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Onto today's topic.

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Today we are going to be
focusing on issues associated

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with oxidation-reduction
half-reactions.

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And I want you think back to my
first lecture,

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where I really talked a lot
about how unifying the concepts

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were that stemmed from the ideas
of Gilbert Newton Lewis,

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because we talked about
acid-base, donor-acceptor,

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electrophile-neucleophile,
and oxidant-reductant all as

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parallel concepts,
with the related theme there

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being what is going on with the
electrons in the system.

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And it is really important to
consider not simply the

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reactions that occur between
molecules, but also the

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reactions that occur when
electrons coming from molecules

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or coming from ions or coming
from metals are connected up in

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some kind of external circuit
that can be interfaced then to

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some real world problems through
the mechanism of

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electrochemistry.
And so we will talk today about

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some of the basic concepts that
you will need.

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This will be the equipment that
you can use to do future

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research to solve the world's
energy problems.

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This is going to be good.
And it is pretty

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straightforward.
I want to give you an example

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of a reaction that can be
decomposed into half-reactions.

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And that is the reaction of
magnesium metal with carbon

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dioxide.
And this will be two

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equivalence of magnesium going
to two magnesium O two plus

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elemental carbon.

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Now, this actually is an
important reaction in that it

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contains one molecule carbon
dioxide, that is really a

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critical energy and environment
molecule that we worry a lot

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about today.
I will emphasize this in a few

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points during my lecture today,
but we want to be able to find

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ways of producing energy that
don't simultaneously produce a

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whole lot of CO two that
goes up into the atmosphere and

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is a greenhouse gas.
And so we will talk about that.

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I would like you to just start
thinking in the background of

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your mind right now about the
molecules, if you can list them,

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that are really important to
energy concerns today.

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With this reaction,
we can decompose this into the

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following half-reactions.
We can see that it can be

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written as carbon dioxide plus
four electrons plus four protons

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going to carbon plus two H two
O.

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And then simultaneously we can
also write two magnesium plus

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two H two O going to magnesium
two equivalence O plus four H

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plus plus four electrons.

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Although the
reaction that one may actually

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carry out would be the one on
top here, that reaction can be

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separated out into the parts
that are associated with

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oxidation and the parts that are
associated with reduction.

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And so, down here in equation
one, we see that CO two

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is undergoing reduction by four
electrons.

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This is a four electron
process.

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And down here,
in equation two,

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what we are seeing is that this
magnesium metal is ultimately

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serving as a source of four
electrons in that process,

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each magnesium going to
magnesium two plus.

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There is implied,
here, that we are going to be

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able to know something also
about oxidation states.

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And so, as we talk about
oxidation-reduction processes,

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I want you to also,
if necessary,

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review about what you know
about the assignment of

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oxidation states in systems.
But by writing this out,

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you can see that on the left
side of equation one,

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electrons are going in.
And on the right side of

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equation two electrons are
coming out.

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And, really,
if you can keep track of what

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the supply of the electrons are
and where the demand is for the

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electrons in the overall system,
you will really have a good

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feeling for how these processes
physically may be taking place.

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And then, up here on top,
this is the sum of those two

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equations, one plus two.
And that is typical of how we

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write oxidation-reduction
half-reactions.

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And when we consider a pair of
half-reactions,

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one of the things that we are
really interested in is the

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potential difference.

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And that is to say if we have
some way, physically,

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of sequestering the two
individual half-reactions into

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different containers,
and if we allow them to

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communicate, but we don't allow
the reaction to proceed to any

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significant extent,
we can measure this.

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We can measure it when the
reaction is not allowed to

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proceed to a significant extent.

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And the reason for that latter
caveat is that when you put into

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communication two sequestered
cells that are each poised,

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one to undergo a reduction and
one to undergo an oxidation,

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then if you put them into
contact and let that reaction

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start to proceed,
you can see that that reaction

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would proceed until such time as
equilibrium is reached.

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And so all throughout that
process, the potential

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difference between the two would
be changing until it stops

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changing.
And when it stops changing,

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we would have reached
equilibrium.

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The two independent
half-reactions at the beginning

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have their own potential,
the degree to which they are

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poised to undergo oxidation or
reduction.

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And then, how do we do this?
Well, you are going to see that

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we can use a thing called a salt
bridge.

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And with a salt bridge,
we can allow for not only

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communication of the electrons
from one of the half-reaction

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cells to the other,
but we are also going to need

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to allow for the movement of
ions.

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This permits ion movement.

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And this requirement for ion
movement stems from the

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necessity to maintain
electroneutrality,

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charge neutrality in each of
the two cells that we have,

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each of the two compartments of
the electrochemical cell that we

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are going to create.
And in reference to

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electrochemical cells,
--

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-- there are lots of different
types of electrochemical cells.

