Acids, Bases And Buffers Pharmaceutical Inorganic chemistry UNIT-II (Part-I)

Pharmaceutical Inorganic chemistry
UNIT-II (Part-I)
Acids, Bases And Buffers
Presented By
Ms. Pooja D. Bhandare
(Assistant Professor)
DADASAHEB BALPANDE COLLEGE OF PHARMACY BESA NAGPUR

Acids, Bases are defined by Four main theories,
1.Traditional theory / concept
2.Arrhenius theory
3.Bronsted and Lowry theory
4.Lewistheory

1. Traditional theory / concept-Acid:
• Acids: are the substances
• Which converts blue litmus paper to red
• Having the PH<7
• Sour taste
• React with bases to form salts and water
• Eg:-Hydrochloric acid(HCl)

1. Traditional theory / concept-Base:
Base: are the substances
• Which converts red litmus paper to blue
• Having the PH>7
• Bitter taste
• React with Acids to form salts and water
• Eg: Sodium Hydroxide (NaOH)

2. Arrhenius theory:
The Swedish chemist Svante Arrhenius proposed the first definition of acids
and bases
• According to Arrhenius Concept, Acid are substance which are capable of
providing
• Hydrogen ions (H+, proton) when dissolved in water and bases are
substances which are capable
• of providing hydroxide ions (OH-, hydroxyl ions) in aqueous solution.
• For example, Hydrochloric acid in the water, HCl undergoes dissociation
reaction to
• produce H+ ion and Cl– ion, as explained below. The concentration of the
H+ ions is increased by
• forming hydronium ion.

• HCl (aq) → H+(aq) + Cl– (aq)
• HCl (aq) + H2O(l) → H3O+(aq) + Cl– (aq)
• Other examples of Arrhenius acids are listed below
• NHO3(aq) + H2O(l) → H3O+ (aq) + NO3–
• In this reaction, nitric acid dissolves in aqueous water to give hydrogen
and nitrate ions.
• Another example of Arrhenius Base is
NaOH + H2O → Na+ + OH- + H2O

Neutralization reaction
Acid react Base
Salt &Water
Eg: Hydrochloric acid react sodium hydroxide
Sodium chloride (Salt) & water
NaOH + HCl------------------- NaCl + H2O

• “Neutralization as the process in which hydrogen ion and hydroxyl ion
combine to form unionized molecule or water”
• NaOH + HCl------------------- NaCl + H2O
• HCl (aq) → H+(aq) + Cl– (aq)
• NaOH + → Na+ + OH-
• H+ + OH- → H2O

Application
1. Non-metallic Oxides are arrhenius acidic in nature
eg CO2, SO2, SO3, N2O5, P4O10
• CO2 + H2O ------------------- H2CO3------ H+ + HCO3
• SO3 + H2O ---------------------H2SO4 -------H+ + HSO4-
2. metallic Oxides are arrhenius bases
eg CaO, Na2O

Limitations of Arrhenius theory:
• The Arrhenius theory is applicable only in aqueous solution; for example, according to the
theory, HCl is an acid in the aqueous solution but not in benzene, even though it donates H+
ion to the benzene. Also, under Arrhenius’s definition, the solution of sodium amide in liquid
• ammonia is not alkaline, even though amide ion deprotonates the ammonia.
• Basisity of Ammonia (No OH-ion)is not explained
• Acidity of BF3,AlCl3 (No H+ ion)is not explained
• Acidity of oxides of P block element (CO2) is not explained
• Basicity of oxides of S block element (Na2o)is not explained
• Neutralization with out absence of solvent is not explained

3.Bronsted -Lowry concept:
Bronsted -Lowry concept: In 1923 the Danish chemist Johannes Nicolaus Bronsted
and the English chemist Thomas Martin Lowry, proposed the theory.
• According to Bronsted-Lowry theory, An acid is any substance (molecular or ionic)
that can donate a proton to any other substance (molecular or ionic) and a base is
any substance that can accept a proton from anyother substance.
• HCl + H2O H3O+ + Cl
• In the above example what is the Bronsted acid? What is the Bronsted base?
In reality, the reaction of HCl with H2O is an equilibrium and occurs in both
directions, although in this case the equilibrium lies far to the right.
HCl + H2O H3O+ + Cl-

• For the reverse reaction Cl- behaves as a Bronsted base and H3O+
behaves as a Bronsted acid.
• The Cl- is called the conjugate base of HCl. Bronsted acids and bases
always exist as conjugate acid-base pairs. Their formulas differ by only
one proton.
• Acid Base conjugate acid conjugate base
HCl + NH3 NH4+ + Cl-

