Ph, buffers & isotonic solution

By; Khalifa M. Asif Y. Asst. Professor Ali-Allana College of Pharmacy, Akkalkuwa
UNIT-V
pH, buffers
&
Isotonic solutions

pH, buffers and Isotonic solutions:
Sorensen’s pH scale, pH determination (electrometric and calorimetric),
applications of buffers, buffer equation, buffer capacity, buffers in
pharmaceutical and biological systems, buffered isotonic solutions.

 For water molecules undergoing ionisation we can write;
 According to law of mass action;
 But, Only one out of 550 water molecules is undergoing ionisation. It
means that [H20] is constant.
Where Kw is called ionic product of water and its value is 1 × 10-14 at
25 °C, I.e.,
𝑯 𝟐 𝑶 ⇋ 𝑯+ + 𝑶𝑯−
𝑲 =
𝑯+
𝑶𝑯−
𝑯 𝟐 𝑶
𝑲 𝑯 𝟐 𝑶 = 𝑯+
𝑶𝑯−
𝑲𝒘 = 𝑯+ 𝑶𝑯−
𝟏 × 𝟏𝟎−𝟏𝟒 = 𝑯+ 𝑶𝑯−

 Pure water has equal concentration of [H+] and [OH-] ions. Thus,
or
 Thus, for neutral water, [H+] ion concentration is 10-7 g ion/litre
 The solution becomes acidic if [H+] > l × 10-7 g ion/litre and
alkaline if [H+] < l × 10-7 g ion/litre
 Reverse is true when solutions [OH-] > l × 10-7 g ion/litre. for
alkaline solutions.
𝟏𝟎−𝟏𝟒
= 𝑯+
𝑯+
𝑯+
= 𝟏𝟎−𝟕

 Hydrogen ion concentration values are complex figures to write and
use in calculations since they contain negative powers of 10.
 In 1909, Sorensen introduced the term pH.
 The pH expresses [H+] of an aqueous solution as a logarithmic
function.
The pH of a solution equals the negative of the logarithm to the
base10 of its hydrogen ion concentration, I.e., pH = -log10 [H+]
 pH offers a convenient mechanism of expressing a wide range of [H+]
in small positive numbers.
 The letter, p in term ‘pH’ stands for German word “ potenz ” (power),
so pH is abbreviation for “power of hydrogen”.

Relation between pH, [H+] and [OH-] concentration in Moles/L at 25 °C Temperature

 The PH scale from 0 to 14 covers all the hydrogen ion concentrations.

 The Electrometric Method is the most accurate of the
methods employed for the determination of Hydrogen Ion
Concentration.
 It is the accepted method for research and laboratory work
requiring pH measurements accurate to 0.1 to 0.001 pH.
1. Electrometric Method
pH determination
Generally pH is determined by following methods;
1. Electrometric Method
2. Colorimetric

Principle
 The basic principle of the electrometric pH measurement is
determination of the activity of the hydrogen ion by potentiometric
measurement using a standard hydrogen electrode and a reference
electrode.
 pH meter detects the change in potential and determine the hydrogen
ion by equation;
𝑬 = 𝑬 𝟎 +
𝟐. 𝟑𝟎𝟑 𝑹𝑻
𝒏
× 𝑭 × log
𝒖𝒏𝒌𝒏𝒐𝒘𝒏 𝑯+
𝒊𝒏𝒕𝒆𝒓𝒏𝒂𝒍 𝑯+
Where,
E = Total potential difference
E0 = Reference potential
R = Gas constant
n = no. of electrons
F = Faraday’s constant

Apparatus:
 pH meter consisting of
potentiometer, a glass electrode,
a reference electrode and a
temperature compensating
device.
 Glass electrode: The sensor
electrode is bulb of special glass
containing a fixed concentration
of HCl and a buffered chloride
solution in contact with an
internal reference electrode.

 The electrodes commonly used are the Hydrogen Electrode, the
Quinhydrone Electrode and the Calomel Electrode.
 The usual combinations are the Hydrogen-Calomel Electrode and
the Quinhydrone-Calomel Electrode Assemblies.

