This presentation is very useful for chemistry ,life science as well as medical students.

1. pH Scale, Buffers,
Redox potential
Ifra
M.Sc. 3rd SEM

2. pH Scale
Danish biochemist ‘ Soren Sorensen’ used a logarithmic scale for
expressing the hydrogen ion concentration. This scale was called pH
scale, where p stands for power and H for Hydrogen ion
concentration. He defined pH as
“The pH of a solution is the negative logarithm of the concentration(in
moles/litre) of hydrogen ions”
pH= ─ log [H+] or log 1/[H+]
pH gives an idea about the acidity or basicity of the solution. The
ionization product of water form the basis of pH scale.

3. Since pure water is neutral, it contains equal concentration of
hydrogen & hydroxyl ions. At a certain temperature, the product of
the concentration of H+ & OH+ ions in pure water is known as ionic
product of water at the temperature. The ionic product of water at
25°C is approx. equal to 1*10-14
In neutral solution [H+]=[OH-]= 10-7 (pH = 7).
In acidic solution, pH value ranges from 0 to 7.
In alkaline(basic) solution, the pH value is between 7 to 14

5. A pH scale shows a range of 0 to 14.
Sometimes the expression pOH is used to denote the basicity of OH-
concentration of a solution.
pH=log 1/[OH-]= ─log [OH-]
pH + pOH = 14
The pH of an aqueous solution can be estimated by using various
indicator, dyes such as phenolphthalein and phenol red. These dyes
ionize at a specific pH to produce coloured ions. In laboratories, the
pH meter.

6. pH OF SOME FLUIDS
FLUID pH
1. Gastric juice 1.0
2. Lemon juice 2.0
3. Tomato juice 4.0
4. Milk , Saliva 6.5
5. Human blood 7.4 – 7.8

7. BUFFERS
SOLUTIONS OF RESERVE ACIDITY AND
RESERVE ALKALINITY

8. BUFFERS
A buffer solution is that which tends to maintain its pH when small
amounts of strong acid or base are added to it. It is a mixture of
weak acid or weak base and their conjugate base or conjugate
acid respectively. It contains a hydrogen ion donor & a hydrogen
ion acceptor form of a weak acids and weak bases. A buffer
system is most effective when the concentration of H+ & H+
acceptor is equal.
Carbonic acid bicarbonate is a common buffering system in blood
plasma. The weak carbonic acid dissociated into H+ & HCO3- as
follows-
H2CO3 → H+ + HCO3-

9. When a small amount of HCl is added to this system, H+ ions are
produce from the acid combine HCO3- ions to form H2CO3. If a small
amount of NaOH is added, the OH produced reacts with H+to form
water molecules. Thus this system soaks the H+ or OH- produced
from strong acid or base & tend to maintain the original pH.
Similarly weak bases & their salts also work as buffer system.
Buffer systems in the organism help in carrying on most of the
biochemical reactions in a narrow pH range of 6 to 8. The blood for
example maintains its constant pH of about 7.4 despite the fact that
it carries a large number & variety of chemicals. Buffer system
provide protection to cells & tissues against sudden change in pH.

11. Buffering in the blood
The pH range of blood is normally in the range of 7.35 to 7.45. If pH decreases
below this range, the symptoms of acidosis appear & death of the animal may
occur at pH 7.8 . This is because the enzymes present in the blood are extremely
sensitive to changes in pH. The major buffer systems of the blood are bicarbonate,
phosphate, haemoglobin & proteins buffers. Haemoglobin is a good buffer because
of its capacity to act as a oxygen acceptor as well as a oxygen donor.
Oxyhaemoglobin( HHbO2) is a stronger acid than than carbonic acid but
haemoglobin(HHb) is weaker acid. When blood is circulated through the
pulmonary veins( in lungs), haemoglobin is converted to oxyhaemoglobin by
absorbing oxygen. Because of its acidic nature it reacts with the bicarbonates
present in blood.
HHbO2 + BHCO3 → BHbO2 + H2CO3 ( B= Na, K, etc)

12. The carbonic acid thus produced is decomposed into carbon dioxide &
water by the enzyme carbonic anhydrase.
H2CO3 → CO2 + H2O
The salt of oxyhaemoglobin formed during reaction with bicarbonate
converted to the salt of haemoglobin by deoxygenation in the tissues,
which then reacts with the carbonic acid produced from carbon
dioxide liberated in the oxidation of carbohydrates, to form
bicarbonate salt and haemoglobin.
BHbO2 → BHb + O2
BHb + H2CO3 → BHCO3 + HHb
The haemoglobin thus liberated goes to the lungs again, where it can
be oxygenated. This cyclic oxidation and deoxygenation of
haemoglobin between lungs and tissues is represented as Henderson
cycle.