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And today we are going to be
discussing two of these in

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particular.
We are going to start out by

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talking about what is called a
Galvanic cell.

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And then, I am going to talk
about how you can use Galvanic

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cells to measure standard
properties of

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oxidation-reduction
half-reactions.

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And then, after we do that,
we are going to talk about the

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properties of redox-active
systems.

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And then, finally,
I will finish up by talking

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about another type of
electrochemical cell,

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and this will be an
electrolytic cell.

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So, let's look at a Galvanic
cell.

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Last week I left my blue chalk
in here.

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Gone.
Oh, well.

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Down one color.
If anyone can find that blue

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chalk for me,
I will be very appreciative.

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As I have been alluding to,
we can set up an

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electrochemical cell.
In this case,

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called the Galvanic cell,
in which we sequestered two

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different half-reactions to the
different compartments.

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And then we are going need for
a control of the communication

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between the two compartments.
And we will do that in the

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following way.
What I am representing,

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here, is a piece of metallic
zinc.

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And the way I have drawn it,
it is supposed to be like a

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simple strip of zinc metal,
to which we could attach some

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alligator clips and can run an
external piece of wire over to

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the other side.
And we can interpose,

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here, some kind of a meter.
This round unit,

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here, could be either a volt
meter or an ammeter to measure,

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respectively,
voltage or current,

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in a system like this.
And then, over on the other

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side, we will have a different
electrode.

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And this electrode will be made
of metallic copper.

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And I have got to have my salt
bridge so that we can maintain

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electroneutrality in these
solutions that we are going to

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have on both sides in the two
separate compartments.

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So, there is my salt bridge.
And this might contain an

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electrolyte, such as potassium
nitrate.

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And that electrolyte would be
suspended in a gelatinous medium

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like agar, for example.
And what we are going to be

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interested in will be the
directionality of the electron

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flow in a system like this.
We are going to be interested

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in the potential that gets set
up when the electrons flow in a

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system like this.
And we are going to want to

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know which side is the anode and
which side is the cathode,

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and what are the equations for
the reactions that are taking

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place at the side that is the
anode and at the side that is

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the cathode.
Let's look at that.

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Let me just point out that if I
had blue, I would be indicating

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the aqueous solution here in
blue.

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Maybe those of you who are
colorblind will think that is

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blue, but anyway.
The idea is that we have an

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aqueous solution here,
on both sides.

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And this one,
over on the right,

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is going to start off with some
concentration of copper two

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ions --

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-- in solution.
So, some molarity of copper two

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will be present in
the solution over here on the

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right-hand side,
where we have the copper

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electrode.
Inside the salt bridge,

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we have potassium cations and
nitrate anions that can go into

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solution on either side to
balance the charged changes that

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are taking place as
oxidation-reduction reactions

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happen on the left-hand side and
on the right-hand side of this

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electrochemical cell,
which is a Galvanic cell.

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And so, what we find is that
you let this go into contact

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briefly.
And you would measure,

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for this cell,
1.1 volts.

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And so, that is the magnitude
of the potential difference

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between the two half-reactions
that are present on both sides

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of the cell.
And, furthermore,

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we would see that the electrons
are starting out over here,

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and they are going this way
around the circuit.

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And that tells us that what we
have on the left electrode at

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zinc is going to be called our
anode.

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That is because at the anode,
--

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-- oxidation is taking place
because metallic zinc is

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undergoing oxidation and
becoming zinc two plus

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and providing,
in so doing,

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two electrons for every zinc
that gets oxidized.

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What you can think of as
happening over here on the side

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that we call the anode is that
you have this piece of zinc

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metal.
And, in order for electrons to

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start coming out to the external
circuit, for every two electrons

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that comes around to the outside
of the circuit,

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you have a single zinc two plus
ion that jumps into

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solution.
A piece of zinc jumps off of

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the surface of the electrode.
That zinc atom leaves as the

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zinc two plus ion.
That supplies two electrons to

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the external circuit.
And, at the same time,

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to balance charge,
we would have to have,

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from the salt bridge entering
solution, two NO three minus,

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two nitrate anions,

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because they are singly
charged, have to come out of the

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salt bridge and into solution
every time a zinc two plus

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00:18:00 --> 00:18:06
drops away from the
electrode and starts flowing

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into the external circuit.
And electrons go around the

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initial voltage when the
reaction has not proceeded to

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any significant extent,
is this 1.1 volts.

