The following content is provided by MIT OpenCourseWare
under a Creative Commons license.
Additional information about our license and MIT
OpenCourseWare in general is available at ocw.mit.edu.
Continuing our discussion of acid-base theory,
I have just drawn here on the board for you the structure of a
popular indicator molecule. Indicator molecules are used
for titrations. And titrations will be of focus
of today's lecture. This is the indicator molecule,
phenolphthalein. It is easier to draw this
molecular structure than it is to spell phenolphthalein,
but you should know how to do both of these things,
or at least recognize them. Now, how do indicator molecules
work? Well, during a titration,
you are mixing chemicals in a way that leads to a changing pH
throughout the experiment. And an indicator molecule like
phenolphthalein is, in fact, a Bronsted acid.
And, if you inspect the structure that I have drawn here
for phenolphthalein, you should be able to begin to
understand just how it is that a molecule like this can serve as
a Bronsted acid. If I write out the formula,
you will see that we have quite a number of carbons and
hydrogens in the molecule. And we have oxygens as well.
Those are the three elements present in a phenolphthalein
molecule. And two of the hydrogens in
this molecule possess Bronsted acidity that makes them valuable
for the function of this system as an indicator molecule.
And what happens is when you have this in solution at low pH,
it is in this form. It is neutral.
And then, as you begin to raise the pH, at some point a base
that has a lone pair of electrons comes along and takes
one of these two protons in a molecule, either this one or
this one. Those two protons are
equivalent. And what I am now going to draw
is a sequence of arrows representing the electronic flow
in the phenolphthalein molecule that occurs when H plus
is transferred over to the base with its lone pair of electrons.
So, I am starting these arrows at pairs of electrons and I am
showing the flow of how they move.
And this is a pretty interesting rearrangement
because it is complicated. It is an electronic
rearrangement that is transmitted through this
molecule such that here in the center, this four coordinate or
sp three hybridized carbon, which is bonded to this
oxygen, has a pair of electrons that migrates to it.
And, if I draw now the product of this reaction,
what you should be able to see quite clearly --
And I am not always explicitly drawing out the lone pairs of
electrons on each of the electronegative oxygen atoms in
the molecule, but the sequence of double
bonds and single bonds in the molecule has been altered in a
profound manner because of the base coming along and removing a
proton from one of these peripheral oxygens,
namely, that one here. So, we have now a singly
negatively charged system. And that charge is balanced,
over here, by the base that has the proton attached to it.
That is the structure-function relationship that is typical of
indicator molecules. And indicator molecules can
either be naturally occurring, as they are in some plants.
Or, alternatively, they can be molecules that are
synthesized to respond to factors like pH in a given way
that might be desired. And, in the case of
phenolphthalein, this change takes place between
8.2 to 10.0 pH units. And the color change associated
with this is colorless to pink.
Looking at the structure of the anionic form of the
phenolphthalein molecule that results from its deprotonation,
what you can see is that this sp three hybridized
carbon atom at the center of the molecule, which in the neutral
form is acting as an insulator, preventing the pi-systems of
the three six-member rings from communicating,
is converted into a three-coordinate sp two
hybridized carbon atom here at the center, --
and provides a conduit for a communication between the
pi-systems of the three six-membered substituted benzene
rings that are in this molecule. And that change in electronic
structure is what leads to the production of this pink
Chromaphor in the anionic deprotonated form of a
phenolphthalein molecule that is important at high pH.
All you have to do, in fact, is to change the
substituents on this molecule in order to get a response at a
different pH value. Bromothymol blue is an
indicator with a structure similar to that of
phenolphthalein, but with a sulfur here in place
of this carbon. And in bromothymol blue,
we also have some isopropyl and bromine substituents on the
aromatic rings. And that modifies its
properties such that it goes from yellow to blue as the pH
rises above 6.0 to about a value of 7.6.
So, if you go to a textbook and look in a table of indicator
molecules, you will be able to select an indicator that is
appropriate for a particular type of titration.
And just to give you a very simple schematic of a classic
type of titration, I am going to show you here --
-- the type of system that you may need to think about in
connection with this next problem set.
And it is a system like this, where we have a flask down here
at the bottom that contains, initially, some volume of a
weak acid.
