DETERMINATION OF Ka OF WEAK ACIDS
Introduction
Acids are substances which provide hydrogen ions. They vary in their ability to ionize from strong
acids, which give up 100% of their hydrogen ions, to extremely weak acids which hardly ionize at all.
The dissociation constant is the value of the equilibrium constant which indicates their acid strength.
Background
The modern Bronsted-Lowry definition of an acid relies on the ability of the compound to donate
hydrogen ions to other substances. When an acid dissolves in water, it donates hydrogen ions to water
molecules to form H3O+ ions. The general form of this reaction, call the ionization reaction, is shown in
Equation One, where HA is the acid and A- its conjugate base after loss of a hydrogen ions. The double
arrow represents a reversible reaction.
HA(aq) + H2O(l) ↔ A-(aq) + H3O+
Equation One
The equilibrium constant expression, Ka, for the reversible ionization of an acid is given in Equation
Two. The square brackets refer to the molar concentrations of the reactants and products.
Equation Two
Not all acids, of course, are equal. The strength of an acid depends on the value of its equilibrium
constant, Ka for Equation One. Strong acids ionize completely in an aqueous solution. The value of Ka
for strong acids is extremely large and Equation One essentially goes to completion – only H3O+ and Aare present in solution. Weak acids, in contrast, ionize only partially in aqueous solutions. The value of
Ka for a weak acid is much less than one and Equation One is reversible – all species (HA, A-, and
H3O+) are present at equilibrium.
Polyprotic acids contain more than one ionizable hydrogen. Ionization of a polyprotic acid occurs in
a stepwise manner, where each step is characterized by its own equilibrium constant, Ka1, Ka2, etc.)The
second reaction (removal of the second acidic hydrogen) always occurs to a much smaller extent that the
first reaction, and so Ka2 is always significantly smaller than the Ka1. Sulfuric acid (H2SO4) and
phosphoric acid (H3PO4) are examples of polyprotic acids.
H2A(aq) + H2O(l) ↔ HA-(aq) + H3O+
Equation Three
HA-(aq) + H2O(l) ↔ A-2(aq) + H3O+
Acid
Iodic
Sulfurous
Acetic
Carbonic
Hypochlorous
Hydrocyanic
Formula
HIO3
H2SO3
HC2H3O2
H2CO3
HClO
HCN
Ka1
1.7 x 10-1
1.7 x 10-2
1.8 x 10-5
4.3 x 10-7
3.0 x 10-8
4.9 x 10-10
Equation Four
Ka2
6.4 x 10-8
5.6 x 10-11
pKa1
0.77
1.77
4.74
6.37
7.52
9.31
pKa2
7.19
10.25
The ionization constant of a weak acid can be determined experimentally by measuring the H3O+
concentration in a dilute aqueous solution of the weak acid. This procedure is most accurate when the
solution contains equal molar amounts of the weak acid and its conjugate base. Consider acetic acid as
an example. Acetic acid, CH3COOH or HC2H3O2, and the acetate ion C2H3O2-, represents a conjugate
acid-base pair. The equilibrium constant expression for ionization of acetic acid is shown in Equation
Five. If the concentrations for acetic acid and acetate ion are equal, then these two terms cancel out in
the equilibrium constant expression, and Equation Five reduces to Equation Six.
CH3COOH(aq) + H2O(l) ↔ CH3COO-(aq) + H3O+(aq)
Ka = [CH3COO-][ H3O+]
[CH3COOH]
Equation Five
Ka = [ H3O+]
Equation Six
In this experiment, solutions are prepared in which the molar concentrations of an unknown acid and
its conjugate base are equal. The pH of these solutions is then equal to the pKa of the acids. The
definition of pKa is closely related to that of pH. Thus, pH = -log[H3O+] and pKa = -logKa. Most of the
unknowns are slats of polyprotic acids that still contain an ionizable hydrogen. Sodium bisulfate
(NaHSO4), for example, is a weak acid salt; it contains Na+ and HSO4- ions. The HSO4- ion is a weak
acid – the equilibrium constant for ionization of HSO4- corresponds to Ka2 for sulfuric acid.
H2SO4(aq) + H2O(l) ↔ HSO4-(aq) + H3O+(aq)
Ka = [HSO4-] H3O+]
Equation Seven
[H2SO4]
HSO4-(aq) + H2O(l) ↔ SO4-2(aq) + H3O+(aq)
Ka = [SO4-2] H3O+]
[HSO4-]
Equation Eight
Experiment Overview
The purpose of this experiment is to determine the pKa values of ionization of two unknown weak
acids. Solutions containing equal molar amounts of the weak acids and their conjugate bases are
prepared by “half-neutralization” of the acid. Their pH values are measured and used to calculate the
pKa value for the unknowns and thus determine their identities. Two trials are run for each unknown
weak acid.
Pre-Lab Questions
Phosphoric acid is a triprotic acid (three ionizable hydrogens). The values of its stepwise ionization
constants are Ka1 = 7.5 x 10-3, Ka2 = 6.2 x 10-8, and Ka3 = 4.2 x 10-13.
1. Write the chemical equation for the first ionization reaction of phosphoric acid with water.
2. Write the equilibrium constant expression (Ka1) for this reaction.
3. What would be the pH of a solution when [H3PO4] = [H2PO4-]
4. Phenolphthalein would not be an appropriate indicator to use to determine Ka1 for phosphoric
acid. Why not? Choose a suitable indicator from the following color chart.