1.Amphoteric: a species that can act as an acid or a base water is an
example of an amphoteric species.
2.Conjugatebase: species that remains after an acid donates its H+.
3.Conjugateacid: species that forms after a base accepts a H+
• 14

• Every Arrhenius Acid is Bronsted Acid
• Every Arrhenius Base is not Bronsted Base e.g., NaOH is Arrhenius base
because it gives
• OH- ion in aqueous solution but not a Bronsted base because it cannot accept
proton.
• Limitations of Bronsted Lowry Concept:
• The protonic definition cannot be used to explain the reactions occurring in
non-protonic solvents such as COCl2, SO2, N2O4, etc.
• Substances like BF3, AlCl3 etc, do not have any hydrogen and hence cannot
give a proton but are known to behave as acids

Lewis Theory
• In 1923 of scientist G.N. Lewis proposed the theory in terms of chemical
structure.
• Lewis Acids:
• Lewis acids accept an electron pair. Lewis Acids are Electrophilic meaning
that they are electron attracting.
• Various species can act as Lewis acids. All cations are Lewis acids since
they are able to accept electrons. (e.g., Cu2+, Fe2+, Fe3+)
• Lewis acids- H+, NH4+, Na+, K+, Cu2+, Al3+, etc.

• Lewis Bases
• Lewis Bases donate an electron pair. Lewis Bases are Nucleophilic meaning that they
“attack” a positive charge with their lone pair. An atom, ion, or molecule with a lonepair of
electrons can thus be a Lewis base.
• Lewis base- NH3, H2O, OH-, Cl-, CN-, S2-, etc.
• (Lewis base) (Lewis acid)
• Boron trifluoride accepts the electron pair, so it is a Lewis acid. Ammonia makes available
(donate) the electron pair, so it is the Lewis base.

Importance of acids and bases in pharmacy
• Acids, bases and their reaction play vital role in pharmacy practice. Some of the
main application of the these are as follows:
• Acid-base neutralization reaction finds use in preparative procedures for the
preparation of suitable salt, and for conversion of certain salts into more suitable
forms.
• Acid-base is used in analytical procedure which is involving acid-base titrations.
• Acids and bases find use as therapeutic agents in the control of and adjustment of
pH of the GI tract, body fluids and urine.

Buffers:
• A buffer is a solution that can resist pH change upon the addition of an
acidic or basic components.
• It is able to neutralize small amounts of added acid or base, thus
maintaining the pH of the solution relatively stable.
• •This is important for processes and/or reactions which require specific
and stable pH ranges.

• Buffers: Buffers are defined as a compound or a mixture of compounds
that resists the pH upon the addition of small quantities of acid or alkali.
Buffer have definite pH value.
• The pH will not change after keeping it for a long period of time. The pH
value altered negligibly by the addition of small quantities of acid or base.
• Buffer action: The resistance to a change in pH is known as buffer action.
So buffers can be added to show buffer action.
• Buffer capacity: The amount of acid/base required to produce a unit
change in pH in a solution is called buffer capacity.

Buffers system:
• A buffer system can be made of a weak acid and its salt or a weak base
and its salt.
• A classic example of a weak acid based buffer is acetic acid
(CH3COOH) and sodium acetate(CH3COONa).
• A common weak base buffer is made of ammonia (NH3) and ammonium
chloride (NH4Cl)

Types of Buffers :
Generally buffers are of two types:
1. Acidic buffers
2. Basic buffers
There are some other buffer system:
3. Two salts acts as acid-base pair. Ex- Potassium hydrogen phosphate and potassium
dihydrogen phosphate.
4. Amphoteric electrolyte. Ex- Solution of glycine.
5. Solution of strong acid and solution of strong base. Ex- Strong HCl with KCl.

1. Acidic Buffers:
• An acidic buffer is a combination of weak acid and its salt with a strong
base. i.e. Weak acid & salt with strong base (conjugate base).
• EXAMPLES:
CH3COOH / CH3COONa
H2CO3 / NaHCO3
H3PO4 / NaH2PO4
HCOOH / HCOONa

2. Basic Buffers:
• A basic buffer is a combination of weak base and its salt with a strong
acid.i.e. Weak base & salt with strong acid (conjugate acid).
• EXAMPLES:
NH4OH / NH4Cl
NH3 / NH4Cl
NH3 / (NH4)2CO3

Mechanism of Buffer action:
Mechanism of Buffer action:
• The resistance of a buffer solution to a change in pH is known as buffer action.
• In a buffer solution, the components interact with each other and produce a
dynamic equilibrium.
• When a small quantity of acid or base is added, the dynamic equilibrium shifts
and nullifies the effect of the addition.

