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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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00:12:03,000 --> 00:12:08,000
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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00:15:01,000 --> 00:15:07,000
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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00:15:21,000 --> 00:15:25,000
electrochemical cell,
which is a Galvanic cell.

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00:15:25,000 --> 00:15:30,000
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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00:15:34,000 --> 00:15:36,000
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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00:16:39,000 --> 00:16:43,000
and providing,
in so doing,

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two electrons for every zinc
that gets oxidized.

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00:16:48,000 --> 00:16:54,000
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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00:17:00,000 --> 00:17:04,000
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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00:17:10,000 --> 00:17:14,000
that comes around to the outside
of the circuit,

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00:17:14,000 --> 00:17:20,000
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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00:17:28,000 --> 00:17:34,000
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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00:17:42,000 --> 00:17:47,000
from the salt bridge entering
solution, two NO three minus,

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00:17:47,000 --> 00:17:50,000
two nitrate anions,

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00:17:50,000 --> 00:17:55,000
because they are singly
charged, have to come out of the

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00:17:55,000 --> 00:18:00,000
salt bridge and into solution
every time a zinc two plus

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00:18:00,000 --> 00:18:06,000
drops away from the
electrode and starts flowing

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00:18:06,000 --> 00:18:11,000
into the external circuit.
And electrons go around the

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00:18:11,000 --> 00:18:15,000
initial voltage when the
reaction has not proceeded to

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00:18:15,000 --> 00:18:18,000
any significant extent,
is this 1.1 volts.

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00:18:18,000 --> 00:18:22,000
So, we have measured the
potential difference between the

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00:18:22,000 --> 00:18:24,000
two half-reactions.
Here is one of these

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00:18:24,000 --> 00:18:27,000
half-reactions.
At the other side,

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00:18:27,000 --> 00:18:31,000
if on the left side we have an
electrode that is the anode,

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00:18:31,000 --> 00:18:35,000
then on the right-hand side we
must have an electrode that is

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the cathode.

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00:18:44,000 --> 00:18:47,000
And we must have reduction
occurring there.

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00:18:47,000 --> 00:18:52,000
And what is getting reduced?
Well, as electrons appear down

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00:18:52,000 --> 00:18:57,000
here at this electrode,
they can encounter a copper two

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00:18:57,000 --> 00:19:03,000
plus ion in solution
that can then become part of the

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00:19:03,000 --> 00:19:08,000
surface of the electrode as a
copper metal atom.

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00:19:08,000 --> 00:19:11,000
In other words,
we are depleting the surface of

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00:19:11,000 --> 00:19:14,000
this piece of metallic zinc over
here.

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00:19:14,000 --> 00:19:16,000
And zinc ions are going into
solution.

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00:19:16,000 --> 00:19:20,000
And, on the right-hand side,
copper ions are coming from

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00:19:20,000 --> 00:19:24,000
solution and becoming
incorporated into the electrode

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00:19:24,000 --> 00:19:27,000
itself.
And so, if you ran this

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00:19:27,000 --> 00:19:31,000
reaction for a while,
and then you took the two

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00:19:31,000 --> 00:19:33,000
pieces of metal,
zinc and copper,

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00:19:33,000 --> 00:19:37,000
that you had used to set up the
Galvanic cell,

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00:19:37,000 --> 00:19:41,000
you could weigh them.
And you could see that the zinc

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00:19:41,000 --> 00:19:44,000
electrode would have gotten
lighter, and the copper

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00:19:44,000 --> 00:19:47,000
electrode would have gotten
heavier.

237
00:19:47,000 --> 00:19:51,000
And so, that is also the basis
for the technique of

238
00:19:51,000 --> 00:19:54,000
electroplating,
where you can take one metal

239
00:19:54,000 --> 00:19:59,000
and cover over the surface of it
with a thin layer of another

240
00:19:59,000 --> 00:20:03,000
metal.
Here we are just putting a new

241
00:20:03,000 --> 00:20:07,000
layer of copper on the surface
of this copper electrode,

242
00:20:07,000 --> 00:20:11,000
and we are depleting what we
call, over on the left,

243
00:20:11,000 --> 00:20:15,000
the sacrificial zinc anode.
And this reaction then,

244
00:20:15,000 --> 00:20:19,000
over on the right-hand side,
is copper two plus plus two

245
00:20:19,000 --> 00:20:24,000
electrons going to copper metal.