Procedure
 Before use, remove electrode from storage solution, rinse, and
blot, dry with a soft tissue paper.
 Calibrate the instrument with standard buffer solution. (Ex:
KCl solution of pH 7.0)
 Once the instrument is calibrated remove the electrode from
standard solution; rinse, blot and dry.
 Dip the electrode in the sample whose pH has to be measured.
 Stir the sample to ensure homogeneity and to minimize CO2
entrainment.
 Note down the reading (pH) from the pH meter.

Colorimetric Determination of the pH

 The basis of colorimetric analysis is the variation in the intensity
of the colour of a solution with changes in concentration (or pH).
 The colour may be due to an inherent property of the constituent
itself (e.g. MnO4− is purple) or it may be due to the formation of
a Coloured compound as the result of the addition of a suitable
reagent (e.g. indicator).
 By comparing the intensity of the colour of a solution of
unknown concentration (or pH) with the intensities of solutions
of known concentrations (or pH), the concentration of an
unknown solution may be determined.
Colorimetric Determination of the pH

A buffer solution is one which resist a change of pH on
addition of an acid or alkali or, on dilution with a solvent.
 The pH of pure water is 7 and the water has no buffer action.
 A buffer solution consists of a mixture of weak acid and its salt or
of a weak base and its salt. To such solution when small amount of
acid or alkali is added no significant change in the pH takes place.
 Buffer solutions can be prepared either by mixing a weak acid
with its salt or a weak base with its salt.
Buffer solution

Types of Buffer Solutions
These are mainly of following two types:
1. Acidic Buffer Solution: The solution having a mixture of weak
acid (e.g., acetic acid) and its salt (e.g., sodium acetate) is known
as acidic buffer.
2. Basic Buffer Solution: The solution having a mixture of weak
base (e, g., ammonium hydroxide) and its salt (e, g. Ammonium
chloride) is a basic buffer.

Properties of buffer solutions
1. The pH of buffer solution remains constant.
 The pH of solution does not change on dilution.
 The pH does not change even after addition of small
quantities of acids or bases.
 The pH of buffer solution does not change on keeping for
long time.
2. The pH of solution remaining constant is useful in number of
chemical reactions.

Buffer action
 Buffer solutions undergo only small changes of pH on the addition
of small amounts of acid or base.
For example,
 If 1 ml of 0.01 M hydrochloric acid is added to 1 litre of pure
water, the pH is reduced to 5.0 from 7.0
 On the other hand, if the acid is added to 0.001 M buffer
solution containing equal quantities of acetic acid and sodium
acetate in water the pH change is only by 0.09 units.
 From this, it is clear that a buffer solution resists change of pH
upon the addition of small quantities of acid or alkali.

Mechanism of buffer action
 If a small quantity of 0.1N HCI is added to the buffer solution, the
base ties up the hydrogen ions released by HCl.
 If an alkali (NaOH) is added to the buffer solution, the free acid of
the buffer solution neutralises the base added
 Thus, the change in PH is resisted by the addition of small
quantities of strong acid.
(𝑯𝑪𝒍 + 𝑯 𝟐 𝑶 ⇋ 𝑯 𝟑 𝑶+ + 𝑪𝒍−)
𝑪𝑯 𝟑 𝑪𝑶𝑶
−
+ 𝑯 𝟑 𝑶+
⇋ 𝑪𝑯 𝟑 𝑪𝑶𝑶𝑯 + 𝑯 𝟐 𝑶
𝑪𝑯 𝟑 𝑪𝑶𝑶𝑯 + 𝑵𝒂𝑶𝑯 ⇋ 𝑪𝑯 𝟑 𝑪𝑶𝑶𝑵𝒂 + 𝑯 𝟐 𝑶

Buffer equation
 Buffer equation may be obtained by considering the effect of
sodium acetate on the ionisation of acetic acid (the salt and the
acid have an ion common with them i.e. Acetate ion, CH3COO-)
The dissociation constant for the acid is given as
𝑪𝑯 𝟑 𝑪𝑶𝑶𝑯 ⇋ 𝑪𝑯 𝟑 𝑪𝑶𝑶
−
+ 𝑯 𝟑 𝑶+
𝑲 𝒂 =
−
𝑯 𝟑 𝑶+
𝑪𝑯 𝟑 𝑪𝑶𝑶𝑯
= 𝟏. 𝟕𝟓 × 𝟏𝟎
− 𝟓

