00:00:00
in this session we're going to be
00:00:02
exploring what happens with
00:00:03
electrochemical equilibrium in
00:00:05
electrochemical environments so thinking
00:00:07
about electrochemistry at equilibrium
00:00:09
our general method for exploring
00:00:12
electrochemistry
00:00:12
is to consider what happens at these
00:00:15
equilibria so firstly we established
00:00:17
this electrochemical equilibrium and
00:00:19
then we disturb this equilibrium by
00:00:21
applying an external potential now an
00:00:24
electrochemical equilibrium is not the
00:00:26
same as a chemical equilibrium but the
00:00:29
definition is we're looking at no net
00:00:30
current flowing across the interface so
00:00:33
if we look at what's going on in this
00:00:34
interface here there will be a
00:00:36
continuous exchange a constant exchange
00:00:37
of electrons but there is no net current
00:00:41
flowing so this is a dynamic equilibrium
00:00:44
so this is different from a chemical
00:00:46
equilibrium it is an electrochemical
00:00:48
equilibrium so let's think about the
00:00:51
different types of cells that we have
00:00:52
that make use of this equilibrium so the
00:00:55
first thing we're looking at is a
00:00:56
galvanic cell so a galvanic cell has a
00:00:59
spontaneous reaction which converts
00:01:01
chemical potential to electrical energy
00:01:03
when the switch for this galvanic cell
00:01:07
is open no current can flow so there's
00:01:10
an electrochemical equilibrium at the
00:01:11
electrode surfaces we have this
00:01:13
continual exchange of electrons that one
00:01:15
and a continual exchange of electrons at
00:01:17
the other but there is no net current
00:01:18
flowing however each electrode is at a
00:01:21
different potential when we close the
00:01:24
switch current is allowed to flow we get
00:01:27
spontaneous oxidation happening at the
00:01:29
anode spontaneous reduction happening at
00:01:31
the cathode and this difference in
00:01:33
potential allows the current to flow
00:01:34
lighting the bulb this particular cell
00:01:37
I've drawn is an example of a cell known
00:01:39
as a Daniell cell it is a standard cell
00:01:42
for electro chemical equilibria it's
00:01:44
well recognized and well understood so a
00:01:46
galvanic cell relies on a spontaneous
00:01:48
chemical process to convert to
00:01:50
electrical energy so what other types of
00:01:53
cell do we have well the electrolytic
00:01:55
cell is the second type these ones have
00:01:57
a non spontaneous reaction and these
00:02:00
ones rely on putting electrical energy
00:02:03
into the cell and it drives a non
00:02:05
spontaneous process the example I've put
00:02:07
here is electrolysis of water so we have
00:02:09
a power supply which applies a potential
00:02:12
difference which forces those fair
00:02:13
to change which creates a reaction at
00:02:17
each surface it raises the potential of
00:02:18
one lowers the potential the other and
00:02:20
drives the reaction forward rechargeable
00:02:23
batteries are an example of something
00:02:26
that behaves as both types of cell so a
00:02:29
rechargeable battery is galvanic on
00:02:31
discharging so it's supplying electrical
00:02:33
energy to the appliance using that
00:02:35
chemical energy to generate the
00:02:37
electrical potential but when we want to
00:02:40
put energy back into it for storage it
00:02:42
becomes electrolytic as we charge it so
00:02:45
it stores that electrical energy as
00:02:46
chemical potential the next phase of
00:02:49
equilibria we want to look at is the
00:02:50
Nernst equation so just a quick recap on
00:02:53
this you covered this in year one I've
00:02:55
linked the video below
00:02:57
this relates electrode potentials to
00:02:59
free energies and there are several ways
00:03:01
to represent this but the main way that
00:03:03
we're most familiar with is this form of
00:03:05
the Nernst equation where we're looking
00:03:07
at how the electrode potential varies
00:03:10
according to the number of electrons
00:03:12
exchanged and the reaction quotient so
00:03:14
it modifies the standard potential for
00:03:16
