The document discusses pH and buffers. It defines pH as the negative log of the hydrogen ion concentration. pH is measured on a scale from 0 to 14, with values below 7 being acidic and above 7 being basic. Buffers resist changes in pH when acids or bases are added by using an equilibrium between a weak acid and its conjugate base or vice versa. Common methods for measuring pH include pH strips, pH indicators, and pH meters. The most important buffer system in the body is the carbonic acid-hydrogen carbonate buffer system, which helps maintain pH levels in body fluids.

1. PH AND
BUFFERS
Speaker :
Manish Verma
B.Sc. MLT 1st
year

2. INTRODUCTION
ph
p means “ Puissance d” and H mean “Hydrogen.
 It is the French word, which means strength/ power
of hydrogen.
 It was introduced by Sorenson in 1909.
Definition : It is defined as the negative log of the hydrogen ion
concentration:
pH = -log [H+]

3. CONT…
 pH is a unit of measure which describes the degree
of acidity or alkalinity (basic) of a solution.
 It is measured on a scale of 0 to 14.
 Low pH values correspond to high concentrations
of H+ and high pH values correspond to low
concentrations of H+.

4. PH VALUE
 The pH value of a substance is directly related to
the ratio of the hydrogen ion and hydroxyl ion
concentrations.
 If the H+ concentration is higher than OH- the
material is acidic.
 If the OH- concentration is higher than H+ the
material is basic.
 7 is neutral, < is acidic, >7 is basic

5. THE PH SCALE
 The pH scale corresponds to the concentration of
hydrogen ions.
 If you take the exponent of the H3O+
concentrations and remove the negative sign you
have the pH of the solution.
 For example pure water H+ ion concentration is 1 x
10^-7 M, therefore the pH would then be 7.

6. MEASUREMENT OF PH
The pH can be measured by:
ph strips
Ph indicators
Ph meter

7. PH STRIPS
 pH strips are pieces of paper that change color
depending on the pH – the acidity or alkalinity – of a
liquid.
 pH strips are Brownish yellow in colour (Whatman
filter); by this method approx value is obtained, which are
less accurate

8. CONT..
 They are used for urine sample and not for blood
or plasma as :
 (1) They are affected by oxidized and reduced
agents
 (2) They are also affected by concentration of
proteins and salts

9. PH INDICATORS
 Definition: Indicators are the organic compounds
of natural or synthetic origin whose colour is
dependent on pH.
 The pH indicators are either weak acids or weak
bases having different colours in undissocisted and
dissociated forms:
HIn <------- H+ + In
-
(Colour-A) (Colour- B)

10. SOME IMPORTANT INDICATORS USED IN A CLINICAL
BIOCHEMISTRY LABORATORY ARE LISTED BELOW:
sr,.
No.
INDICATOR Ph range Colour in
acidic ph
Colour in
basic ph
1 Phenophthalein 9.3-10.5 colourless pink
2 Methyl orange 3.1-4.6 red yellow
3 Bromophenol blue 3.0-4.6 yellow blue
4 Methyl red 4.4-6.2 Red yellow
5 Phenol red 6.8 – 8.4 yellow red
6 Litmus 4.5-8.3 red Blue

11. PH METER
 The pH meter is a laboratory equipment which used to
measure acidity or alkalinity of a solution
 The pH meter measures the concentration of
hydrogen ions [H+] using an ion-sensitive electrode.
 It is the most reliable and convenient method for
measuring ph.

12. CONSTRUCTION
 A typical pH meter consists of a special measuring
probe (a glass electrode) connected to an
electronic meter that measures and displays the pH
reading.
 At the bottom of the probe there is a bulb, the bulb
is a sensitive part of a probe that contains the
sensor.
 Never touch the bulb by hand and clean it with the
help of an absorbent tissue paper with very soft
hands, being careful not to rub the tissue against
the glass bulb in order to avoid creating static
charges.
 Ph meter consists of mainly two electrodes :
 Glass electrode
 Reference electrode

