Buffer

1. Experiment No. ___
Aim- Preparation of buffers and estimation of pH.

2. Definition
Buffers are compounds or the mixture of compounds that resist change of pH in
the solution upon the addition of a small amount of acid or alkali.
It is a mixture of weak acid and its salt with a strong base or a mixture of a weak
base and its salt with a strong acid.
Examples- Acetic acid with sodium acetate
Ammonium hydroxide with ammonium chloride
The pH of a buffer depends upon the pK value of its acids and the ratio of the salt
to acid concentration. For every buffer, pK of the acid is constant therefore, its pH is
dependent on the ratio of salt to acid.
pH= pK + log [salt]
[acid]

3. Buffer Capacity
The magnitude of resistance offered by the buffer to change in pH on addition of an
acid or alkali is called the buffer capacity or buffer value.
It is also defined as the gram equivalent of acid or alkali required by one litre of a
buffer to undergo a change of pH by 1.
A buffer requiring more acid or alkali is said to have more buffer capacity. It is
dependent on the molar concentration of the buffer and is usually maximum when
the salt to acid ratio is one i.e. present in equal concentrations.

4. Henderson-Hasselbalch Equation
The Henderson-Hasselbalch equation provides a relationship between the pH of
acids (in aqueous solutions) and their pKa (acid dissociation constant).
It is useful for estimating the pH of a buffer solution.
It is used to calculate the isoelectric point of proteins (point at which protein
neither accept nor yield proton).
Derivation to be written on plain side
The Henderson-Hasselbalch Equation is derived as:-
HI [H+] + [I-]
Ka= [H+][I-]
[HI]
[H+]= Ka [HI]
[I-]

5. taking negative log of both sides, we get
- log [H+]= - log Ka – log [I-]
[HI]
pH= pKa + log [salt] Henderson- Hasselbalch
[acid] equation

6. Preparation of Buffers to be written on plain side
Take two clean dry test tubes ‘A’ and ‘B’ and add all the reagents as given below in
the table.
Mix the contents thoroughly and buffer solution A and B are ready in tube A of
pH ≈ 5.8 and tube B of pH ≈ 7.2.
S/N Reagent Tube A pH= 5.8 Tube B pH= 7.2
1. M/15 Na2HPO4 2.0 mL 14.4 mL
2. M/15 KH2PO4 23.0 mL 5.6 mL

7. pH
It is defined as the negative log of the hydrogen ion concentration. It is a unit of
measure which describes the measure of acidity or alkalinity of a solution.
pH = - log [H+]
It was introduced by Sorenson in 1909 and is measured on a scale of 0 to 14.
Measurement of pH
By pH paper- Take a small piece of broad range pH paper of range 2-10 and
moisten it with buffer A prepared before. Wipe the excess of the solution with
filter paper. Find the pH of solution by comparing with the color chart. Repeat the
process with buffer B.
Narrow range pH paper can also be used.

8. By indicators- They are the substances which change their color with variation in
pH of the solution in which they are present.
They are usually weak acids or bases having different colors in ionized and
unionized forms.
An indicator which is a weak acid is undissociated in acid form and gives the ‘acid’
color.
In presence of dilute alkali, it will form a salt which will dissociate freely and
display the ‘alkaline’ color.
At intermediate pH, the color will be a mixture of an ‘acid’ and an ‘alkaline’ color.
Examples- Methyl red indicator- 4.4 red in pH undissociated acid form and 5.8
yellow in pH dissociated alkaline form.
Phenol red indicator- 6.8 yellow in pH undissociated acid form and 8.4
red in pH dissociated alkaline form.

9. By universal indicator- This is a broad range indicator. Take 2 mL of the buffer
solution A and B in two separate test tubes and add a few drops of universal
indicator to each. Mix and compare the color with color chart to find the
approximate pH.

10. By pH scale- A measure of the acidity or the alkalinity of a solution is measured on
a pH scale of 0 to 14. The midpoint of 7.0 on the pH scale represents neutrality i.e.
a "neutral" solution is neither acid nor alkaline. Numbers below 7.0 indicate acidity
whereas numbers greater than 7.0 indicate alkalinity.

11. By pH strips- pH strips are brownish yellow piece of paper that change color depending
on the pH. Approximate value is obtained which makes it less accurate for the
estimation. They are used for urine samples and not for blood or plasma as they are
affected by oxidizing and reducing agents and also by concentration of proteins and
salts present in the sample.

12. By pH meter- It is an electrical device that is used to measure hydrogen ion
activity in a solution.
Principle- It is based on the principle of measurement of electromotive force
(e.m.f.) generated between the two electrodes due to the difference in [H+]
concentration.
If two different metal electrodes are connected, the difference in their electrode
potential can be measured as e.m.f. (net difference of the two electrode
potentials).
One of the electrodes is a standard hydrogen electrode of known potential, so the
electrode potential of the unknown electrode sensitive to hydrogen ion
concentration can be measured. This is the basis of pH meter, i.e. the electrode
potential generated by H+ concentration in an unknown solution is measured
against a standard hydrogen electrode potential.

