this point in your chemistry
career, you should be able to predict the products of chemical
reactions, the states of the products, and whether the reaction will
occur spontaneously at any given set of conditions. You should
even be able to determine the rate
at which the reactants are consumed and predict the amount of time it
take to produce a given amount of product. While this is
extremely useful information, it only applies to a limited set of
reactions, namely those that occur in one direction only:
A + B
C + D
the reactants A and B
sufficient energy and the proper geometry to form the products C and
What about a reaction in which C and D now become reactants in the
direction and form the products A and B?
C + D A + B
when A and B were
mixed, the reaction proceeds in the forward direction to produce C and
D. However, as time progresses, the concentration of C and D
increases causing an increase in the rate of the reverse
reaction. Concurrent with this increased rate of the reverse
reaction is a reduction of the forward rate due to the decrease in the
concentration of A and B. At some point, the rate of the forward
and reverse reactions will become the same and we will reach a state of
A + B C + D
state of dynamic equilibrium
does not mean that the forward and reverse reactions have
stopped. Molecules of A and B are still reacting to form C and D
and molecules of C and D are reacting to form A and B. However,
since the rate of the forward and reverse reactions is the same, it
will appear that nothing is happening. As such, all quantifiable
physical and chemical properties such has pH,
color, and concentration will remain constant.
which a moles of A react with b moles of B to produce c moles of C and
moles of D,
aA + bB cC + dD
specify an equilibrium
constant, Kequil (same as Kc), that
relates the concentration of all product and reactant species,
[A], [B], [C], and [D] are
the molar concentration of all species present at equilibrium.
The exponents, a, b, c, and d represent the stoichiometric coefficients
from the balance chemical reaction. Kequil is
a constant for all
conditions at a given temperature (normally 25°C unless otherwise
purpose of this experiment is
to familiarize you with the concept of an equilibrium reaction.
The ionization of a weak acid or weak base is a typical example of an
equilibrium process. Consider the reversible ionization of the
classic weak acid, acetic acid:
+ H2O H3O+(aq)
reversible arrows tell us
ionization reaction does not go to completion. Sometime after
acetic acid (CH3CO2H or HAc) is mixed with water,
reverse of the ionization process (combination) will begin to occur as
concentrations of the hydronium ion (H3O+ or H
+) and acetate ion (CH3CO2- or
Ac -) increase. At some time, the opposing reactions
will be occurring at the same rate and the concentrations of all
reactants and products will remain constant. Once we have reached
this state of dynamic equilibrium, we can define the equilibrium
dilute aqueous solutions, the
concentration of H2O is essentially constant at 55.5
M. Since it is a
constant, we can rearrange the equilibrium equation and define an new
constant for the ionization of weak acids, Ka:
calculate the ionization
constant, Ka, for acetic acid, it is necessary to
experimentally determine the equilibrium concentrations of H+,
Ac-, and HAc. Method:
Based on the discussion above, if we want
to determine the Ka
for any weak acid (HA), we need to determine the equilibrium
concentration of H+,
A- , and HA. The most straight forward of these is [H+
], because we know that the pH = -log[H+]. So if we
measure the pH of the equilibrium solution, we will not only know the
concentration of the hydrogen ion, [H+], but the
concentration of the weak acid's conjugate, [A-], as
well. As an example, let's assume that the pH of this solution
2.37. This means that the value for both [H+] and [A-]
[H+] = [A-] = 10-pH = 10-2.37
= 4.27x10-3 M
we still need to
determine the equilibrium concentration of HA. Unfortunately,
this is difficult to determine since most methods of analysis will
change the concentration of the HA and cause the equilibrium to
shift. Since we cannot directly determine the [HA], we need to
find the initial concentration of HA. To do this we need to
neutralize all of the HA present by titrating it with a strong base of
known concentration. As the H+ from the weak acid is
neutralized by the strong base, the equilibrium will shift
to the right generating more H+. This process will
continue as the strong base is added until all of the HA has been
converted to H + and A- (equivalence
point). This is no longer
an equilibrium solution, it only contains A-(aq),
Na +(aq), and H2O(l).
For example, if 23.6 mL of 0.321 M NaOH were required to neutralize
mL of the HA solution, then the initial concentration of HA would have
MHA = MNaOH
VNaOH / VHA
MHA = 0.321 x
23.6 / 50.0
MHA = 0.152 M
can calculate the
equilibrium concentration of HA, by subtracting the equilibrium [H+]
concentration from the initial HA concentration:
[HA] = 0.152 - 4.27x10-3
= 0.148 M
have all of the
necessary to calculate the Ka for our weak acid!