our real reaction conditions so remember
00:03:19
the reaction quotient whenever we're
00:03:21
looking at equilibria is a product of
00:03:23
the activity to the right hand side
00:03:24
divided by the products the activity to
00:03:26
the left hand side you remember doing
00:03:28
this as a products over reactants but
00:03:31
when we're dealing with equilibria we
00:03:32
don't really have products and reactants
00:03:34
so we need to look at the different
00:03:36
sides of the equation we normally use
00:03:40
concentrations rather than activities
00:03:41
assuming that the standard activity is
00:03:43
unity this only applies at very low
00:03:45
concentrations however but it means that
00:03:47
we end up with Q being unitless which
00:03:50
makes life a lot easier
00:03:52
so let's think about our reaction
00:03:54
quotient here if we have our general
00:03:56
reaction remember we have the products
00:03:58
of the right hand side divided by the
00:04:00
products of the left hand side and at
00:04:02
low concentrations activities are
00:04:04
approximately equal to concentration so
00:04:06
we can use this approximation this
00:04:08
doesn't tell the whole story of course
00:04:10
we need to look at the half equations
00:04:12
there's reduction half equations to find
00:04:14
the number of electrons that are being
00:04:15
transferred this allows us to establish
00:04:17
the value of n and therefore use the
00:04:19
Nernst equation
00:04:22
whenever we're thinking of using these
00:04:25
reactions we need to consider what
00:04:26
phases were working within so we're what
00:04:29
wondering whether we're looking at
00:04:30
solids liquids or gases so the first
00:04:33
thing to do is start with the half-cell
00:04:35
reactions so by convention we always
00:04:37
write these as reductions so if we
00:04:39
consider the phenomenon of the
00:04:41
electrolysis of water to release oxygen
00:04:43
gas we have liquid and gas present so
00:04:47
have liquid water we have gaseous oxygen
00:04:49
and we have aqueous hydrogen so how do
00:04:52
we treat the reaction quotient how do we
00:04:55
consider a concentration when we have
00:04:57
bulk liquid and bulk gas but whenever
00:04:59
we're thinking of the solvent remember
00:05:01
we're thinking about activities the
00:05:03
activity can be taken as a unity because
00:05:05
it is the solvent and the activity
00:05:07
doesn't change significantly as part of
00:05:09
the reaction so because it's not
00:05:11
changing we can accept it cancels out as
00:05:13
one for the gas we want to consider the
00:05:17
partial pressure so the partial pressure
00:05:19
of oxygen since the gases were almost
00:05:21
always working with atmospheric pressure
00:05:23
and there are being evolved at
00:05:26
atmospheric pressure and they are pure
00:05:27
gases at the point of evolution we can
00:05:30
also take these as being unity so this
00:05:34
allows us to simply consider this in
00:05:35
terms of the aqueous terms now we need
00:05:37
to make sure we look at our pressures
00:05:39
look at our concentrations to make sure
00:05:40
that that still applies but almost
00:05:43
always our gases and our solvents will
00:05:45
be considered under standard conditions
00:05:48
so let's apply the Nernst equation to a
00:05:51
full cell so using this we can predict
00:05:54
the variation of cell potentials the
00:05:56
spontaneous direction of reaction free
00:05:57
energy change okay this is a recap from
00:05:59
what you've done before so let's think
00:06:01
about the technique that we're going to
00:06:02
use so the first thing we need to do is
00:06:04
we need to write down the cell so we're
00:06:06
going to use the copper hydrogen cell
00:06:09
that I've detailed here the first thing
00:06:12
we need to do is you need to write down
00:06:13
the cell remembering to balance our half
00:06:15
cells so if we look at our two half
00:06:17
cells we can see that we don't have an
00:06:21
equal number of electrons so firstly we
00:06:23
need to balance the electron term we're
00:06:24
going to do this by multiplying the
00:06:26
second one by two so we get to H+
00:06:30
