pH by KK SAHU sir

pH
By
KAUSHAL KUMAR SAHU
Assistant Professor (Ad Hoc)
Department of Biotechnology
Govt. Digvijay Autonomous P. G. College
Raj-Nandgaon ( C. G. )

Synopsis
 Introduction
 History
 Acid & Base
 Ionization of water
 Definitions of pH
(1) Mathematical Definition
(2) pH
(3) pOH
 Buffer solution
(1) Types
(2) Buffer action
(3) Biological buffer systems

 Henderson – Hasselbalch Equation
 Measurement of pH
(1) pH Scale
(2) pH indicators
(3) pH meter
 pH in human body and nature
 Importance
 Conclusion
 Reference

INTRODUCTION
 pH is a measure of acidity and basicity of a solution.
 pH is a defined on the negative log at the base 10 of hydrogen ion
concentration of a solution.
pH = - Log 10 [H+]
 Low PH indicates a high concentration of Hydronium ion (H3O+), while a
high pH indicates a low connection.
 The negative of the logarithm matches the number of places behind the
decimal point. for ex, 01 molar HCL should be near pH 1 and 0.001 molar
HC L should be near PH4.
 Pure water is neutral and can be considered either a very weak acid or very
break base (center of the pH scale) giving it a pH of 7 (at 250C)
 Solution with a pH less than 7 (at 250 C) are said to be acidic and solution
with a pH greater that 7 are said to be basic.

HISTORY
 Concept of Acid and Base was give by Arrhenius (1884) Bronsted &
Lowry and Lewis (1923)
 In 1906 Max Cremer discovered that the difference between liquids could
be studied by blowing a thin bubble of glass and placing one liquid inside it
and another outside it created potential that could be measured.
 In 1906 Frits Haber and Zygmut klemsiewcz discovered that glass bulb
could he used as to measure hydrogen ion activity and that this followed
logarithmic function.
 In 1909 Sorenson 1st introduced the concept pH.
 In 1936 First Commercial pH meter was made by Arhold Orville Beakman.
 In 1970 Jenco Electronics (Taiwan) designed and manufactured the 1st
portable digital pH meter.

Acid and Base
Earlier Classification
According to earlier classification :-
Acid :- It is a substance having sour taste, Turns blue litmus to red.
liberates hydrogen reacting with active Metals and a base. Ex. Citric acid
Base :- Is a substance having bitten test slippery in touch, turns red litmus to
blue and reads with acids. Ex. Ash
There definitions of acid and base have many exception and limited
applications.
Arrhenius concept of Acids and Base :-
According to Arrhenius (1884)
Acid :- is a substance which provides H+ in an aqueous solution.
Ex : HCl H+ + Cl –
(a q) (a q)
H2SO4 H+ + HSO4
-
(a q) (a q)
Base :- Its is a substance which provides hydroxyl (OH) ion in an
aqueous solution.
Ex. : NaOH Na+ (a q) + OH-
(a q)
Ca(OH)2 Ca2 +(
a q) + OH-
(a q)

Bronsted Lowery Concept :-
According to Bronsted and lowery.
Acid :- Acid is a Substance that Donets a proton
HCl + H2O H3O +(a q) + Cl –(a q)
Base :- Base is a substance that accept OH- ion.
NH3 + H2O NH4 + + OH -
(CO2)3 + H2O HCO3 - + OH -
H2O can behave as a base as well as an acid hence it is called
amphiprotic as amphigoric substance.
Lewis Concept :-
According to Lewis (1993)
Acid :- acid is a substance (atom, ion, or molecule) which is capable to accept a
pats of electrons.
Base :- Base is a substance which is capable to donate alone pair or electrons.
H3N : + BF3 H3N BF3

IONIZATION OF WATER
Ionization of water is the reversible of the process results in the
formation of Hydrogen and Hydroxide ion . However the used the term
Hydrogen ion and the symbol H+ it must be understood that "Bare" Hydrogen
ions i.e. proton do not exist in water. Hydrogen ion like most other ions are
always Hydrated. Hydrated form of the H+ ion is called the Hydronium ion .
This is often designated H3O+ but actually each H+ is closely surrounded by
several H2O molecules, The number depending on the temperature .
The ionization of water according to the equation.
H2O H+OH-

