ionization The topic is pic by physical pharmacy book.pptx

Ionization
Presented by: M Abubakar Zubair

Ionization
 It is a process in which atom or molecule is converted
into ions by the addition or removal of charged particle.
 E.g ionization of water into H+
and OH-
ion

Theory of ionization
 It was presented by Arrhenius in 1887 which
consist of following postulates
 The substances called electrolytes are believed to
contain electrically charged particles called ions.
 These charges are positive for H+ ion or ions
derived from metals and negative for the ions
derived from non-metals.
 Number of electrical charges carried by an ion is
equal to the valency of corresponding atom.

 Molecules of electrolytes (acids, bases and salts) dissociate
into oppositely charged ions on dissolution in water,
 e.g. + + −
𝑁𝑎𝐶𝑙 ⇌ 𝑁𝑎 𝐶𝑙
 𝐻𝐶𝑙 ⇌ 𝐻 + + −
𝐶𝑙
 𝑁𝑎𝑂𝐻 ⇌ 𝑁𝑎+ + −
𝑂𝐻
The number of positive and negative charges on the ions must be equal so
that the solution as a whole remains neutral.
 In a solution, the ions are in a state of disorderly or random motion. Upon
colliding they may combine to give unionized molecules

 Thus ionization is a reversible process in which the solution
contains ions of electrolyte together with unionized molecules.
 𝐻2𝑆𝑂4( ) 2 + ( ) +
𝑎𝑞 ⇌ 𝐻 𝑎𝑞 𝑆𝑂4
2
− ( )
𝑎𝑞
 The extent of ionization or the degree of ionization depends upon the
nature of electrolyte.
 Strong electrolytes such as HCl etc. ionize completely in water.
 Weak electrolytes such as acetic acid (CH3COOH) ionize only slightly
 Ionization is not affected by electric current.

Types of Solvents:
 On the basis of accepting or donating proton, solvents can
be classified into four groups.
i. Protophillic
ii. Protogenic
iii. Amphiprotic
iv. Aprotic

PROTOPHILLIC:
 Any solvent that can accept proton (H+ ) from solute is
called Protophillic.
 For example, liquids like (acetones, ether & liquid
ammonia) are Protophillic solvents.
 Also called basic solvents.

PROTOGENIC:
 Those solvents which can donate proton (H+ ) are called
Protogenic solvents.
 They are usually acids in nature.
 E.g. H2SO4, HCl, CH3COOH, etc.

AMPHIPROTIC:
 Those solvents which can donate or accept the proton
can be grouped as amphiprotic solvents.
 This group contains H2O, alcohols.

APROTIC:
 Those solvents which cannot donate or accept protons are
called as Aprotic.
 They are used to study acidic and basic reactions of other
compounds.
 Hydrocarbons are grouped in this class of solvents.

Law of Mass Action:
 Law of mass action is stated by Guldberg and Waage. It
defines the influence of the concentration of the reactants on
the rate of reaction.
 Law of mass action states that the rate at which substance
reacts is proportional to its active mass and the rate of
chemical reaction is proportional to the product of the active
masses of the reactants.
 Active mass is the number of moles per liter.
 It is represented by placing the chemical formula of the
substance in square brackets. For example, HCl is represented
as [HCl].

Equilibrium
The state at which two opposing forces or actions are
balancing each other is called equilibrium.
 Chemical Equilibrium:
The state of reversible reaction when the two opposing
reactions occur at the same rate and the concentration of
reactants and products do not change with time is called
chemical equilibrium

Law of Chemical Equilibrium:
 When the above stated law of mass action is applied to a
reaction in equilibrium, the result is termed as the law of
chemical equilibrium.
 For example, the reaction,
𝑎𝐴 + + D
𝑏
𝐵
↔ 𝑐𝐶 𝑑

 The law of chemical equilibrium states that the product
of molar concentration of the products raised to the
power equal to its co-efficient, divided by the product
of the molar concentration of the reactants raised to its
co-efficient, is constant at constant temperature and is
termed as equilibrium constant.
 Kc= d
Where “Kc” is called constant of law of mass action, and
is equilibrium constant for a specific chemical reaction.