 If sodium acetate is added to the acetic acid solution, it ionises to
produce acetate ions as;
 and it results in the increase in the concentration of CH3COO-
momentarily.
 To re-establish the constant Ka = 1.75 × 10-5, the hydrogen ion
term in the numerator is instantaneously decreased.
 This results in the increase of concentration of CH3COOH in the
denominator i.e. The reaction
is favored. Thus the constant Ka remains unaltered.
𝐂𝐇 𝟑 𝐂𝐎𝐎𝐍𝐚 ⇋ 𝐂𝐇 𝟑 𝐂𝐎𝐎
−
+ 𝐍𝐚
+
−
+ 𝑯 𝟑 𝑶+
⇋ 𝑪𝑯 𝟑 𝑪𝑶𝑶𝑯 + 𝑯 𝟐 𝑶

 The pH of the final solution (i.e. the buffer solution) is obtained from
the above equation.
 By rearrangement
 Since CH3COOH ionises only slightly (i.e. to an extent of 0.0000175
mole/liter), the concentration of CH3COOH may be considered to
represent the total concentration of acid in the solution and can be
replaced by [Acid].
 The acetate ion (CH3COO-) is contributed entirely by sodium acetate
as the contribution by the acid is negligible, Therefore [CH3COO-]
may replaced by [salt]. Therefore;
𝑯 𝟑 𝑶+
= 𝑲 𝒂
𝑪𝑯 𝟑 𝑪𝑶𝑶𝑯
−
𝑯 𝟑 𝑶+
= 𝑲 𝒂
𝒂𝒄𝒊𝒅
𝒔𝒂𝒍𝒕

Expressing in logarithmic form,
log [H3O+] = log Ka + log [acid] - log [salt]
Reversing the sign
- log [H3O+] = - log Ka - log [acid] + log [Salt]
or
This is known as Henderson Hasselbalch equation for a weak acid and
its salt.
pH = pKa+ log
salt
acid

 To know the effectiveness of a buffer on a quantitative basis, the term
buffer capacity (ß) is used. First introduced by van Slyke in 1922.
Buffer capacity is defined as the amount of acid or base that
must be added to the 1 liter buffer to produce a 1 unit change of pH.
Hence,
 Where the d[B] is gram equivalent of strong acid or base added and
d[H] is pH change due to addition of acid or base.
 The larger the buffer capacity, the more resistant the buffer is to
changes in pH.
 The capacity of a given buffer is determined by the concentrations of
acid and salt and also by the ratio of acid/salt or base/salt.
β =
d B
d pH
Buffer capacity

1. Ratio of Salt/Acid or Base
2. Total buffer concentration
3. Temperature
4. Ionic strength
Factors affecting Buffer capacity

Role of buffers in pharmacy
The buffers play an important role in pharmaceutical preparations to
ensure stable pH conditions for the medicinally active compound.
1. Solubility: Solubility of compounds can be frequently controlled by
providing a medium of suitable pH. The required pH is adjusted by
buffers.
E.g. Many inorganic salts such as salts of Fe+3, phosphates,
borates become soluble in acid media; but precipitate in alkaline
media.
2. Colour: Colour of many natural dyes, present in fluid extract or of
certain synthetic drugs has been found to be pH dependent. e.g., red
colour of cherry and raspberry syrups has been maintained at acidic
pH which becomes pale yellow to nearly colorless at alkaline pH.

3. Stability of certain compounds :
e.g.
i. Adrenaline with dissolved oxygen in presence may undergo
reductions with alkaline medium. So its solution for
injection is to be buffered to most stable pH range 2.5 - 5.
i. Ascorbic acid and penicillin are unstable in an alkaline pH.
Some compounds have been found to be structurally unstable within
certain PH ranges usually autoxidation, giving rise to insoluble solids
or gases,
e.g.,
i. Sodium thiosulphate and sodium polysulphide preparations
have to be stored at alkaline conditions to prevent
separation of sulphur.
ii. Nitrites become brown in acid media because of formation
of Coloured nitrogen oxides.