246
00:20:24,000 --> 00:20:30,000
So, that is a prototypical
example of a Galvanic cell.

247
00:20:30,000 --> 00:20:36,000
And in setting that up and
thinking about what is happening

248
00:20:36,000 --> 00:20:42,000
there, in terms of depletion of
one electrode so that the other

249
00:20:42,000 --> 00:20:48,000
one can be increased in mass,
you may be aware that you know

250
00:20:48,000 --> 00:20:53,000
of places where this principle
is actually used.

251
00:20:53,000 --> 00:20:56,000
In construction,
for example,

252
00:20:56,000 --> 00:21:03,000
there are large steel bridges
that are constructed.

253
00:21:03,000 --> 00:21:07,000
And one does not want those to
oxidize away and to become

254
00:21:07,000 --> 00:21:11,000
fragile in their structure,
so that bridges break down and

255
00:21:11,000 --> 00:21:15,000
bridges are not safe anymore,
shall we say.

256
00:21:15,000 --> 00:21:19,000
And so, what is often done
there is large pieces of zinc,

257
00:21:19,000 --> 00:21:23,000
actually, are put into
electrical contact with the

258
00:21:23,000 --> 00:21:27,000
metal of the bridge so that it
is the zinc that actually gets

259
00:21:27,000 --> 00:21:32,000
oxidized away as a sacrificial
anode, instead of the bridge

260
00:21:32,000 --> 00:21:36,000
itself.
The same for any of you who are

261
00:21:36,000 --> 00:21:40,000
sailors or boaters in general.
You probably know that zinc

262
00:21:40,000 --> 00:21:45,000
anodes are present usually in
several places on boats so that

263
00:21:45,000 --> 00:21:49,000
electrolysis does not occur and
corrode away the metal parts of

264
00:21:49,000 --> 00:21:52,000
your boat.
Instead, the zinc sacrificial

265
00:21:52,000 --> 00:21:56,000
anode is oxidized away and is a
source of electrons.

266
00:21:56,000 --> 00:22:00,000
So, that is an important
principle.

267
00:22:10,000 --> 00:22:13,000
Once you have started looking
at some different kinds of

268
00:22:13,000 --> 00:22:16,000
oxidation-reduction
half-reactions,

269
00:22:16,000 --> 00:22:20,000
you are going to want to know
just how oxidizing is something

270
00:22:20,000 --> 00:22:22,000
or just how reducing is
something.

271
00:22:22,000 --> 00:22:26,000
And, in order to do this,
we need to pick one particular

272
00:22:26,000 --> 00:22:31,000
half-reaction that will serve as
our universal standard.

273
00:22:31,000 --> 00:22:34,000
And, to that half-reaction,
we will compare everything

274
00:22:34,000 --> 00:22:37,000
else.
And so, that brings us to the

275
00:22:37,000 --> 00:22:40,000
discussion of standard reduction
potentials, or standard

276
00:22:40,000 --> 00:22:42,000
potentials.

277
00:22:57,000 --> 00:23:00,000
For this, I am going to make
another Galvanic cell.

278
00:23:20,000 --> 00:23:23,000
And, concerning this Galvanic
cell, we will have a number of

279
00:23:23,000 --> 00:23:26,000
the same questions that we have
had before.

280
00:23:26,000 --> 00:23:30,000
But the one that we are going
to be using on the right-hand

281
00:23:30,000 --> 00:23:33,000
side is going to be very
special.

282
00:23:33,000 --> 00:23:39,000
Here we indicate our water
line, again, in green.

283
00:23:39,000 --> 00:23:47,000
And I am going to put in a salt
bridge into the system.

284
00:23:47,000 --> 00:23:56,000
And I will have an electrode,
here, on the left and another

285
00:23:56,000 --> 00:24:02,000
electrode over here on the
right.