 Characteristics of Equilibrium Constant:
 Its value remains constant at a given temperature irrespective
of the direction of approach.
 The value of the equilibrium constant remains constant at given
temperature and pressure irrespective of the concentration of the
reactants and products.
 The value of equilibrium constant depends on the nature and
temperature of the reaction but it remains unaffected in the presence or
absence of catalyst.
 It gives information about the reaction proceeding in a particular
direction at a given temperature.

PH
It is the negative log of the hydrogen ion concentration.
pH = -log [H+ ] ACID BASE BALANCE
 pH is a unit of measure which describes the degree of acidity
or alkalinity (basic) of a solution.
 It is measured on a scale of 0 to 14.
Low pH values correspond to high concentrations of H+ and high pH
values correspond to low concentrations of H+.

PH VALUE
The pH value of a substance is directly related to the ratio of the
hydrogen ion and hydroxyl ion concentrations.
 If the H+ concentration is higher than OH- the material is acidic.
 If the OH- concentration is higher than H+ the material is basic.
 7 is neutral, < is acidic, >7 is basic
THE PH SCALE
The pH scale corresponds to the concentration of hydrogen ions.
For example, pure water H+ ion concentration is 1 x 10^- 7
M,
therefore the pH would then be 7.

Acid
Any compound which forms H ions in solution (proton donors)
⁺
eg: Carbonic acid releases H ions
⁺
Base
Any compound which combines with H ions in solution (proton acceptors)
⁺
eg: Bicarbonate (HCO3 ) accepts H+ ions
⁻
ACID–BASE BALANCE
Normal pH : 7.35-7.45
Acidosis
Physiological state resulting from abnormally low plasma
Alkalosis
Physiological state resulting from abnormally high plasma
Acidemia: plasma pH < 7.35
Alkalemia: plasma pH > 7.45

How to Figure it out?
If you were given HCl with a concentration of 0.003M, what would be the pH?
pH = -log [H+]
pH = -log[0.003M]
pH = 2.52 pH will not have a unit!!!
How does pH and pOH relate?
pH and pOH relate in the following way
pH + pOH = 14
pOH is the measure of the [OH-], in a solution.
What would be the pOH of a solution with a pH of 5.75?
pOH = 14 - 5.75
pOH = 8.25

Sample Problem
What is the pH of a solution with a [OH-] = 2.35 X 10-9
?
1 st determine pOH
pOH = -log [OH-]
pOH = -log[2.35X10-9
M]
pOH = 8.63
2nd
determine pH
pH = 14 – pOH
pH = 14- 8.63
pH = 5.37 (acidic)

MEASUREMENT OF PH
PH METER
The pH meter is a laboratory equipment which used to measure acidity
or alkalinity of a solution
The pH meter measures the concentration of hydrogen ions [H+ ]
using an ion-sensitive electrode.
It is the most reliable and convenient method for measuring ph.

pH Indicators:
Indicators may be considered as weak acid or weak basis that
act like buffers and also exhibit color changes as their degree of
dissociation varies with pH.
E.g. methyl red shows its fuel alkaline color yellow at pH about 6
and full acid color red at pH 4.
So indicator offers a convenient way of calorimetric method of
determining the pH of solution.
The dissociation of an acid indicator can be expressed as:
𝐾𝑖𝑛 = [𝐻3𝑂+
][ −] [ ]
𝑙𝑛 𝐻𝑙𝑛
HIn is unionized form and In is the ionize form, unionize form gives acid color
and ionize form gives basic colour. When an acid is added to the solution of
indicator, the H+ ion concentration increases and HIn predomination and give
acid colour. When a base is added [H3O + ] is reduced and more ionized form is
produce so the color changes.
𝑝𝐻 = + [ ] [ ]
𝑝𝐾𝑖𝑛 𝐵𝑎𝑠𝑒 𝐴𝑐𝑖𝑑

pKa:
pKa of an acid is the –ve log of its acid dissociation constant.
It describes the strength of an acid
Just as pH that is used to find H+ of a solution,
pKa can be used to describe the dissociation constant of a weak acid.
The higher the pKa, the weaker is the acid. E.g.
 pKa of CH3COOH is 4.76.
 pKa of HNO3 is –1.45.
 pKa of (COOH)2 is 1.27.
 pKa of CF3COOH is 0.25.