Indicator
1
2
3
4
pH
5
6
Phenolphthalein
C
C
C
C
C
C
C
Methyl Red
R
R
R
R
O
O
Y
Y
Orange IV
O
O
P
P
Y
Y
Y
Y
7
8
9
10
11
R
R
Y
Y
Y
Y
Y
Y
Pink Pink
C = colorless, R = red, O = orange, Y = yellow, P = peach
5. What would be the pH of a solution prepared by combining equal quantities of NaH2PO4 and
Na2HPO4? Explain with an equation.
6. Sufficient strong acid is added to a solution containing NaHPO4 to neutralize one-half of it.
What will be the pH of this solution? Explain.
Materials
Phenolphthalein, 1.0%, 0.10M NaOH, Unknown weak acids A – E (about 0.5g each), DI water,
electronic balance, 250mL beaker, 250mL Erlenmeyer flask, 100mL graduated cylinder, pH meter,
stirring rod, wash bottle with DI water, weigh boats, pipets.
Safety
Acids and bases are skin and eye irritants. Wear goggles. Inform your teacher and clean up all acid and
base spills immediately with the appropriate neutralizing agent. Phenolphthalein is an alcohol based
indicator and thus flammable. Keep the solution away from the flames. Wash hands thoroughly with
soap and water before leaving the lab.
Procedure
1. Obtain an unknown weak acid and record the code of the unknown in your logbook data table.
2. Using one of your numbered weigh boats, measure out a small quantity (0.15 to 0.20g) of the
unknown in each weighing dish. It is NOT necessary to know the exact mass of each sample.
3. Using the graduated cylinder, precisely measure out 50.0mL deionized water in 250mL beaker.
4. Transfer the sample from weigh boat #1 (Trial 1) to the water in the beaker and stir to dissolve.
5. Using a graduated cylinder, precisely transfer 25.0mL of the acid solution prepared in step 4 to
an Erlenmeyer flask.
6. Add three drops of phenolphthalein solution to the acid solution in the Erlenmeyer flask.
7. Using the Beral pipet, add sodium hydroxide dropwise to the flask. Gently swirl the flask while
adding the sodium hydroxide solution.
8. Continue adding the sodium hydroxide dropwise and swirling the solution until a faint pink color
persists throughout the solution for at least five seconds. This is called the endpoint. Note: A
pink color develops immediately after the base is added, but fades quickly once the solution is
swirled. When nearing the endpoint, the pink color begins to fade more slowly. Proceed
cautiously when nearing the endpoint so as to not overshoot it and have to start over.
Note: At this point the solution in the beaker contains exactly one half of the original amount of the
acid, essentially all of which is in the acid form, HA. The Erlenmeyer flask contains an equal amount
of the conjugate base A- obtained by neutralization.
9. Pour the contents of the Erlenmeyer flask back into the beaker. Pour the solution back and forth
a few times to mix. Note: It is important to transfer all of the solution from the Erlenmeyer flask
back into the beaker.
10. Using a pH meter, measure the pH of the solution in the beaker, which now contains equal molar
amounts of the acid and its conjugate base. Record the pH in the logbook data table.
11. Dispose of the beaker contents by pouring down the drain with the water running and rinse both
the beaker and the Erlenmeyer flask with distilled water. Dry the beaker with a paper towel.
12. Repeat steps 3-11 using the sample from weighing dish #2 (Trial 2).
13. Repeat steps 1-12 using a second, different unknown weak acid.
Data
Should include the headings of Unknown Acid Code, Trial #, pH
Calculations and Analysis
1. For each unknown, average the pH readings for Trials #1 and #2 and calculate the average pKa
value for each unknown acid. Show your work. Make a Calculations Table with average pH,
pKa, and identity of unknown.
2. Comment on the precision (reproducibility) of the pKa determinations. Describe sources of
experimental error and their likely effect on the measured pKa (pH) values.
3. The following table lists the identities of the unknowns in this experiment. Calculate the pKa
value for each acid in the table. Show your work.
Weak Acid
Potassium hydrogen sulfate
Formula
KHSO4
Ka
Ka2 of H2SO4 = 1.0 x10-2
Potassium hydrogen phthalate
KHC8H4O4
Ka2 of H2C8H4O4 = 3.9 x10-6
Potassium hydrogen tartrate
KHC4H4O6
Ka2 of H2C4H4O6 = 4.6 x10-5
pKa
4. Compare the experimental pKa with the literature values reported in Question 3. Determine the
probable identity of each unknown you uses and enter your answers in the Calculation table.
5. Write a chemical equation for the ionization of each weak acid in the list of unknowns
(Questions #3).
6. Why is it not necessary to know the exact mass of each acid sample (step 2 in the procedure)?
7. Why is it not necessary to know the exact concentration of the sodium hydroxide solution used in
step 7 of the procedure?
8. Why is it necessary to measure the exact volume of the deionized water used to dissolve the acid
(step 3), as well as the exact volume of solution transferred from the beaker to the Erlenmeyer
flask (step 5)?
LAB REPORT
•
•
•
•
•
•
Title page
Prelab
Data Table
Calculations Table
Calculations 1-8 – show your work!
One lab report – everyone submits their own calculations table and section…..
Lab report is due__________________________________ initial the parts you did.