 Mechanism of Action of acidic
buffers:
• Consider a buffer system of CH3COOH (Weak electrolyte) and CH3COONa (Strong
electrolyte). There will be a large concentration of Na+ ions, CH3COONa – ions,
and undissociated CH3COOH molecules.
When an acid is added:
• If a strong acid (HCl) is added in CH3COOH / CH3COONa buffer, the changes that
will occur may be represented as:
CH3COONa Na + COO H + Cl
CH3COOH
• The hydrogen ions yielded by the HCl are quickly removed as unionized acetic acid,
and the hydrogen ion concentration is therefore only slightly affected (because acetic
acid produced is very weak as compared to HCl added).
-
- +
+

When a base is added:
• If a strong base (NaOH) is added in NH4OH / NH4Cl buffer, the changes
that will occur may be represented as:
CH3COOH CH3COO + H OH + Na NaOH
H2O
• The hydroxyl ions yielded by the NaOH are therefore removed as water. The
supply of hydrogen ions needed for this purpose being constantly provided
by the dissociation of acetic acid..
+
+
-
-

 Buffer equation-Henderson-Hasselbalch equation:
• The buffer equation is also known as Henderson-Hasselbalch equation,
with the help of this equation it is possible to calculate the pH of a buffer
solution of known concentration or to makebuffer solution of known pH.
• Two separate equations are obtained for each type of buffer, acidic and
basic
• pH of acidic buffer: The hydrogen ion concentration obtained from the
dissociation of weak acid HA is given by equation,

HA H+ + A-
𝐾𝑎 =
[H+][A−]
[HA]
Ka = equilibrium constant
[H+] = 𝐾𝑎
[HA]
[A−]
Taking logarithms of both sides of the equation & multiplying throughout by -1
gives
-log[H+] = -log𝐾𝑎-log
[HA]
[A−]
pH= pKa + log
[A−]
[HA]
𝑝𝐻 = 𝑝𝐾𝑎 + log
[congugated base]
[Acid]

pH of an alkaline buffer: The ionization of a weak base BOH is given by,
BOH B+ + OH-
𝐾b =
[B+][OH−]
[BOH]
Kb = equilibrium constant
[OH−] = 𝐾b
[BOH]
[B+]
Taking logarithms of both sides of the equation & multiplying throughout by -1
gives
-log[OH−] = -log𝐾𝑎-log
[BOH]
[B+]
pH= pKa + log
[B+]
[BOH]
𝑝𝐻 = 𝑝𝐾𝑎 + log
[Congugated acisd]
Base]

Buffer capacity:
• Buffer capacity may also be defined as “The maximum amount of either
strong acid or strong base that can be added before a significant change in
the pH will occur”.
• The maximum amount of strong acid that can be added is equal to the
amount of conjugate base present in the buffer whereas the maximum
amount of base that can be added is equal to the amount of weak acid
present in the buffer.

• Buffer capacity is depend on the factors:
1. The concentration of the acid or base component of the buffer (Direct relation)
2. The pH of the buffer
• Buffer can act best at pH = pKa and buffering range is pH = pKa +1
• Or It may be defined as the moles of strong acid or strong base required to change
the pH of 1000 ml of buffer solution by one unit.
• The magnitude of the resistance of a buffer to pH changes is referred to as the
buffer capacity, β.
• Where, ΔB is the small increment in gram equivalents (g Eq)/liter of strong base
added to the buffer solution and ΔpH: change in a pH

Standard Buffer Solutions:
• The standard buffer solutions of pH ranging from 1.2-10 are possible to
prepare by appropriate combinations of 0.2N HCl or 0.2N NaOH or 0.2M
solution of potassium hydrogen phthalate, potassium dihydrogen phosphate,
boric acid, potassium chloride.
• Standard buffers with pH range:
Buffer pH
Hydrochloric acid buffer 1.2-2.2
Acid Phthalate buffer 2.2-4.0
Neutralised phthalate buffer 4.2-5.8
Phosphate buffer 5.8-8.0
Alkaline Borate buffer 8-10

 Preparation of Buffer Solutions:
Weak acid with pKa = desired pH should be selected
↓
Ratio of salt and acid needed should be calculated using buffer equation
↓
Individual concentration of buffer salt and acid determined.
↓
Ingredients are dissolved in carbon dioxide free water
↓
Buffer capacity of 0.01-0.1 is adequate
↓
Concentration of 0.05-0.5 M is sufficient
↓
Allowed to establish equilibrium
↓
pH verified.