13. Standard buffer reagent
of pH 4.0, 7.0, 9.0 and
10.0 at 25 oC
make a simple bottle and write only
pH and temp
Parts of pH meter

14. pH measuring electrodes
• Standard Hydrogen Electrode- it consists of a platinum plate coated with
platinum black (to absorb hydrogen) and dipped in a hydrogen ion solution of
known concentration (1.0 M HCl). The hydrogen gas needs to be constantly
passed into it at 1 atmospheric pressure for the maintenance of H+
concentration. The electrode potential of this is taken as zero.
When a similar electrode is dipped into a solution of unknown H+ concentrated
and connected by a bridge of KCl solution, the electrode potential of the
unknown can be measured. KCl solution is used to allow the movement of ions
and not cross mixing of both the solutions.
Thus, the electrode potential E is related to pH by the equation:
pH= E/0.0591

15. The standard hydrogen electrode is quite difficult and cumbersome to use.
Therefore, calomel electrode is used as a standard electrode (secondary
reference electrode).
For the measurement of pH of the unknown solution, a glass electrode is
used.
• Calomel Electrode- It is a mercury-mercurous chloride electrode dipped in a
solution of KCl through a porous plug. The mercury is in contact with a
platinum wire whose potential varies with the concentration of KCl (0.1 M -
1.0 M or saturated KCl). At 1.0 M KCl solution at 25 oC, its potential is equal
0.2444 Volts.
The potential (V) at the existing temperature is given by equation:
V= 0.2444- 0.0024 (t-25) where t is the actual temperature in oC.

16. Calomel electrode

17. • Glass Electrode- This is used as a [H+] measuring electrode. It consists of H+
selective glass membrane which covers the tip of the glass electrode. This
membrane is very fragile and delicate. Inside the glass is a silver wire covered
with AgCl paste and dipped in 0.1 M HCl solution.
• Combined Electrode- A combined electrode having both the reference
electrode and the measuring electrode is used. The ph of the unknown solution
at 25 oC is given by the equation:
pH= (E-K)/0.0591
where K is constant and E is the measured potential. The response of
this electrode is linear from 0-10. It can further extend up to pH 12 by the help
of calibration.

18. Glass electrode Combined electrode

19. Procedure-
A pH meter is a combination of a combined electrode and an electronic
measuring and display device. They usually incorporate a temperature correction
also.
Standard buffers of pH 4.0, 7.0, 9.0 and 10.0 are used.
Unknown buffers/solutions are used for the measurement of pH.
• Switch on the instrument.
• Note the room temperature and adjust the temperature correction with the
temperature probe if required. Wash the electrode with distilled water and
gently wipe with soft tissue paper.
• Take standard buffers of pH 4.0, 7.0, 9.0 and 10.0 in 250 mL bottles or as
instructed on the bottle. They should be sufficient so that the electrode
bulb can dip into them. Similarly take the unknown buffer/solution in
another beaker which is to be measured.

20. • Calibration- It is the process of adjusting or standardizing the equipment to
make its functioning more precise.
Calibrate the pH meter with standard buffer solutions. First calibrate with
the pH 7.0 (zero point calibration). Then calibrate with the buffer of pH 4.0
(if the pH of the unknown solution is expected to be in acidic range) or with
a buffer of pH 9.0 and pH 10.0 (if the pH of the unknown solution is
expected to be in the alkaline range).
For calibrating, carefully lift the electrode of the instrument and wipe its tip
gently with a soft tissue paper to remove the excess water. Carefully lower
the electrode into the buffer solution (pH 7.0) so that the bulb dips into the
solution. Adjust the instrument with coarse and fine adjustments present to
show a pH value of 7.0 on the display device of the instrument. Then
remove the electrode gently and wash with distilled water. Gently wipe with
soft tissue paper. Now calibrate with other buffers given. Recheck the
calibration once again.

21. • Now after washing the electrode again with water, dip the electrode into the
unknown solution. Allow the instrument to read the pH value and note it only
when the reading is stable.
• Remove the unknown solution and wash the electrode with distilled water and
wipe it softly with the tissue paper as done before the start of the calibration.
• Switch off the instrument.
Precautions-
 Glass electrode is fragile and must be handled with care. It should not be left to
dry but kept dipped in KCl solution bottle to prevent dehydration of pH sensing
membrane and dysfunction of the electrode.
 Temperature correction is must before the calibration with the unknown
buffers.
 After every solution change, the electrode should be washed with distilled
water and wiped gently with soft tissue paper.