+ - e - going to H - and the standard
00:06:38
cell potential is unchanged by this
00:06:41
operation it's still going to be the
00:06:42
same potential even if we've doubled up
00:06:45
okay so that gives us something we can
00:06:48
equate we now subtract this from the
00:06:50
copper equation and what we end up with
00:06:53
is the overall cell reaction which
00:06:55
allows us to identify our left hand and
00:06:57
right hand side so our overall cell
00:07:00
equation is copper two-plus thus HG goes
00:07:06
to copper solid + 2 h plus and our
00:07:14
standard cell potential is simply the
00:07:16
difference between the two which in this
00:07:17
case is not point three four - no point
00:07:20
naught which is not point of three four
00:07:22
volts okay so that gives us our overall
00:07:26
cell reactions where we have our right
00:07:28
hand side and our left hand side
00:07:35
so this means that our reaction quotient
00:07:38
is going to be the concentrations of the
00:07:44
right-hand side so remember each class
00:07:49
will be squared because we've got that
00:07:50
two they're divided by concentration of
00:07:53
Cu two plus and the concentration of H
00:07:57
two now looking at this we've got copper
00:08:01
solid so the activity of this is going
00:08:03
to be unity so it's not going to be
00:08:05
changed by the reaction it's going to be
00:08:07
a prop it's going to be essentially
00:08:08
constant so we don't need to consider
00:08:10
that because it goes to one age two
00:08:13
let's say we're working under standard
00:08:14
conditions so we expect one atmosphere
00:08:17
of pressure so that would be unity as
00:08:19
well so that can disappear as well so
00:08:22
that tells us everything we want to deal
00:08:23
with great our ourselves at the standard
00:08:28
conditions well let's just say that
00:08:30
would consider our concentration of
00:08:32
protons to be 1 mole per diem cubed and
00:08:37
we consider the concentration of copper
00:08:39
two plus two B naught 0.2 moles per diem
00:08:42
cube so we're no longer operating under
00:08:46
standard conditions so we need to apply
00:08:48
our Nernst equation to adjust the
00:08:50
potential so we need to find out what
00:08:52
the reaction quotient is well
00:08:53
fortunately the proton concentration is
00:08:56
1 but the copper concentration has
00:08:58
changed so Q becomes 1 squared over not
00:09:03
0.2 which is equal to 5 and this allows
00:09:07
us to apply our Nernst equation our
00:09:10
Nernst equation then becomes
00:09:13
e is equal to not 0.34 volts minus R T
00:09:20
so 8.314 joules per Kelvin per mole
00:09:26
times our temperature resets at 25
00:09:29
Celsius that's 298.15 kelvin divide that
00:09:35
by n F so the number of electrons from
00:09:38
here we've got two electrons two
00:09:41
electrons times the Faraday constant
00:09:43
which is nine point six four eight zero
00:09:46
times tenth and four coulombs per mole
00:09:53
and we multiply that by log 5 okay let's
00:09:57
do some quick approximating which is not
00:10:00
0.34 volts first let us consider the yet
00:10:03
units so our Kelvin cancel I'm per mole
00:10:08
cancel and then we're ended end up with
00:10:11
a Coulomb which if you remember a
00:10:13
Coulomb is an amp second okay so that's
00:10:17
going to give us our units that we're
00:10:18
dealing with we're working with SI units
00:10:20
throughout so we don't need to do any
00:10:21
conversions so minus eight point three
00:10:24
one four times two nine eight
00:10:26
that's about just over eight times just
00:10:28
under 300 so you should end up with
00:10:30
something that's around about two
00:10:32
thousand four hundred eight times three
00:10:34
hundred two thousand four hundred divide
00:10:36
that by two times nine point six four
00:10:39
eight zero times tenth four which is
00:10:41
going to be about eighteen nineteen
00:10:46
times 10 to the 4 okay log 5
00:10:53
well 19 times 10 to the 4 that's a
00:10:55
products button in nineteen nineteen
00:10:58
hundred times ten to two so this is not
00:11:01
quite 1.4 times that so it's going to be
00:11:05
about what no point three four volts
