Buffer for physical pharmacy

1. UNIT: 8 BUFFERED AND ISOTONIC SOLUTIONS
BUFFERS:
 Buffers are compounds or mixture of compounds that, by their presence in solution
resists change in PH upon the addition of small quantities of acid or alkali.
 The resistance to change in PH is known as buffer action.
 If a small amount of strong acid or base is added to water or a solution of sodium
chloride the PH is altered considerably; such system have no buffer action.
BUFFER ACTION:
 The ability of certain solutions to resists change in their PH upon addition of an
acid or base is known as buffer action.
 Let us make an example of a solution of sodium chloride in water. Its PH value is7
 Addition of even 1ml of 1N HCL to 1 litre of sodium chloride solution lowers its
PH value from 7 to about 3.similarly the addition of even 1ml of 1N NaOH
solution to 1 litre of Nacl solution raises its PH to about 11.sodium chloride
solution is not a buffer.
 Now, let us take an example of a mixture of a weak acid (acetic acid) and its salt
(sodium acetate).
 Acetic acid is very slightly disassociated in solution while sodium acetate being a
salt is almost completely disassociated.
 The mixture thus consists of CH3COOH molecules as well as CH3COO- and Na+
If a strong acid is added to the mixture,the H+ ions supplied by the acid are
immediately taken up by CH3COO- ions to form very slightly disassociated
CH3COOH.
H+ +CH3COO- = CH3COOH
 The hydrogen ion is neutralized by acetate ion present in the mixture and there is
very little change in the PH value of mixture.
 On the other hand, if a strong base is added, the OH- ions supplied by the base are
neutralized by acetic acid present in the mixture and again there is a very little
change in the PH of the solution.
OH- +CH3COOH = CH3 COO- + H2O
BUFFER EQUATION
Buffer equation for a weak acid and its salt
 Let us take an example of the effect of addition of sodium acetate on the
ionization of acetic acid. Both sodium acetate and acetic acid have ion in
common between them,i.e CH3COO-
 The disassociation constant for the acid is given by

2.  CH3COOH + H2O = CH3COO- +H3O
Ka = [CH3COO-] [H3O]
[CH3COOH]
 If sodium acetate is added to the acetic acid solution it ionizes to give acetate
ion.
 CH3COONa = CH3COO- + Na+
This causes a momentary increase in the concentration of CH3 COO – in the
solution. To re-establish the disassociation constant for the solution, the
hydrogen ion term in the numerator is instantaneously decreased.
This results in the increase of the concentration of CH3COOH in the
denominator.
 A mixture of a weak base and its salt also behave in a similar manner.
 Let us take an example of a mixture containing equimolar solutions of
ammonium hydroxide and its largely disassociated salt, ammonium chloride.
 The mixture contains undisassociated NH4 OH as well as NH4
+ and Cl- ions.
 If a strong acid is added to this mixture, the H+ ions supplied by the acid are
neutralized by the base NH4OH.
 H+ + NH4 OH = H2O + NH4+
 On the other hand if a strong base is added, the OH- are neutralized by NH4
+
ions forming very slightly disassociated NH4OH.
 OH- + NH4 =NH4 OH
 In both the cases there is very little change in the PH.
 The pH of a buffer solution and the change in pH upon the addition of an acid or
base can be calculated by use of the buffer equation.
 This expression is developed by considering the effect of a salt on the ionization
of a weak acid when the salt and the acid have ion in common.
 H3O + CH3COO- = CH3COOH + H2O
 The ionization of acetic acid is expressed upon the addition of the common
acetate ion to the solution and this is known as common ion effect.
 The pH of the buffer solution can be obtained by rearranging the above equation
for disassociation constant.
 [ H3O] = [CH3COOH]
Ka [CH3COO-]
 [ H3O] = [Acid]
Ka [Salt]
 Log[H3O] =LogKa +Log[Acid]-Log[salt]
 Reversing the sign we get,
 -Log[H3 O]=-LogKa - Log [Acid] + Log [salt]
 PH = Pka + log [salt]/ [Acid]

