Buffer chemistry

1. BI11.2: Preparation of buffers and pH
meter

2. pH & its significance

3. • All biochemical reactions in vivo and in vitro
are greatly influenced by hydrogen ions of
surrounding medium.
• The convenient way of expressing hydrogen ions
concentration is by the term pH which is defined
as the negative logarithm of hydrogen ion
concentration.
• Enzymes are optimally active at a particular H+
ion concentration.

4. • Enzymes are optimally active at a particular H+
ion concentration.
• Water is dissociated to H+
and OH-
ions.
pH = Log [H+
]
• Eg: pH of a solution having hydrogen ion
concentration 0.00000001 (10-7
) is 7.

5. • The pH values of some important biological
fluids are as follow.
Body fluids pH
Pancreatic juice 8.8
Bile 7.6
Blood (at 38°C) 7.35
Saliva & human milk 6.7
Urine 6.0
Gastric juice 1.77

6. What is buffer ?

7. • The H+
ion concentration varies very little in any
biochemical fluid or environment.
• This variation is arrested by some bases or acids
which respectively absorb or donate H+
ions
depending upon situation.
• This phenomenon is called Buffering.

8. • A buffer is a mixture of a weak acid and its alkali
salts or a weak base and its acid salts.
• The buffer prevents a change in the pH on the
addition of an acid or a base.
• eg: Bicarbonic acid buffer
On addition of strong acid such as HCl bicarbonate
reacts with the acid forming carbonic acid which is
a weak acid.

9. HCl + NaHCO3 NaCl + H2CO3

10. • The capacity of the buffer decreases as the ratio
deviates from 1.
• In general buffers should be used at a pH ± 1
from pk.
• If the ratio is beyond 50:1 or 1:50, the buffer
system is considered to have lost its buffering
capacity.

11. • In aqueous solutions the pH ranges from 0-14.
• Molar concentration of H+
and OH-
ions in pure
water is 1 X 10-7
mol/L.
• With water, at neutral point
[H+
] = [OH-
] = 1 X 10-7
mol/L
• Therefore; pH at neutral point is :
–log [10-7
mol/L], so neutral pH = 7.
• More than 7 pH indicates the solution is
alkaline.

12. • Less than 7 pH indicates the solution is acidic.
• The pH 0 would be given by 1M HCl.
• The pH 14 would be given by 1M NaOH.
• Many biochemical substances posses functional
groups that are weak acids or weak bases eg: -
COOH (Carboxylic acid), phosphates present on
proteins, nucleic acids, coenzymes, intermediary
metabolities.
• Their dissociation behavior influences pH and
the structure and function of the entire molecule.

13. • Their dissociation behavior influences pH and
the structure and function of the entire molecule.
• The pH of a buffer solution can be calculated by
Handerson Hasselbalch equation.
𝑝𝐻 =𝑝𝐾𝑎 + 𝐿𝑜𝑔
𝑆𝑎𝑙𝑡
𝑎𝑐𝑖𝑑

14. Determination of pH

15.  By using indicator:
• Indicators are the compounds which change the
color with changes in the pH of the solution to
which they are added.
• They are weak acids of weak bases.
• In unionized forms, the indicators show one
color while in their ionized forms they have
different colors.

16. • The color of the solution in presence of an
indicator which in turn depends upon the relative
proportions of ionized and unionized forms of
the indicators which in turn depends on the
hydrogen ion concentration.
• For such indicators there is a definite pH range
in which it is present as mixture of the ionized
and unionized forms.

