Buffers and Isotonic Solutions – Pharmaceutical Inorganic Chemistry B. Pharma 1st Semester

Buffers and Isotonic Solutions

Contents

• Definitions and concept of pH, pH scale and buffers

• Different buffer system and its compositions

• Pharmaceutical applications of pH and buffers

• Buffer action and mechanism of buffer action

• Buffer capacity and its importance

• Effect of pH and buffers on tissue irritation

• Isotonic solutions and its importance

• Paratonic solutions and its adverse effects on the
physiology

• Different methods of adjustment of tonicity

• Henderson-Hasselbalch equations for acidic and basic
buffers

• Applications of buffer equations

• Pharmaceutical buffer systems and the preparation of
pharmaceutical buffer system

• Physiological buffer systems and their importance

Learning
Objectives

At the end of this
lecture, student will be able to

• Explain the concept and importance of pH and pH scale

• Describe the importance of pH scale

• Explain the concept and composition of buffers and buffer
systems

• Describe the applications of pH and buffers

• Explain the buffer action and its mechanism

• Describe buffer capacity and maximum buffer capacity

• Discuss the relationship between buffer capacity and pH on
tissue irritation

• Explain the concept of isotonic solutions and its
importance in the physiological systems

• Explain the concept of paratonic solutions and its
importance in the pharmaceutical formulations

• Describe the different methods of adjustment of tonicity

• Discuss the buffer equation for weak acid and weak base

• Describe the application of buffer equation

• Explain the buffer equation for weak acid or base and its
salt

• Discuss the application of buffer equation

• Describe pharmaceutical buffer system and its importance

• Describe the method of preparation of pharmaceutical
buffers

• Explain the various biological buffer system and its
importance

pH and
Sorensen’s pH Scale

• In thermodynamic terms, pH is defined as negative
logarithm of activity of hydronium ions

• Sorensen defined pH as the logarithm of the reciprocal of
the hydrogen ion concentration

• Mathematically pH is expressed as:

pH=log 1/[H3O+] ………..(1)

Equation (1) may be rearranged as

pH=log1 – log[H3O+]…………(2)

Since log 1 is zero, equation (2) can be written as

pH= -log[H3O+]

• pH may be defined as negative logarithm of hydrogen ion
concentration

• Sorensen established the term pH, to represent hydrogen
ion potential

• Term p is used to express the negative logarithm.

• Concentration of H3O+ is expressed
as molarity, moles/liter etc.

• Solutions are stated as weakly acidic or strongly alkaline
and the extent of acidity of a solution may be explained by Sorensen’s scale

Sorensen’s
pH Scale

• Based on the pH values and different concentrations of H3O+
ions, a scale is devised and named after Sorensen, who developed it

• The scale starts with a zero pH, i.e., hydrogen ion
concentration is 1, which means the solution is strongly acidic

• At the other end of the scale, pH is 14, i.e., hydrogen
ion concentration is 10-14 strongly alkaline

• The central point pH is the scale is 7.0, because [H3O+]
is equalto [OH], i.e., hydrogen ion concentration is 10-7

• pH=7 means neutral

• The region with pH values below 7.0 is designated as
acidic and above PH 7.0 is designated as basic (or alkaline)

Applications
of pH

• Enhancing solubility

• Increasing stability

• Improving purity

• Optimizing biological activity

• Comforting the body

• Storage of products

Definition
and Applications of Buffers

• Buffers are compounds or mixtures of compounds that, by
their presence in solution, resist changes in pH upon the addition of small
quantities of acid or alkali

• The resistance to change in pH is known as buffer action

• Different characteristic properties of buffers are-

– they have a definite pH value

– pH value of buffers does not alter either on keeping for
long periods or on dilution

– The pH value of the buffer is very slightly altered by the
addition of small quantities of acids or alkalis

Applications
of Buffers

• The applications of buffers are-

– Enhancing solubility

– Increasing stability

– Improving purity

– Optimizing biological activity

– Comforting the body

Buffer
Systems- Composition and Examples

• A combination of two or more compounds is used in the
preparation of buffer solutions as described below-

-Weak acid and its conjugate base, i.e., the salt of weak
acid with a strong base. Example, a solution containing acetic acid and sodium
acetate

– Weak base and its conjugate acid, i.e., the salt of weak
base with a strong acid. Example, a solution containing ammonium hydroxide and
ammonium chloride

– Two salts can act as an acid-base pair. Example, a
solution of monobasic potassium phosphate (KH2PO4) and
dibasic potassium phosphate (K2HPO4)

– Amphoteric electrolytes act as buffer systems. Example is
the solution of glycine