 Buffers in pharmaceutical systems or
Application of buffer:
1. Solubility enhancement: The pH of the pharmaceutical formulations are adjusted to an
optimum value so that the drug remains solubilised though out its shelf-life and not precipitated
out.Eg. Sodium salicylate (Asprin) precipitates as salicylic acid in acidic environment.
2. Increasing stability: To prevent hydrolysis and for maximum stability, the pH of the
medium should be adjusted suitably. Eg. Vitamins
3. Improving purity: The purity of proteins can be identified from its solubility at their
isoelectric point as they are least soluble at this point. The isoelectric pH can be maintained
using suitable buffers . Eg. Insulin precipitates from aqueous solution at pH 5.0-6.0.
4. Optimising biological activity: Enzymes have maximum activity at definite pH values.
Hence buffer of desired pH is added to the preparation.

5. Comforting the body: The pH of the formulations that are administered to different tissues of the body should
be optimum to avoid irritation (eyes), haemolysis (blood) or burning sensation (abraded surface). The pH of the
preparation must be added with suitable amount of buffers to match with the pH of the physiological fluid.
Eg: buffer in various dosage forms
Dosage Form Application Buffer used
Solids (Tablets,
Capsules, powder)
Control pH for release.
Reduce gastric irritation
Citrate buffer,
Phosphate buffer
Semisolids (Creams,
Ointments)
Stability Citric acid and
sodium citrate
Parenteral Products pH maintenance (pH
high: tissue, necrosis
pH low: pain)
Citrates, glutamate,
acetate, phthalate
Opthalamic products Drug solubility and
stability
Borates, carbonates,
phosphates

Stability of buffers
• Treat buffer solution with care.
• The typical shelf-life for commercial technical buffer is 2 years unopened and 3-6
month open.
• The typical shelf life of alkaline buffer is 1 month open.
• Alkaline buffer = change with dissolved air and form carbonic acid which decreased
pH of the solution.
• Maintain at 25°C
• Preserved in colored bottle.
• Storage condition maintain for inhibition of growth of microorganism.

Buffered isotonic solution:
• Isotonic buffered solution is defined as a solution which maintains the isotonicity
and the pHas that of the body fluids. Isotonic buffer solution should be compatible
with the body fluids for the following reasons.
• Blood and lacrimal fluids are in vivo buffer systems. Any solution that comes in
contact with these fluids should be buffered to a desired pH, so that these are
compatible with the body fluids.
• Some solutions are meant for the application on delicate membranes of the body.
Such solutions may cause haemolysis, tissue irritation, necrosis and tissue toxicity.
In such cases, solutions must be just to the same osmotic pressure and tonicity as that
of the body fluids.

• Osmosis is the diffusion of solvent through a semi-permeable membrane.
• Water always flows from lower solute concentration [dilute solution] to higher
solute concentration until a balance is produced
• Osmotic pressure is the force that cause this diffusion .
• Tonicity is a measure of the osmotic pressure of two solutions separated by a
semi-permeable membrane.

Types of Buffer Isotonic solution

1. Isotonic Solutions:
• Isotonic solutions are those solutions which produce the same osmotic pressure as that of the cell
contents in question, without net gain or loss of both solutions, provided the cell membraneis
impermeable to the solutes.
• Isotonic solutions are iso-osmotic as well as isotonic with the cells and membranes. Some of the
standard isotonic solutions are:
0.9% w/v Normal saline (sodium chloride) solution
5.0% w/v Dextrose solution
2.0% w/v Boric acid solution
These solutions do not cause swelling or shrinking of tissues when applied. Therefore, discomfort
would not be caused when instilled into the eyes, nasal tract and when injected intoblood or other body
fluids.

• In the human body, different types of cell membranes are available. All are
not having same- level of permeability to a single substance. For example,
red blood cell membrane and mucous lining of the eye are not the same.
• Therefore, isotonic solutions of 0.9% w/v sodium chloride also need not
necessarily be isotonic with respect to all the living membranes, but many
of them are roughly isotonic.