00:11:09
minus twenty four hundred by 1900 is
00:11:12
going to be a but one point four times
00:11:14
ten to the minus two times log 5
00:11:19
remember this is asking what power of e
00:11:22
do we need to get five e is about two
00:11:24
point seven so it's going to be about
00:11:26
one in a bit it's not going to be
00:11:27
squared let's say that this is going to
00:11:29
be about one point five so one point
00:11:32
four times one point five gives us not
00:11:35
0.34 volts ten to the minus two - not
00:11:39
point not one point four times one point
00:11:41
five is about two point one times seven
00:11:46
minus two so get not point not to one
00:11:49
and remember we need to think about the
00:11:53
units so joules per amp per second which
00:11:57
is the definition of a volt so these
00:11:59
units are congruent so they add up
00:12:01
subtract this we end up with about nor
00:12:04
point three two volts which gives the
00:12:08
overall cell potential using the Nernst
00:12:10
equation adjusting for non-standard
00:12:12
conditions we can go through and we can
00:12:15
just check with a calculator of course
00:12:17
just checking with a calculator we end
00:12:18
up with that goes to at two four seven
00:12:22
eight this one goes to nineteen point
00:12:26
two seven six
00:12:30
times 10 to the 4 which means we end up
00:12:33
with not points not 1 2 8 4 log 5 which
00:12:46
gives us hmm comes out at minus naught
00:12:50
point naught 2 1 so we get the same
00:12:53
value with the calculator ok so this
00:12:56
tells us how we fix our equilibrium cell
00:12:59
potential to give us the overall cell
00:13:02
potential for this particular equation
00:13:03
and gives us the direction of
00:13:06
spontaneous change so it's a positive
00:13:08
value so it remains a positive value
00:13:11
which tells us that the direction of
00:13:12
spontaneous change is to go to form this
00:13:15
copper solid once again make sure you
00:13:18
revisit your first-year material that
00:13:19
you can follow this derivation we now
00:13:24
need to consider what we mean when we
00:13:25
say electrochemical versus chemical
00:13:27
equilibria electrochemical equilibria as
00:13:29
we said has no net current and are the
00:13:32
conditions for measuring cell potential
00:13:34
but electrochemical equilibrium has to
00:13:36
be present otherwise we're not able to
00:13:37
measure a cell potential when we set up
00:13:39
our cell like this our voltmeter has
00:13:41
have a very high internal resistance
00:13:43
that we don't get any current flowing so
00:13:45
we don't allow current to flow therefore
00:13:47
we've got electrochemical equilibria at
00:13:49
both electrodes that allows us to
00:13:51
measure that cell potential chemical
00:13:53
equilibrium however requires in a
00:13:55
slightly different definition it
00:13:57
requires that the Delta G for the entire
00:13:59
process is zero so remember that Delta G
00:14:03
is minus NFE that means that the cell
00:14:06
potential overall has to be zero but if
00:14:09
we allow current to flow the cell
00:14:11
potential is clearly not zero so we need
00:14:13
to reach a different state so a galvanic
00:14:16
cell at equilibrium remember galvanic
00:14:18
cell is where we have a spontaneous
00:14:19
current flowing a chemical equilibrium
00:14:21
the cell potential will be zero which
00:14:24
means Q equals K so the reaction
00:14:27
quotient is the equilibrium constant we
00:14:29
can use this equilibrium constant
00:14:31
because it's related to the standard
00:14:34
cell potential anyway what this allows
00:14:36
us to do is it allows us to predict an
00:14:38
equilibrium constant from a measured
00:14:40
standard
00:14:41
seneschal so we can get at the
00:14:42
equilibrium constant by using the cell
00:14:44
potential which is not measured at
00:14:46
chemical equilibrium so what does this
00:14:49
mean what we can find kay let's consider
00:14:52
an equilibrium such as this one where we
00:14:54
have the solvation of silver bromide
00:14:56
it's a sparingly soluble salt so p.m. it
00:14:59
being able to measure this solubility
00:15:01