3.  This is known as buffer equation or the Henderson – Hasselbalch equation for a
weak acid and its salt.
 The buffer equation is a useful equation (expression) used in the preparation of
buffered pharmaceutical solutions.
 It is satisfactory for calculation within the PH range 4 to 10.
 In actual condition there is small difference in the PH of the solution, since the
activity coefficient of the components varies with the concentration taking into
consideration the activity coefficient.
 PH = Pka + log [salt]/ [Acid] x LogÝAC
 Where ÝAC indicates the activity coefficient of the common ion.
 Numerical:1 Calculate the PH of the buffer solution consisting of 0.1M each of acetic
acid and sodium acetate [ Pka of acetic acid=4.76]
Solution:
Concentration of acetic acid [Acid] = 0.1M
Concentration of sodium acetate [salt]= 0.1M
Pka of acetic acid=4.76
According to the Henderson – Hasselbalch equation,
PH = Pka + log [salt]/ [Acid]
PH = 4.76+ log [0.1]/[0.1]
PH = 4.76
Hence, the PH of the buffer solution = 4.76
Buffer equation for weak base and its salt
Although buffers are not usually prepared using weak bases and their salt because of the poor
stability, certain pharmaceutical solutions such as solutions of ephedrine base and ephedrine
hydrochloride which are mixture of weak bases and their salts are often encountered.
The buffer equation for solution of weak bases and their corresponding salt can be derive in a
similar manner to that of weak acid buffers.
[OH-] = Kb [base]/ [salt] ………… (i)
Now since ionic product of water,
Kw = H3O+ X OH-
OH- = Kw / H3O+ ……………..(ii)
Expressing in logarithmic terms,
Log Kw – Log [H3O+] = Logkb + Log base – Log [salt]

4. Rearranging the given terms;
– Log [H3O+] = - Log Kw+ LogKa + Log base – Log [salt]
PH = Pkw – Pkb + Log [base]/[salt]
Numerical 2: Estimate the pH of a solution containing 0.10 mole of Ephedrine and 0.01 mole
of Ephedrine hydrochloride per litre of solutions. [Pkb of Ephedrine is 4.64]
Solution:
Concentration of Ephedrine [Base] = 0.10 mole
Concentration of Ephedrine Hydrochloride [salt] = 0.01 Mole
Pkb of Ephedrine = 4.64
Ionic product of water (Pkw) = 14
For a weak base and its salt;
PH = Pkw – Pkb + Log [base]/[salt]
PH = 14 – 4.64 + Log [0.10] / [0.01]
PH = 14 – 4.65 + 1
PH = 10.36
BUFFER CAPACITY
 The magnitude of the resistance of a buffer to pH changes is referred to as buffer
capacity. It is also known as buffer efficiency, buffer index and buffer value.
 Koppel, Spiro, Van slyke introduced the concept of buffer capacity and defined it as the
ratio of the increasement of strong base to the small changes in pH brought about by this
addition. For the present discussion the approximate formula can be used.
 β = ∆ B
∆PH
 In which delta has its usual meaning a finite change and ∆B is the small increasement in
gram equivalent /litre of strong base added to the buffer solution to produce a pH change
of ∆ pH.