286
00:24:07,000 --> 00:24:10,000
We will connect up these
electrode using alligator clips.

287
00:24:10,000 --> 00:24:14,000
We will need an external
circuit with a meter,

288
00:24:14,000 --> 00:24:17,000
so we will be able to compare
the potential of the two

289
00:24:17,000 --> 00:24:21,000
half-reactions of interest.
But then, over here on the

290
00:24:21,000 --> 00:24:25,000
right, this electrode is going
to be a little different.

291
00:24:25,000 --> 00:24:29,000
Because this one is going to
have something like a test-tube

292
00:24:29,000 --> 00:24:35,000
inverted over it --
-- because the reaction that we

293
00:24:35,000 --> 00:24:39,000
will talk about here,
in fact, does involve a gas.

294
00:24:39,000 --> 00:24:45,000
And that gas is hydrogen,
another one of our important

295
00:24:45,000 --> 00:24:51,000
energy molecules to be talking
about in the context of today's

296
00:24:51,000 --> 00:24:55,000
lecture.
At the end, I would like to see

297
00:24:55,000 --> 00:25:01,000
if you have a list of energy
molecules collected from today's

298
00:25:01,000 --> 00:25:05,000
lecture.
But also, we are going to

299
00:25:05,000 --> 00:25:08,000
choose, as the electrode
material, here,

300
00:25:08,000 --> 00:25:11,000
this metal piece of electrode
material.

301
00:25:11,000 --> 00:25:17,000
This is going to be platinum.
And the reason we are going to

302
00:25:17,000 --> 00:25:20,000
choose platinum,
a very common electrode

303
00:25:20,000 --> 00:25:26,000
material, is that platinum is a
so-called noble metal.

304
00:25:31,000 --> 00:25:35,000
And that means that it is quite
impervious to most chemical

305
00:25:35,000 --> 00:25:37,000
reactions.
Over there, we were talking

306
00:25:37,000 --> 00:25:41,000
about zinc and copper electrodes
that themselves do change during

307
00:25:41,000 --> 00:25:43,000
the reaction.
And over here,

308
00:25:43,000 --> 00:25:47,000
we want to have a piece of
metal that can provide the

309
00:25:47,000 --> 00:25:51,000
valuable function of giving us
electrical contact of the

310
00:25:51,000 --> 00:25:55,000
chemical reactions on the two
sides, but an electrode that

311
00:25:55,000 --> 00:25:59,000
will remain clean and unchanged
at its surface throughout the

312
00:25:59,000 --> 00:26:04,000
course of the reaction.
So, we use a platinum electrode

313
00:26:04,000 --> 00:26:07,000
over here.
And you can see that with this

314
00:26:07,000 --> 00:26:12,000
inverted test tube through which
the wire passes and with some

315
00:26:12,000 --> 00:26:16,000
kind of a provision for a side
arm attachment here,

316
00:26:16,000 --> 00:26:20,000
we can hook up a cylinder of H
two gas.

317
00:26:20,000 --> 00:26:25,000
And we can start bubbling H two
over this electrode,

318
00:26:25,000 --> 00:26:30,000
like this, so that it actually
will have bubbles of H two

319
00:26:30,000 --> 00:26:35,000
coming out like that.
In fact, this solution over

320
00:26:35,000 --> 00:26:40,000
here will be saturated with H
two.

321
00:26:40,000 --> 00:26:46,000
And a further consideration for
our reference electrode,

322
00:26:46,000 --> 00:26:52,000
which is going to be known as
the standard hydrogen electrode,

323
00:26:52,000 --> 00:26:54,000
or SHE.

324
00:27:06,000 --> 00:27:12,000
That is, the electrode against
which everything else is going

325
00:27:12,000 --> 00:27:15,000
to be compared.
And it also has,

326
00:27:15,000 --> 00:27:20,000
in the solution here,
1.0 molar H three O plus.

