+ || reference electrode
Initially the left cell contains only Fe2+
...
Once
the first drop of Ce3+ has been added, the potential of the left cell is controlled
by the ratio of oxidized and reduced iron according to the Nernst equation
E =
...
059 log
[Fe2+ ]
[Fe3+ ]
When the equivalence point is reached, the Fe2+ will have been totally consumed
(the large equilibrium constant ensures that this will be so), and the potential
will then be controlled by the concentration ratio of Ce3+ /Ce4+
...
If one works
out the actual cell potentials for various concentrations of all these species, the
resulting titration curve (Fig
...
The end point is found not by measuring a particular cell voltage, but
by finding what volume of titrant gives the steepest part of the curve
...
H2 (g, 1atm) | Pt | H+ (? M ) || reference electrode
4
THE NERNST EQUATION
25
1
...
1
...
8
Initially, the solution is mostly Fe2+
...
0
...
4
0
10
20
30
volume of Ce4+ solution added
Figure 9: Curve for the potentiometric titration of FeSO4 with Ce(SO4 )2
glass hydration gel layers
external (test)
solution
internal
soution
0
...
In 1914 it was discovered that a thin glass membrane enclosing a solution
of HCl can produce a potential that varies with [H+ ] in about the same way
as that of the hydrogen electrode
...
10) are manufactured
in huge numbers for both laboratory and field measurements
...
The potential of a glass electrode is given by a form of the Nernst equation
very similar to that of an ordinary hydrogen electrode, but of course without
4
THE NERNST EQUATION
26
the H2 :
RT
ln[aH+ + constant]
F
The reason a glass membrane would behave in this way was not understood
until around 1970
...
These sodium ions diffuse to
whichever side of the membrane has the lower H+ concentration, where they
remain mostly confined to the surface of the glass, which has a porous, gelatinous
nature
...
Emembrane = constant +
Ion-selective electrodes The function of the membrane in the glass electrode is to allow hydrogen ions to pass through and thus change its potential,
while preventing other cations from doing the same thing1
...
Since about 1970, various other membranes have been developed which show similar selectivities to certain other
ions
...
Membrane potentials
You may recall the phenomena of osmosis and osmotic pressure that are observed
when two solutions having different solute concentrations are separated by a
thin film or membrane whose porosity allows small ions and molecules to diffuse
through, but which holds back larger particles
...
This will produce a charge imbalance between the two solutions, with
the original solution having the charge sign of the larger ion
...
This potential difference is usually called a membrane
potential or Donnan potential
...
11 shows a simple system containing the potassium salt of a protein on
one side of a membrane, and potassium chloride on the other
...
The value of this potential difference can be expressed by a
1 This selectivity is never perfect; most glass electrodes will respond to moderate concentrations of sodium ions, and to high concentrations of some others
...
This
membrane potential can be observed by
introducing a pair of Pt electrodes
...
The membrane potential can be expressed in terms of the ratio
of either the K+ or Cl− ion activities:
∆Φ =
{K + }right
{Cl− }left
RT
RT
ln
ln
=
+}
nF
{K left
nF
{Cl− }right
(14)
The membrane surrounding most living cells contains sites or “channels” through
which K+ ions are selectively transported so that the concentration of K+ inside the
cell is 10-30 times that of the intracellular fluid
...
059 log10
1
= −70 mV
20
which is consistent with observed values
...
The metabolic
processes governing this action are often referred to as “ion pumps”
...
The normal potential difference between the inner and outer parts of nerve cells is
about −70 mv as estimated above
...
This has the effect of temporarily
opening the Na+ channel; the influx of these ions causes the membrane potential of
the adjacent portion of the nerve to collapse, leading to an effect that is transmitted
along the length of the nerve
...
5
5
BATTERIES AND FUEL CELLS
29
Batteries and fuel cells
An electrochemical cell which operates spontaneously can deliver an amount of
work to the surroundings whose upper limit (in the case of reversible operation)
is equal to the fall in free energy as the cell reaction proceeds
...