And that weak acid might be, for example,
acetic acid, as we talked about last time.
And we will say, for example,
that this could be 0.1 molar CH three COOH.
That is acetic acid. And the calculation that we
left off with at the end of last hour is one that we will get to
solving today that had to do with what is the pH of a tenth
molar acetic acid solution in water?
Or, in other words, what is the pH of this solution
down here containing the acetic acid at the beginning of this
titration that we are going to carry out.
We might have a weak acid down here.
We might also have our indicator present,
so that we will see a color change when we pass through a
particular pH. And we are going to make the pH
rise by putting into this burette here,
which will allow us to control a dropwise addition of a strong
base solution that we are going to add in and slowly neutralize
our acetic acid that is down here at the bottom.
So, this is strong base. And, for example,
that might be tenth molar sodium hydroxide.
So, there is a typical titration.
And, if you think about it, because we are adding one
solution to another, we are going to have a
constantly changing volume in this bottom solution throughout
the experiment. And so ultimately what we are
going to be interested in, in our questions such as the
following, as the volume increases --
And we will give the volume, for the purposes of today's
lecture, the symbol the lower case letter m -- As the volume
increases, what happens to the pH?
We know that we are using a strong base to neutralize an
acid, so the pH is going to start out low and somehow is
going to rise. But what if we wanted to
predict the mathematical form of the rise in pH as a function of
this increasing volume of the solution?
Then we would have to have some equations that would be useful
for describing this physical property as a function of this
increasing volume. This is what we often want to
be able to do in chemistry. We want to be able to describe
properties that change either as a function of time or as a
function of some other variable, here volume,
in a titration. And we want to be able to see
if we can come up with a set of equations that would predict
that change as a function of this variable that is changing.
And, in order to do that, here, I am going to first point
out that we can put together some titration equations.
And, ultimately, we are going to want to
generate a mathematical model for this titration.
And then, if you are experimentalist,
you want to go ahead and take that mathematical model and
compare it to actual experimental data.
But let's work on getting a model generated,
here. First we are going to have a
set of chemical equations. These correspond to the acid
HA, which is acetic acid in our example, in water reacting to
give some concentration of H three O plus plus the
anion A minus, which is acetate ion,
in this particular case. We have a strong base,
which is sodium hydroxide. And when you put sodium
hydroxide in water, because it is a strong base,
it ionizes completely. And I am signifying that with a
single forward arrow rather than with an equilibrium arrow.
This going into water is going to be Na plus and
hydroxide, OH minus. Now, at certain points in a
titration curve, you would not need to
necessarily go much further than this in order to have enough
information to start solving for the pH at a particular value of
a volume m. But at some points along that
titration curve, that is not enough information.
And that is because water itself can sometimes get into
the equation. Here is an interesting
equation, which is H two O plus H two O going to H three O plus
plus O minus.
That is a conceivable reaction that could be occurring even in
pure water, but the value of the equilibrium constant for that
reaction is very small. And I will draw that as an
equilibrium to distinguish it from the irreversible ionization
of sodium hydroxide in aqueous solution.
I have a pair of equilibria that I have written.
I have one forward reaction. And I want to point out,
too, that this third equation I have written here has a special
name, which is autoprotolysis.
I will tell you more about the equilibrium constant associated
with autoprotolysis in a moment. But because this equilibrium
constant is very small, meaning that there is very
little H three O plus and OH minus present in
pure water, because of that pure water itself is not a very good
electrolyte. Meaning pure water itself is
not a very good conductor of electricity, because the
autoprotolysis reaction lies mostly over here to the left in
pure water. But at certain points along a
titration curve, such as the one we are going to
want to develop on this panel here, we are going to need to
include autoprotolysis in order to get the right answer because
it is there, it can occur. Let's continue.
Another type of equation that we can use to put this into some
kind of a mathematical footing is the balance of charge in a
solution like the one that we are talking about in that beaker
to which we are adding sodium hydroxide solution.
And the charge balance consideration tells us that the
concentration of hydronium ion plus the concentration of sodium
ion must be equal to the concentration of hydroxide plus
the concentration of acetate ion, A minus.