327
00:27:20,000 --> 00:27:23,000
So, indeed, it is a very

328
00:27:23,000 --> 00:27:29,000
strongly acidic medium.
And what we are measuring,

329
00:27:29,000 --> 00:27:35,000
over on the right-hand side,
is a half reaction that may

330
00:27:35,000 --> 00:27:41,000
correspond either to two H plus
plus two electrons going to

331
00:27:41,000 --> 00:27:44,000
H two.

332
00:27:44,000 --> 00:27:49,000
And that will be the case if
this side is the cathode.

333
00:27:49,000 --> 00:27:53,000
But if this side turns out to
be the anode,

334
00:27:53,000 --> 00:28:00,000
then we would be measuring the
opposite reaction.

335
00:28:00,000 --> 00:28:03,000
And I will talk more about that
in a second.

336
00:28:03,000 --> 00:28:08,000
Let's say, for example,
that our electrode over here

337
00:28:08,000 --> 00:28:11,000
is, in fact, made of zinc.

338
00:28:19,000 --> 00:28:22,000
If our electrode over here is
made of zinc.

339
00:28:22,000 --> 00:28:27,000
And we may have zinc two plus
in solution,

340
00:28:27,000 --> 00:28:32,000
here, and we may again have
something like potassium nitrate

341
00:28:32,000 --> 00:28:37,000
or some other
electrolyte present in our salt

342
00:28:37,000 --> 00:28:41,000
bridge.
Then what we can do is put

343
00:28:41,000 --> 00:28:43,000
these things briefly into
contact.

344
00:28:43,000 --> 00:28:48,000
And we want to know two things.
We want to know what the

345
00:28:48,000 --> 00:28:52,000
magnitude of the potential
difference is and what the

346
00:28:52,000 --> 00:28:57,000
direction in which the electrons
are flowing is so that we know

347
00:28:57,000 --> 00:29:02,000
which one is the cathode and
which one is the anode.

348
00:29:02,000 --> 00:29:07,000
And, in this particular case,
the direction of electron flow

349
00:29:07,000 --> 00:29:10,000
is, as before,
away from the zinc.

350
00:29:10,000 --> 00:29:16,000
Once again, zinc atoms are
jumping off the surface of the

351
00:29:16,000 --> 00:29:20,000
electrode as zinc two plus
ions.

352
00:29:20,000 --> 00:29:25,000
And, every time that happens,
two electrons go into the

353
00:29:25,000 --> 00:29:30,000
external circuit and come around
here.

354
00:29:30,000 --> 00:29:34,000
And the chemical reaction that
occurs here, in this case,

355
00:29:34,000 --> 00:29:38,000
is that those two electrons
that came from one of the zinc

356
00:29:38,000 --> 00:29:43,000
atoms react, as shown here,
with two H plus ions to make

357
00:29:43,000 --> 00:29:45,000
H two.

358
00:29:45,000 --> 00:29:49,000
And so, we know the direction of
electron flow,

359
00:29:49,000 --> 00:29:53,000
we know the exact reactions
that are taking place on the

360
00:29:53,000 --> 00:29:57,000
left and on the right.
And we know that,

361
00:29:57,000 --> 00:30:02,000
for this particular choice of
substances, our anode is on the

362
00:30:02,000 --> 00:30:06,000
left, and our cathode is on the
right.

363
00:30:06,000 --> 00:30:09,000
And let me come over here.

364
00:30:17,000 --> 00:30:21,000
There is another way that I can
write this cell.

365
00:30:21,000 --> 00:30:27,000
And so, I would like to
introduce cell notation to you.

366
00:30:27,000 --> 00:30:32,000
And that will be like this.
We will have zinc.

367
00:30:32,000 --> 00:30:41,000
And then, a solid line.
And then, zinc two plus.

368
00:30:41,000 --> 00:30:48,000
And then, a solid double
vertical line.

369
00:30:48,000 --> 00:30:57,000
And then, H plus,
another solid vertical line,

370
00:30:57,000 --> 00:31:02,000
H two.
And then another solid line.

371
00:31:02,000 --> 00:31:07,000
And then platinum.

372
00:31:07,000 --> 00:31:10,000
That is the way I set that up.