5.  According to the equation, the buffer capacity of a solution has a value of 1 when addition
of 1 gram equivalent of strong base (or acid) to 1 litre of buffer solution results in a
change of 1 pH unit.
 The significance of this index will be appreciated better when it is applied to the
calculation of the capacity of a buffer solution.
APPROXIMATE CALCULATION OF BUFFER CAPACITY
 Consider an acetate buffer containing 0.1 Mole each of acetic acid and sodium acetate in
1 litre of solution.
 To this added 0.01 Mole portions of sodium hydroxide. When the first increments of
sodium hydroxide is added the concentration of sodium acetate, the [salt] term in the
buffer equation increases by 0.01 mole/litre and the acetic acid concentration [acid]
decreases proportionately because each increments of base converts 0.01 mole of acetic
acid into 0.01 mole of sodium acetate according to the reaction
 HAc + NaOH = NaAc + H2O
[0.1-0.01] [0.01] [0.1 + 0.01]
The change in concentration of the salt and the acid by the addition of a base are
represented in the buffer equation, by using modified form;
PH = Pka + log [salt] + [Base]
[Acid]- [Base]
Before the addition of the first portion of sodium hydroxide the PH of the buffer
solution is
PH = 4.76 + Log 0.1 + 0 / 0.1-0 = 4.76
 The buffer capacity is not a fixed value for a given buffer system, but instead depends
upon the amount of base added.
 The buffer capacity changes as the ratio Log [salt/acid] increases with added base.
 With the addition of more sodium hydroxide, the buffer capacity decreases rapidly and
when the sufficient has been added to convert the acid rapidly into sodium ions and
acetate ions the solution no longer passes an acid reverses.
 The buffer system has its greatest capacity before any base is added and therefore
according to the equation,PH = Pka
 The buffer capacity is also influenced by an increase in the total concentration of the
buffer constituents because obviously, because a great concentration of salt and acid
provide a greater alkaline and acid reverse.
 A more exact equation for buffer capacity is given below;
 The buffer capacity calculated is only approximate. It gives the average buffer capacity
over the increasement of base added.
 Koppel, spiro and van slyke develop more exact equation.

6.  Β = 2.3 C Ka [H3O+]
(Ka + [H3O+] )2
Where C is the total buffer concentration that is the sum of the molar concentration of the
acid and salt.
MAXIMUM BUFFER CAPACITY
A buffer solution containing a weak acid and its salt has a maximum buffer capacity βMax
When PH is equal to the Pka value for the weak acid or in equivalent terms [ H3O] equals
pka
Substituting [H3O] for Ka in the above experiment we get,,
βMax = 2.303/4 x C
βMax = 0.576 C
Numerical 3:
A buffer solution contained 0.1 M each of acetic acid and sodium acetate and its PH was
4.76 .To this 0.01 moles of sodium hydroxide was added and pH of the resultant solution
was 4.85. Calculate the buffer capacity β.
Solution:
Change in pH (∆ pH) = 4.85 – 4.75 = 0.09
Quantity of sodium hydroxide added (∆ B) = 0.01 mole’
Buffer capacity,β
β = ∆B
∆Ph
β = 0.01/0.09 = 0.11
BILOGICAL BUFFERS
 The PH of the blood is maintained at about 7.4 by primary buffer components in the
plasma and secondary buffer components in the erythrocytes.
 The plasma contains carbonic acid- carbonate and acid /alkali sodium salt of phosphoric
acid while erythrocytes contains two buffer system Oxyhaemogloblin-hemoglobin and
acid alkali potassium salt of phosphoric acid.
 Values of buffer capacity of the blood ranging from 0.025 to 0.039 gram equivalents per
PH unit have been reported in the literature.
 The physiological pH range of the blood is 7.0 to 7.8.
 When PH of the blood decreases or rises beyond this range, life is in danger.
 The pH of the lacrimal fluid or tears is about 7.4 with a range of 7.0 to 7.8.
 Tears have found to have a great degree of buffer capacity and dilution up to 1:15 with
neutral distilled water is possible before an appreciable change in PH is noticed.