17. Indicator pH range color change
Litmus 4.5-8.8 Red-blue
Thymol blue (acid range) 1.2-2.8 Red-yellow
Thymol blue (alkaline range) 8.0-9.6 Yellow-blue
Methyl orange 3.0-4.4 Red-yellow
Congo red 3.0-4.5 Blue-red
Bromophelon blue 3.0-4.6 Yellow-blue
Bromophelon green 3.8-5.4 Yellow-blue
Phenolphthalein 8.3-10.0 Colorless-red
Methyl red 4.4-6.2 Red-yellow
Bromocresol purple 5.2-6.8 Yellow-purple
Cresol red 7.2-8.8 Yellow-red
Bromothymol blue 3.0-4.6 Yellow-blue

18.  By using pH paper:
• The indicator paper given to you is accompanied
by a color chart which shows different colors
which the indicator exhibits at different pH
values.
• Take a strip of indicator paper & moist it with
small drop of solution whose pH has to be
determined.

19. • Remove the excess fluid adhering to the
indicator paper and compare with the color chart
of the indicator and thus determine the pH.
• Do not let the paper dry before comparing with
indicator.

20.  By using pH meter:
• The most accurate method for the routine
measurement of pH is by pH meter in which a
change in [H+
] is measured as a change in
electrical potential.
• If a metal rod is palaced in a solution of its salts,
it acquires potential.

21. • If 2 different metals are dipped into the solutions
of their own salts, the difference can be
measured or calculated from two separate
potentials.
• A standard electrode is thus required against
which the potential of all other electrodes can be
compared.

22. • This is the standard hydrogen electrode
consisting of platinum rod dipped in an aqueous
solution with the given H+
activity in which
hydrogen gas is bubbled continuously at 1
atmosphere pressure.
• But as this is too cumbersome to be used as a
reference electrode routine use, other secondary
reference electrodes of known potential in
relation to the standard hydrogen electrodes are
used.

23. • The most commonly used secondary electrode is
the calomel electrode consisting of mercury-
mercurous chloride in contact with a saturated
solution of KCl. Its potential is pH independent.

24. • In the pH meter the most commonly used pH
dependent unit is the glass electrode. Certain
types of borosilicate glass are permeable to H+
but not to other cation and anions.
• Therefore; if a thin glass membrane separates 2
solutions of different pH, a potential difference
is generated across the membrane, the
magnitude of which is given by the equation

25. • Where E = Potential, R= gas constant, T=
absolute temperature, F= Faraday constant,
[H+
]I= concentration on the inside which is
fixed (0.1 N HCl), [H+
]o= concentration on the
outside (test sample)
E

26. • The voltage measured by such a system is
primarily the difference between that of the glass
and the reference electrodes and it is linearly
related to the pH of the test solution.
• The system therefore consists of the glass
electrode in contact with the solution to be
measured, the calomel reference electrode, a
KCl bride (the KCL should flow slowly into the
sample) and the measuring device (meter).

27. • These are designed so that pH 7 gives a zero
potential. High resistances are used so that very
little current is drawn from circuit (a large
current flow could change the ion concentration)

28.  Precautions
• The glass electrode is fragile and must be
handled with care.
• It must nit be left dry; usually dipped in 3M
KCl.
• The temperature compenstion dial must be set
before it is calibrated as potential produced is
dependent on temperature.

29. • The meter must be calibrated first with a
standard buffer pH 7, and then with standard of
pH 4 (if the test sample is expected to be acidic)
or with a standard of pH 9-10 (if the test sample
is expected to be alkalic).

31.  Setting of instrument
• Keep the SELECTOR in zero position
• Switch on the instrument and wait for 10
minutes.
• Connect the cleaned electrode the proper
terminals.
• Adjust ZERO control to the pointer read 7 Ph.

32.  Steps for pH measurement using digital pH
meter
• The pH instrument is calibrated with
commercially prepared buffers of known pH
values; 3 buffers are used pH -4 (acidic), -7
(neutral), and pH 10 (alkaline).
• Caliberation of the system in verified using the
same buffers as sample buffers (buffer of
unknown pH).

33. • Upon verifying the caliberation, the sample
buffer is estimated for its pH.
• The electrode should be carefully washed after
each pH determination.

35. Presented by: Dr. Krattika Singhal