Because what I have done here in the charge balance
consideration is I have said okay, what are all the possible
charged species in solution? I can make a list of them.
It is those four things. And I know that the number of
positive charges must equal the number of negative charges
because this is overall a neutral beaker in which we are
putting things. But everything that you put in
is charge neutral, so the positives have to equal
the negatives. That is a very nice limiting
equation that helps us understand how the different
competing chemical reactions and equilibria will all settle down
and arrive at the physical result, which will be our
observable. And then, in addition to that,
we have a mass balance consideration.
And in this mass balance consideration,
we will be able to write that our HA initial value --
-- that is square brackets, once again, denoting
concentration -- will be, at any point in time,
equal to the actual value of HA plus the concentration of A
minus. And that just says that when
you have acetic acid and you put it into water,
it ionizes partly to hydronium ion and acetate ion.
And that which does not ionize is unionized HA.
And so, the ionized A minus concentration plus the
unionized HA concentration is equal to the sum total of the
acetic acid that is present in solution.
And, similarly, we have another mass balance
equation, which is just that the sodium hydroxide initial is
equal to our sodium ion concentration because all of it
is ionized. That is a little simpler than
the situation with the acetic acid.
And then, in addition to those equations, we have our
equilibrium equations.
And we are going to have two of these because there are two
equilibria that I wrote over there under titration equations.
The first of these is going to have the equilibrium constant
Ka, and that is going to be H three O plus times A
minus, which is acetate in the case of
acetic acid, all over HA. That is the expression for the
equilibrium of acetic acid ionization in dilute aqueous
solution. And I will give you a value for
this in a moment. I won't write it right here.
Well, why don't I just do that. It is 1.8x10^-5 for this
specific example of acetic acid. And then, if you were working
with some other weak acid, you would be able to go to a
table and look up the Ka value for it, if it has been measured
and reported and tabulated. And then, we have Kw.
And this is the equilibrium expression for the
autoprotolysis reaction that I mentioned over there,
the other equilibrium, which is H three O plus times
OH minus. And that is equal to 1x10^-14,
so a very small equilibrium constant.
And some of you may be wondering why water is not
appearing in an expression like this one or like this one.
And the reason for this is that in the expressions that I am
deriving here, we are using concentrations in
place of activities for these species.
In the derivation of the equilibrium constants,
the activity of pure water at very high concentration is
approximated as one. And so, this is simplified as
shown here. And having put all these
equations together, what we are going to want to do
is take the following approach. We are looking at the
expression for electro-neutrality here,
the equal number of positively and negatively charged species
in solution. And I want to get this
expression rewritten in terms of things that I know at any point
along the titration curve, with the exception of the one
thing that I want to know, the pH.
And the pH we can calculate easily if we know the H three O
plus concentration. So, I am going to be seeking to
get Na plus, OH minus,
and A minus rewritten in terms of H three O plus
and constants using the equations that I have up
here on the board.
First, let's attempt to do that for A minus using this
Ka expression. And I am going to rearrange it.
And it is going to be equal to Ka times HA divided by H three O
plus. But I still have a problem with
this expression for A minus because I am going to
need to replace HA with something that will allow me to
get everything in terms of just H three O plus and
constants. And so, if I look over here at
the mass balance equation, I can use this one and replace
HA with HA initial minus A minus.
Now I have A minus written in terms of an initial
concentration of HA itself and constants.
I need to rearrange this a little bit.
I am going to multiply through by H three O plus so
that I can write A minus times H three O plus is equal to Ka
times HA initial minus Ka times A minus.
And hopefully I have done that
right. And, if so, I will be able to
say that A minus is going to be equal to Ka times HA initial all
over H three O plus.
Does that look good? Plus Ka, thank you.
H3O+ plus Ka.
And let's bring that down here.
Because now, I would like to go ahead and
make this substitution for the charge.
We have our A minus. The other ones are a little
easier. We now have this equation H
three O plus. And we were going to replace Na
plus by NaOH initial.
And we are going to recognize that we can use the equation for
autoprotolysis to replace hydroxide with Kw divided by H
three O plus. And then now we have the A
minus being replaced by Ka HA initial all over H
three O plus plus Ka.