But wait! What if we don't know the
concentration of the
we used to titrate the weak acid? No problem...we will
Standardization is a process of comparing an unknown against a known or
In this case, we will titrate a known quantity of standard acid with
unknown base. Using our M1
V1 = M2V2 relationship, we will be
to determine the exact concentration of our base.
with all standardization
procedures, the real problem is picking an appropriate standard.
A primary standard is a substance that is readily available in a pure
form (<0.02% impurities), it is stable, easy to dry, is not
hydroscopic, and should have a fairly high
equivalent weight to minimize the consequences of errors in mass
determination. We are fortunate that such a standard exists for
our situation, the mono potassium salt of the organic di-acid, phthalic
acid (KHC8H4O 4, or KHP, mw = 204.223
example, if we dissolve 1.000
g of KHP in 50 mL of water and titrate this solution with 31.6 mL of
our unknown base, what is the molarity of our base? First we need
to remember that at the equivalence point (where the indicator changes
color); the moles of KHP equal the moles of NaOH:
moles KHP =
moles NaOH =
1.000g / 204.223g/mol = 0.0049 moles
31.6 mL of our base
0.0049 moles, the molarity of our base is:
0.0049 moles / 0.0316 liters = 0.155 M
that we know the
our base, we can titrate our unknown weak acid to determine its initial
and use the pH meter to determine the equilibrium [H +],
and [A-]. With these measurements, it is a simple
matter to calculate the Ka
for any weak acid. [Pssst....there are also other ways of
determining the Ka
for a weak acid, but
that is a story for another day.]
Although they seem simple, many people initially have trouble with
titrations. There is a good deal of eye-hand coordination
involved and a lot of small errors than can creep in to ruin your
experiment. The following are some tips that should help you be
- Make sure you have studied the video, before coming to lab: An
Overview of Titrations
- Add your base to the burette over the sink. If you try to
add it in the burette clamp, some might spill into your acid sample and
you would have to start over.
- Don't forget to put the indicator in your sample.
- Make sure you setup your burette so the tip is below the top of the beaker.
- Set your magnetic stirrer as fast as possible, BUT no
splashing. If any of the acid splashes on the side, you need to
wash it down with distilled water.
- Make sure you have no bubbles in the tip of your burette.
This probably causes 75% of the problems students have with getting
titrations to be reproducible.
- When you see that the pink color start to persist, slow the addition of base to a drop at a time.
- To get the best equivalence points, you will need to 'cut'
drops. Barely open your stopcock and let less than a drop form on
the tip. Then use your distilled water bottle to squirt it into
- Your burette is a direct read delivery burette. This means,
'What You See Is What You Get'. If you start at 0.00mL and stop
at 23.56mL, you used 23.56mL.
- Very important: read all burette readings to 0.01mL!
Remember, you always read to one place past what is marked on the
measuring device. Also make sure you look directly at the burette
at eye level, do not look down or up to read the meniscus, this will
cause parallax errors. The following figure shows an initial
volume of 9.62mL and a final volume of 24.16mL:
Preparing the Sodium Hydroxide
- Clean and dry a 600 mL beaker .
- Weigh out approximately 6
grams (to 0.001 g) of sodium hydroxide pellets directly to the 600 mL
beaker. NOTE: Handle the
sodium hydroxide pellets with care. Sodium
Hydroxide is very hydroscopic and can cause
burns if it comes in contact with your skin. Be sure to use
weighing boats or weighing paper to determine the mass and to deliver
hydroxide pellets to the beaker. When removing the sodium
pellets from the reagent container replace the cap of the container as
as possible. Clean up any spilled sodium hydroxide pellets
- Add approximately 400 mL of
distilled water to the beaker and stir the solution until the
pellets have completely dissolved.
- Put this beaker of sodium
hydroxide on a paper towelS so it does not mare the bench top.
Make sure you immediately clean up any spills.
Titration of NaOH Solution:
- Obtain a 50 mL burrett, close the stopcock and fill it to
the top with distilled water. Open the stopcock and allow all of
the water to drain.
- Close the stopcock and fill your burrett with 50 mL of your NaOH solution (from
above) so that the
solution comes in contact with the entire inner surface of the burrett.
- Open the stopcock and allow all of the NaOH to drain through the tip.
- Fill the burrett to the top with the NaOH. Open the
all the way to flush all bubbles out of the tip. When all bubbles have
been flushed out (it may take several tries), close the stopcock and
refill the burrett.
- Read the bottom of the meniscus and record the initial reading
to the nearest 0.01 mL. The Teflon stopcock should turn smoothly with a
little resistance. If the stopcock is too loose, tighten it a
little, otherwise the solution will leak around the stopcock and the
titration will be for naught. A
leaking burrett is a major cause of error in titrations.