product so essentially the equilibrium
00:15:04
product for this dissolution becomes
00:15:06
very tricky because we have a very low
00:15:08
solubility so how do we predict it well
00:15:11
once again we can use electrochemistry
00:15:13
so we need electrochemical potentials to
00:15:16
generate a reaction and we can work
00:15:19
backwards from that to determine our
00:15:20
equilibria my constant for this
00:15:22
dissolution so these are the two cells
00:15:24
that we're interested in and we just
00:15:26
apply the same rules as we did before we
00:15:29
combine needs to make the original
00:15:30
equation and identify the standard
00:15:33
electrode potential for that reaction
00:15:34
now the key thing is that the overall
00:15:36
reaction is not a redox process this
00:15:40
cell is purely hypothetical so the cell
00:15:42
isn't actually real but we use it for
00:15:45
the purposes of this investigation to do
00:15:50
this we have our two cells we have our
00:15:52
first cell and our second cell so all
00:15:55
we're going to do is we're going to
00:15:55
subtract the second one from the first
00:15:57
one this should generate our overall
00:16:00
cell potential so if we do 1 minus 2 we
00:16:04
end up with our hypothetical cell
00:16:06
becoming silver bromide minus silver
00:16:14
plus
00:16:17
going to a bromide line now the electron
00:16:20
terms balance so we can use these
00:16:21
equations directly remember we've got
00:16:23
the first one subtracting a second one
00:16:25
which gives us an a standard cell
00:16:28
potential of minus not 0.788 volts we
00:16:35
can rearrange this equation to give us
00:16:36
our silver bromide into silver plus
00:16:42
bromide okay now we just apply this
00:16:45
relationship here to find our
00:16:47
equilibrium cell potential so our East
00:16:50
and 'red is RT over NF l-- okay let's
00:16:54
rearrange this multiply both sides by NF
00:16:56
so log K is e standard NF divided by RT
00:17:07
and once again this is just simple case
00:17:12
of plugging things in we have a single
00:17:13
electron in the process we end up with 1
00:17:16
times I fired a constant which is nine
00:17:18
point six four eight zero times ten to
00:17:21
the four coulombs per mole multiplied by
00:17:26
our cell potential times minus not 0.788
00:17:30
volts divided by RT which is 8.314
00:17:35
joules per Kelvin per mole times 298.15
00:17:41
kelvin
00:17:43
okay apply the same approximations as we
00:17:46
did before remember 8.314 times two two
00:17:49
nine eight it's approximately 2,400 so
00:17:52
let's say something like naught point
00:17:54
seven eighty eight times this will give
00:17:55
us seven point five
00:17:57
times 10 to the 4 what are our units
00:18:00
gonna be let's quickly cancel some units
00:18:02
kelvins cancel Kelvin mole mole remember
00:18:07
we said that a volt
00:18:09
it's Joule per amp per second and a
00:18:13
Coulomb is an amp second we end up with
00:18:17
our amp seconds cancel our joules cancel
00:18:23
and we end up with a unitless term which
00:18:25
is what we need for our log K to work
00:18:27
okay
00:18:28
75 times 10 to the 4 which is
00:18:31
approximately 75,000 divided by 2400
00:18:35
cancel zeroes 750 divided by 24 is going
00:18:40
to be well there are 30 25 and 75 so
00:18:43
it's gonna be slightly more than 30
00:18:44
let's say it's approximately equal to 31
00:18:47
forgive me we've lost a minus sign we
00:18:49
need to keep that minus sign in - 31
00:18:52
so our okay that we get at the end is
00:18:55
going to be equal to e to the power
00:18:57
negative 31 which is going to be an
00:19:02
exceptionally small number indeed we can
00:19:05
apply a calculator again as we did
00:19:06
before and what we find is we end up
00:19:08
with if you multiply the top row out we
00:19:11
end up with minus seven point six three
00:19:15
times ten 10 to the four divide it by
00:19:19
two four seven eight which when we work
00:19:22
that through we end up with we end up
00:19:25
with - give me a - yes - 30 point seven
00:19:31
eight so we get approximately the same
00:19:33
number what I'm trying to show here is
00:19:35
through approximation we can get fairly