4. Patient comfort: Injectable and preparations for internal or external
use become irritating if their pH is different greatly from that normal
for the particular tissues involved. An extremely acid or alkaline pH
must be avoided because of tissue damage.
5. Activity of medicinal compounds: Optimum pH conditions for
activity of medicinal compounds have to be maintained.
e.g.
i. Buffering methylamine with sodium dihydrogen phosphate.
ii. Adjustment of the pH of sodium hypochlorite to lower
values tends to increase the germicidal effectiveness of the
preparation.

 Usually for the preparation of buffer solutions, weak bases and their
salts are not used since many bases are volatile and their solutions
have high temperature coefficients.
 KCI is a neutral salt and is added to adjust the ionic strength of the
buffer system. Some pharmaceutical buffers are;
Pharmaceutical buffers
Buffers pH
HCI and KCl 1.2 to 2.2
HCI and potassium hydrogen phthalate 2.2 to 5.0
NaOH and potassium hydrogen phthalate 4.2 to 5.8
NaOH and potassium dihydrogen phosphate 5.8 to 8.0
Boric acid, NaOH and KCl 8.0 to 10.0

 Various buffer solutions used in the pharmaceutical formulations are
as follows;
Buffers pH
Acetic acid / sodium acetate 3.8 to 5.6
Phosphate acid/ sodium phosphate 5.0 to 8.0
Citric acid/ sodium citrate 1.2 to 6.6
Boric acid/ sodium borate 7.8 to 10.6

Biologic Buffers
 Blood pH is maintained at about 7.4 by the buffers present in plasma
as well as by the buffers in the erythrocytes.
 The plasma contain carbonic acid/bicarbonate and acid/alkali sodium
salts of phosphoric acid as buffers.
 In the erythrocytes, the buffers are haemoglobin/oxyhaemoglobin and
acid/alkali potassium salts of phosphoric acid.
Biological fluids pH
Blood 7.4 – 7.5
Tear 7.0 – 8.0
Urine 4.5 – 8.0
Gastric juice 1.5 – 3.5
Bile 6.0 – 8.5
Saliva 5.4 – 7.5
 pH of some physiological
fluids are given in table;

The important biological buffer systems are the dihydrogen phosphate system and
the carbonic acid system.
1. The Phosphate Buffer System:
 The phosphate buffer system operates in the internal fluid of all cells.
 This buffer system consists of dihydrogen phosphate ions (H2PO4
-) as
hydrogen ion donor (acid) and hydrogen phosphate ions (HPO4
-2) as
hydrogen-ion acceptor (base).
 These two ions are in equilibrium with each other as indicated by the
chemical equation given below.
 If additional hydrogen ions enter the cellular fluid, they are consumed in
the reaction with HPO4
-2, and the equilibrium shifts to the left. If additional
hydroxide ions enter the cellular fluid, they react with H2PO4
-, producing
HPO4
-2 , and shifting the equilibrium to the right.
𝑯 𝟐 𝑷𝑶 𝟒(𝒂𝒒)
−
⇋ 𝑯(𝒂𝒒)
+
+ 𝑯𝑷𝑶 𝟒(𝒂𝒒)
−

2. The Carbonic Acid System:
 Another biological fluid in which a buffer plays an important role in
maintaining PH is blood plasma.
 In blood plasma, the carbonic acid and hydrogen carbonate ion equilibrium
buffers the PH.
 In this buffer, carbonic acid (H2CO3) is the hydrogen ion donor (acid) and
hydrogen carbonate ion (HCO3
-) is hydrogen-ion acceptor (base). The
simultaneous equilibrium reaction is shown below.
 This buffer functions in the same way as the phosphate buffer.
 Additional H+ is consumed by HCO3
- and additional OH- is consumed by
H2CO3.
𝑯 𝟐 𝑪𝑶 𝟑(𝒂𝒒) ⇋ 𝑯(𝒂𝒒)
+
+ 𝑯𝑪𝑶 𝟑(𝒂𝒒)
−