And at this point, you have actually done a really
nice thing, because you have generated a general titration
equation for the titration of interest in this particular
case. You can see that at any point
along that curve for a given volume m, we can now stop and
walk along in one milliliter increments.
And, in each spot along that curve, we could go ahead and
solve this equation to get the value of H three O plus.
And then we would just have to
take the negative log of that, and we would have the pH at
that position. And we are interested if this
type of a model for the mathematical form of the
titration curve bears any resemblance to what one sees in
reality. And, if so, one could then be,
at least for the time being, until your theory no longer
fits, satisfied that you have accounted for the various
equilibria that could be present in a system such as the one we
are describing here. This happens to be a cubic
equation, and it has, therefore, three roots.
And when we solve it at each of these points along the titration
equation, because of the physically realistic quantities
that we are going to be interested in,
only one of the roots is positive.
We are talking about H three O plus concentration.
It cannot be a negative concentration.
The positive root is the one we want, and then we can convert
that to pH and can see what kind of result we get.
And so, I will show you how we can do that using a tool
available on Athena.
First of all, let me tell you that for the
purposes of doing this, I am going to replace these
various quantities by some simpler symbols,
here. I am going to go x plus a,
a is our NaOH initial, is equal to (b over x) plus (c
times d) over (x plus d).
So, I am going to use those simpler symbols for this
equation. I could rearrange it and show
you how it is a cubic equation, you set it equal to zero,
and then you solve, and you get the roots.
Just like a quadratic equation but with another term.
And, instead of doing that, we are just going to leave it
in this form, which is simpler to look at.
And we are just going to use that.
And so, let's go ahead and define A.
That is the initial concentration of sodium
hydroxide. We are going to go 0.1.
And remember I need to get this in terms of the quantity m here,
which is the volume. And so, I am going to take m,
minus 20 over m. And the reason I am doing is
this is I am saying that we have an 0.1 molar solution of sodium
hydroxide that we are adding. But there is a dilution factor
that you have to take care of. So, at every value m we are
saying that we are starting with a volume of 20 in the acetic
acid flask at the bottom there at the beginning.
And then we are starting to add volume units of sodium
hydroxide. And so the volume,
m, is increasing. That is how I take care of
that, there. And then our value of b is Kw,
so that is 1 E minus 14.
That is for autoprotolysis. And then we have c,
that is our initial concentration of acetic acid.
And we happen to have a dilution factor here,
too, but it is simpler, just 20 divided by m.
And there is that. And then d, what did I say that
d was? d is our equilibrium constant
for HA, 1.8 E minus 5.
There is our acidity constant for acetic acids.
Now I have the four things in there that I need,
and those are four things that I know because of the initial
conditions that I am defining for this problem.
And then I am going to define my equation as being equal to x
plus a is equal to (b over x) plus (c times d) all over (x
plus d). There is my equation.
And this nice calculator, all I have to do,
in principle, is tell it what the value of m
is. And then it is going to solve
that cubic equation, give me back the three roots.
I take the positive one and I take the negative log of it,
and then I know the pH at that value of m.
I can develop a whole titration curve based on the equations
that we have been discussing. And so let's choose a value of
m here to start with, and I am going to choose 30.
We are starting at 20. At the very beginning of the
titration, we are at 20. Let's see what happens at 30
given the particular value of m. Here is what I do.
There may be other ways to do this.
I am defining this thing called the solution set.
Because the solving of that equation is going to give us
three values. And then I want to be able to
pick the one that I want. We would use sol-set is equal
to, we are going to solve that equation, and we want x,
which is our H three O plus concentration.
And sometimes it chokes. Let's see what happens.
Sometimes it is a little slow. But it takes a lot longer if
you actually try to work it all out by hand.
I have a picture of an image of this that I grabbed.
I can show you if this fails on me.
Don't fail on me. There we go.
You can see it got the answer. And so here is a positive root,
and then there are two negative roots to this cubic equation in
x. And I want to find out the pH.
When we are up to 30 volume units, m equals 30.
And so, I will define p here as equal to the negative log base
10 of the first of the answers in my solution set,
sol-set one, which is the positive x value.
And, bingo, there is the pH, 4.745 whatever dot,
dot, dot. That is now the pH at a volume
of 30. Now, you figure I am starting
with 20 volume units of 0.1 molar acetic acid.
