TVS AP Chemistry Lab
Determination of the Ka of a Weak Acid
Introduction
Understanding the chemistry of acids and bases is critical to gaining a full comprehension
of equilibrium. As we have seen, any acid base reaction is actually a competition
between two acids or two bases, with the more powerful acid (or base) driving the
position of equilibrium. For example, consider the reaction:
HCl(aq) + OH-(aq)  H2O(l) + Cl-(aq) K=1 x 1014
Since the K value is so large, it is evident that HCl is a more powerful acid than H2O.
Similarly, OH- must be a better base the Cl-.
Since the relative donating ability of acids will ultimately determine the equilibrium
constant for an acid base reaction, we need a quantitative method to determine how
powerful an acid is. Traditionally, the way this is done is to compare several different
acids to a common base, with that common base being water. Thus, if one acid is a
better donator to water than another, it is a more potent (or stronger) acid.
The chemical reaction occurring when the donation occurs is:
HA(aq) + H2O(l)  H3O+(aq) + AHere, HA is a generic acid (i.e. A could be Cl, C2H3O2, etc….). From what we have
learned from equilibrium, the equilibrium constant expression for the reaction would be:
[H O+ ][A" ]
K= 3
[HA]
Remember that since water is a pure liquid of constant concentration, it does not appear
in the equilibrium constant expression. This constant is called the acid dissociation
constant and is given the label Ka. The larger the value of this constant, the better donator
! measure this constant, we need to determine the
the acid is. Now, in order to
concentration of the hydronium ion, H3O+. From the stoichiometry of the reaction, the
concentration of hydronium is equal to the concentration of A-. Thus, if we know these
concentrations and the initial concentration of the acid, the value of Kq is easily
determined.
Traditionally, there are two methods to determine the concentration of the hydronium ion.
One procedure involves the use of an indicator, which is a substance that turns various
colors when exposed to different H3O+ concentrations. There are a literally hundreds of
different indicators one can choose, but the problem with all of them is that a definite
color change usually occurs as the concentration of hydronium changes by a factor of
100. Therefore the indicator method is a little too inexact to use to obtain a precise value
of Ka. In contrast, we will use a pH meter, which is a device that uses redox reactions to
precisely measure the hydronium ion concentration of a solution to at least two
significant figures.
This lab has three parts and will be completed as a class effort. In the first part of the lab,
we will examine how a small amount of strong acid can drastically change the pH of a
solution of distilled water. Then, each lab group will prepare an acetic acid solution of
known concentration and use the pH meter to measure the amount of H3O+ in solution.
Then, commercial vinegar will be tested, and the pH measured will be used to determine
the concentration of that solution. Finally, A known amount of strong acid (HCl) will be
added to an acetic acid solution, and the pH calculated.
Procedure:
A) Illustrating the effects of Strong Acids on the pH of distilled water.
To see how the pH meter is operated, take a large test tube and add 20 ml of distilled
water to it. Observe how the instructor calibrates the pH meter measure the pH of the
distilled water solution. Then add 1 drop of a strong acid solution and observe the new
pH of the solution. Record both pH’s.
B) Determining the Ka of acetic acid
Each lab group will have been assigned a concentration of acetic acid solution to prepare
from a stock 2.00 M solution. Be sure in your procedure that you have outlined exactly
how much of the 2.00 M solution you will need to dilute to 100 ml to make the solution
of desired concentration (this will be checked before the lab and is a significant portion of
the procedure grade). Go to the buret filled with the 2.00 M soln. and place the amount
needed to make your soln. into a 100 ml volumetric flask. Using distilled water and a
pipet, fill the flask to the mark and mix well. After all groups have made their solutions,
fill a test tube approximately 1/4 full of solution and measure the pH. Create a table of
pH for each concentration of solution made by the class.
C) Determination of the Concentration of Commercial Vinegar
Commercial vinegar is typically sold as a 5% soln of acetic acid. Fill a test tube 1/4 full
of a vinegar solution and measure its pH. Record this pH in the data table.
D) Observing the effect of a strong acid on the dissociation of a weak acid.
As a class, take one of the solutions made in part B of the experiment and measure out
10.0 ml in a graduated cylinder (it is necessary to measure the pH of the soln?). Add 10.0
ml of standardized HCl solution and measure the pH.
Calculations:
A) In Part A you want to show two calculations. First, calculate the concentration of
H3O+ (hereafter referred to as H+) in the distilled water. Then calculate the
concentration of H+ after adding the drop of strong acid. Assuming that the total
volume of the solution doesn’t change by the addition of the one drop of acid,
how many moles of H+ were contained in the one drop.
B) In Part B, you want to calculate the Ka of acetic acid. From the pH of the solution
you should be able to calculate both the concentration of acetate and acetic acid at
equilibrium. Then use the definition of Ka to obtain its value. You must show the
calculation for your individual solution that you have made, but you can just
report the other Ka’s obtained from other trials. Average these results and
compare to the literature value with a percent error calculation.
You should show the calculation of the percent ionization of each trial (again just
show the calculation for your trial; report the rest) Recall that the definition of
percent ionization is:
[H + ]
%ion =
x100
[HA]0
C) Use the average Ka value obtained in part B to calculate the concentration of
commercial vinegar from the pH value. Commercial vinegar is stated to be 5.0 %
be weight, meaning that 100. g of vinegar solution contains 5.0 g of acetic acid.
Assuming the density!of such a solution is 1.00 g/ml, calculate the expected
concentration acetic acid in vinegar in mol/L. Finish this section with a percent
error calculation.
D) Based on the listed concentration of the HCl solution, calculate its expected pH
when it is diluted from 10.0 ml to 20.0 ml. Compare this pH to measured pH of
the acetic acid-HCl mixture. In addition, use RICE to calculate the percent
ionization of the acetic acid in the mixture, and compare this value to the one
obtained in part B.
Theory:
Your theory section should address the following questions in paragraph format:
1) Explain why the percent ionization of the acetic acid increased as the initial
concentration of the acid decreased.
2) Does one have to consider the donating ability of a weak acid when it is mixed
with a strong acid? Explain
3) What is one thing you could do to change the Ka value of acetic acid? Could this
factor have any bearing on your percent error in part B of the lab? Explain
Sources of Error: None
Complete the Lab with a standard conclusion stating the value of Ka obtained for
the acid, with the percent error.