– Solutions of strong acids and solutions of strong bases
exhibit buffer action by virtue of relatively high concentration of hydronium
ions and hydroxyl ions. For example, hydrochloric acid buffers cover the range
of 1.2 to 2.2, which include potassium chloride

• The solutions of drugs themselves manifest buffer action,
however their buffer capacities are low. A few examples are:

– Ephedrine in acidic media forms a salt of ephedrine
hydrochloride, which acts as buffer system similar to weak base and its salt
with strong acid

– When salicylic acid is stored in a soft glass bottle,
sodium ions in the container react with salicylic acid and forms sodium
salicylate. The solution behaves as a buffer similar to a weak acid and its
salt with strong base

Buffer
Action-Mechanisms

• Buffer action of acid buffer

• The ionization equation of a mixture of weak acid and its
salt, example, acetic acid and sodium acetate can be given as:

Strong electrolyte:

                                                                         
H2O

CH3COONa ———à
Na+ + CH3COO …………(completely ionized)

Weak acid:

CH3COOH + H2O ßà
H3O+ + CH3COO ……….(slightly
ionized)

• Therefore, the solution contains very few H3O+
ions, but has an excess of sodium ions and acetate ions

• When a small amount of acid is added, the H3O+
ions present in the solution reacts with CH3COO as:

H3O+ + CH3COO à
CH3COOH + H2O

• Since added free H3O+ are not
available, pH does not change

• When a small amount of base is added, the hydroxyl ions
furnished by the base are neutralised by acetic acid as:

OH + CH3COOH à
CH3COO + H2O

• Since added free OHions are not available, pH
does not change

• Thus buffer action is maintained when a small amount of
acid or base is added

Buffer action of
alkaline buffer

• Buffer action of a mixture of ammonium hydroxide and ammonium
chloride is considered and the equation can be given as :

• Strong electrolyte:

                                                                           
H2O

NH4Cl ———à
NH4+ + Cl ……..(completely ionized)

Weak base:

                                                                                
H2O

NH4OH ß——à NH4+
+ OH ……..(slightly ionized)

• The solution contains very few OH ions, but
has an excess of ammonium ions and chloride ions

• When a small amount of acid is added, the H3O+
ions obtained from acid react with NH4OH as:

H3O+ + NH4OH ßàNH4+
+2H2O

• Since added free H3O+ ions are not
available, pH does not change

• When a strong base is added, the hydroxyl ions furnished
by the base are neutralized by NH4+ as:

OH + NH4+ ßà
NH4OH

• Since added free OH ions are not available, pH
does not change

Buffer
Capacity

• Buffer capacity (also known as buffer efficiency, buffer
index or buffer value) is defined as the ratio of the increment of strong base
(or acid) to the small change in pH brought about by this addition

• In other words, the magnitude of the resistance of a
buffer to pH change is referred to as the buffer capacity

• Buffer capacity, β, is mathematically expressed as-

β= ΔB / ΔpH

where, Δ is a finite change, and ΔB is the small increment
in gram equivalents/liter of strong base added to the buffer solution to
produce a change of ΔpH

• Van Slyke’s equation for buffer capacity is represented
as:

β= 2.303C Ka[H3O+]
/ (Ka +[H3O+]) 2

Where C is the total buffer concentration

• Maximum buffer capacity can be given by βmax=0.576C

Influence
of Buffer Capacity and pH On Tissue Irritation

• If the buffer capacity is kept low, then the pH of the
solutions for introduction into the eye may vary from 4.5 to 11.5 without
marked pain or damage

• pH range of non-irritation cannot be established, but it
depends upon the buffer capacity of the buffer employed in the formulations

• Tissue irritation, due to large pH differences between the
solution being administered and the physiologic environment in which it is
used, will be minimal when:

– The lower is the buffer capacity of the solution

– The smaller is the volume used for a given concentration

– The larger the volume and buffer capacity of the
physiologic fluid

Buffered
Isotonic Solutions

• The solutions which have the same salt concentration and
the same osmotic pressure as the red blood cell contents is said to be isotonic
with blood

• Examples of isotonic
solutions
are- 0.9% w/v sodium chloride solution, 5%w/v dextrose solution
and 2% w/v boric acid solution

Buffered isotonic
solution
is defined as a solution which maintains the isotonicity and the
pH as that of the body fluids

Hypertonic
solutions
are those solutions containing the solute in higher concentration
than that is required for isotonic solutions. Examples- 2% w/v sodium chloride
solution, 10% w/v dextrose solution etc.

• When red blood cells are suspended in a hypertonic
solution, the water within the cells passes out through the cell membranes in
an attempt to dilute the surrounding salt solution. This outward passage of
water causes the cells to shrink and becomes wrinkled or crenated (crenulation)

Hypotonic solutions
are those solutions containing the solute in lower concentration than that is
required for isotonic solutions. Examples- 0.2% w/v sodium chloride solution,
3% w/v dextrose solution etc.