13. CONT…
 A typical modern pH probe is a combination electrode,
which combines both the glass and reference electrodes
into one body. The combination electrode consists of the
following parts (see the drawing):
 1. a sensing part of electrode, a bulb made from a
specific glass
 2. internal electrode, usually silver chloride electrode or
calomel electrode
 3. internal solution, usually a pH=7 buffered solution of
0.1 mol/L KCl for pH electrodes
 4. when using the silver chloride electrode, a small
amount of AgCl can precipitate inside the glass electrode
 5. reference electrode, usually the same type as 2
 6. reference internal solution, usually 0.1 mol/L KCl
 7. junction with studied solution, usually made from
ceramics or capillary with asbestos or quartz fiber.
 8. body of electrode, made from non-conductive glass or
plastics.

14. WORKING :
 (1) Solution being tested;
 (2) Glass electrode, consisting of
 (3) a thin layer of silica glass containing
metal salts, inside which there is a potassium
chloride solution
 (4) and an internal electrode
 (5) made from silver/silver chloride.
 (6) Hydrogen ions formed in the test solution
interact with the outer surface of the glass.
(7) Hydrogen ions formed in the potassium
chloride solution interact with the inside
surface of the glass.
 (8) The meter measures the difference in
voltage between the two sides of the glass
and converts this "potential difference" into a
pH reading.
 (9) Reference electrode acts as a baseline or
reference for the measurement—or you can
think of it as simply completing the circuit.

15. ELECTRODE OR PH PROBE

16. MEASURING THE PH OF A
SOLUTION
 Rinse the electrode with distilled water and dry it.
 Place the electrode in the solution of unknown
pH.
 Turn the functions selector from the Standby
position to the pH position.
 Read the pH of the solution on the meter’s scale
or the screen.
 Turn the functions selector again to the Standby
position.

17. CALIBRATION OF PH METER
 For very precise work the pH meter should be
calibrated before each measurement. For normal
use calibration should be performed at the
beginning of each day.
 The reason for this is that the glass electrode does
not give a reproducible e.m.f. over longer periods of
time.
 Calibration should be performed with at least two
standard buffer solutions that span the range of pH
values to be measured. For general purposes
buffers at pH 4.00 and pH 10.00 are acceptable.

19. BUFFER
A buffer solution is a solution which resists changes in pH
when a small amount of acid or base is added.
Typically a mixture of a weak acid and a salt of its
conjugate base or weak base and a salt of its conjugate
acid.
The resistive action is the result of equilibrium between the
weak acid (HA) and its conjugate base (A-) or vice versa.
HA(aq) + H2O(l) → H3O+
(aq) + A-
(aq)

20. TYPES OF BUFFERS
Two types :
 ACIDIC BUFFERS –
Solution of a mixture of a weak acid and a salt of this
weak acid with a strong base.
E.g. CH3COOH + CH3COONa
( weak acid ) ( Salt )
 BASIC BUFFERS –
Solution of a mixture of a weak base and a salt of this
weak base with a strong acid.
e.g. NH4OH + NH4Cl
( Weak base) ( Salt)

21. HOW BUFFERS WORK
 Equilibrium between acid and base.
 Example: ACETATE BUFFER
 CH3COOH  CH3COO- + H+
 If more H+ is added to this solution, it simply shifts
the equilibrium to the left, absorbing H+, so the [H+]
remains unchanged.
 If H+ is removed (e.g. by adding OH-) then the
equilibrium shifts to the right, releasing H+ to keep
the pH constant

22. Buffer with equal
concentrations of
conjugate base and acid
OH-H3O+
Buffer after addition of H3O+
H2O + CH3COOH H3O+ +CH3COO-
Buffer after addition of OH-
CH3COOH + OH - H2O + CH3COO-

23. LIMITS OF THE WORKING RANGE OF A
BUFFER
 Consider the previous example:
 CH3COOH  CH3COO- + H+
 If too much H+ is added, the equilibrium is shifted all the
way to the left, and there is no longer any more CH3COO-
to “absorb” H+.
 At that point the solution no longer resists change in pH;
it is useless as a buffer.
 In order for a buffer to work well the concentration of the
acid/base and its salt must be much higher than the
strong acid/base added.