Applications of pH in Pharmacy:
1. DRUG SOLUBILITY:
As many drugs are either weak acids or bases, so their
solubility is affected by the change in pH.
• The weakly acidic drugs (e.g. aspirin) are more soluble
in alkaline solution because in alkaline solution they are
converted into salt form, which is more soluble than in
acidic form conversely the weakly basic drugs are more
soluble in acidic solution.
• And if the pH is lowered then acidic drug will be
precipitated. So for the solubility of acidic drug pH
should be higher than 7 & for basic drug lower than 7

2. DRUG STABILITY:
• Drugs are only stable at a certain pH. So for this purpose
the pH of drugs must be constant otherwise the
ionization of drug may take place, which decreases the
therapeutic effect of the drug.
• For example, cocaine HCl (Salt) is for two months at pH
5.7 But during this period pH is decreased to 4.2 So
stability is decreased.
• If pH is increased from 5.7 to 6, its decomposition into
its ions is decreased & stability increased.

DRUG ACTIVITY:
The activity of drug also depends upon their pH and its
ionized or unionized form.
For example, mandelic acid, benzoic acid etc. are acidic
drugs.
They are more effective in unionized form than ionized
form. So such acidic drugs need acidic medium (low pH) to
get unionized form & for better effectiveness

. 4. DRUG ABSORPTION:
According to pH partition theory:
“The acidic drugs are absorbed by the stomach (with low
pH) and basic drugs are absorbed by intestine (having high
pH).”
On the basis of this theory, we can conclude that for the
absorption of an acidic drug (e.g. aspirin) the medium
should have high concentration of hydrogen ions (low pH)
& for the absorption of basic drug. (e.g. paracetamol) the
absorption medium should have low concentration of
hydrogen ions (having high pH).

5. ENZYME ACTIVITY:
The activity of enzymes present in our body is influenced
by pH. Some enzymes are active in acidic medium, some
in alkaline and other in neutral medium. And beyond that
limit enzymes become inactive and may be destroyed.
E.g. pepsin works at 1.5 – 1.6

6. THERAPEUTIC EFFICACY:
Solutions to be applied to tissues or administrative
parenterally are liable to cause irritation if their pH is
greater than the relevant body fluid e.g. pH of ophthalmic
and nasal solution should be maintained.
7. PERMEABILITY OF THE DRUG THROUGH BIOLOGICAL
MEMBRANES:
This depends on the extent of ionization of drugs. The pH of
solution can affect the extent of ionization of weakly acid /
weakly basic drug. A drug in the non-ionized form is more
permeable than its ionized form.

Definition
“Buffers are compounds or mixtures of compounds that by their
presence in the solution resist changes in the pH upon the
addition of small quantities of acid or alkali.”

Necessity of a buffer system:
Sometimes it is necessary that a solution of a definite pH be prepared
and stored. The preservation of such a solution is even more difficult
than its preparation. If solution meets air, it will absorb CO2 and
becomes acidic. On the other hand, if solution is stored in a glass
bottle, alkaline impurities from the glass may alter its pH. Due to
these reasons, pharmaceutical solutions are buffered as the buffer
solutions can maintain pH at some constant value when even small
amounts of acid or base are added.