373
00:31:10,000 --> 00:31:15,000
We could, actually,
add another solid line here and

374
00:31:15,000 --> 00:31:20,000
put platinum over here.
The thing that I want you to

375
00:31:20,000 --> 00:31:26,000
recognize about this notation
for Galvanic cells is that the

376
00:31:26,000 --> 00:31:31,000
solid single lines represent
interfaces, direct contacts

377
00:31:31,000 --> 00:31:36,000
between things.
Here, it would be a platinum

378
00:31:36,000 --> 00:31:39,000
solid electrode
connected with solid zinc,

379
00:31:39,000 --> 00:31:43,000
possibly.
And then, the way I drew it

380
00:31:43,000 --> 00:31:47,000
over there, it would just be
solid zinc, no platinum on the

381
00:31:47,000 --> 00:31:49,000
left.
And then this solid line

382
00:31:49,000 --> 00:31:53,000
represents a solid liquid
interface because the electrode

383
00:31:53,000 --> 00:31:57,000
is dipped into a solution that
contains zinc two

384
00:31:57,000 --> 00:32:00,000
ions.
And that is in communication

385
00:32:00,000 --> 00:32:04,000
with the other half of this
electrochemical cell by a salt

386
00:32:04,000 --> 00:32:08,000
bridge.
And the salt bridge is

387
00:32:08,000 --> 00:32:11,000
represented by the double
vertical line.

388
00:32:11,000 --> 00:32:14,000
And then we have,
in solution,

389
00:32:14,000 --> 00:32:17,000
protons.
And then, a solid liquid.

390
00:32:17,000 --> 00:32:20,000
Well, actually,
in this case,

391
00:32:20,000 --> 00:32:25,000
a liquid gas interface to the
gaseous H two.

392
00:32:25,000 --> 00:32:30,000
And H two is also
dissolved.

393
00:32:30,000 --> 00:32:34,000
In order for this to be used as
a reference electrode,

394
00:32:34,000 --> 00:32:38,000
we have to pick standard
conditions.

395
00:32:38,000 --> 00:32:43,000
And so, we are usually talking
about one atmosphere of hydrogen

396
00:32:43,000 --> 00:32:47,000
for the standard hydrogen
electrode.

397
00:32:47,000 --> 00:32:52,000
And we are talking about a
concentration of 1.0 molar of

398
00:32:52,000 --> 00:32:56,000
our strong acid.
And so, those are our standard

399
00:32:56,000 --> 00:33:01,000
conditions, along with 25
degrees C.

400
00:33:01,000 --> 00:33:05,000
And then this solution that
contains the protons and the

401
00:33:05,000 --> 00:33:10,000
hydrogen gas is in a solution
solid interface with the solid

402
00:33:10,000 --> 00:33:14,000
platinum electrode that is not
going to be undergoing any

403
00:33:14,000 --> 00:33:17,000
change.
So, these are the types of cell

404
00:33:17,000 --> 00:33:21,000
notations that you will
encounter when looking into

405
00:33:21,000 --> 00:33:25,000
electrochemistry.
And we found out that the left

406
00:33:25,000 --> 00:33:30,000
side, where the zinc metal
is going to zinc two

407
00:33:30,000 --> 00:33:34,000
plus,
is our anode.

408
00:33:34,000 --> 00:33:39,000
And also, the potential
difference, when we measure it

409
00:33:39,000 --> 00:33:45,000
using our voltmeter right here
by turning on the contact very

410
00:33:45,000 --> 00:33:51,000
briefly, is around 0.763 volts.
So, that is the magnitude of

411
00:33:51,000 --> 00:33:56,000
our potential difference.
And this is equal to delta E

412
00:33:56,000 --> 00:34:02,000
zero cell.
When I write E zero,

413
00:34:02,000 --> 00:34:08,000
that means standard potential.
And this is written as a delta

414
00:34:08,000 --> 00:34:12,000
here because that is the
difference between the

415
00:34:12,000 --> 00:34:17,000
potentials for the two
half-reactions that we have

416
00:34:17,000 --> 00:34:20,000
written up there.
But it is a very simple

417
00:34:20,000 --> 00:34:26,000
difference precisely because the
standard hydrogen electrode is

418
00:34:26,000 --> 00:34:32,000
our reference electrode.
Over here we want to write the

419
00:34:32,000 --> 00:34:36,000
definition of standard cell
potential.