As the reaction continues the free energy of the
system falls, so as time goes by less energy remains to be recovered
...
A battery is a practical adaptation of this arrangment and usually consists of
a number of cells connected in series so as to attain the desired output voltage
...
2;
this was used, among other things, for telegraphy and railroad signalling during
the nineteenth century, when batteries provided the only practical source of
electrical power
...
The latter
is also called a storage cell, which more aptly describes its ability to convert
electrical energy into chemical energy and then re-supply it as electrical energy
on demand
...
A primary cell,
as expemplified by an ordinary flashlight battery, cannot be recharged with any
efficiency, so the amount of energy it can deliver is limited to that obtainable
from the reactants that were placed in it at the time of manufacture
...
The cell is represented by
Pb(s) | PbSO4 (s) | H2 SO4 (aq) | PbSO4 (s), PbO2 (s) | Pb(s)
and the net cell reaction
Pb(s) + PbO2 (s) + 2 H2 SO4 (aq) −→ 2 PbSO4 (s) + 2 H2 O
The reaction proceeds to the right during discharge and to the left during charging
...
The fuel cell
Conventional batteries supply electrical energy from the chemical reactants
stored within them; when these reactants are consumed, the battery is “dead”
...
In this case the
reactants can be thought of as “fuel” to drive the cell, hence the term fuel cell
...
At the time, it was
already known that water could be decomposed into hydrogen and oxygen by
electrolysis; Grove tried recombining the two gases in a simple appratus, and
discovered what he called “reverse electrolysis”:
H2 (g) −→ 2 H+ + 2 e−
1
+
+ 2 e− −→ H2 O
2 O2 (g) + 2 H
E◦ = 0 v
E ◦ = 1
...
23 v
(15)
It was not until 1959 that the first working hydrogen-oxygen fuel cell was developed
...
One reason for the interest in fuel cells is that they offer a far more efficient
way of utilizing chemical energy than does conventional thermal conversion
...
If the hydrogen were simply burned
in oxygen, the heat obtainable would be −∆H ◦ = 242 kJ mol−1 , but no more
than about half of this heat can be converted into work2 , so the output would
be 121 kJ mol−1 or less
...
Coating the electrode
with a suitable catalytic material is almost always necessary to obtain useable
output currents, but good catalysts are mostly very expensive substances such
as platinum, so that the resulting cells are too costly for most practical uses
...
6
Electrochemical Corrosion
Corrosion is the destructive attack of a metal by chemical or electrochemical
reaction with its environment
...
This is the reason that considerable energy (and
expense) must go into the extraction of a metal from its ore
...
To do so, the metal must lose electrons, and this requires
the presence of an electron acceptor or oxidizing agent
...
The special characteristic of most corrosion processes is that the oxidation
and reduction steps occur at separate locations on the metal
...
In this sense the system can be regarded as
an electrochemical cell in which the anodic process is something like
Fe(s) −→ Fe2+ (aq) + 2 e−
and the cathodic steps can be any of
O2 + 2 H2 O + 4 e− −→ 4 OH−
2 This
limit is a consequence of the Second Law of Thermodynamics
...
= (1 − Thigh )/Tlow
...
3 A noble metal is one that appears toward the bottom of a listing of standard reduction
potentials as in Table 3
...
6, you can see that such metals serve as lower-free energy
sinks which can accept electrons from higher (less noble) metals
...
Figure 14: Electrochemical corrosion of a nail showing anodic and cathodic
regions
H+ + e− −→
2+
M
+ 2e
−
1
2
H2 (g)
−→ M(s)
where M is an more noble metal
...
Which parts of the metal serve as anodes and cathodes can depend on many
factors, as can be seen from the irregular corrosion patterns that are commonly
observed
...
14
...