Types of buffers: Generally, buffers are of two types;
Acidic buffers
Basic buffers
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
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

Phosphate Buffers (Double salt buffers):
Besides the two general types of buffers (i.e., acidic & basic), a third appears to exist.
This is buffer system composed of two salts:
Monobasic potassium phosphate (KH2PO4)
Dibasic potassium phosphate (K2HPO4).
Buffer action
The resistance of a buffer solution to a change in pH is known as buffer action.

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, CH3COO – ions, and un-
dissociated 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:
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 CH3COOH / CH3COONa buffer, the changes that will
occur may be represented as:
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.

Mechanism of Action of basic buffers:
Consider a buffer system of NH4OH (Weak electrolyte) and NH4Cl (Strong electrolyte).
There will be a large concentration of NH4
+
ions, Cl
-
ions, and un- dissociated NH4OH
molecules. When an acid is added
If a strong acid (HCl) is added in NH4OH / NH4Cl buffer, the changes that will occur
may be represented as:
The hydrogen ions yielded by the Hcl are therefore removed as water. The supply of
OH ions needed for this is constantly provided by the ammonium hydroxide.
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:
The hydroxyl ions yielded by the NaOH are therefore quickly removed as unionized
ammonium hydroxide and the pH of solution is only slightly affected.

Mechanism of Action of phosphate buffers:
In KH2PO4 / K2HPO4 buffer system, H2PO4
-
serves as weak acid and HPO4
-2
serves as
conjugate base.
When hydronium ions are added, then
HPO4
-2
+ H3O+ H
→ 2PO4
-
+ H2O
When hydroxyl ions are added to this buffer, the following reaction takes place;
H2PO4
-
+ OH
-
HPO
→ 4
-2
+ H2O

BUFFER EQUATION (Henderson – Hasselbalch equation)
For Acid Buffers:
The pH of acid buffer can be calculated from the dissociation constant, Ka of the weak
acid and the concentrations of the acid and salt used.
The dissociation expression of the weak acid can be represented as:
HA H
↔ +
+ A-
Ka = [H+
] [A
-
] / [HA] Or
[H+] = Ka [HA] / [A-
] ------------- (1)
A weak acid is only slightly dissociated, and its dissociation is further depressed by
the addition of the salt (HA) which provides A-
ion (common ion effect) as a result the
equilibrium concentration of the unionized acid is nearly equal to the initial
concentration of the acid. The equilibrium concentration of A-
is assumed to be equal
to the initial concentration of the salt added since it is completely dissociated.
Therefore, in above equation (1), we represent concentration of A-
by salt
concentration.

[H+] = Ka. [Acid] / [Salt] --------- (2)
Taking log on both sides, we get:
log[H+] = logKa + log [Acid] / [Salt]
multiplying both sides by –ve sign,
-log[H+] = -logKa - log [Acid] / [Salt]
As -log[H+] = pH & -logKa = pka
pH = pka - log[Acid] / [Salt]
OR
pH = pka + log[Salt] / [Acid] ---------- (3)
Eq. (3) is called as Henderson – Hasselbalch equation.
It helps in calculating the pH value of buffer solution, if the concentrations
of acid as well as that of the salt are known.

For Basic Buffers:
Buffer equation for basic buffer can be calculated in same way as that for acidic
buffers. Consider a basic buffer composed of a mixture of weak base (BOH) and its salt
(BA). The dissociation constant for base can be written as,
BOH B
↔ +
+ OH-
Kb = [B+
] [OH-
] / [BOH] OR
[OH-
] = Kb [BOH] / [B+] ------------- (1)
A weak base is only slightly dissociated, and its dissociation is further depressed by the
addition of the salt (BA) which provides B+
ion (common ion effect) as a result the
equilibrium concentration of the unionized base is nearly equal to the initial
concentration of the base. The equilibrium concentration of B+
is assumed to be equal
to the initial concentration of the salt added since it is completely dissociated.
Therefore, in above equation (1), we represent concentration of B+
by salt
concentration.
[OH-
] = Kb. [Base] / [Salt] --------- (2)
Taking log on both sides,
we get log[OH-
] = logKb + log [Base] / [Salt]
multiplying both sides by –ve sign,
-log[OH-
] = -logKb - log [Base] / [Salt]
As -log[OH-] = pOH & -logKb = pkb
pOH = pkb – log [Base] / [Salt] Or