14. Haemoglobin as buffer
CO2 +H2O
H2CO3
HHb
HCO3- + H+
Hb
Carbonic anhydrase
HCO3-
Cl- Cl-
CO2
ERYTHROCYTESPLASMA

15. Renal regulation of blood pH
PHOSPHATE BUFFER
Renal tubular cell
Na+
HCO3- + H+
H2CO3
CO2 + H2O
Na2HPO4 pH=7.4
Na+ NaHPO4-
H+
NaH2PO4 pH=4.5
EXCRETED
Na+
HCO3-
BLOOD TUBULAR LUMEN

16. PROTIEN BUFFER SYSTEM
RENAL TUBULAR CELL
Glutamine
Glutaminase
Glutamate
Na+
HCO3- + H+
CA
H2CO3
CO2 + H2O
BLOOD
Na+
HCO3-
TUBULAR LUMEN
NH3
Na+
H+
NH4+
EXCRETED

17. pH=7.4
CO2
(H2CO3)
HCO3-
Lungs
(CO2 exhaled)
Metabolism
(CO2 generated)
Kidneys
(HCO3- generated
,H* lost)
Erythrocytes
(CO2
transported ,
HCO3-
generated)
pH Regulation In Blood

18. DISEASES OCCUR DUE TO pH
DISTURBANCE
ALKALOSIS – It is a rise in pH
(a)METABOLIC- due to increase in
bicarbonate.
(b)RESPIRATORY- due to decrease in
carbonic acid.
(c)Occur due to vomiting, anemia,
hypokalemia and at high altitude.
(d)Compensated by hypoventilation
and HCO3- excretion by kidney.
ACIDOSIS – It is a decline in
pH
(a)METABOLIC- due to decrease in
bicarbonate.
(b)RESPIRATORY- due to an increase
in carbonic acid .
(c)Occur due to diabetes, heart, liver,
lungs and kidney problems.
(d)Compensated by hyperventilation
and HCO3- retained by kidney.

19. Other Examples
Buffering Agent pKa Useful pH Range
CITRIC ACID 3.13,4.76,6.40 2.1- 7.4
ACETIC ACID 4.8 3.8-5.8
KH2PO4 7.2 6.2-8.2
CHES 9.3 8.3-10.3
BORATE 9.24 8.25-10.25

20. USES
• In determining pHof unknown solutions.
• In studying the rate of chemical reactions.
• In the manufacture of ethyl alcohol from molasses (pH 5-6.8).
• In paper manufacture , leather tanning etc.
• In preparing cultures in biological specimens.

21. pH of the buffer system
pH of the buffer solution can be calculated if the composition of the mixture
as well as the ionization constant of the weak electrolyte is known. For example
in a buffer mixture of acetic acid and sodium acetate, the pH can be
determined if the ionization constant of acetic acid is known. The formula for
such a determination can be derived as follows-

23. Hendorson-Hassel Equation
1. For acid , pH=pKa + (salt)/(acid)
2. For base , pOH=pKb + (salt)/(base)
3. pH + pOH= pKw =14
4. In blood, pH=pKa +log(base)/(acid)= pKa+log(HCO3-)/(H2CO3)

24. REDOX POTENTIAL
The process of electron transfer accompanied with the oxidation – reduction of
the system. A compound losing electron is oxidized while a compound gaining it
is reduced. For example ferrous ion is oxidized to ferric ion is oxidized to ferric
ion by losing one electron & vice-versa. The quantitative measure of the affinity
of a compound to lose or gain electron is the redox potential.
Fe++ → Fe+++ + e-

25. A redox system can be compared to a dry electric cell. In a cell, the e- are transferred
through a wire from one electrode to the other. This generates an electric an electric
current. The capacity to gain or lose e- in such a system is electrode potential. It can
be measured through a standard hydrogen potential of zero, at 1N concentration & 1
atmospheric pressure. The electrode potential of a reducing- oxidizing system can
also measured in a similar way. Electrode potential in this case will be the redox
potential.
Redox potential of an organic compound can be measured in the laboratory by using
a standard platinum electrode.
E = E• + RT/nF In [oxidant]/[reductant]
where E= redox potential
E•= Redox potential of mixture containing equimolal concentration of
oxidant and reductant.

26. R= Gas constant
T= absolute temperature
F= Faraday number= 96500 coloumbs
n= number of e-
This equation is called Peter’s equation.
Under normal conditions of temperature i.e 30◦C, valency change of 2 &
converting into log10 [oxidant]/[reductant]

27. Redox potential of hydrogen involving system
E=E• + RT/nF In[oxidant]/[reductant]+ RT/nF In[H+]