NORTH ALLEGHEHY SENIOR HIGH SCHOOL
AP Chemistry
DETERMINATION OF THE EQUIVALENT MASS AND pKa OF AN UNKNOWN ACID
INTRODUCTION
In this experiment you will determine the equivalent mass of an unknown acid, that is, the mass of the acid
that supplies one mole of hydrogen ions. The acid, a solid crystalline substance, will be weighed out and
titrated with a standard solution of sodium hydroxide. From the moles of base used and the mass of the acid,
you will be able to determine the equivalent mass of the acid. Next you will plot the titration curve of the
acid, with pH on the vertical axis and the volume of NaOH on the horizontal axis. From this graph you will
be able to determine the value of the equilibrium constant for the dissociation of the acid.
Acids are substances that contain ionizable hydrogen atoms within the molecule. Strong acids ionize totally,
weak acids partially. The value of Ka, the equilibrium constant for the dissociation of the acid, is an indication
of the strength of the acid. We can also speak of the pKa, the -log(Ka), as an indication of acid strength.
An acid may contain one or more ionizable hydrogen atoms in the molecule. The equivalent mass of an
acid is the mass that provides one mole of hydrogen ions. It can be calculated from the molar mass
divided by the number of ionizable hydrogen atoms in a molecule. For example: hydrochloric acid, HCl, contains one ionizable hydrogen atom; the molar mass is 36.45 g/mole; the equivalent mass is also 36.45 g/mole.
Sulfuric acid, H2SO4, contains 2 ionizable hydrogen atoms; the molar mass is 98.07 g/mole; the equivalent
mass is 49.04 g/mole. Thus, 36.45 g of HCl or 49.04 g of H2SO4 would provide you with one mole of H+
ions.
The equivalent mass may be determined by titrating an acid with a standard solution of NaOH. Since one
mole of NaOH will react with one mole of hydrogen ion, at the equivalence point the following relation
holds:
Vbase x Mbase = moles base used = moles H+ present in acid
Equiv. Massacid = grams acid/moles H+
The concentration of the NaOH solution must be accurately known. To "standardize" the NaOH, that is to
find its exact molarity, the NaOH is titrated against a solid acid, potassium hydrogen phthalate, sometimes
abbreviated KHP. The KHP is chosen because it is easily dried and weighed, and has a relatively high equivalent mass. The formula of KHP is:
(KHC8H4O4)
It contains one ionizable H+. The titration can be followed using phenolphthalein as an indicator.
A graph of pH vs. mL of NaOH added can be drawn by carefully following the titration with a pH meter.
There should be a significant change in pH in the vicinity of the equivalence point. Note that the equivalence
point will probably NOT be at pH 7, but will be on the basic side. The value of the equilibrium constant for
the dissociation of the acid can be obtained from the graph.
If we represent the dissociation of the acid as:
HA + H2O
H3O+ + A-
the equilibrium expression is:
Ka = [H3O+][A-]/[HA]
When the acid is half-neutralized, [HA] = [A-], so these terms cancel in the above equation, and Ka = [H3O+].
Therefore, when the acid is half-neutralized, the pH = pKa. The point where pH = pKa can be found from
the graph below.
Figure 1 : pH during titration of a monoprotic acid with NaOH where:
A = Volume of NaOH at equivalence point
B = Volume at half-neutralization
C = pH when half-neutralized or pKa
PROCEDURE
Part 1: Preparation and Standardization of Approximately 0.12 M NaOH
Since solid NaOH rapidly absorbs both H2O and CO2, a solution of exact molarity cannot be
prepared by weighing the solid and diluting it to volume. Instead, you must prepare a solution of
approximately the desired concentration, and find its exact molarity by titrating it against a standard
substance.
In an effort to eliminate waste, each 0.12 M NaOH solution below will be shared by TWO GROUPS.
Thus, pick a partner group to make your base solution and share that solution throughout the lab.