24. •HANDERSON HASSELBALCH
EQUATION
 Lawrence Joseph Henderson wrote an equation, in
1908, describing the use of carbonic acid as a buffer
solution.
 Karl Albert Hasselbalch later re-expressed that formula
in logarithmic terms, resulting in the
Henderson–Hasselbalch equation.

25.  Describes the derivation of pH as a measure of
acidity in biological and chemical systems.
 The equation is also useful for estimating the pH of
a buffer solution.
 It is widely used to calculate the isoelectric point of
proteins( point at which protein neither accept nor yield
proton) .

26. Ka =
[H+] [Ac-]
[HAc]
take the -log on both sides
The Henderson-Hasselbalch Equation derivation
-log Ka = -log [H+] -log
[Ac-]
[HAc]
pH = pKa + log
[Ac-]
[HAc]
= pKa + log
[Proton acceptor]
[Proton donor]
HAc H+ + Ac-
pKa = pH -log [Ac-]
[HAc]
apply p(x) = -log(x)
and finally solve for pH…

27. BUFFER CAPACITY OR BUFFER ACTION
 It is defined as the number of moles of an acid or a base
required to be added to one litre of the buffer solution so as
to change its pH by one.
no. of moles of the acid or base
 Buffer capacity = added to 1 litre of buffer
change in pH
NOTE : Buffer capacity of a buffer is maximum when the
concentration of the weak acid and its salt or weak
base and its salt are equal.

28. CONT..
 - The greater the buffer capacity the less the pH
changes upon addition of H+ or OH-
-
Choose a buffer whose pKa is closest to the desired
pH.
pH should be within pKa ± 1

29.  The buffer range is the pH range over which the buffer is
effective.
 Buffer range is related to the ratio of buffer component
concentrations.
[HA]
[A-] The closer is to 1, the more effective the buffer.
 If one component is more than 10 times the other,
buffering action is poor. Since log10 = 1, buffers have a
usable range within ± 1 pH unit of the pKa of the acid
component.

30. USES OF BUFFERS
 Buffers In cosmetics.
 Buffers In saops.
 Buffers in bacteriological studies.
 Calibration of pH meters.
 Control of pH in industrial reaction.
 Wine making.
 pH balanced shampoos and deodorants.

31. BUFFER SYSTEM IN
BODY FLUIDS

32. THE CARBONIC ACID HYDROGENCARBONATE BUFFER SYSTEM
• The carbonic acid-hydrogen carbonate ion buffer is
the most important buffer system.
• Carbonic acid, H2CO3, acts as the weak acid
• Hydrogen carbonate, HCO3
-, acts as the conjugate
base
• Increase in H+(aq) ions is removed by HCO3
-(aq)
• The equilibrium shifts to the left and most of the
H+(aq) ions are removed

33. CONT..
 The small concentration of H+(aq) ions reacts with
the OH-(aq) ions
 H2CO3 dissociates, shifting the equilibrium to the
right, restoring most of the H+(aq) ions
 Any increase in OH-(aq) ions is removed by H2CO3

34. 34
PHOSPHATE BUFFER SYSTEM
 The phosphate buffer system (HPO4
2-/H2PO4
-)
plays a role in plasma and erythrocytes.
 H2PO4
- + H2O ↔ H3O+ + HPO4
2-
 Any acid reacts with monohydrogen phosphate
to form dihydrogen phosphate
dihydrogen phosphate monohydrogen phosphate
 H2PO4
- + H2O ← HPO4
2- + H3O+
 The base is neutralized by dihydrogen phosphate
dihydrogen phosphate monohydrogen phosphate
 H2PO4
- + OH- → HPO4
2- + H3O+

35. 35
PROTEINS AS A BUFFER
 Proteins contain – COO- groups, which, like acetate ions
(CH3COO-), can act as proton acceptors.
 Proteins also contain – NH3
+ groups, which, like
ammonium ions (NH4
+), can donate protons.
 If acid comes into blood, hydronium ions can be
neutralized by the – COO- groups
 - COO- + H3O+ → - COOH + H2O
 If base is added, it can be neutralized by the – NH3
+
groups
- NH3
+ + OH- → - NH2 + H2O