Significance of Henderson – Hasselbalch equation:
By this equation, the pH of a buffer solution can be calculated from the initial
concentrations of the weak acid and the salt provided when ka is given. However, the
Henderson equation for a basic buffer will give pOH, and so pH can be calculated as;
pkw = pH + pOH
or pH = pkw – pOH
Where pkw is a –Ve log of water ion product
pH = 14 – PoH
Also, the dissociation constant of a weak acid (pka) or a weak base (pkb) can be
calculated by measuring the pH of a buffer solution containing equimolar
concentrations of the acid (or base) and the salt.

BUFFER CAPACITY
The buffer capacity of a buffer solution is “a measure of its magnitude of its
resistance to change in the pH on an addition of an acid or a base.”
Buffer capacity is also referred as buffer index, buffer value, buffer efficiency or
buffer coefficient.
The buffer capacity represented by ‘β’ may also be defined as:
“The ratio of the increment (amount added) of strong acid or base to the small
change in pH (ΔpH) brought about by this addition”.
β = ΔA or ΔB / ΔpH
Where, ΔA or ΔB represents the small increment (in gram equivalents / liter of strong
acid or base added) to the buffer to bring about a pH change of ΔpH.
According to the above equation, a solution has a buffer capacity of 1 when one liter
of it requires one gram equivalent of a strong acid or base to change the pH by one
unit. So, smaller the pH change in a solution upon the addition of an acid or base,
greater is the buffer capacity and vice versa.

APPLICATIONS OF BUFFERS:
In biological systems:
The pH of blood is maintained at about 7.4 by two buffer systems. That are;
a) primary buffers : these are present in plasma. The plasma contains; carbonic
acid / carbonate & acid /alkali sodium salt of phosphoric acid.
b) secondary buffers: these are present in erythrocytes which are; oxy-haemoglobin /
hemoglobin & acid / alkali potassium salts of phosphoric acid.
In pharmaceutical systems:
Buffers are widely used in the field of pharmacy as ingredients in most of the
pharmaceutical formulations in order to adjust the pH of the product to that required
for maximum stability. In parenteral preparations (i.e., injections): In case of
parenteral preparations, pH should be considered carefully as large deviations of pH
may lead to serious consequences. The ideal pH of a parenteral product is 7.4, which
is pH of blood. The most commonly used buffers in parenteral products (injections) are
acetate, phosphate, citrate and glutamate.

In ophthalmic preparations
(i.e., eye preparations): Buffers are generally used in ophthalmic preparations to
maintain the pH within the physiological pH range of lacrimal fluid (i.e., eye fluid).
The lacrimal fluid has a pH in range 7 – 8 , but it has good buffering capacity and can
tolerate preparations having pH values between 3.5 – 10.5 with little discomfort.
Outside this range (i.e., 3.5 – 10.5), increase lacrimation may occur with other
complications.
The buffering agents most used in ophthalmic preparations include borate, carbonate
and phosphates.
In ointments and creams:
Topical products (which are used on skins) such as ointments and creams are also
buffered to ensure stability of the formulation. The most used buffers in ointments
and creams are citric acid / its salts & phosphoric acid / its salts.

ISOTONIC SOLUTIONS
"When two solutions have same osmotic pressure and salt concentration are said to be
isotonic solutions”. Iso (same) and tonic (concentration).
Physiologically, isotonic solutions are solutions having the same osmotic pressure as that of
the body fluids when separated by a biological membrane. Biological fluids including blood
and lachrymal fluid normally have an osmotic pressure corresponding to that of 0.9% w/v
solution of sodium chloride. Thus 0.9% solution of sodium chloride is said to be isotonic with
the physiological fluids.
In addition, to carrying out pH adjustments, pharmaceutical solutions that are meant for
application to delicate membranes of the body should also be adjusted to approximately
the same osmotic pressure as that of body fluids. Isotonic solutions cause no swelling or
contraction of the tissues with which they come in contact, and produce no discomfort
when instilled in the eye, nasal tract, blood or other body tissues. Isotonic sodium chloride
is a familiar pharmaceutical example of such a preparation.