You will share your standardization data as well.
a. Obtain a 1 liter bottle and rinse with distilled water.
b. Measure approximately 18 mL of 6 M NaOH and dilute it to about 1 L (near the rounded top of the
one liter plastic bottle). Mix well. This will make approximately a 0.12 M NaOH solution.
c. Using a milligram balance, obtain between 0.4 and 0.6 g of dry potassium hydrogen phthalate (KHP)
using wax weighing paper. Wash the KHP into a 125 mL Erlenmeyer Flask using a wash bottle. Add
about 40 mL of distilled water and swirl until dissolved.
d. Obtain a buret. Rinse it with tap water and then at least twice with portions of your NaOH solution.
Allow some of the rinse NaOH to run through the tip.
e. Fill the buret with your NaOH solution. Open the stopcock briefly to make sure the NaOH fills the
tip.
f. Record the initial volume of the NaOH on the buret (for this part it doesn’t have to start at 0 mL.
Just make sure you read the beginning volume. These burets should be read to the 0.01 mL).
g. Add 3 drops phenolphthalein to the acid in the flask and then titrate with the NaOH until the first
trace of pink color persists for 30 seconds.
Tips: 1 – Place a sheet of white paper underneath the flask to help see the pink color.
2 – Constantly swirl the flask while adding the NaOH.
3 – Using a tiny bit of water, rinse the sides of the flask to ensure all of the NaOH
makes it into the flask.
h. Record final volume on buret.
i. Complete a second trial, repeating steps c-h. If you use slightly more acid on the second trial, this trial
should be more rapid since you will know approximately how much NaOH you can safely add before
you get to the endpoint.
j. Trade data with the lab group sharing your NaOH.
Part 2: Determining the pKa of the Unknown Acid using a pH meter.
a.
b.
c.
d.
Weigh a 0.3 – 0.4 g sample of the unknown acid.
Using a 250 mL Erlenmeyer Flask, dissolve the sample in 100 mL of distilled water.
Add 3 drops phenolphthalein.
Using the setup described in class, submerge the pH electrode in your solution and record the initial
pH of the solution.
e. Make sure your buret volume begins at 0 mL.
f. Titrate your unknown acid with the standard base, recording the volume of base added and the pH of
the system after the addition. Please use the following tips:
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



At the beginning of the titration, the pH won’t change much (buffer region). We will begin by
adding three or four mL of base at a time, recording the pH after each addition (record base
added to 0.01 mL from your buret).
When the pink starts to appear for longer periods of time, you should decrease the base
increments to one or two mL at a time.
We are trying to reach the endpoint using the phenolphthalein, so as the pink lasts longer and
longer, add less and less base, to the point where you will be adding it dropwise. Continue to
record pH.
When you are beyond the equivalence point, resume adding base every one or two mL at a
time.
Get at least three data points at a pH above 10 to smooth out your curve.
Part 3: Getting more data for the equivalence point - no pH meter (if time allows).
a. Taking about the same amount of unknown acid in part two, dissolve your acid in a 125 mL
Erlenmeyer Flask and add 3 drops of phenolphthalein.
b. Titrate to the phenolphthalein endpoint.
c. Share this data with your lab group partners (if you get a chance to complete this).
CALCULATIONS
1. Determine the concentration of the standardized NaOH used in the titration by the average of the
concentrations calculated from at least four trials.
2. Determine the equivalent mass (molar mass) of your unknown acid. If you were able to do part 3 this
data should be included in this part.
3. Using your titration graph, determine the pKa of your unknown acid (show this on the graph).
4. Based upon the results of #2 and #3 above, determine the identity of your unknown acid from the
list provided. Please explain your choice.
QUESTIONS
1. What is equivalent mass? How is equivalent mass and molar mass related for a monoprotic acid?
Diprotic acid?
2. Why must the KHP and the acid samples be dried before use? If the KHP was not dried before the
standardization of NaOH, how would it affect the calculated value of the concentration of NaOH?
3. Why is the equivalence point of a weak acid not at a pH = 7?
4. How accurately can you read a buret? Why is it better to titrate with a volume of base that fills most
of your buret?