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00:18:18 --> 00:18:22
So, we have measured the
potential difference between the

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00:18:22 --> 00:18:24
two half-reactions.
Here is one of these

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half-reactions.
At the other side,

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if on the left side we have an
electrode that is the anode,

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then on the right-hand side we
must have an electrode that is

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the cathode.

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And we must have reduction
occurring there.

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And what is getting reduced?
Well, as electrons appear down

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here at this electrode,
they can encounter a copper two

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plus ion in solution
that can then become part of the

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surface of the electrode as a
copper metal atom.

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In other words,
we are depleting the surface of

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this piece of metallic zinc over
here.

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And zinc ions are going into
solution.

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And, on the right-hand side,
copper ions are coming from

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solution and becoming
incorporated into the electrode

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itself.
And so, if you ran this

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00:19:27 --> 00:19:31
reaction for a while,
and then you took the two

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pieces of metal,
zinc and copper,

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00:19:33 --> 00:19:37
that you had used to set up the
Galvanic cell,

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you could weigh them.
And you could see that the zinc

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00:19:41 --> 00:19:44
electrode would have gotten
lighter, and the copper

249
00:19:44 --> 00:19:47
electrode would have gotten
heavier.

250
00:19:47 --> 00:19:51
And so, that is also the basis
for the technique of

251
00:19:51 --> 00:19:54
electroplating,
where you can take one metal

252
00:19:54 --> 00:19:59
and cover over the surface of it
with a thin layer of another

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00:19:59 --> 00:20:03
metal.
Here we are just putting a new

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00:20:03 --> 00:20:07
layer of copper on the surface
of this copper electrode,

255
00:20:07 --> 00:20:11
and we are depleting what we
call, over on the left,

256
00:20:11 --> 00:20:15
the sacrificial zinc anode.
And this reaction then,

257
00:20:15 --> 00:20:19
over on the right-hand side,
is copper two plus plus two

258
00:20:19 --> 00:20:24
electrons going to copper metal.

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00:20:24 --> 00:20:30
So, that is a prototypical
example of a Galvanic cell.

260
00:20:30 --> 00:20:36
And in setting that up and
thinking about what is happening

261
00:20:36 --> 00:20:42
there, in terms of depletion of
one electrode so that the other

262
00:20:42 --> 00:20:48
one can be increased in mass,
you may be aware that you know

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00:20:48 --> 00:20:53
of places where this principle
is actually used.

264
00:20:53 --> 00:20:56
In construction,
for example,

265
00:20:56 --> 00:21:03
there are large steel bridges
that are constructed.

266
00:21:03 --> 00:21:07
And one does not want those to
oxidize away and to become

267
00:21:07 --> 00:21:11
fragile in their structure,
so that bridges break down and

268
00:21:11 --> 00:21:15
bridges are not safe anymore,
shall we say.

269
00:21:15 --> 00:21:19
And so, what is often done
there is large pieces of zinc,

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00:21:19 --> 00:21:23
actually, are put into
electrical contact with the

271
00:21:23 --> 00:21:27
metal of the bridge so that it
is the zinc that actually gets

272
00:21:27 --> 00:21:32
oxidized away as a sacrificial
anode, instead of the bridge

273
00:21:32 --> 00:21:36
itself.
The same for any of you who are

274
00:21:36 --> 00:21:40
sailors or boaters in general.
You probably know that zinc

275
00:21:40 --> 00:21:45
anodes are present usually in
several places on boats so that

276
00:21:45 --> 00:21:49
electrolysis does not occur and
corrode away the metal parts of

277
00:21:49 --> 00:21:52
your boat.
Instead, the zinc sacrificial

278
00:21:52 --> 00:21:56
anode is oxidized away and is a
source of electrons.

279
00:21:56 --> 00:22:00
So, that is an important
principle.

280
00:22:00 --> 00:22:10


281
00:22:10 --> 00:22:13
Once you have started looking
at some different kinds of

282
00:22:13 --> 00:22:16
oxidation-reduction
half-reactions,

283
00:22:16 --> 00:22:20
you are going to want to know
just how oxidizing is something

284
00:22:20 --> 00:22:22
or just how reducing is
something.

285
00:22:22 --> 00:22:26
And, in order to do this,
we need to pick one particular

286
00:22:26 --> 00:22:31
half-reaction that will serve as
our universal standard.

287
00:22:31 --> 00:22:34
And, to that half-reaction,
we will compare everything

288
00:22:34 --> 00:22:37
else.
And so, that brings us to the

289
00:22:37 --> 00:22:40
discussion of standard reduction
potentials, or standard

290
00:22:40 --> 00:22:42
potentials.