7.  Similarly, the average PH of urine is about 6.0 with a range of 4.5 to 7.8 When PH of urine
is altered beyond this range, remedial action is taken by the kidneys in the form of
retention or excretion of hydrogen ions in order to maintain the PH within the range.
PHARMACEUTICAL BUFFERS
 Buffers are widely used in the field of pharmacy or ingredients in pharmaceutical
formulations either to adjust the PH of the product to that required for maximum stability
or for maintaining the PH of the product.
 In the field of scientific research, buffers have generally been used to control PH within a
certain range so that, for Example rates of reaction can be investigated.
BUFFERS IN TABLET FORMULATION
 Buffers have been used in tablets and capsules to control the PH in the microenvironment
surrounding the drug particles.
 This is especially helpful in cases of drugs where the absorption is dissolution rate limited
from the unbuffered formulations.
 Buffers have also been employed in formulations containing acidic drugs to reduce
gastric irritation.
 Buffering agents that have been used in solid oral dosage forms include antacids such as
sodium bicarbonate, magnesium carbonate and sodium citrate.
BUFFERS IN OPTHALMIC PREPRATION
 Buffers are generally used in ophthalmic preparations to maintain the PH within the
physiological PH range of the lacrimal fluid.
 The lacrimal fluid has good buffering capacity and solutions with PH values between 3.5
and 10.5 can usually be tolerated with little discomfort.
 Outside this PH range irritation of the eye accompanied by increase lacrimation occur.
 Ideally ophthalmic preparations should be formulated at physiological PH but often this
PH is not the ideal one for best solubility and/or stability of the drug.
 Most of the ophthalmic drugs are weakly acidic or basic. At low PH this drug
disassociated and goes into the solution and remains stable.
 However at this PH the therapeutic effect is lower since only the undisassociated form of
the drug is able to penetrate the lipoidal membrane.
 Buffers are therefore added to adjust the pH to a value that is best with regard to the
solubility and stability of the drug and which is well tolerated by the eye.
 When solution is instilled into the eye, the pH slowly rises to that of tear solution and the
weak acid or base (drugs) gets converted into the undisassociated form which provides
the maximum therapeutic effect.

8.  The buff,ering agents most commonly used in ophthalmic preparations include borate,
phosphate and carbonate buffers.
 These preparations are also made isotonic to prevent discomfort and injury to the surface
of the eye.
BUFFERS IN PARENTERAL PREPRATIONS
 Consideration of pH is important in case of parenteral products since highly alkaline PH
(above 9) can cause tissue necrosis which an acidic pH [below 3] can result in extreme
pain at site of injection
 The ideal pH of a parenteral product is 7.4, the pH of blood. However blood being a good
buffer itself, the pH of small volume parenteral is not necessarily required to be at
physiological pH.

 The pH selected for such product is generally a compromise between the stability and
solubility of the medicament as well as the irritancy of the preparations.
 Buffers are usually added for adjusting the PH of parenteral products to a suitable value.
 The buffer capacity of small-volume parenteral is however kept low so that PH can be
adjusted by the blood-buffer system.
 The most commonly used buffers in parenteral products are acetate, phosphate, and
citrate and glutamate buffers.
BUFFERS IN CREAM AND OINTMENTS
 Topical products usually have a tendency to undergoes change in PH during storage
which may adversely affect the stability of the product.
 Buffers are therefore included in such preparation to maintain the stability of the product.
 The most commonly used buffers in creams and ointments include citric acid and its salt
or phosphoric acid and its salt.
PREPRATION OF PHARMACEUTICAL BUFFER
 Buffer solutions are used frequently in pharmaceutical practice, particularly in the
formulation of ophthalmic solutions.
 They also find application in the colorimetric determination of PH and research studies in
which PH must be held constant.
Gifford suggested two stock solutions, one containing boric acid and other
monohydrated sodium carbonate, which when mixed in various proportions, yeild buffer
solutions. PH values ranges from 5 to 9.

