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Determination of Ka of several weak acids
OBJECTIVE
To determine the acid ionization constant, KC , for acetic acid and an unknown monoprotic acid
by using indicators and by using a pH meter.
CONCEPT TO BE TESTED
The value of Ka , for a weak acid can be determined from the concentrations of the species
present at equilibrium.
TEXT REFERENCES
(1) Whitten, K, W., Davis, R. E., and Peck, L. General Chemistry, 5th ed., Saunders College
Publishing, Philadelphia, 1996, Sections 11.1.–11.3, Chapters 18 and 19 (2) Safety, pp1–3, (3)
Laboratory Techniques Section K and L and Figure L.1.
INTRODUCTION
The equation, H 2 O(l) + HA(aq) H 3 O
(aq) + A
- (aq), describes the equilibrium conditions
when a weak acid, HA, dissolves in water and ionizes to form H 3 O
and K ions. The ratio of the
algebraic product of the concentrations of the ions to the concentration of the unionized acid is
constant and is given the symbol, Kc.
HA
H O A
K
3 A
(Eq. 1)
K (^) c applies to equilibrium conditions of the ionization reaction. If the acid is relatively strong, the
extent of ionization is reflected in a relatively large value for Kc. A large value of Kc is the result
of large concentrations of H 3 O
and A
- ions divided by the smaller concentration of unionized
HA molecules, A weaker acid has smaller concentrations of H 3 O
and A ions and a larger
concentration of HA molecules. Its K (^) c value will be smaller. The values of Kc for different acids
tell us about their relative strengths as acids.
There are many important weak acids. For example, vinegar is a dilute solution of acetic acid,
CH 3 COOH. Carbonic acid, H 2 CO 3 , is the weak acid formed by CO 2 and H 2 O in carbonated bever‐
ages, in blood, and in many other systems.
The value of Kc, for an acid can be calculated from the measured values of the concentration of
the ionized and un‐ionized species present in solution at equilibrium.
Example 1: The equilibrium concentrations for a given acid were found to be as follows.
HA H 3 O
Equil. conc. 1.0 x 10
- M 1.3 x 10 - M 1.3 x 10 - M
- 3 x 10 1.3x 10
HA
H O A
K
K (^) a = 1.7 x 10
The [H 3 O
] in an aqueous solution can be expressed as pH, which is defined as
pH = ‐log [H 3 O
] (Eq. 2)
Equilibrium constants can also be expressed in a similar way.
pK (^) a = ‐ [log Ka ] (Eq. 3)
By taking the negative of the common logarithm of Equation 1, a useful relationship between
pK (^) a and pH called the Henderson‐Hasselbach equation can be derived.
HA
A log a
pH pK
(Eq. 4)
If a solution of a weak acid is titrated to its endpoint, the anion, A
- , species is present in the
form of a salt. If this solution and the acid are mixed to give a 50:50 mixture then [HA] = [A
solution that contains the weak acid, HA, and the anion, A
- , of this weak acid from its salt is a
buffer system. In a buffer solution in which [A
- ] = [HA], the measured pH of the solution equals
to the pK (^) a value. These conditions offer a convenient way to determine the pK (^) a of an acid from
pH measurements. K (^) a , is determined from pK (^) a (Eq. 3).
A. The Use of a pH Meter to Measure pH (See webpage for electrode calibration instructions)
A pH meter gives more accurate measurements of pH than do indicators, but it uses more
expensive equipment and must be calibrated, pH meters are generally similar in their operation
Your instructor will provide some instructions for the use of the computer based pH system.
A pH meter consists of a voltage measuring device and two electrodes. A reference electrode,
usually a calomel electrode, provides a constant potential while glass electrode generates a
potential that is proportional to the pH of the solution. When the two electrodes are placed in a
solution, the voltage is displayed on a meter that is calibrated in pH units.
The pH meter is calibrated against a standard solution of known pH before use. Care must be
used in handing the electrodes. The thin membrane that is permeable to H 3 O
ions is very
C. Determination of Ka Values for an Unknown Monoprotic Acid Using a pH Meter
Step 8. Repeat Part B Steps 2 through 7 using one of the unknown weak monoprotic acid
solutions as you did with acetic acid.
D. Data Analysis
Step 9. Plot your data as pH (vertical axis) versus mL of NaOH solution added
Step 10. From your graphs, determine the exact volume of standard NaOH solution
required to reach the equivalence point and then calculate the concentration of
the acetic acid in your original solution. (Remember: MAVA = MBV (^) B)
Step 11. From your graphs, determine the value of Ka for CH 3 COOH and for the unknown
acid. (Remember: the pK (^) a = pH at the half‐way point of the titration.)