However, practically all metallic surfaces that have been exposed to
the atmosphere are coated with a thin film of the metal oxide, which tends to
shield the metal from the electrolyte and thus prevent corrosion4
...
The fact that such sites are usually
hidden from view accounts for much of the difficulty in detecting and controlling
corrosion
...
4 Metals such as aluminum and stainless steels form extremely tough and adherent oxide
films that afford extraordinary corrosion resistance
...
The tendency of oxygen-deprived locations to become anodic is the cause of
many commonly-observed patterns of corrosion
...
15a)
begins when corrosion hollows out a narrow hole, or pit, in the metal
...
Fig
...
Rusted-out cars and bathroom stains
...
You will also
have noticed that once corrosion starts, it tends to feed on itself
...
The high
pH produced in these cathodic regions tends to destroy the protective oxide film, and
may even soften or weaken paint films, so that these sites can become anodic
...
A very common cause of corrosion is having two dissimilar metals in contact, as
might occur near a fastener or at a weld joint
...
Moisture and conductive salts on the outside surfaces acts as an external conductor,
effectively short-circuiting the cell and producing very rapid corrosion; this is why cars
rust out so quickly in places where salt is placed on roads to melt ice
...
For example, in homes where copper tubing is used for plumbing, there is
always a small amount of dissolved Cu2+ in the water
...
In the case
of chrome bathroom sink fittings, this leads to the formation of Cr3+ salts which
precipitate as greenish stains
...
Sacrificial coating of more active metal
protects exposed area of base by
supplying electtrons, keeping it cathodic
...
The most obvious strategy is to stop
both processes by coating the object with a paint or other protective coating
...
A more sophisticated strategy is to maintain a continual negative electrical
charge on a metal, so that its dissolution as positive ions is inhibited
...
The source of electrons can be an external direct current power supply (commonly used to protect oil pipelines and other buried
structures), or it can be the corrosion of another, more active metal
...
16a
...
Dissolution of this sacrificial coating leaves behind electrons which
concentrate in the iron, making it cathodic and thus inhibiting its dissolution
...
The common tin-plated can (Fig
...
As long as
the tin coating remains intact, all is well, but exposure of even a tiny part of
the underlying iron to the moist atmosphere initiates corrosion
...
7
Electrolytic cells
Electrolysis refers to the decomposition of a substance by an electric current
...
” One pole of a battery of copper-zinc cells was connected
to the spoon, and the other was connected to platinum wire which dipped
into the melt
...
The potash appeared to be a conductor in a high degree, and as long
as the communication was preserved, a most intense light was exhibited
at the negative wire, and a column of flame, which seemed to be owing to
the development of combustible matter, arose from the point of contact
...
In
another experiment, Davy observed “small globules having a high metallic
lustre, precisely similar in visible characters to quicksilver, some of which
burnt with explosion and bright flame, as soon as they were formed, and
others remained, and were merely tarnished, and finally covered by a white
film which formed on their surfaces
...
Thus if a solution of
nickel chloride undergoes electrolysis at platinum electrodes, the reactions are
cathode:
anode:
net reaction:
Ni2+ + 2 e− −→ Ni(s)
2 Cl− −→ Cl2 (g) + 2 e−
Ni2+ + 2 Cl− −→ Ni(s) + Cl2 (g)
E ◦ = −0
...
36 v
E ◦ = 1
...
The free energy is supplied
in the form of electrical work done on the system by the outside world (the
surroundings)
...
Electrolysis involving water
If we substitute sodium chloride for nickel chloride, dihydrogen is produced at
the cathode instead of sodium:
7
ELECTROLYTIC CELLS
cathode:
anode:
net reaction:
H2 O + 2 e− −→ H2 (g) + 2 OH−
Cl− −→ 1 Cl2 (g) + e−
2
Na+ + Cl− −→ Na(l) + 1 Cl2 (g)
2
36
E = +0
...