9.  Sorensen proposed a mixture of the salts of sodium phosphate for buffer solutions of PH
6 to 8.
 Sodium chloride is added to each buffer mixture to make it isotonic with body fluids.
 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 mixture isotonic.
 It is used in the ophthalmic solutions in the PH range of 7 to 9.
The clark-lubs mixtures their corresponding PH ranges
are as follows
 Hcl and Kcl,PH 1.2 to 2.2
 Hcl and potassium hydrogen pthalate, PH 2.2 to 4.0.
 NaOH and potassium hydrogen phthalate, PH4.2 to 5.8.
 NaOH and KH2PO4,PH 5.8 TO 8.0
 H3BO3,NaOH and KCl,PH 8.0 TO 10.0
The following steps should be helpful in the development
of a new buffer
 Select an acid having a Pka of approximately equal to the PH at which the buffer is to be
used. This will ensure the maximum buffer capacity.
 From the buffer equation, calculate the ratio of salt and weak acid required to obtain the
desire PH.The buffer equation is satisfactory for approximate calculation within the PH
range of 4 to 10.
 Consider the individual concentration of the buffer salt and acid needed to obtain a
suitable buffer capacity. A concentration of 0.05 to 0.5 M is usually sufficient, and a
buffer capacity of 0.01 to 0.1 is generally adequate.
 Other factors of some importance in the choice of a pharmaceutical buffer include
availability of chemicals, sterility of the final solution, stability of the drug and buffer on
aging, cost of materials and freedom from toxicity.
 Finally determine the PH and buffer capacity of the completed buffer solution using a
reliable PH meter. In some cases, sufficient accuracy is obtained by the use of PH papers.

10. BUFFERED ISOTONIC SOLUTIONS
 Two solutions are said to be iso-osmotic or isotonic if they exert the same osmotic
pressure when separated by a semi-permeable membrane.
 Physiologically, isotonic solutions are solutions having the same osmotic pressure as that
of the body fluids when separated by a biological membrane.
 Body fluids including blood and lachrymal fluid normally have an osmotic pressure
corresponding to that of 0.9% solution of sodium chloride.
 Thus 0.9% solution of sodium chloride is said to be isotonic with the physiological fluid.
 Solutions with osmotic pressure lower than that of the body fluids or of 0.9% sodium
chloride solutions are commonly referred to as hypotonic and those having a higher
osmotic pressure are termed as hypertonic.
 if red blood cells are suspended in a 2% sodium chloride solution, the water with in the
cells passes through cell membrane.
 In an attempt to dilute the surrounding salt solution until the salt concentration on both
sides of the erythrocytes membrane are identical.
 This outside passage of water causes the cells to shrink and become wrinkled and
crenated. The salt solution in these instances is said to be hypertonic with respect to blood
cells contents.
 Finally if the blood is mixed with 0.2% sodium chloride solutions or with distilled water,
water enters the blood cells causing them to swell and finally brust, with the liberation of
hemoglobin. This phenomenon is known as haemolysis, and the weak salt solution or
water is said to be hypotonic with respect to blood.
MEASUREMENT OF TONICITY
 HAEMOLYTIC METHOD
 COLLIGATIVE METHOD

11. HAEMOLYTIC METHOD:
 In this method red blood cells are suspended in the solution whose tonicity is to be
determined. If the solution causes shrinkage of the cells, they are said to be hypertonic
and if they cause haemolysis they are said to be hypotonic.
 Quantitative measurements are possible using this method based on the facts that
hypotonic solution liberates Oxyhaemogloblin in direct proportion to the number of cell
haemolysed.
COLLIGATIVE METHOD
 It has been determined that solution having same tonicity exhibit similar behavior with
respect to their colligative properties such as lowering of vapour pressure, depression in
freezing point,etc.
 Hence isotonicity of a solution may be determined by determining its colligative
properties.
CALCULATION OF TONICITY USING Liso VALUES
 The depression in freezing point of solution of weak as well as strong electrolytes may be
given by the equation :
 ∆Tf = kf C
 Where ∆Tf is the depression in freezing point, Kf is the molal depression constant and c
is the molar concentration.
 However it has been found that the actual freezing point depression is always greater than
the value obtained by using the above equation. This is actually due to the deviation from
ideal behavior of solutions.
 In order to compensate for this deviation, a factor [ i] known as vant Hoff factor is
introduced.
 In other words,[ i] represents the number of times greater that the colligative effect is for
a real solution( of an electrolyte or non – electrolyte) than for an ideal solution.
The equation thus becomes: ∆Tf =ikf C
Substituting ikf with a new factor L we get ;
∆Tf =LC
L = ∆Tf / c