0.9% w/v solution of NaCl has essentially the same concentration and
osmotic pressure as the RBCs contents. Thus, it is said to be isotonic
with blood, and is also known as " Physiological Saline solution", or
"Normal saline solution"
If RBCs are suspended in 0.9% solution of NaCl, they undergo no change
in their size and shape due to isotonicity.

HYPERTONIC SOLUTION
"A hyper-tonic solution is one that has greater concentration than reference solution
(i.e., RBCs Contents)."
A hyper-tonic solution has greater osmotic pressure than that of reference solution. If
RBCs are suspended in 2% w/v solution of NaCl (i.e., hypertonic solution), then water
present within the RBCs will come out (i.e., due to osmosis, from dilute RBCs fluid to
concentrated hypertonic solution) into the surroundings to dilute the NaCl solution
(hypertonic solution). This exit of water from RBCs causes their shrinkage and RBCs
become wrinkled in shape. This shrinkage of RBCs is known as "Plasmolysis".

HYPOTONIC SOLUTION
"A hypo tonic solution is one that has lower concentration than reference
solution (i.e., RBCs contents). A hypo-tonic solution has lower osmotic
pressure than that of reference solution. If RBCs are suspended in 0.1 % w/v
solution of NaCl (i.e., hypotonic solution), then water from this solution will
enter the RBCs (i.e., due to osmosis, from dilute hypotonic solution to RBCs
fluid) to dilute the fluid within the RBCs causing their swelling, which may
later result in rupturing of RBCs and release of haemoglobin. This rupturing
of RBCs is known as "Haemolysis”

ISO-OSMOTIC OR ISOSMOTIC SOLUTIONS
"Solutions having the same osmotic pressure but not necessarily the same
concentration are said to be iso-osmotic or isosmotic solutions."
E.g., 2% w/v solution of boric acid has the same osmotic pressure as 0.9% w/v solution
of NaCI but not the same concentration. So, both solutions are iso-osmotic but not iso-
tonic.

Measurement Of Tonicity
The tonicity of solutions may be determined by one of the following two methods:
1) Haemolytic Method .
2) Colligative Method
Haemolytic method:
The acting principle of this method is the observation of the effect of various
solutions of drugs on the appearance of RBCs when suspended in those solutions. If,
there is no change in size and shape of RBCs when immersed in test solution on
observing with microscope, then this solution is isotonic to the blood. This method can
be made more accurate by using a hematocrit, which is a centrifuge head in which a
graduated capillary tube is held in each of the two arms.

One capillary tube (tube A) is filled with blood diluted with 5 ml of 0.9%
w/v NaCI (isotonic solution).
The other capillary tube (tube B) is filled with blood diluted with an
equal volume i.e., 5ml of test solution.
Both tubes are centrifuged (i.e., rotated at high speed). After
centrifuge, the blood cells are concentrated at one end of the capillary
tubes and the volume occupied by the cells (i.e., PCV -Packed Cell
Volume) is measured.
Finally, the PCV of test solution tube (tube B) is compared with PCV of
isotonic solution tube (tube A), and following inferences are made.

RESULTS:
If PCV of test solution (tube B) is same as that of tube A, then test solution is regarded
as isotonic.
If RBCs volume (i.e., PCV) of tube is more than that of tube A, then test solution is
regarded as hypotonic solution (increase in PCV is due to swelling of RBCs, which
occurs in case of hypotonic solution).
If RBCs volume (i.e., PCV) of tube is less than that of tube A, then test solution is
regarded as hypertonic solution (decrease in PCV is due to shrinkage of RBCs, which
occurs in case of hypertonic solution).