420
00:35:07,000 --> 00:35:14,000
And that is a delta E zero for
your cell is equal to delta E

421
00:35:14,000 --> 00:35:21,000
zero for your cathode minus--
Sorry, this is not delta here.

422
00:35:21,000 --> 00:35:28,000
Just E zero for your cathode
minus E zero for your anode.

423
00:35:40,000 --> 00:35:44,000
But E zero for the
standard hydrogen electrode,

424
00:35:44,000 --> 00:35:49,000
because this is our reference
electrode, is equal to zero at

425
00:35:49,000 --> 00:35:51,000
all temperatures.

426
00:35:56,000 --> 00:36:02,000
And that is by definition.
So, the standard potential for

427
00:36:02,000 --> 00:36:08,000
that set of conditions that
constitutes our standard

428
00:36:08,000 --> 00:36:14,000
hydrogen electrode is taken as
the zero of potential for

429
00:36:14,000 --> 00:36:21,000
comparison with any other type
of half cell that you might be

430
00:36:21,000 --> 00:36:25,000
able to consider.
And so, we can see further

431
00:36:25,000 --> 00:36:32,000
that, in this particular case,
where we have our cathode as

432
00:36:32,000 --> 00:36:39,000
the standard hydrogen electrode,
we have zero and minus E zero

433
00:36:39,000 --> 00:36:45,000
for the anode,
which we measured as 0

434
00:36:45,000 --> 00:36:48,000
volts.

435
00:36:53,000 --> 00:36:57,000
That means that for the
reaction zinc going to zinc two

436
00:36:57,000 --> 00:37:01,000
plus plus two electrons,

437
00:37:01,000 --> 00:37:06,000
which is our anodic reaction,
we have a standard potential of

438
00:37:06,000 --> 00:37:09,000
-0.763 volts.
And generally,

439
00:37:09,000 --> 00:37:12,000
what you are going to be
interested in,

440
00:37:12,000 --> 00:37:17,000
as you consider different kinds
of substances from throughout

441
00:37:17,000 --> 00:37:21,000
the periodic table with
reference to their ability to

442
00:37:21,000 --> 00:37:25,000
take place in
oxidation-reduction reactions,

443
00:37:25,000 --> 00:37:29,000
is you are going to want to
know what your standard

444
00:37:29,000 --> 00:37:33,000
potential is.
You are not going to always

445
00:37:33,000 --> 00:37:37,000
have a system that you want to
consider that is under standard

446
00:37:37,000 --> 00:37:40,000
conditions.
These standard potentials

447
00:37:40,000 --> 00:37:43,000
always are referenced to some
standard conditions,

448
00:37:43,000 --> 00:37:47,000
as I mentioned specifically for
the standard hydrogen electrode.

449
00:37:47,000 --> 00:37:51,000
And so, part of next lecture on
Monday is we are going to show

450
00:37:51,000 --> 00:37:54,000
you how to handle systems that
are not under standard

451
00:37:54,000 --> 00:37:58,000
conditions because then you can
handle some real practical

452
00:37:58,000 --> 00:38:02,000
problems.
But let's look at a different

453
00:38:02,000 --> 00:38:03,000
type of cell.

454
00:38:10,000 --> 00:38:13,000
Let's write down the following
cell.

455
00:38:40,000 --> 00:38:43,000
This new cell that I have
written the cell notation for,

456
00:38:43,000 --> 00:38:48,000
instead of drawing up the whole
diagram of this Galvanic cell,

457
00:38:48,000 --> 00:38:52,000
happens to be one in which we
have platinum electrodes on both

458
00:38:52,000 --> 00:38:54,000
sides.
Here we have the silver |

459
00:38:54,000 --> 00:38:58,000
silver plus
redox couple on the left-hand

460
00:38:58,000 --> 00:39:02,000
side.
And we have the redox couple of

461
00:39:02,000 --> 00:39:06,000
H plus with H two
on the right-hand side.