Proceeds to only a very slight extent at 250C only about 1 out of every 10
million molecules in pure water is ionized at any instant. Although water has
only a very slight tendency to iodize the products H2 and OH2 have very
profound biological effects. For this reason we must be able to express the extent
of ionization of water quantitatively.
We can do this by writing the expression for the equilibrium constant for the
reversible reaction (A)
[H+] [OH-]
Keq =
[H2O]
We can now simplify this expression since the concentrating of H2O is
relatively very high (it is equal to the number of grams of H2O in I.L. divided
by the gram molecular weight. or 1000/18 = 55.5 M) and thus is essentially
constant in relation to the very low concentrations of H+ and OH- ions in pure
water at 250 C. namely , 1 x 10-7 M . Accordingly we can substitute 55.5 in the
equilibrium constant expression to yield.
[H+] [ OH-]
Keq =
55.5
55.5 Keq = [H+] [OH-]
The value for Keq has been carefully estimated from electrical
conductivity measurements of water [only the ions] arising from the dissociation
of H2O can carry current in pure water) and found to be 1.8 x 10-16 at 280 C.
Substituting this value for Keq in the above equation gives.

(55.5) (1.8x10-16) = [H+] [ OH-]
99.9 x 10-16 = [H+] [ OH-]
1.0 x 10-14 = [H+] [ OH-]
The symbol Kw is used to designate the product 33.3 Keq and we have the relationship.
Kw = 1.0 X 10-14 = [H+] [OH-]
Kw, called the ion product of water, has the value 1.0 X 10-14 at 250 C. What the means is
that the product [H+][OH-] in aqueous solution at 250C always equals the fixed number
1X10-14 when there are exactly equal concentrations of both H+ and OH-, as in pure
water, the solution is said to be neutral. Under these conditions the concentrations of H+
and OH- can be calculated from the ion product of water as follows.
Kw = [H+] [OH-] = [H+]2
Solving for H+ gives
Furthermore, the ion product of water says that whenever the concentration of H+ ions is
greater than 1 X 10-7 M, the concentration of OH- must become less than 1 X 10-7 M.
and vice versa. Thus when the concentration of H+ is very high as. In a solution of
hydrochloric acid, the OH- concentration must be very low, since the product of their
concentrations must be 1 X 10-14.Conversely, when the concentration of OH- is very
high as in a solution of sodium hydroxide, the concentration of H+ must be very low.
Thus from the ion product of water we can calculate the H+ concentration if we know the
OH- concentration, or vice versa.

Handerson equation
 Ph of a buffer solution can be calculated with the help of
HANDERSON EQUATION.
 For this consider a buffer of weak acid HA and its salt.
 HA↔H+ +A-
Ka=[H+] [A-]/[HA]
Ka=dissociation constant of acid or [H+]=Ka [HA]/[A]
Salt is completely ionized while due presence of excess A-
from the salt .
The dissociation of weak acid will be depressed more due to
common ion effect.

.
 Or [H+]=Ka [acid]/[salt]
taking log value
log10[H+]= log10ka+log10[acid]/[salt]
-log10[H+]= -log10ka-log10[acid]/[salt]
-log10[H+]= -log10ka+log10[salt]/[acid]
Or pH=pKa + log [salt]/[acid]
This is HANDERSON EQUATION.

Definition of pH :-
1. Mathematical definition :-
pH is defined as the decimal logarithm of hydrogen ion activity in a
solution. pH is a dimensionless quantity.
pH = -Log10 (aH+) = log10(1H+)
(aH+) = activity of hydrogen ion
The reason for this definition is that a H + is a property of a single
ion. which can only he measured experimentally by meant of an ion-
selective electrode which responds according to Nernst equation, to H+
activity.

pH is commonly measured by means of a combined glass electrode.
Which measures the potential difference, or electromotive force E
between an electrode sensitive to the hydrogen ion activity and a
reference electrode, such as a a calomel electrode or silver chloride
electrode. The combined glass electrode ideally follows the Nernst
equation :
RT E0-E
E = E 0 + In (a H+ ) ; p H =
n F 2.303 RT/F
Where E = measured potential,
E0 = Standard electrode potential
F = Faraday Constant and
N = the number of electrons transferred

pH
“pH is defined as the negative log at the base 10 of hydrogen ion
concentration of solution.”
pH = - log 10 [H+]
POH
POH is some times used as measures of the concentration of
hydroxide ions, (OH-). pOH is not measured independently but is derived
from pH. The concentration of hydroxide ions in water is related to the
concentration of hydrogen ions by
pOH + pH = pKw
pOH = pKw - pH
So at room temperature
pOH = 14 – pH

PH Indicators :-
A PH indicator is a Organic chemical compound that is added
in small amount to a solution so that the pH of the solution can be
determined visually
Theory :-
PH indicator are frequently weak acids or weak bases The
general reaction formulated as follows .
HInd + H2O H30+ + Ind -
Hydronium ion Conjugate atc base
Here Hind stands for the acid from and Ind- for the conjugate base of
the indictor. It is ten ratio of these that determines of color of the
solution and that connects the color to the pH value For pH indicators
that are weak protolytes we can write Henderson Hasselbalch equation
for them.