2. Hypertonic Solutions:
• Hypertonic solutions are defined as those solutions containing the solute in
higher concentration than that is required for isotonic solutions. Some
hypertonic solutions are:
• 2.0% w/v Normal saline (sodium chloride) solution (concentration > 0.9 %
w/v).
• 10.0 % w/v Dextrose solution (concentration > 5.0% w/v).
• 3.0 % w/v Boric acid solution (concentration > 2.0% w/v).

• When red blood cells are suspended in a 2.0 % w/v solution of sodium
chloride, the water within the cells passes out through the cell membranes
in an attempt to dilute the surrounding salt solution.
• This process continues until the salt concentrations on both sides of the
erythrocyte membrane are equal.
• Thus outward passage of water causes the cells to shrink and becomes
wrinkled or crenated. Such a salt solution is said to be hypertonic with
respect to blood.

3. Hypotonic Solution:
• Hypotonic solutions are defined as those solutions containing the Solute in lower
concentration than that is required for isotonic solutions
• Some hypotonic solutions are:
0.2% w/v Normal saline (sodium chloride) solution (concentration < 2.0 % w/v).
• When blood cells are suspended in a 0.2 % w/v solution of sodium chloride (or in
distilled water), water enters the blood cells causing them to swell and finally burst
with the liberation of haemoglobin.
• This process is known as haemolysis. Such a weak salt solution is said to be
hypotonic with respect to blood

Measurement of Tonicity
• Isotonicity value is defined as the concentration of an aqueous sodium chloride
solution having same colligative properties as the solution in blood.
• Apart from sodium chloride, a number of chemicals and drugs are also included
in the formulations. These ingredients also contribute to the tonicity of the
solution. Therefore, methods are needed for verifying the tonicity and adjusting
the tonicity

1. Hemolytic method:
• Red blood cells (RBCs) are suspended in various drug solutions and the
swelling of RBCs is observed bursting, shrinking and wrinkling of the
blood cells.
• In hypotonic solutions, oxyhemoglobin is released, which is in direct
proportion to the number of cells hemolyzed.
• In hypertonic solutions, the cells shrink and become wrinkled or crenated
• In isotonic solutions, the cells do not change their morphology. This
method is used for the determination of isotonicity value.

2. Cryoscopic method or depression of freezing
point:
• Colligative properties of solutions are helpful in determining the isotonicity
values. Among them, freezing point depression is extensively applied.
• Water has the freezing point of 0 °C. When substances such as sodium chloride
are added to water, the freezing point of water decreases.
• The depression of the freezing point (ΔTf) of blood and tears is 0.52 °C.
Therefore, the value of 0.9 % w/w NaCl solution should also be -0.52 °C. Such
a solution shows same osmotic pressure as that of the blood. Hence, the
functions of RBC and tissues do not alter.

Methods of adjusting the tonicity:
• Normally, solution dosage forms contain drugs of desired dose and several
excipients needed for formulation. In order to render such solutions isotonic,
sodium chloride, dextrose, etc. are added. Several methods are available for
adjusting the tonicity.
• Osmotic pressure is not a readily measurable quantity, but freezing point
depression (another colligative property) is more easily measured.

Class I methods:
In this type, sodium chloride or other substances are added to the solution in
sufficient quantity to make it isotonic. Then the preparation is brought to its
final volume withan isotonic or a buffered isotonic diluting solution.
These methods are of two types:
Cryoscopic method
Sodium chloride equivalent method.

Class II methods:
• In this type, water is added in sufficient quantity make the preparation
isotonic. Then the preparation is brought to its volume with an isotonic or a
buffered isotonic diluting solution.
• These methods are of two types:
White-Vincent method
Sprowls method.

1. Cryoscopic Method of Adjusting the Tonicity:
Principle:
• Water has the freezing point of 0 °C. Blood contains a number of substances such as carbonic acid,
carbonates, salts of phosphoric acid and hemoglobin. As a result, the depression in the freezing point
of the blood is -0.52 °C.
• When substances such as sodium chloride are added to water, the freezing point of water decreases.
The extent of depression in the freezing point depends on the concentration of the added substance.
• For example, sodium chloride at 1 % w v solution decreases the freezing point of water to - 0.58°. In
order to make the drug solution isotonic, the freezing point depression of the solution must be
maintained at-0.52°.

• Initially the drug solution is prepared whose depression in the freezing
point (ΔTf) is known. The remaining (ΔTf) value is adjusted by adding
additional substances such as sodium chloride.
• For the purpose of calculate, the freezing point depression of a number of
drugs are determined experimentally or theoretically a concentration of 10
% w/v (or sometimes 0.5 % w/v). Similarly the freezing point depression
values of 1 w/v solution of sodium chloride and other general ingredients
are also determined.