Buffered isotonic solution
 Tonicity is a property of a solution about a membrane, and is equal to
the sum of the concentrations of the solutes which have the capacity
to exert an osmotic force across that membrane.
 Tonicity depends on solute permeability.
If a semi-permeable membrane is used to separate solutions of
different solute concentrations, a phenomenon known as osmosis
occurs to establish concentration equilibrium.
 The pressure driving this movement is called osmotic pressure and is
governed by the number of particles of solute in solution.
 Tonicity is generally classified in three types;
1. Hypertonicity
2. Hypotonicity
3. Isotonicity

Hypertonicity:
A solution having higher osmotic pressure than the body
fluids (or 0.9% NaCl solution) is known as hypertonic solution.
 These solutions draw water from the body tissues to dilute and
establish equilibrium.
 An animal cell in a hypertonic environment is surrounded by a
higher concentration of impermeable solute than exists in the
inside of the cell causing it to shrink.

For example,
 If 2.0% NaCl solution is added to blood, osmotic pressure directs a net
movement of water out of the cell causing it to shrink (the shape of the cell
becomes distorted) and wrinkled, as water leaves the cell. This movement is
continued until the concentrations of salt on both sides of the membrane are
identical.
 Hence, 2.0% NaCl solution is hypertonic with the blood,

Hypotonicity:
A solution with low osmotic pressure than body fluids is
known as hypotonic solution.
 Administration of a hypotonic solution produces swelling of tissues
as water is pulled to the biological cells (tissues or blood cells) from
dilute the hypertonic solution.
 The effects of administering a hypotonic solution are generally
more severe than with hypertonic solutions, since ruptured cells can
never be repaired.
 Hypotonic solutions show opposite effect compare to hypertonic
solutions that the net movement of water is into the cell causing
them to swell.

For example,
 If 0.2% NaCl solution is added to blood, the cells get swelled and
burst.
 Therefore, O. 2% NaCl solution is hypotonic with respect to the
blood.

Isotonicity:
The solution that have the same osmotic pressure as that of
body fluids are said to be isotonic with the body fluid.
 Body fluids such as blood and tears have osmotic pressure
corresponding to that of 0.9. % NaCl or 5% dextrose aqueous
solution thus, a 0.9% NaCl or 5% dextrose solution is called as
isosmotic or isotonic.
 The term isotonic means equal tone, and is used interchangeably
with isosmotic regarding specific body fluids.
 A cell in an isotonic environment is in a state of equilibrium with its
surroundings with respect to osmotic pressure. When the amount of
impermeable solute is same on the inside and outside of the cell,
osmotic pressure becomes equal.

For example,
 On addition of 0.9 g NaCl/100 mL (0.9%) into blood, the cells
retain their normal size.

Methods Used to Determine Tonicity Value:
Many chemicals and drugs are used in the pharmaceutical
formulations. These substances contribute to the tonicity of the
solution. Hence methods are needed to verify the tonicity and adjust
isotonicity.
Two of the methods used to determine tonicity value are as follows.
1. Hemolytic method
2. Cryoscopic method

1) Hemolytic method:
 Isotonicity value is calculated by using hemolytic method in
which the effect of various solutions of drug is observed on the
appearance of red blood cells suspended in solutions.
 In this method, RBC's are suspended in various solutions and
the appearance of RBC's is observed for swelling, bursting,
shrinking and wrinkling of the blood cells.
 In hypotonic solutions, oxyhaemoglobin released is
proportional to number of cells hemolyzed; in case of
hypertonic solutions, the cells shrink and become wrinkled
where as in case of isotonic solutions the cells do not change
their morphology.

2) Cryoscopic method:
 Isotonicity values can be determined from the colligative
properties of the solutions.
 For this purpose, freezing point depression property is most
extensively used.
 The freezing point of water is 0º C, and when any substance
such as NaCl is added to it the freezing point of water decreases.
 The freezing point depression ∆Tf of blood is -0.52º C.
 Hence the ∆Tf value of the drug solution must be - 0.52º C. This
solution shows osmotic pressure equal to the blood and hence
RBC’s morphology found to be unchanged.

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