• When red blood cells are suspended in a 0.2% w/v solution
of sodium chloride, water enters the blood cells causing them to swell and
finally burst with the liberation of haemoglobin. This process is known as
haemolysis

Methods of
Adjusting Tonicity

• Class I methods-
sodium chloride or some other substance is added to the solution of the drug to
lower the freezing point of the solution to -0.520C, thus make the
solution isotonic with body fluids

• Two methods under class I are: Cryoscopic method and
Sodium chloride equivalent method

• Class II methods-
water is added to the drug in sufficient amount to form an isotonic solution.
The preparation is then brought to its final volume with an isotonic or a
buffered isotonic solution

• The two methods under class II are White-Vincent method
and Sprowls method

• Cryoscopic method-
the method involves the depression of freezing point by adding sodium chloride

• Sodium chloride equivalent method- the tonicic equivalent
or sodium chloride equivalent of a drug is the amount of sodium chloride that
is equivalent to 1g of the drug

• White- Vincent method

• Sprowls method

Buffer
Equation- Henderson- Hasselbalch Equation

• Buffer equation is also known as Henderson-Hasselbalch
equation

• Buffer equation is developed based on the effect of salt
on the ionization of a weak acid, when the salt and acid have a common ion

• An acid buffer, acetic acid and sodium acetate, is
considered for deriving the buffer equation

• The ionization equilibrium equation for weak acid (acetic
acid) may be shown as:

Weak acid:

CH3COOH + H2O ßàH3O+
+ CH3COO……..slightly ionized

• Applying the Law of Mass Action, the acid dissociation constant
(Ka) is written as:

Ka= [H3O] [CH3COO] / [CH3COOH]
=1.75X10-5 ………..(1)

• When sodium acetate is added to acetic acid, equation (1)
is momentarily disturbed

• As salt also supplies the acetate ion, the term [CH3COO]
in the numerator increases

• To re-establish the constant Ka, the [H3O]
in the numerator decreases and the equilibrium is shifted in the direction
shown below:

CH3COO + H3O+ à
H2O + CH3COOH

• Common ion [CH3COO] repressed
ionization of acetic acid and this is an example of common ion effect

• The pH of the final solution may be obtained by
rearranging equation (1)-

[H3O+]=Ka [CH3COOH] / [CH3COO]
…………..(2)

• Acid is weak and ionizes slightly, [CH3COOH]
may remain unaltered, hence, [CH3COOH] = [acid]

• Salt is completely ionized, the entire [CH3COO]
may be obtained directly from the salt and be written as [salt], hence, [CH3COO]=[salt]

• Substituting the values for [acid] and [salt] in equation
(2) gives:

[H3O+]=Ka[acid] / [salt] ……………..(3)

• Taking logarithms of equation (3) and reversing the signs:

-log [H3O]= -log Ka-log [acid] + log [salt]………….(4)

• Substituting, pH= -log [H3O] and pKa= -log Ka in equation
(4):

pH= pKa +log[salt] / [acid] …………..(5)

• Equation (5) is known as buffer equation or
HendersonHasselbalch’s equation for acid buffer

• Henderson-Hasselbalch equation for a weak base and its
salt is given by-

pH= pKw – pKb + log[𝐁𝐚𝐬𝐞]
/ [
𝐒𝐚𝐥𝐭]
………..(6)

Applications
of Buffer Equation

• For the preparation of a specified pH solution

• To calculate the pH of an unknown solution

• To predict the drug absorption based on the ionized and
unionized fraction of the drug molecules

• Determination of pKa based on the pH of the solutions

• Prediction of solubility based on pH of different
pharmaceutical solutions

• Selection of a suitable salt form

Pharmaceutical
Buffer System

• Pharmaceutical buffer systems are important in the
formulation of ophthalmic and parenteral drug delivery systems

• Two stock solutions suggested by Gifford, one containing
boric acid and the other monohydrated sodium carbonate, which when mixed in
various proportions, yield buffer solutions withpH values from 5 to 9

• Sorensen proposed a mixture of the salts of sodium
phosphate for buffer solutions of pH values 6 to 8

• A buffer system suggested by Palitzsch and modified by
Hind and Goyan consists of boric acid, sodium borate and sufficient sodium
chloride to make the mixtures isotonic

• The Clark-Lubs mixtures and their corresponding pH ranges
are as follows-HCl and KCl, pH 1.2 to 2.2