The equation derived from the acidity constant state that when
pH equals the pKa value of indicator both species are present in 1:1 ratio
If pH is above the pka value the concentration of the conjugate base is
greater than the concentration of the acid and the color associated
associated. with the conjugate base dominates. If pH is below the pKa
value the converse is true.
Indicators Color PH Transition range
Acid Base
1- Methyl 1 Orange Red Orange 3.1-4.4
2- Bromo Phenol blue Yellow Purple 3.0-4.6
3- Bromo thymol blue Yellow Blue 6.8 - 7.6
4- Congo red Blue Red 3.0-5.2
5- Indigo carmine Blue Yellow 11.6-14.0
6- Litmus Red Blue 4.5-8.3
8- Alzerine red Yellow Red 5.0-6.8
9- Phonal red Yellow Blue 6.8 - 8.4
10- Bromo cresol green Yellow Blue 4.0-5.6

Uses :-
 In titration
 In analytic chemistry and biology experiments to determine the extent of a
chemical reaction.
pH scale :-
 Scale pH is a means by which we can measure the pH of a solution.
 Solution with a pH less than 7 considered as a acidic solution and solution
with a PH greater than 7 are said to be basic.
 It is discovered by Sorenson in 1909
 pH scale is Logarithmic it is not arithmetic.
 Kw the ion product of water, is the basis for the pH scale
Neutral
Acidic 7 Basic
X
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Uses :-
1- In measuring the pH of a solution

Uses :-
1- In measuring the pH of a solution
pH Meter :-
 A pH meter is an electronic instrument used to measure the pH of a solution.
 An electric potential develops when one liquid is brought in contact with another
one but a membrane is needed to keep such liquids apart.
 A PH meter measures essentially the electro chemical potential between a known
liquid inside a thin glass bulb (glass electrode) and an unknown liquid outside.
The glass electrode allows agile and small H+ interact with the glass electrode
which measure the potential or hydrogen.

 Another reference electrode is used to measure the electrode voltage.
Since a small leakage of on take place in the reference electrode a pH
meter should not be used in measuring liquid of low conductivity and
small container should not be used. The pH meter measures the
electrical potential (follow the drawing closewise from the meter)
between the mercuric chloride of the reference electode and its
potassium chloride lic the unknown liquid. It solution in side the glass
electrode, and the potential between thea solution and the silver
electode. But only the potential between the unknown liquid and the
solution inside the glass electrode change from sample to sample. So
other potentials can be calibrated out of the equiation. The calomel
reference electrode consists of a glass tube with a potassium chloride
(KCl) electrolyte which is in intimate contact with a mercuric chloride
element at the end of a KCL element, it is a fragile construction joined
by a liquid junction tip made of porous ceramic or similar material.
This kind of electrode is not easily polsoned by heavy metals and
sodium . The glass electrode consists of a sturdy glass tube with a thin
glass

bulb welded to it. Inside is a known solution of potassium chloride (KCl) bufered
at a pH of 7.0 A silver elec trode with a silver chloride tip makes contact with the
inside solution. to minimise electronic interference, the probe is sheelded by a
foil shield, often found inside the galss electrode Most modern pH metres also
have a thermistor temperature probe which allows for authomatic temperature
correction, since pH varies somewhat with temperature.
Naturally accuring PH meter :-
Many plants or plant parts contain chemicals from the naturally colored
anthocynine family of compounds. They are red in acidic solution and blue in
basic. Extracting anthocynin from red cabbage leaves or skin of lemon to form a
crude acid base indicator is popular in introductory chemistry demonstration.
Anthocynine can be extracted from a multitude of colored plants or
plant parts from leaves (red cabbage) flowers (Poppy etc.) berries (blue berries,
black current) and stem.