462
00:39:06,000 --> 00:39:09,000
So, the right-hand side
corresponds to the standard

463
00:39:09,000 --> 00:39:13,000
hydrogen electrode,
and the left-hand side is a new

464
00:39:13,000 --> 00:39:16,000
redox couple,
or a new half-reaction,

465
00:39:16,000 --> 00:39:20,000
that we want to compare to the
standard hydrogen electrode so

466
00:39:20,000 --> 00:39:23,000
that we will know,
on an absolute scale,

467
00:39:23,000 --> 00:39:27,000
where it falls relative to all
the other half-reactions we

468
00:39:27,000 --> 00:39:31,000
might want to measure.
And so, we want to know,

469
00:39:31,000 --> 00:39:34,000
which way do the electrons
flow?

470
00:39:39,000 --> 00:39:43,000
And, in this case,
it turns out that the direction

471
00:39:43,000 --> 00:39:48,000
of electron flow is this way.
So, this is the opposite of

472
00:39:48,000 --> 00:39:53,000
what we talked about in the case
of comparing zinc to the

473
00:39:53,000 --> 00:39:57,000
standard hydrogen electrode.
Electron flow has been

474
00:39:57,000 --> 00:40:00,000
reversed.
What that means is that this

475
00:40:00,000 --> 00:40:05,000
over here, the SHE is now our
anode.

476
00:40:05,000 --> 00:40:11,000
And now the silver-silver plus
electrode is our

477
00:40:11,000 --> 00:40:14,000
cathode.
And we want to know not only

478
00:40:14,000 --> 00:40:20,000
the direction of electron flow,
the direction of electron flow

479
00:40:20,000 --> 00:40:26,000
tells us what is doing the
reduction and what is doing the

480
00:40:26,000 --> 00:40:32,000
oxidation, but we also want to
know the magnitude.

481
00:40:32,000 --> 00:40:37,000
And this one turns out to be,
this delta E cell,

482
00:40:37,000 --> 00:40:43,000
is equal to 0.8 volts.
And, by the definition of

483
00:40:43,000 --> 00:40:49,000
standard cell potentials that I
gave you over there,

484
00:40:49,000 --> 00:40:55,000
you can see that what we are
getting now, because the

485
00:40:55,000 --> 00:41:00,000
electron flow is reversed,
our sign is reversed,

486
00:41:00,000 --> 00:41:06,000
and so our E zero for the
reaction Ag plus plus an

487
00:41:06,000 --> 00:41:12,000
electron going to silver is

488
00:41:12,000 --> 00:41:18,000
equal to 0.8 volts,
positive.

489
00:41:18,000 --> 00:41:22,000
Notice that the zinc-zinc two
plus couple was

490
00:41:22,000 --> 00:41:25,000
negative with respect to the
standard hydrogen electrode,

491
00:41:25,000 --> 00:41:29,000
but because the electron flow
is reversed for silver plus,

492
00:41:29,000 --> 00:41:32,000
silver redox couple,
we now have a positive

493
00:41:32,000 --> 00:41:37,000
potential relative to the
standard hydrogen electrode.

494
00:41:45,000 --> 00:41:51,000
What would like to arrive at is
a nice big table where we look

495
00:41:51,000 --> 00:41:54,000
at standard potential --

496
00:42:01,000 --> 00:42:06,000
-- with reference to this
standard hydrogen electrode,

497
00:42:06,000 --> 00:42:12,000
which is our zero of potential.
And you will be able to find a

498
00:42:12,000 --> 00:42:17,000
table like this in your book.
What we found is that zinc is

499
00:42:17,000 --> 00:42:22,000
down here at -0.763,
so that was zinc-zinc two plus

500
00:42:22,000 --> 00:42:28,000
at a negative
potential relative to hydrogen

501
00:42:28,000 --> 00:42:33,000
plus electrons.
We found that up here at +0.8,

502
00:42:33,000 --> 00:42:39,000
we have the silver-silver plus
redox couple.

