Weak Acid Titration
Introduction
In this lab you will be titrating a weak acid (C2H4O2 – acetic acid), with a strong base
NaOH, while recording the pH. From the collected data a titration curve will be plotted.
You will compare this titration curve with the one you previously generated in the Acid
Base Titration Lab (strong acid / strong base).
Most substances that are acidic in water are actually weak acids. Because weak acids
dissociate only partially in aqueous solution, an equilibrium is formed between the acid
and its ions. The ionization equilibrium is given by:
HX(aq) ↔ H+(aq) + X-(aq)
where X- is the conjugate base.
The equilibrium constant is then:
Ka= { [H+][X-] / [HX] }
the smaller the value for Ka, the weaker the acid. Weaker acids dissociate less ([H+] is
smaller compared to [HX]) and therefore have a less drastic effect on pH.
On the other hand, strong acids such as HCl dissociate almost completely in water – as
was discussed in the previous Acid Base Titration Lab.
For the titration plot the graph of pH versus volume of base added. Locate and label the
end point. This is assumed to correspond to the mid-point of the vertical portion of the
curve, where pH is increasing rapidly. From the graph read the volume of base need to
the reach the end point.
There are a number of differences between the titration curves for a strong acid versus the
weak acid.
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The weak-acid solution has a higher initial pH.
The pH rises more rapidly at the start, but less rapidly near the end point.
The pH at the equivalence point does not equal 7.00 (pH > 7.00) for the weak acid
titration.
Note these differences when comparing the weak acid (C2H4O2) / strong base NaOH,
titration against the titration curve you previously generated in the Acid Base Titration
Lab (strong acid / strong base).
Degree of Ionization:
The degree of ionization is the fraction of molecules of a weak electrolyte that reacts with
water to give ions. We are interested the fraction of [H+] relative the initial concentration
of acetic acid. You will be asked to calculate this value for your titration.
Common Ion Effect:
The common ion effect is the shift in an ionic equilibrium produced by the addition of a
solute that provides an ion common to the equilibrium. This is fundamental Le Chatlier’s
principle in action. You will perform a series of calculations addressing the common ion
effect.
Weak Acid Titration
Procedure
**NOTE: Use instructions provide by the instructor – DO NOT USE INSTRUCTIONS
IN CHEMLAB! To remove the instructions on the screen, and free-up more working
area, perform the following operation: click on the OPTIONS tab; then click on LAB
ONLY. The instructions “disappear” and all of the area is now lab space.**
Titration of weak acid:
1. Obtain a 100 ml beaker from the equipment menu.
2. Select the beaker and add 30 ml of 0.100 M C2H4O2 (acetic acid) solution.
3. With the beaker selected, add the pH meter (from the equipment menu). Record this
information in Observations.
4. Turn on the collection of titration data for the beaker by selecting the "collect titration
data" menu from the procedures menu. Note only one set of data can be collected per
lab so you will be asked to drop the previous data set.
5. Add a 50 ml burette to the lab.
6. Fill the burette with 50 ml of 0.100 M NaOH solution.
7. Open the Titration data window by selecting the "View titration data" from the
Procedures menu.
8. Add NaOH, quickly at first (1ml increments) then slowly (.1ml increments) as pH
starts to rise more rapidly.
9. Continue adding slowly until the pH is observed to level off then add NaOH in larger
increments again.
10. Continue adding NaOH until an excess of 20ml have been added after the leveling
off.
11. Locate the end point of the titration (ml of base added and pH).
12. Add a title to the Titration curve.
13. Copy the Titration curve from the Titration data window to Word. Print and include
with your report.
Weak Acid Titration
Observations
Name:__________
Sect:____________
Data
Initial pH of acetic acid in water =
__________
From the end point, (mid point of the steepest part of titration curve), determine the
Volume (ml) of base (NaOH) added:
__________mL
Does the volume of base delivered to neutralize the acid make sense? Provide an
explanation.
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
______________________________________________________________________
Calculations
1. The degree of ionization is defined as the fraction of molecules of a weak electrolyte
that reacts with water to give ions. Calculate the degree of ionization (amount of [H+]
generated from the dissociation of acetic acid in water) of acetic acid in your initial
30mL sample of 0.100 M C2H4O2 solution, where:
degree of ionization of acetic acid = [H+ due to dissociation of acetic acid in
water] / [initial concentration acetic acid]
[H+ due to dissociation of acetic acid in water] = antilog[-pH]
so,
degree of ionization of acetic acid = antilog[-pH] / [0.100 M] C2H4O2
________[M]
2. Calculate the expected pH of the following solutions - use the provided example for
guidance:
Example:
0.10 M H3PO4 with Ka1 = 6.9 X 10-3
Solution: H2O + H3PO4 ↔ H3O+ + H2PO4
Ka1 = {[H3O+][H2PO4]}/[H3PO4] = X2/(0.10 – X) ≈ X2/0.10
X = [H3O+] = 2.6 X 10-2 M
pH = -log[H3O+] = 1.59
2-A
0.10 M NH3 with kb = 1.8 X 10-5
pH = _______
2-B
0.10 M NH4NO3 with ka = kw/kb (NH3)
pH = _______
3. The common-ion effect is defined as the shift in an ionic equilibrium produced by the
addition of a solute that provides an ion common to the equilibrium. Compare the
following two solutions:
Solution 1: 0.100 M C2H4O2
Solution 2: 5mL 0.100 M C2H4O2 + 5mL 0.100 M HCl
The [H3O+] ions from the strong acid in solution 2 force the equilibrium C2H4O2 ↔ H2O
+ C2H3O2- almost completely to the left. As a result, the pH is determined by the H3O+
ions from the strong acid. Calculate the pH of Solution 1 (ka = 1.738 x 10-5) and 2
(assume HCl dominate for calculation)
pH Solution 1= _______
pH Solution 2= _______
4. Examine the common-ion effect for 50mL of 0.100 M NH3 + 50mL of 0.100 M
NH4NO3. The solution contains a substantial amount of NH4+, an ion common the
the equilibrium NH3 + H2O ↔ NH4+ + OH-. As a result, the equilibrium is shifted to
the left, an the pH is less than it would be in the absence of NH4+ ions. Calculate the
pH of this solution (kb = 1.8 X 10-5).
pH = _______
Discussion Write a comprehensive, grammatically correct, paragraph addressing the
differences between the titration curves for a strong acid (Acid Base Titration Lab) versus
the weak acid.