291
00:22:42 --> 00:22:57


292
00:22:57 --> 00:23:00
For this, I am going to make
another Galvanic cell.

293
00:23:00 --> 00:23:20


294
00:23:20 --> 00:23:23
And, concerning this Galvanic
cell, we will have a number of

295
00:23:23 --> 00:23:26
the same questions that we have
had before.

296
00:23:26 --> 00:23:30
But the one that we are going
to be using on the right-hand

297
00:23:30 --> 00:23:33
side is going to be very
special.

298
00:23:33 --> 00:23:39
Here we indicate our water
line, again, in green.

299
00:23:39 --> 00:23:47
And I am going to put in a salt
bridge into the system.

300
00:23:47 --> 00:23:56
And I will have an electrode,
here, on the left and another

301
00:23:56 --> 00:24:02
electrode over here on the
right.

302
00:24:02 --> 00:24:07


303
00:24:07 --> 00:24:10
We will connect up these
electrode using alligator clips.

304
00:24:10 --> 00:24:14
We will need an external
circuit with a meter,

305
00:24:14 --> 00:24:17
so we will be able to compare
the potential of the two

306
00:24:17 --> 00:24:21
half-reactions of interest.
But then, over here on the

307
00:24:21 --> 00:24:25
right, this electrode is going
to be a little different.

308
00:24:25 --> 00:24:29
Because this one is going to
have something like a test-tube

309
00:24:29 --> 00:24:35
inverted over it --
-- because the reaction that we

310
00:24:35 --> 00:24:39
will talk about here,
in fact, does involve a gas.

311
00:24:39 --> 00:24:45
And that gas is hydrogen,
another one of our important

312
00:24:45 --> 00:24:51
energy molecules to be talking
about in the context of today's

313
00:24:51 --> 00:24:55
lecture.
At the end, I would like to see

314
00:24:55 --> 00:25:01
if you have a list of energy
molecules collected from today's

315
00:25:01 --> 00:25:05
lecture.
But also, we are going to

316
00:25:05 --> 00:25:08
choose, as the electrode
material, here,

317
00:25:08 --> 00:25:11
this metal piece of electrode
material.

318
00:25:11 --> 00:25:17
This is going to be platinum.
And the reason we are going to

319
00:25:17 --> 00:25:20
choose platinum,
a very common electrode

320
00:25:20 --> 00:25:26
material, is that platinum is a
so-called noble metal.

321
00:25:26 --> 00:25:31


322
00:25:31 --> 00:25:35
And that means that it is quite
impervious to most chemical

323
00:25:35 --> 00:25:37
reactions.
Over there, we were talking

324
00:25:37 --> 00:25:41
about zinc and copper electrodes
that themselves do change during

325
00:25:41 --> 00:25:43
the reaction.
And over here,

326
00:25:43 --> 00:25:47
we want to have a piece of
metal that can provide the

327
00:25:47 --> 00:25:51
valuable function of giving us
electrical contact of the

328
00:25:51 --> 00:25:55
chemical reactions on the two
sides, but an electrode that

329
00:25:55 --> 00:25:59
will remain clean and unchanged
at its surface throughout the

330
00:25:59 --> 00:26:04
course of the reaction.
So, we use a platinum electrode

331
00:26:04 --> 00:26:07
over here.
And you can see that with this

332
00:26:07 --> 00:26:12
inverted test tube through which
the wire passes and with some

333
00:26:12 --> 00:26:16
kind of a provision for a side
arm attachment here,

334
00:26:16 --> 00:26:20
we can hook up a cylinder of H
two gas.

335
00:26:20 --> 00:26:25
And we can start bubbling H two
over this electrode,

336
00:26:25 --> 00:26:30
like this, so that it actually
will have bubbles of H two

337
00:26:30 --> 00:26:35
coming out like that.
In fact, this solution over

338
00:26:35 --> 00:26:40
here will be saturated with H
two.

339
00:26:40 --> 00:26:46
And a further consideration for
our reference electrode,

340
00:26:46 --> 00:26:52
which is going to be known as
the standard hydrogen electrode,

341
00:26:52 --> 00:26:54
or SHE.

342
00:26:54 --> 00:27:06


343
00:27:06 --> 00:27:12
That is, the electrode against
which everything else is going

344
00:27:12 --> 00:27:15
to be compared.
And it also has,

345
00:27:15 --> 00:27:20
in the solution here,
1.0 molar H three O plus.

346
00:27:20 --> 00:27:23
So, indeed, it is a very

347
00:27:23 --> 00:27:29
strongly acidic medium.
And what we are measuring,

348
00:27:29 --> 00:27:35
over on the right-hand side,
is a half reaction that may

349
00:27:35 --> 00:27:41
correspond either to two H plus
plus two electrons going to

350
00:27:41 --> 00:27:44
H two.

