UNIVERSITI TUNKU ABDUL RAHMAN
Faculty of Science
Bachelor of Science (Hons) Chemistry
Year 1 Trimester 1
UDEC1224 – Chemistry Laboratory II
Name : Chang Ling Ying (1901683),
Kirthinee Jegatheesan (1902731)
Experiment : Experiment 4
Title : Verification of the Benderson-Hasselbalch Equation
using Buffer Solutions
Date of Experiment : 15 July 2019
Date of Submission : 22 July 2019
Lecturer : Dr. Ooi Mei Lee
Practical group : P (3)
Title
Verification of the Benderson-Hasselbalch Equation using Buffer Solutions
Objectives
To study the validity of Henderson-Hasselbalch equation to a buffer system.
To determine the buffer capacity of a buffer system.
Introduction
Buffer solutions are one of the most important chemical reagents used in chemical
researches. Their usefulness stems mostly from their ability to resist changes in pH. Buffer
solutions are usually formed by the combination of a weak acid with its anion or a weak base
with its cation, that is a conjugate acid-base pair. The Henderson-Hasselbalch equation,
which was developed by the American biological chemist L. J. Henderson and the Swedish
physiologist K. A. Hasselbalch, is useful for estimating the pH of a buffer solution and
finding the equilibrium pH in an acid-base reaction. The equation can be used to determine
the amount of acid and conjugate base needed to make a buffer solution of a certain pH. The
formula for the Henderson-Hasselbalch equation is given by:
pH= p K a + log ¿ ¿
pKa = acid dissociation constant
[A-] = concentration of the conjugate base
[HA] = concentration of the starting acid
In this experiment, we are going to determine the validity of this equation and thus,
calculating the buffer capacity of a buffer system using this equation.
Materials
Standard buffer solution (pH = 4.00), 1 M aqueous sodium hydroxide, 1 M aqueous acetic
acid
Glasswares and Apparatus
Burettes, 100 cm3 beaker, pH meter
Experimental Procedures
1. The pH meter is standardized with the standard buffer solution, pH = 4.00.
2. 50 cm3 of 1 M acetic acid is pipetted into a 100 cm 3 beaker and 1 cm3 of 1 M sodium
hydroxide is then added using a pipette.
3. The solution is mixed thoroughly and the pH of this solution is measured.
4. The pH measurements are repeated for further 1 cm3 addition each up to a total of 5
cm3 of 1 M sodium hydroxide.
5. The pH of the solution is then continued to record for further 5 cm3 additions each up
to a total of 45 cm3 of 1 M sodium hydroxide.
6. The pH measurements are repeated again for further 1 cm3 addition each up to a total
of 50 cm3 of 1 M sodium hydroxide.
Results
Volume of 1M Volume of 1M [SALT] / [ACID] log10 ( [SALT] /
CH3COOH (aq) (x) NaOH (aq) (y) pH = y / (x-y) [ACID] )
50 1 2.37 0.0204 -1.6902
50 2 3.05 0.0417 -1.3802
50 3 3.33 0.0638 -1.1950
50 4 3.54 0.0870 -1.0607
50 5 3.54 0.1111 -0.9542
50 10 3.92 0.2500 -0.6021
50 15 4.18 0.4286 -0.3680
50 20 4.37 0.6667 -0.1761
50 25 4.56 1.0000 0
50 30 4.75 1.5000 0.1761
50 35 4.97 2.3333 0.3680
50 40 5.23 4.0000 0.6021
50 45 5.66 9.0000 0.9542
50 46 5.79 11.5000 1.0607
50 47 5.96 15.6667 1.1950
50 48 6.20 24.0000 1.3802
50 49 6.66 49.0000 1.6902
50 50 11.08 ∞ ∞
Graph of pH vs log10([SALT]/[ACID])
7
0
-2 -1.5 -1 -0.5 0 0.5 1 1.5 2
Discussion :
The Henderson-Hasselbalch approximation allows us one method to approximate the pH of a
buffer solution. The basic equation is as follows:
pH≈pKa+log10[A−][HA](1)
We have straightforward calculations for strong acids and bases, but the computations behind
buffers are rather complex and time consuming. By using the fact that weak acids and bases
barely ionize, allowing us to approximate the pH of buffer solutions using initial
concentrations. Though the approximation has a few restrictions, it simplifies a lengthy
calculation into a simple equation derived from K . The magnitude of the equilibrium constant
for an ionization reaction can be used to determine the relative strengths of acids and bases. The
equilibrium constant for acid dissociation is
The equilibrium constant for base dissociation is
In order to obtain Henderson Hasselbalch equation, rearrange the following equation:
Buffer capacity of a buffered solution is defined as in terms of the amount of protons or
hydroxide ions it can absorb without a significant change in buffer. A buffer with a large
capacity contains large concentration of the buffering components and thus can absorb a
relatively large amount of protons or hydroxide ions and show little pH change. The pH of a
buffer system is determined by the ratio [A-]/[HA]. The capacity of a buffer system is
determined by the magnitude of [A-] and [HA]. Buffer capacity is also depends on the
concentration of the solution. The greater the concentration of a weak acid and its conjugate
base, the greater the buffer capacity. Buffer capacity may expressed as
A buffer solution has the maximum buffer capacity when the ratio or = 1 for
acidic buffer and basic buffer respectively. In such a solution pH = pKa for acidic buffers and
pOH= pKb for basic buffers. A solution can act as buffer only if the ratio of salt to acid or
base is between 0.1 and 10. Thus, the pH of an acidic buffer can have the range from pKa - 1
or pKa + 1. Similarly, a basic buffer can act as buffer in the pOH range from pKb – 1 to pKb +
1.
Conclusion :
The pH of the buffer system remains relatively unaffected by the addition of slight amounts
of acid or base. Acids have a pH below 7, so an acidic buffer solution simply has a pH less
than 7. Using our results and knowing the concentration of the conjugate base and weak acid, the
pH of the acidic buffer system can be calculated. The same is true for an alkaline buffer system if the
conjugate acid and weak base concentrations are known.In addition, the more concentrated the buffer
pair, the greater the buffer capacity of the system, which means better resistance to pH changes.
References :
(1) http://1chemistry.blogspot.com/2012/03/verification-of-henderson-hasseblach.html
(2) "Henderson-Hasselbalch Equation: Its History and Limitations," Henry N. Po and N.
M. Senozan, J. Chem. Educ., 2001, 78 (11), p 1499
(3) National Academy of Sciences (U.S.) Biographical memoirs. City of Washington,
1945. Vol. XXIII, 2d memoir.
(4) Petrucci, et al. General Chemistry: Principles & Modern Applications. 9th ed. Upper
Saddle River, New Jersey 2007.