00:19:37
close to the actual value we're
00:19:39
expecting but what happens if we apply
00:19:42
the Nernst equation two half-cells
00:19:43
so once again let's revisit the copper -
00:19:46
cell we spoke a little bit about what's
00:19:49
the sound of one hand clapping remember
00:19:51
we can't measure these things in
00:19:52
isolation so every half cell is measured
00:19:55
relative to that standard hydrogen
00:19:56
electrode but what does the standard
00:20:00
electrode potential actually mean for a
00:20:01
half cell you know what what meaning do
00:20:04
we ascribe it well fundamentally it's a
00:20:06
balance point it's not saying that it is
00:20:09
not 0.34 volt
00:20:10
to drive copper in this direction it's
00:20:15
saying that the copper 1/2 cell is 0.34
00:20:18
volts more positive than the standard
00:20:20
hydrogen electrode it's just a relative
00:20:22
measurement so it shows the potential
00:20:25
that we would need to apply to switch
00:20:27
from galvanic to electrolytic cell
00:20:29
behavior if we want to look at how other
00:20:33
half cells compare we need to think
00:20:35
about free energies we can pretty much
00:20:37
measure anything we wish but it's
00:20:40
important that we consider free energies
00:20:42
it's not always possible to directly
00:20:44
compare electrode potentials remember
00:20:46
that not all half cells can be directly
00:20:48
measured so we use this relative
00:20:50
comparison between half cells to
00:20:52
determine the standard electrode
00:20:53
potential for hypothetical half cells
00:20:56
because not all of them can be directly
00:20:58
measured so let's think about the direct
00:21:00
reduction of our iron 3 to learn most
00:21:02
electron processes our single electron
00:21:04
or pair of electrons if we think about
00:21:06
iron three-plus while we can either add
00:21:08
one electron to become iron two-plus or
00:21:10
four iron two-plus we can add two
00:21:12
electrons to become iron metal now we
00:21:14
can't combine these because the
00:21:16
electrons don't cancel
00:21:18
we can't equate these we can't multiply
00:21:20
them up nothing's going to cancel out
00:21:22
because the electron terms don't work we
00:21:25
can't double this first ronk then we end
00:21:26
up with two Fe three-plus and 2 Fe 2
00:21:28
plus we would just end up with 3 iron
00:21:30
species in our final equation so because
00:21:33
the electrons don't cancel we have to
00:21:35
use free energies we convert these
00:21:37
electrode potentials into free energies
00:21:39
for each process so what we do is we
00:21:41
simply add both of them together to
00:21:43
obtain the overall equation and then we
00:21:45
add the free energies together so you
00:21:48
can see if we add these together the
00:21:49
iron 2 terms cancel out we gain an
00:21:51
electron and that gives us the overall
00:21:54
cell equation here okay but what do we
00:21:57
do in terms of free energies well we
00:21:59
just simply work out the free energy for
00:22:00
each of them so the free energy for the
00:22:04
first equation simply becomes one one
00:22:06
electron times the Faraday constant
00:22:07
times 0.77 1 the second one becomes 2
00:22:11
times minus 0.4 4 times the Faraday
00:22:14
constant which gives us our final value
00:22:17
of + naught point 109
00:22:19
time's the faraday constant which gives
00:22:22
a final cell potential when we work
00:22:24
backwards we simply apply this equation
00:22:26
in Reverse and we end up with a cell
00:22:28
potential of negative naught point
00:22:29
naught 3 6 so this allows us to
00:22:31
determine the cell potential of any
00:22:33
electrode provided we can establish how
00:22:36
we can put it together from the existing
00:22:38
measurable electrodes this is similar to
00:22:40
the process I spoke about before with
00:22:42
the law of independent migration and the
00:22:44
similar principle that you've applied
00:22:45
with Hess's law and any other situation
00:22:47
where you use a number of known
00:22:49
quantities to find the unknown to close
00:22:51
the loop so we have all of these
00:22:53
electrode potentials that we can