351
00:27:44 --> 00:27:49
And that will be the case if
this side is the cathode.

352
00:27:49 --> 00:27:53
But if this side turns out to
be the anode,

353
00:27:53 --> 00:28:00
then we would be measuring the
opposite reaction.

354
00:28:00 --> 00:28:03
And I will talk more about that
in a second.

355
00:28:03 --> 00:28:08
Let's say, for example,
that our electrode over here

356
00:28:08 --> 00:28:11
is, in fact, made of zinc.

357
00:28:11 --> 00:28:19


358
00:28:19 --> 00:28:22
If our electrode over here is
made of zinc.

359
00:28:22 --> 00:28:27
And we may have zinc two plus
in solution,

360
00:28:27 --> 00:28:32
here, and we may again have
something like potassium nitrate

361
00:28:32 --> 00:28:37
or some other
electrolyte present in our salt

362
00:28:37 --> 00:28:41
bridge.
Then what we can do is put

363
00:28:41 --> 00:28:43
these things briefly into
contact.

364
00:28:43 --> 00:28:48
And we want to know two things.
We want to know what the

365
00:28:48 --> 00:28:52
magnitude of the potential
difference is and what the

366
00:28:52 --> 00:28:57
direction in which the electrons
are flowing is so that we know

367
00:28:57 --> 00:29:02
which one is the cathode and
which one is the anode.

368
00:29:02 --> 00:29:07
And, in this particular case,
the direction of electron flow

369
00:29:07 --> 00:29:10
is, as before,
away from the zinc.

370
00:29:10 --> 00:29:16
Once again, zinc atoms are
jumping off the surface of the

371
00:29:16 --> 00:29:20
electrode as zinc two plus
ions.

372
00:29:20 --> 00:29:25
And, every time that happens,
two electrons go into the

373
00:29:25 --> 00:29:30
external circuit and come around
here.

374
00:29:30 --> 00:29:34
And the chemical reaction that
occurs here, in this case,

375
00:29:34 --> 00:29:38
is that those two electrons
that came from one of the zinc

376
00:29:38 --> 00:29:43
atoms react, as shown here,
with two H plus ions to make

377
00:29:43 --> 00:29:45
H two.

378
00:29:45 --> 00:29:49
And so, we know the direction of
electron flow,

379
00:29:49 --> 00:29:53
we know the exact reactions
that are taking place on the

380
00:29:53 --> 00:29:57
left and on the right.
And we know that,

381
00:29:57 --> 00:30:02
for this particular choice of
substances, our anode is on the

382
00:30:02 --> 00:30:06
left, and our cathode is on the
right.

383
00:30:06 --> 00:30:09
And let me come over here.

384
00:30:09 --> 00:30:17


385
00:30:17 --> 00:30:21
There is another way that I can
write this cell.

386
00:30:21 --> 00:30:27
And so, I would like to
introduce cell notation to you.

387
00:30:27 --> 00:30:32
And that will be like this.
We will have zinc.

388
00:30:32 --> 00:30:41
And then, a solid line.
And then, zinc two plus.

389
00:30:41 --> 00:30:48
And then, a solid double
vertical line.

390
00:30:48 --> 00:30:57
And then, H plus,
another solid vertical line,

391
00:30:57 --> 00:31:02
H two.
And then another solid line.

392
00:31:02 --> 00:31:07
And then platinum.

393
00:31:07 --> 00:31:10
That is the way I set that up.

394
00:31:10 --> 00:31:15
We could, actually,
add another solid line here and

395
00:31:15 --> 00:31:20
put platinum over here.
The thing that I want you to

396
00:31:20 --> 00:31:26
recognize about this notation
for Galvanic cells is that the

397
00:31:26 --> 00:31:31
solid single lines represent
interfaces, direct contacts

398
00:31:31 --> 00:31:36
between things.
Here, it would be a platinum

399
00:31:36 --> 00:31:39
solid electrode
connected with solid zinc,

400
00:31:39 --> 00:31:43
possibly.
And then, the way I drew it

401
00:31:43 --> 00:31:47
over there, it would just be
solid zinc, no platinum on the

402
00:31:47 --> 00:31:49
left.
And then this solid line

403
00:31:49 --> 00:31:53
represents a solid liquid
interface because the electrode

404
00:31:53 --> 00:31:57
is dipped into a solution that
contains zinc two

405
00:31:57 --> 00:32:00
ions.
And that is in communication

406
00:32:00 --> 00:32:04
with the other half of this
electrochemical cell by a salt

407
00:32:04 --> 00:32:08
bridge.
And the salt bridge is

408
00:32:08 --> 00:32:11
represented by the double
vertical line.