– HCl and potassium hydrogen phthalate, pH 4.2 to 5.8

• The Clark-Lubs mixtures and their corresponding pH ranges

– NaOH and potassium hydrogen phthalate, pH 2.2 to 4.0

– NaOH and KH2PO4, pH 5.8 to 8.0

– H3BO3, NaOH, and KCl, pH 8.0 to 10.0

Preparation
of Pharmaceutical Buffer Solutions

• The sequence of steps involved in the preparation of
buffers is as follows:

– A weak acid should be selected, which is having pKa value
approximately equal to the desired pH of the solution. This ensures maximum
buffer capacity

– From the buffer equation, the ratio of the salt and acid
needed for obtaining a suitable buffer capacity should be calculated. For a pH
range from 4 to 10, the buffer equation is satisfactory

-The individual concentrations of the buffer salt and acid
(or base), should be determined for a desired buffer capacity. A concentration
of 0.05 to 0.5 M is sufficient

– The buffer capacity of 0.01 to 0.1 is generally adequate

– The ingredients should be dissolved in carbon dioxide free
water and allowed to remain for some time to establish equilibrium condition

– The pH of the solution should be verified by a suitable
means, pH meter or pH indicator paper

• The procedure remains same for the preparation of basic
buffers

Physiological
(Biological) Buffer System

• Three important biological buffer systems are

1. Blood

– The pH of the blood is maintained at a pH of 7.4 by the
primary buffers in the plasma and secondary buffers in the erythrocytes

• The plasma contains carbonic acid/bicarbonate and
acid/alkali sodium salts of phosphoric acid as buffer

• In the erythrocytes the two buffer systems are
haemoglobin/oxyhaemoglobin and acid/alkali potassium salts of phosphoric acid

• The buffer equation for the carbonic acid/bicarbonate
buffer of the blood is:

pH=6.1+log [HCO3]/[H2CO3]

Where [H2CO3] represents the
concentration of CO2 present as H2CO3
dissolved in the blood.

2. Lacrimal Fluid or
Tears

• The pH of tears is about 7.4 with a range of 7 to8 or
slightly higher

• Lacrimal fluids have been found to have a great degree of
buffer capacity, allowing dilution of 1:15 with neutral distilled water

3. Urine

• The urine of a normal adult has a pH averaging about 6
units

• It may be as low as 4.5 or as high as 7.8

• When the pH of the urine is below normal values, hydrogen
ions are excreted by the kidneys

• When the urine is above pH 7.4, hydrogen ions are retained
by action of the kidneys in order to return the pH to its normal range of
values

Summary

pH- Negative
logarithmic value of hydronium ion concentration

pH scale-
Devised by Sorensen, helps in determination of acidity and basicity of a
chemical substance

Buffers- These
are mixture of compounds that resist the change in pH

• A combination of two or more compounds are used in the
preparation of buffer solutions

• The combination in buffer solutions are a weak acid and
its conjugate base or a weak base and its conjugate acid

• The typical example of drug as buffer systems are-
ephedrine / ephedrine hydrochloride and salicylic acid and sodium salicylate

• Buffer action and
its mechanism-
The action by which buffers resist the change in pH. It acts
by common ion effect

• Buffer capacity-
The intensity of buffer action is buffer capacity

• Buffer capacity, pH and tissue irritation-Low buffer
capacity and physiological pH will reduce tissue irritation

• Isotonic solutions-
Solutions having the same tonicity as that of the body fluids

• Paratonic
solutions-
Solutions having more tonicity (hypertonic or hypotonic) than
that of the body fluids

• Methods of adjustment of tonicity- Different methods are
cryscopic method, sodium chloride equivalent method, whiteVincent and sprowls
method

• Buffer equation –
Explained by Henderson-Hasselbalch and is given by:

pH= pKa +log[acid] / [salt] …………..(weak acid)

pH= pKw- pKb + log [𝐁𝐚𝐬𝐞]
/ [
𝐒𝐚𝐥𝐭]
…………..(weak base)

• Buffer equations are used for the preparation and
stabilization of different pharmaceutical preparation

• Pharmaceutical
buffer system-
Different buffer systems proposed by various scientist, can
be prepared in lab scale

• Typical examples of pharmaceuticalbuffer systems are: NaOH
and potassium hydrogen phthalate, pH 2.2 to 4.0

– NaOH and KH2PO4, pH 5.8 to 8.0

-H3BO3, NaOH, and KCl, pH 8.0 to 10.0

• Method of
preparation of pharmaceutical buffer-
Different step by step methods are
followed by choosing a suitable weak acid or a weak base for the preparation of
a buffer system

• Biological buffer
systems –
Different biological buffer systems are:

– Blood

– Lachrymal fluid

– Urine

 

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