36 v
The reason that sodium is not a product of this reaction is best understood by
locating the couples Na+ /Na and H2 O/H2 ,OH− in Fig
...
Reduction
of Na+ (E ◦ = 2
...
You will recall that water can be oxidized as well as reduced, so if we replace
the chloride ion with an anion such as nitrate or sulfate that is much more
difficult to oxidize, the water is oxidized intead
...
41 v ([OH− ] = 10−7 M )
E ◦ = −0
...
23 v
Faraday’s laws of electrolysis
One mole of electric charge (96,500 coulombs), when passed through a cell, will
discharge half a mole of a divalent metal ion such as Cu2+
...
The weights of substances formed at an electrode during electrolysis are
directly proportional to the quantity of electricity that passes through the
electrolyte
...
The weights of different substances formed by the passage of the same
quantity of electricity are proportional to the equivalent weight of each
substance
...
Thus one mole of V3+ corresponds to three equivalents of this species, and will
require three faradays of charge to deposit it as metallic vanadium
...
7
ELECTROLYTIC CELLS
37
+
Ð
Cl2 out
H2 out
caustic
spent brine
Na+
brine feed
ClÑ ® Cl2
Na+
H2 O ®
H2 + OHÐ
ion-selective
polymer membrane
water
Figure 17: Membrane cell for industrial chloralkali production
Industrial electrolytic processes
For many industrial-scale operations involving the oxidation or reduction of both
inorganic and organic substances, and especially for the production of the more
active metals such as sodium, calcium, magnesium, and aluminum, the most
cost-effective reducing agent is electrons supplied by an external power source
...
The chloralkali industry
...
Because the reduction
potential of Na+ is much higher than that of water, the latter substance undergoes decomposition at the cathode, yielding hydrogen gas and OH−
...
36 v
−
...
7 v
0v
i
ii
iii
iv
A comparison of the E ◦ s would lead us to predict that reactions (ii) and (iv)
would predominate and that (iii) would be unimportant
...
The net reaction for
the chlorination of brine is thus
2 NaCl(aq) + 2 H2 O −→ 2 NaOH + Cl2 (g) + H2 (g)
7
ELECTROLYTIC CELLS
38
positive bus bar
carbon anode
lowered into melt at rate
of about 2-3 cm per day
frozen crust of
electrolyte and alumina
carbon lining
molten electrolyte consisting of alumina
Al2O3 and cryolite 3NaF¥AlF3
molten aluminum
product
negative collector plate
Figure 18: Hall-H´rault cell for the production of aluminum
e
Since chlorine reacts with both OH− and H2 , it is necessary to physically separate the anode and cathode compartments
...
A small amount of this mercury would normally find its way into the
plant’s waste stream, and this has resulted serious pollution of many major
river systems and estuaries and devastation of their fisheries
...
Electrolytic refining of aluminum
Aluminum is present in most rocks, makes up 8% of the earth’s crust, and is
potentially the world’s most abundant metal
...
For the same reason, aluminum cannot be isolated by
electrolysis of aqueous solutions of its compounds, since the water would be
electrolyzed preferentially
...
S
...
The Hall-H´rault process takes advantage of the principle that the melting
e
point of a substance is reduced by admixture with another substance with which
it forms a homogeneous phase
...
The anodes of the cell are made of carbon (actually a mixture of
pitch and coal), and this plays a direct role in the process; the carbon gets
7
ELECTROLYTIC CELLS
39
oxidized (by the oxide ions left over from the reduction of Al3+ ) to CO, and the
free energy of this reaction helps drive the aluminum reduction, lowering the
voltage that must be applied and thus reducing the power consumption
...
Since aluminum cells commonly operate at about 100,000 amperes, even a slight
reduction in voltage can result in a large saving of power
...
c 1994 by Stephen K
...
March 8, 1994
Title: electrochemistry
Description: well ..defined...notes..on electochemistry
Description: well ..defined...notes..on electochemistry