409
00:32:11 --> 00:32:14
And then we have,
in solution,

410
00:32:14 --> 00:32:17
protons.
And then, a solid liquid.

411
00:32:17 --> 00:32:20
Well, actually,
in this case,

412
00:32:20 --> 00:32:25
a liquid gas interface to the
gaseous H two.

413
00:32:25 --> 00:32:30
And H two is also
dissolved.

414
00:32:30 --> 00:32:34
In order for this to be used as
a reference electrode,

415
00:32:34 --> 00:32:38
we have to pick standard
conditions.

416
00:32:38 --> 00:32:43
And so, we are usually talking
about one atmosphere of hydrogen

417
00:32:43 --> 00:32:47
for the standard hydrogen
electrode.

418
00:32:47 --> 00:32:52
And we are talking about a
concentration of 1.0 molar of

419
00:32:52 --> 00:32:56
our strong acid.
And so, those are our standard

420
00:32:56 --> 00:33:01
conditions, along with 25
degrees C.

421
00:33:01 --> 00:33:05
And then this solution that
contains the protons and the

422
00:33:05 --> 00:33:10
hydrogen gas is in a solution
solid interface with the solid

423
00:33:10 --> 00:33:14
platinum electrode that is not
going to be undergoing any

424
00:33:14 --> 00:33:17
change.
So, these are the types of cell

425
00:33:17 --> 00:33:21
notations that you will
encounter when looking into

426
00:33:21 --> 00:33:25
electrochemistry.
And we found out that the left

427
00:33:25 --> 00:33:30
side, where the zinc metal
is going to zinc two

428
00:33:30 --> 00:33:34
plus,
is our anode.

429
00:33:34 --> 00:33:39
And also, the potential
difference, when we measure it

430
00:33:39 --> 00:33:45
using our voltmeter right here
by turning on the contact very

431
00:33:45 --> 00:33:51
briefly, is around 0.763 volts.
So, that is the magnitude of

432
00:33:51 --> 00:33:56
our potential difference.
And this is equal to delta E

433
00:33:56 --> 00:34:02
zero cell.
When I write E zero,

434
00:34:02 --> 00:34:08
that means standard potential.
And this is written as a delta

435
00:34:08 --> 00:34:12
here because that is the
difference between the

436
00:34:12 --> 00:34:17
potentials for the two
half-reactions that we have

437
00:34:17 --> 00:34:20
written up there.
But it is a very simple

438
00:34:20 --> 00:34:26
difference precisely because the
standard hydrogen electrode is

439
00:34:26 --> 00:34:32
our reference electrode.
Over here we want to write the

440
00:34:32 --> 00:34:36
definition of standard cell
potential.

441
00:34:36 --> 00:35:07


442
00:35:07 --> 00:35:14
And that is a delta E zero for
your cell is equal to delta E

443
00:35:14 --> 00:35:21
zero for your cathode minus--
Sorry, this is not delta here.

444
00:35:21 --> 00:35:28
Just E zero for your cathode
minus E zero for your anode.

445
00:35:28 --> 00:35:35


446
00:35:35 --> 00:35:40


447
00:35:40 --> 00:35:44
But E zero for the
standard hydrogen electrode,

448
00:35:44 --> 00:35:49
because this is our reference
electrode, is equal to zero at

449
00:35:49 --> 00:35:51
all temperatures.

450
00:35:51 --> 00:35:56


451
00:35:56 --> 00:36:02
And that is by definition.
So, the standard potential for

452
00:36:02 --> 00:36:08
that set of conditions that
constitutes our standard

453
00:36:08 --> 00:36:14
hydrogen electrode is taken as
the zero of potential for

454
00:36:14 --> 00:36:21
comparison with any other type
of half cell that you might be

455
00:36:21 --> 00:36:25
able to consider.
And so, we can see further

456
00:36:25 --> 00:36:32
that, in this particular case,
where we have our cathode as

457
00:36:32 --> 00:36:39
the standard hydrogen electrode,
we have zero and minus E zero

458
00:36:39 --> 00:36:45
for the anode,
which we measured as 0

459
00:36:45 --> 00:36:48
volts.