00:22:54
determine remember we said that a
00:22:57
standard electrode potential is simply a
00:22:58
relative measurement it's saying that
00:23:00
something is however much more positive
00:23:02
or negative than the hydrogen electrode
00:23:05
but this means it can be helpful to have
00:23:07
a visual aid whatever we think of visual
00:23:09
aids we think of drawing a graph the way
00:23:11
that we typically visualize electrode
00:23:13
potentials is to plot current against
00:23:15
potential and think about where we're
00:23:16
starting where we're going from so when
00:23:18
we're thinking about electrochemical
00:23:19
equilibrium if we've got a current
00:23:21
potential graph is zero so everything
00:23:24
becomes single dimension in this
00:23:26
particular visualization so we're going
00:23:28
to plot our standard electrode
00:23:29
potentials at I equals 0 our standard
00:23:31
hydrogen electrode by definition is at 0
00:23:34
so this is assuming standard conditions
00:23:37
where the activity of hydrogen is 1 the
00:23:38
partial pressure of hydrogen is 1 let's
00:23:41
look at the copper electrode that we
00:23:43
spoke about so once again let's say
00:23:45
we've got an activity of 1 this is
00:23:47
standard conditions remember and we have
00:23:50
our standard potential marked at 0.34
00:23:53
volts our silver electrode again
00:23:57
standard conditions has a more positive
00:24:00
potential so what we're saying is that
00:24:03
both of these are positive relative the
00:24:06
standard hydrogen electrode but we're
00:24:08
saying now that our silver chloride is
00:24:11
0.46 volts more positive than our copper
00:24:15
likewise our copper is 0.46 volts more
00:24:18
negative than our silver electrode these
00:24:21
are simply just relative measures and
00:24:23
this graph helps us visualize them if we
00:24:27
change our concentration from standard
00:24:28
conditions we'll get a different
00:24:30
potential so if we take the silver
00:24:32
chloride and we
00:24:33
juice the concentration we find we
00:24:36
reduce it to 1 millimolar we find that
00:24:39
applying the Nernst equation we end up
00:24:42
with a drop in our cell potential to
00:24:43
0.62 volts this is still more positive
00:24:47
than our copper electrode so our overall
00:24:50
cell potential we would find by simply
00:24:52
finding the difference between the two
00:24:54
in terms of trying to predict what's
00:24:56
going on here if we draw a graph like
00:24:58
this one way to remember it is that
00:25:01
whatever is on the right is the species
00:25:03
being reduced so in this cell we would
00:25:05
expect the silver cation to be reduced
00:25:07
while the species on the left we would
00:25:10
expect copper metal to be oxidized to
00:25:13
copper two-plus and this gives us a
00:25:16
simple way of picturing what's going on
00:25:19
if we then allow a current to flow we
00:25:22
can then apply an external voltage to
00:25:24
drive a reaction in a particular
00:25:25
direction that we wish so if we apply a
00:25:28
higher voltage remember this raises and
00:25:32
lowers the electrode potentials and
00:25:33
drives the reaction a different way to
00:25:37
summarize electrode potentials it's
00:25:39
always helpful to have a visual aid our
00:25:41
standard cell potential for any system
00:25:43
will never change it's measured under
00:25:45
standard conditions and fundamentally
00:25:48
free energy still govern all processes
00:25:50
only the free energy can be used to
00:25:52
predict the direction of spontaneity so
00:25:54
whenever we have our cell we would need
00:25:56
to formally convert to a free energy to
00:25:58
determine the direction of spontaneous
00:25:59
change and concentration has a big
00:26:02
effect on cell potentials so
00:26:04
everything's under standard conditions
00:26:05
but the minute we change that
00:26:06
concentration we get a different cell
00:26:08
potential these visual representations
00:26:11
can be really helpful to work out what's
00:26:13
going on because sometimes a quick
00:26:14
sketch can allow us to just
00:26:16
discombobulated the mathematics going on