503
00:42:39,000 --> 00:42:44,000
These are thermodynamic
quantities, so you can look on a

504
00:42:44,000 --> 00:42:49,000
table of standard reduction
potentials and you can tell

505
00:42:49,000 --> 00:42:54,000
which direction electrons will
flow if you set up cells that

506
00:42:54,000 --> 00:42:59,000
involve those redox
half-reactions.

507
00:42:59,000 --> 00:43:04,000
Another electrode that we used
today was copper two plus

508
00:43:04,000 --> 00:43:09,000
combining with two electrons to
give copper metal.

509
00:43:09,000 --> 00:43:13,000
It turns out that one is

510
00:43:13,000 --> 00:43:17,000
positive also by about positive
0.3 volts.

511
00:43:17,000 --> 00:43:23,000
You see hydrogen is here and it
reduces silver-silver plus.

512
00:43:23,000 --> 00:43:27,000
Hydrogen reduces silver plus to

513
00:43:27,000 --> 00:43:32,000
silver because it is up there.
Zinc is down here.

514
00:43:32,000 --> 00:43:37,000
Zinc also serves as an anode
with respect to any one of those

515
00:43:37,000 --> 00:43:41,000
three because all of those three
are at a potential positive

516
00:43:41,000 --> 00:43:45,000
relative to zinc on this scale
of standard reduction

517
00:43:45,000 --> 00:43:49,000
potentials.
You can also have some other

518
00:43:49,000 --> 00:43:53,000
potentials, like way down here
at about -2.7 volts is the

519
00:43:53,000 --> 00:43:57,000
sodium-sodium plus
redox couple,

520
00:43:57,000 --> 00:44:02,000
way down there at -2.7
That is why sodium is so much

521
00:44:02,000 --> 00:44:05,000
fun to heave into a body of
water.

522
00:44:05,000 --> 00:44:09,000
I mean, it is fantastic.
You get reduction of the

523
00:44:09,000 --> 00:44:13,000
protons to make hydrogen,
which then explodes.

524
00:44:13,000 --> 00:44:17,000
So, this is fantastic.
And some of you may know that

525
00:44:17,000 --> 00:44:21,000
there is an annual,
and I am not recommending that

526
00:44:21,000 --> 00:44:25,000
you do this, by the way.
You can see how negative the

527
00:44:25,000 --> 00:44:32,000
potential is down here.
Metallic sodium is a very

528
00:44:32,000 --> 00:44:35,000
strong reducing agent,
indeed.

529
00:44:35,000 --> 00:44:42,000
A very important reaction is up
here at about +1.23 volts

530
00:44:42,000 --> 00:44:48,000
relative to the standard
hydrogen electrode.

531
00:44:48,000 --> 00:44:56,000
And this important reaction is
oxygen plus four H pus plus four

532
00:44:56,000 --> 00:45:03,000
electrons going to two H two O.

533
00:45:11,000 --> 00:45:14,000
So, this +1.23 volts is very
important.

534
00:45:14,000 --> 00:45:18,000
In the time remaining,
I am not going to be able to

535
00:45:18,000 --> 00:45:23,000
tell you about electrolysis.
I believe next hour I will

536
00:45:23,000 --> 00:45:27,000
start off by talking about
electrolysis.

537
00:45:27,000 --> 00:45:31,000
And, if you can do electrolysis
using oxidizing and reducing

538
00:45:31,000 --> 00:45:35,000
equivalents that derive from
photovoltaic cells,

539
00:45:35,000 --> 00:45:40,000
so you are converting sunlight
into separated electron whole

540
00:45:40,000 --> 00:45:45,000
pairs, you can drive a reaction
like this the other way and

541
00:45:45,000 --> 00:45:50,000
learn how to make oxygen from
water, and ultimately also

542
00:45:50,000 --> 00:45:54,000
hydrogen from water.
We will talk about that a

543
00:45:54,272 --> 00:45:57,000
little bit next time.
Have a nice weekend.