Derivation:
Freezing point depression (ΔTf) of blood is 0.52°C. Since the drug solution must be isotonic, it must have ΔTf,
same as that of the blood, i.e. ΔTf = 0.52°C.
Total drug solution ΔTf = ΔTf of drug + ΔTf adjusting substance ---------- (1)
Freezing point depression (ΔTf) of the total drug solution = 0.52°C
ΔTf value of the drug = x X ΔTf of 1 % drug solution = d
Where,
x = drug concentration in the preparation, g/100 mL
ΔTf for adjusting solution = w X a
Where,
W = weight of the adjusting substance, g/100 mL
a = ΔTf of the adjusting substance (sodium chloride), 1% (=0.58)
For an isotonic solution, equation (1) is substituted by the terms mentioned above. 0.52° = d + wa

The % w/v of adjusting substance needed is:
W= (0.52-d)/a = (0.52-d)/0.58 --------- (2)
Equation (2) is valid, if 1 % drug solution is specified. For any given percentage strength of
medicament (PSM), equation (2) may be modified as:
W= [0.52- (PSM x d)]/ 0.58 ----------- (3)
Thus, the desired concentration of adjusting substance is calculated and added in order to make the
drug preparation isotonic with blood. Each solute exerts its effect on the freezing point, although
others are present.
Hence, if two or more substances are present, a sum of their freezing point depression should be
considered.
• Advantage:
• Determination of depression in the freezing point is much simpler and more convenient.

2. Sodium Chloride Equivalent Method:
• Tonicity equivalent or sodium chloride equivalent method is used to adjust the tonicity of
pharmaceutical solutions.
• Sodium chloride equivalent (D) of a drug is the amount of sodium chloride that is
equivalent to 1 g of the drug. In this definition, equivalent refers to sodium chloride
concentration having the same osmotic effect as that of the drug. In the absence of
available data, the E value of a new drug can be calculated from equation (4).
• E= [17 X Liso]/M -------------- (4)
• Where,
• M = molecular mass, AMU
• Liso = freezing point depression constant of the drug.

Method:
The percent of sodium chloride required for adjusting isotonicity can be calculated
using equation (5).
PSA = 0.9 - (PSM E of medicament) ----------- (5)
Where,
PSM= Percent strength of medicament
PSA = Percent of sodium chloride for adjustment of isotonicity.
Equation (5) is used to calculate the amount of adjusting substance (sodium
chloride) required for making the solution isotonic. It is valid for 100 mL solution.

1. White-Vincent Method
• This method involves the addition of water to the given amount of drug to
make isotonic solution, followed by the addition of some other isotonic
solution (e.g. 0.9% NaCl) to make the final volume.
• White Vincent, from their study of need of pH adjustment in addition to
tonicity of ophthalmic solution, developed an equation The volume of
water that should be added in given amount of drug to make isotonic
solution is calculated by using:
𝑉 = 𝑊 × 𝐸 × 111.1
Where,
• V = volume of water needed to make isotonic solution
• W = given weight of drug in grams
• E = NaCl equivalent value of drug
• 111.1 = constant

• NUMERICAL: Make 50 ml isotonic solution from 0.5 gm of boric acid.
E value of boric acid is 0.50. Solution:
• Given amount of boric acid = 0.5 gm
• Required volume = 50 ml
• E value of boric acid = 0.50
• Fistly, we calculate the amount of the water that should be added in 0.5
gm of boric acid to make isotonic solution by using formula,
𝑉 = 𝑊 × 𝐸 × 111.1 𝑉 = 0.5 × 0.5 × 111.1 = 27.8 𝑚𝑙
So, 0.5 gm of boric acid is added in 27.8 ml of water to male isotonic
solution. But, final volume that is required is 50ml. so, remaining 22.2ml of
some other isotonic solution (e.g. 0.9% NaCl) are added to make up final
volume 50 ml.

2. Sprowl Method:
• Sprowls method is a simplification of White-Vincent method in which
values of V for the drug of fixed weight (0.3 g) are computed and
construed.
• This is commonly used for ophthalmic and parental solutions.
𝑉 = 𝑊 × 𝐸1% × 111.1 𝑜𝑟 𝑉 = 33.33 × 𝐸1%

Acids, Bases And Buffers Pharmaceutical Inorganic chemistry UNIT-II (Part-I)

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