460
00:36:48 --> 00:36:53


461
00:36:53 --> 00:36:57
That means that for the
reaction zinc going to zinc two

462
00:36:57 --> 00:37:01
plus plus two electrons,

463
00:37:01 --> 00:37:06
which is our anodic reaction,
we have a standard potential of

464
00:37:06 --> 00:37:09
-0.763 volts.
And generally,

465
00:37:09 --> 00:37:12
what you are going to be
interested in,

466
00:37:12 --> 00:37:17
as you consider different kinds
of substances from throughout

467
00:37:17 --> 00:37:21
the periodic table with
reference to their ability to

468
00:37:21 --> 00:37:25
take place in
oxidation-reduction reactions,

469
00:37:25 --> 00:37:29
is you are going to want to
know what your standard

470
00:37:29 --> 00:37:33
potential is.
You are not going to always

471
00:37:33 --> 00:37:37
have a system that you want to
consider that is under standard

472
00:37:37 --> 00:37:40
conditions.
These standard potentials

473
00:37:40 --> 00:37:43
always are referenced to some
standard conditions,

474
00:37:43 --> 00:37:47
as I mentioned specifically for
the standard hydrogen electrode.

475
00:37:47 --> 00:37:51
And so, part of next lecture on
Monday is we are going to show

476
00:37:51 --> 00:37:54
you how to handle systems that
are not under standard

477
00:37:54 --> 00:37:58
conditions because then you can
handle some real practical

478
00:37:58 --> 00:38:02
problems.
But let's look at a different

479
00:38:02 --> 00:38:03
type of cell.

480
00:38:03 --> 00:38:10


481
00:38:10 --> 00:38:13
Let's write down the following
cell.

482
00:38:13 --> 00:38:40


483
00:38:40 --> 00:38:43
This new cell that I have
written the cell notation for,

484
00:38:43 --> 00:38:48
instead of drawing up the whole
diagram of this Galvanic cell,

485
00:38:48 --> 00:38:52
happens to be one in which we
have platinum electrodes on both

486
00:38:52 --> 00:38:54
sides.
Here we have the silver |

487
00:38:54 --> 00:38:58
silver plus
redox couple on the left-hand

488
00:38:58 --> 00:39:02
side.
And we have the redox couple of

489
00:39:02 --> 00:39:06
H plus with H two
on the right-hand side.

490
00:39:06 --> 00:39:09
So, the right-hand side
corresponds to the standard

491
00:39:09 --> 00:39:13
hydrogen electrode,
and the left-hand side is a new

492
00:39:13 --> 00:39:16
redox couple,
or a new half-reaction,

493
00:39:16 --> 00:39:20
that we want to compare to the
standard hydrogen electrode so

494
00:39:20 --> 00:39:23
that we will know,
on an absolute scale,

495
00:39:23 --> 00:39:27
where it falls relative to all
the other half-reactions we

496
00:39:27 --> 00:39:31
might want to measure.
And so, we want to know,

497
00:39:31 --> 00:39:34
which way do the electrons
flow?

498
00:39:34 --> 00:39:39


499
00:39:39 --> 00:39:43
And, in this case,
it turns out that the direction

500
00:39:43 --> 00:39:48
of electron flow is this way.
So, this is the opposite of

501
00:39:48 --> 00:39:53
what we talked about in the case
of comparing zinc to the

502
00:39:53 --> 00:39:57
standard hydrogen electrode.
Electron flow has been

503
00:39:57 --> 00:40:00
reversed.
What that means is that this

504
00:40:00 --> 00:40:05
over here, the SHE is now our
anode.

505
00:40:05 --> 00:40:11
And now the silver-silver plus
electrode is our

506
00:40:11 --> 00:40:14
cathode.
And we want to know not only

507
00:40:14 --> 00:40:20
the direction of electron flow,
the direction of electron flow

508
00:40:20 --> 00:40:26
tells us what is doing the
reduction and what is doing the

509
00:40:26 --> 00:40:32
oxidation, but we also want to
know the magnitude.

510
00:40:32 --> 00:40:37
And this one turns out to be,
this delta E cell,

511
00:40:37 --> 00:40:43
is equal to 0.8 volts.
And, by the definition of

512
00:40:43 --> 00:40:49
standard cell potentials that I
gave you over there,

513
00:40:49 --> 00:40:55
you can see that what we are
getting now, because the

514
00:40:55 --> 00:41:00
electron flow is reversed,
our sign is reversed,

515
00:41:00 --> 00:41:06
and so our E zero for the
reaction Ag plus plus an

516
00:41:06 --> 00:41:12
electron going to silver is

517
00:41:12 --> 00:41:18
equal to 0.8 volts,
positive.

518
00:41:18 --> 00:41:22
Notice that the zinc-zinc two
plus couple was

519
00:41:22 --> 00:41:25
negative with respect to the
standard hydrogen electrode,

520
00:41:25 --> 00:41:29
but because the electron flow
is reversed for silver plus,

521
00:41:29 --> 00:41:32
silver redox couple,
we now have a positive

522
00:41:32 --> 00:41:37
potential relative to the
standard hydrogen electrode.