And by the time m equals 30 volume units,
that means I have added ten volume units of strong base,
enough to quench half of my acetic acid.
And what if I had just only quenched none of it?
I can change m now. Once you have this in here,
you don't have to retype these things.
I can go back here and say let m be 20.
And, if I do that, now I am actually at the zero
point, before I add any strong base.
And there is the pH now,
The answer to the expression that we had at the very end of
class last time that was going to tell us what is the pH of
tenth molar acetic acid, there is the answer,
2.87, because that is what we have when we start.
Now you can see that you could just go through and use that to
get the pH at any value of m along this titration curve.
And when you do that, let's show you what the result
is, now what you see here is a set of data.
I start at 20. And I'm going up in single
volume units, 21, 22, 23, and 24,
I am showing what the pH is at each of those choices of m.
And I have gone all the way from 20 to 60.
And you take that data set --
And then you can take a file, like the one I just showed you,
which just had values of m and corresponding values of pH.
And you can go ahead and just plot that thing.
And, when you do that, you get the form of the
titration curve here, which is the classic form of a
titration curve for titrating a weak acid with a strong base.
And you see that here, m was 20, and we let it
increment through all the way up to 60.
And what you see initially, and you are going to need to
study this one and the other cases of titration curves that
are in that chapter of your book, is how the little subtle
features differ from one case to another.
They do differ. A much simpler titration curve
than this one is obtained if you titrate a strong acid with a
strong base. This is titration of a weak
acid with a strong base. And you see that initially,
we have that 2.87 initial pH value, that it kind of shoots up
quickly here at the beginning. There is a steep rise,
initially. And then it kind of levels off.
And in this region here, the pH is not changing very
rapidly at all as you are adding more sodium hydroxide.
Until you get out here, and the pH just shoots up.
And you will notice that it is around where we get to 40.
Because we started with 20 mLs of tenth molar acetic acid.
And, when we get out to m equals 40, we have added 20 mLs
of tenth molar sodium hydroxide. At that point the pH just
shoots up. And then over here it turns
over and just goes kind of slowly for a minute.
And it just approaches an asymptotic limiting value over
here on the high pH side of the scale.
One thing you will notice here is that at this point where we
have added the same number of equivalence of sodium hydroxide
molecules as we had acetic acid molecules to start with,
this occurs over here at greater than 8.0 pH units rather
than at the neutral value of
And I would like you to think about why that might be.
And that will be something that will be covered in recitation,
but it is important. If we had derived here today
the expression for the titration of a strong acid with a strong
base, then at 40 we would be right at 7.0,
but when we were using this weak acid, acetic acid,
it comes up here over 8.0 pH units.
This point where we have added as many strong base molecules as
we started with, weak acid molecules,
is called the equivalence or the stoichiometric point in the
titration. Moreover, where would your
indicator molecule change color? At what point,
at what value of m, for example?
It would be close to there. It would be because,
in the case of phenolphthalein, it changes between right in
here. It might change right around
maybe 41 or so. m equals 41 might be where the
color change from colorless to pink occurs, for the reasons
that we have discussed. And then right here,
it turns out at this value of m equals 30, the pH is equal to
the pKa of acetic acid. And that is a relationship that
I would like you to think about. And it is a relationship that
comes into play very strongly when we talk about the
preparation of buffer solutions according to the
Henderson-Hasselbach equation. At m equals 30,
the pH value of the titration is the pKa of acetic acid,
which is right here at about 4.7 or so.
We did that one a moment ago. We found out what that value
was. And, in plus or minus 1.0 pH
units of that, we have the part of the
titration curve that is called the buffering region.
In here we have generated a buffer in situ.
And buffers help the pH not to change too much when acids or
bases are added in solution. That is what buffers do.
They try to help you keep in a constant pH range because lots
of things are dependent upon a particular pH for proper
functioning. And you know that this is
important in the world's oceans, it is important in our
environment, and it is important in our bodies.
So, the pH considerations for aqueous solutions are quite
important. And I hope that today I have
been able to give you a nice way of thinking about some of the
math that goes behind a titration curve.
Have a nice Halloween.