12.  The L value can be obtained from the value of freezing point depression of solutions of
representing compounds of a given ionic type at a concentration c that is isotonic with the
body fluids. This specific value of L is represented as ‘Liso’
 The ‘Liso’ value is therefore defined as the specific value of L which is equal to ikf at a
concentration of drug that is isotonic with body fluids.
 The Liso value for a 0.9 % sodium chloride solution whose freezing point depression is
0.520 c and therefore isotonic with body fluids is :
Liso = ∆Tf / c = 0.52/ 0.154
=3.4
 Where 0.154 represents the molar concentration of 0.9 % sodium chloride solution.
Compounds of similar ionic types have similar Liso values.
METHODS OF ADJUSTING ISOTONICITY
 In order to render a hypotonic solution, isotonic with body fluids, substances such as
sodium chloride and dextrose are added to it.
 Alternatively, water is added to drug substances in sufficient amount to form an isotonic
solution which is further diluted with an isotonic or a buffered isotonic solution to give
the final volume.
The following methods are generally followed for adjustment of tonicity
 Cycroscpic orFreezing point depressionmethod
 Sodium chloride equivalent method
 White-Vincent method
 Sprowls Method
1) Cycroscpic orFreezing point depressionmethod:
 Body fluids such as plasma and lachrymal secretions have a freezing point of – 0.520 C
by virtue of different solute present in them. Hence All solutions which freeze at – 0.520
C will be isotonic with these fluids.
 Depression in freezing point being an additive property, the following formula is used
for the calculation of the quantity of a substances required to make solutions isotonic
with physiological fluids.
% W / V of adjusting substances = [0.52 – a] / b

13.  Where, a represents the depression in freezing point due to unadjusted solutions
or substance ; b represents the depression in freezing point of 1 % w/v of
adjusting substances.
2) Sodium chloride equivalent method: The sodium chloride equivalent or tonic
equivalent of a drug is the amount of sodium chloride that is equivalent to 1 gram of the drug
osmotically. The sodium chloride equivalents E can easily be in the literature. Alternatively, the
value can be calculated from the Liso values using the following equation.
E = 17 Liso / M
Where M is the molecular weight of the drug.
Preparation of isotonic solution using the sodium chloride equivalent values ‘E’ simply involve
multiplying the quantity of each drug in the prescription by its sodium chloride equivalent and
subtracting these values from the concentration of sodium chloride that is isotonic with body
fluids, i.e 0.9 g / 100 ml
Thus, quantity of sodium chloride required to render 100 ml solutions containing x gram of
drug isotonic
= 0.9 – (xE)
3) White-Vincent method
This method involves the addition of sufficient quantity of water to a drug in order to prepare an
isotonic solution. An isotonic or a buffered isotonic solution is then added to this drug solution to
give the final volume. The volume of the water required for a particular quantity of drug to
prepare an isotonic solution can be calculated from the following equation:
V= w x E x 111.1
Where v is the volume in milliliters of an isotonic solution that can be prepared by dissolving w
gm of drug in water; E is the sodium chloride equivalent of the drug and 111.1 is the constant
representing the volume in milliliters of isotonic solution obtained by dissolving 1 gm of sodium
chloride in water.
4) Sprowls Method
A further modification of white – Vincent method is the Sprowls method which uses tables
listing the volume V of isotonic solution that can be prepared by mixing 0.3 g of a drug in water

14. B.pharm 2nd semesterPurbanchal University
PREPRAED BY: Prakash Babu Dahal