523
00:41:37 --> 00:41:45


524
00:41:45 --> 00:41:51
What would like to arrive at is
a nice big table where we look

525
00:41:51 --> 00:41:54
at standard potential --

526
00:41:54 --> 00:42:01


527
00:42:01 --> 00:42:06
-- with reference to this
standard hydrogen electrode,

528
00:42:06 --> 00:42:12
which is our zero of potential.
And you will be able to find a

529
00:42:12 --> 00:42:17
table like this in your book.
What we found is that zinc is

530
00:42:17 --> 00:42:22
down here at -0.763,
so that was zinc-zinc two plus

531
00:42:22 --> 00:42:28
at a negative
potential relative to hydrogen

532
00:42:28 --> 00:42:33
plus electrons.
We found that up here at +0.8,

533
00:42:33 --> 00:42:39
we have the silver-silver plus
redox couple.

534
00:42:39 --> 00:42:44
These are thermodynamic
quantities, so you can look on a

535
00:42:44 --> 00:42:49
table of standard reduction
potentials and you can tell

536
00:42:49 --> 00:42:54
which direction electrons will
flow if you set up cells that

537
00:42:54 --> 00:42:59
involve those redox
half-reactions.

538
00:42:59 --> 00:43:04
Another electrode that we used
today was copper two plus

539
00:43:04 --> 00:43:09
combining with two electrons to
give copper metal.

540
00:43:09 --> 00:43:13
It turns out that one is

541
00:43:13 --> 00:43:17
positive also by about positive
0.3 volts.

542
00:43:17 --> 00:43:23
You see hydrogen is here and it
reduces silver-silver plus.

543
00:43:23 --> 00:43:27
Hydrogen reduces silver plus to

544
00:43:27 --> 00:43:32
silver because it is up there.
Zinc is down here.

545
00:43:32 --> 00:43:37
Zinc also serves as an anode
with respect to any one of those

546
00:43:37 --> 00:43:41
three because all of those three
are at a potential positive

547
00:43:41 --> 00:43:45
relative to zinc on this scale
of standard reduction

548
00:43:45 --> 00:43:49
potentials.
You can also have some other

549
00:43:49 --> 00:43:53
potentials, like way down here
at about -2.7 volts is the

550
00:43:53 --> 00:43:57
sodium-sodium plus
redox couple,

551
00:43:57 --> 00:44:02
way down there at -2.7
That is why sodium is so much

552
00:44:02 --> 00:44:05
fun to heave into a body of
water.

553
00:44:05 --> 00:44:09
I mean, it is fantastic.
You get reduction of the

554
00:44:09 --> 00:44:13
protons to make hydrogen,
which then explodes.

555
00:44:13 --> 00:44:17
So, this is fantastic.
And some of you may know that

556
00:44:17 --> 00:44:21
there is an annual,
and I am not recommending that

557
00:44:21 --> 00:44:25
you do this, by the way.
You can see how negative the

558
00:44:25 --> 00:44:32
potential is down here.
Metallic sodium is a very

559
00:44:32 --> 00:44:35
strong reducing agent,
indeed.

560
00:44:35 --> 00:44:42
A very important reaction is up
here at about +1.23 volts

561
00:44:42 --> 00:44:48
relative to the standard
hydrogen electrode.

562
00:44:48 --> 00:44:56
And this important reaction is
oxygen plus four H pus plus four

563
00:44:56 --> 00:45:03
electrons going to two H two O.

564
00:45:03 --> 00:45:06


565
00:45:06 --> 00:45:11


566
00:45:11 --> 00:45:14
So, this +1.23 volts is very
important.

567
00:45:14 --> 00:45:18
In the time remaining,
I am not going to be able to

568
00:45:18 --> 00:45:23
tell you about electrolysis.
I believe next hour I will

569
00:45:23 --> 00:45:27
start off by talking about
electrolysis.

570
00:45:27 --> 00:45:31
And, if you can do electrolysis
using oxidizing and reducing

571
00:45:31 --> 00:45:35
equivalents that derive from
photovoltaic cells,

572
00:45:35 --> 00:45:40
so you are converting sunlight
into separated electron whole

573
00:45:40 --> 00:45:45
pairs, you can drive a reaction
like this the other way and

574
00:45:45 --> 00:45:50
learn how to make oxygen from
water, and ultimately also

575
00:45:50 --> 00:45:54
hydrogen from water.
We will talk about that a

576
00:45:54.272 --> 45:57
little bit next time.
Have a nice weekend.