- For more information about the proper use
of a burrett and titration procedures, click on this hyper link: "How
Read A Burrett" or "Everything you wanted to know about titrations but
afraid to ask".
- Use a repipetter to deliver 10.00 mL of the standard HCl
into a clean 150 mL beaker. Wash down the sides of the beaker
the wash bottle. The addition of water at this stage has no
the total amount of acid already present in the beaker. Be sure
to record the concentration of the standard HCl!
- Add 2 drops of phenolphthalein indicator to the acid
solution. The solution should remain colorless.
- Add a magnetic stir bar to the beaker and place on the heating
stir plate. Adjust the stirring rate to obtain a vortex without
of the solution splattering on the sides of the beaker. Avoid spilling any of the beaker contents.
Any loss of sample would render the titration worthless.
- Rinse down the inside of the beaker occasionally and continue,
slowly adding NaOH until the first permanent, faint pink color persists
for at least 30 seconds. At this point the titration is complete
Read the final volume of NaOH and record to the nearest 0.01 mL.
Remember these are 'direct read' burretts, what you see is what you
get! Some of you may have used other burretts where they start at
50 mL and go to zero, so you subtract your reading from 50. DO NOT do that with these burretts!
- Calculate the normality of the NaOH solution.
- Repeat Steps #7-12 with 20.00 mL of standard HCl. Fill the burrett and
read it before each titration if there is a possibility of
before a titration is complete.
- Repeat Steps #7-12 with 30.00 mL of standard HCl.
- The results from these three titrations should agree within
± 0.005 N, otherwise titrate a fourth sample and throw out the
value which does not agree. Average the results and use this
value as the concentration of your STANDARDIZED BASE. Do
not discard the base which you have just standardized or clean the
burrett. This standardized NaOH solution will be
used to determine the concentration of an unknown acid solution in the
of Initial Weak
- Pick one of the of the
unknown weak acids and be sure to record its number.
- Using the repipetter,
20.00 mL of this acid to a clean 100 mL Erlenmeyer flask.
Use distilled water to make sure all of the acid has been rinsed off
sides of the flask.
- Add 2 drops of
phenolphthalein indicator to the flask.
- Put a magnet stirring bar in
the flask and set the stir plate to a moderate rate (avoid splashing).
- Fill your 50 mL burette with
your standardized sodium hydroxide solution.
- Titrate the unknown weak
acid with your standardized sodium hydroxide solution until a faint
pink color remains for 30 seconds.
- Repeat Steps 2 - 6, for two
additional 20 mL aliquots of the same unknown weak acid and average
- Once you are satisfied with
your results you may pour any remaining sodium hydroxide solution down
the sink with copious water.
Hydrogen Ion Concentration:
- Add approximately 20 mL of
the unknown acid to a clean 50 mL beaker.
- Remove the pH electrode from
its buffer solution and rinse it off with distilled water.
- Submerge the pH electrode
into the unknown acid and wait 10-30 seconds for the meter to stabilize.
- Record the pH and the
temperature of the unknown acid solution.
- Sodium hydroxide solutions can etch glass and leave white rings
on the bench top. So, it is important to thoroughly clean any
glassware that has come in contact with the sodium hydroxide.
- There is a special procedure for cleaning your burette.
First, open the stopcock and completely drain the sodium hydroxide from
your burette. Then fill it with Burette rinse (squeeze bottle
with black tape) and allow it to completely drain. Finally, fill
your burette with distilled water and allow it to drain
completely. Be sure to leave the stopcock open after you are done.
- Carefully clean the whole bench top including the area around the
- From your titration data,
calculate the molarity of your sodium hydroxide solution.
- How reproducible were your
results for the molarity?
- From your titration of the
weak acid, calculate the it's initial concentration.
- How reproducible were your
results for this concentration?
- Knowing the initial
concentration of the weak acid and its pH, calculate the Ka
for this acid.
- In the back of your textbook
is a list of Ka values for several weak acids. Keep in
mind that as careful as you were with your titrations there are several
factors outside of your control so your answer may be a factor of
2, 3, even 5 off from the value in the back of the book. So use
some common sense when deciding which unknown acid you had.
- The Ka values in
the back of your textbook have been very accurately determined using a
process similar to the one you have used. One major difference is
water used in the analysis is freshly boiled before use. Why was
this done? Don't tell me to make the water pure, I want to know
- Phenolphthalein was used as
the indicator in this experiment. What is the role of the
indicator? Why was phenolphthalein used and not some other
indicator (see your textbook for other indicators)?
(Updated 8/2/13 by C.R. Snelling)