Buffer Solution :-
Buffer in a solution in which :-
1- pH value is definite
2- PH does not change on dilution as on keeping for sometime.
3- An a adding acid or base in less quantity change in pH is negliable.
Such solution are called buffer solution.
Types of buffer solution :-
1- Acidic buffer Solution - Acidic buffer are formed with weak acid and its
salt solution.
ex. Acetic acid + Sodium Acetate
2- Basic Buffer Solution - Basic buffer are mixture of weak base and its
salt solution.
NH4Cl + NH4OH
Buffer action of buffer solution :-
3- Buffer action of acidic buffer
ex . CH3COOH and CH3 COONa
Ionization of sodium acetate will produce CH3COO- ion in large
quantity while H+ will be less because of weak acid and common ion
effect.

CH3COONa CH3COO- + Na+
CH3COOH CH3COO- + H+
On adding HCl, H+ ions produced will combine with acetate ions to
form CH3COOH which will ionize very less, so its pH will not change
H + + CH3 COO- CH3COOH (Weak Acid)
(Very less ionization)
On adding a drop of NaOH + OH- ion produced will combine with H+ of acetic
aicde and pH remains constant.
OH- + CH3COOH H2O + CH3COO

Buffer action of Basic Buffer :-
Ex. NH4Cl + NH4OH –
In this solution NH4+ will very high due to ionization of NH4Cl (Strong
electrolyte) which will depress the ionization of weak base NH4OH (Common
ion effect)
NH4Cl NH4 + + Cl- (Complete ionization)
NH4OH NH4 + + OH (Very less ionization)
On adding a drop of HCl, H+ will combine with OH of NH4OH and PH
remains constant
H + + NH4OH H2O + NH4+
On adding a drop of NaOH the produced OH- will combine with
NH4 + to form NH4OH (weak base)
OH- + NH4 + NH4oH (weak base)
Biological Buffer Systems
 The Phosphate Buffer Systems
 The Bicarbonate Buffer System
 The Protein Buffer Systems
 The amino acids buffer system
 The Hemoglobin Buffer Systems

Importance :-
1- In laboratory :_
In study of velocity of chemical reactions buffer are used.
2- Qualitative analysis :-
In removal of Po4 buffer CH3COONa + CH3COOH is used
3- Industries :-
Production of alcohol by fermentation In manufacturing of sugar, Paper
and in electroplating industries.
4- In Organisms :-
In organisms cellular activity depends on enzyme and activity of
enzyme depends on certain pH.
Handerson Equation :-
pH of a buffer solution can be calculated with the help of Handerson’s
equation for this consider a buffer of weak acid HA and its salt.
HA H+ + A-

1- pH in human body :-
1- Gastric acid - 1
2- Lysosomes - 4.5
3- Granules of chromaffin cells - 5.5
4- Human skin - 5.5
5- Neutral H20 at 3rc - 6.81
6- Urine - 6.0
7- Cytosol - 7.2
8- Blood - 7.34 - 7.45 + 7.4
9- Mitochondrial matrix - 7.5
10- Pancreas Secretions - 8.1
11- Sweat - 4 - 5.5
12- Tears - 7.1 - 7.4

Acid base homeostasis :-
The pH of different cellular components body fluids and organism
usually tightly regulated in a process called acid base homeostasis.
Acidosis :-
The pH of blood plasma of severely diabetic person is often lower than
the normal value of 7.4 this condition is called acidosis.
Alkaloses :-
When the PH of blood is higher the normal condition it is known as
alkalosis.
Ionization of Water
Ionization of water is the reversible process results in the formation of
hydrogen and hydroxide ions. However, when we use the term “Hydrogen ion”
and symbol H+ it must be understood that “bare” hydrogen ions i.e., protons, do
not exist in water, hydrogen ions, like H+ ion is called hydronium ion. This is
often designated H3O+ but actually each H+ is closely surrounded by several
H2O mol ecules, the number depending on the temperature.
The ionization of water according to equation
H2O H+ + OH-
Proceeds to only a very slight extent ; at 250C only about 1 out

IMPORTANCE :-
 In daily life
 In Microbial word
 In Agriculture
 In Medical
 In Food Industries
 In Milk Production
 In Bear Production
 In Human Body

Reference
Books Author
Fundamental of Biochemistry Dr. J.L. Jain,
Chemistry –Principle of Biochemistry Lehringer (Nelson & Cox)
Websites : www.wikipedia.com

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