Colligative Method:
It has been determined that solutions having same tonicity exhibit similar behavior
with respect to their colligative properties such as lowering of vapor pressure,
depression in freezing point, etc. Hence, tonicity of a solution may be determined by
determining its colligative properties. For making isotonic solutions, the quantities of
substances to be added may be calculated by following methods:
Based on molecular concentration
Based on freezing point data
Based on sodium chloride equivalent (E) value
White-Vincent method

Based on molecular concentration
1% molecular concentration: If one gram molecule (i.e., one-gram molecular
weight) of a substance is dissolved in 100 ml of water, the resulting solution
will be of 1% molecular concentration. For example, the molecular weight of
boric acid is 62, so if 62 gms (i.e. one-gram molecular weight) of boric acid is
dissolved in 100 ml of water, the resulting solution will have 1% molecular
concentration. For non-ionizing substance, an aqueous solution having 1%
molecular concentration, depresses the freezing point to - 18.6 C°, and
freezing point of plasma is - 0.52 C°. So, by using this information, we can
calculate the molecular concentration of blood plasma as follows;
1)A depression of - 18.6 C° in freezing point of solution is due to = 1% Molecular
concentration A depression of 1 C° in freezing point of solution is due to = 1 /
- 18.6 % Molecular concentration
2) A depression of - 0.52C° in freezing point of plasma is due to = 1 /- 18.6 x -
0.52 = 0.03% Molecular concentration
So, molecular concentration of plasma is 0.03%. Therefore, any solution having
the molecular concentration of 0.03% will be isotonic with blood (having the
same concentration - which means isotonic).

The formula for calculating the w/v percent of ionizing and non-
ionizing substances required to make isotonic solutions with blood
plasma is as follows;
• For non-ionizing substances W / V % of substance required =
0.03% x Gram molecular weight
• For ionizing substances W/V % of substance required = 0.03% x
Gram molecular weight / no. of ions yielded by the molecule

Based on freezing point data
Body fluids such as blood plasma and lachrymal secretions have a freezing point of -
0.52 C° due to different solutes present in them. 0.9% solution of NaCI (isotonic
solution) also has freezing point of - 0.52 C . Hence, all solutions which freeze at -
0.52 C° will be isotonic with these fluids.
Adjustment to the tonicity of solutions is simplified if the freezing point of 1% solution
of substance whose tonicity is to be adjusted (i.e., unadjusted solution) and freezing
point of 1% solution of adjusting substance are known.
The quantity of the adjusting substance needed for making the solution isotonic with
blood may be calculated from the general formula given below;
Amount of adjusting substance required = - 0.52 -a /b
Were,
a = freezing point of 1% solution of un-adjusted solution
b = freezing point of 1% solution of adjusting substance

Based on Sodium Chloride equivalent (E)
Sodium chloride equivalent i.e. (E) of a drug is defined as; "The grams of Sodium
chloride that will produce the same osmotic effect as 1 gm of that drug." For
example, potassium chloride has sodium chloride equivalent (E) value of 0.76 gm
NaCI / gm of KCI-. This means 0.76 gm of NaCI produce the same osmotic effect as 1
gm of KCI. To make a solution of a particular drug isotonic with blood plasma, the
sodium chloride equivalent value (E) of that drug is noted from the reference table.
This E value is multiplied with the %age of the drug solution and result so obtained is
subtracted from 0.9%. The difference in value so obtained is the amount of NaCI
needed to adjust the tonicity of the solution to isotonic value.
Amount of NaCI required = 0.9% - {%age of solution x E)

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%
NaCI) to make the final volume.
The volume of water that should be added in given amount of drug to make isotonic
solution is calculated by using following formula;
V = W x E x 111.1
Where, V = volume of water needed to make isotonic solution
W = given weight of drug in grams
E = NaCI equivalent value of drug
111.1 = constant

Pharmaceutical Buffers: The buffer action of drugs is very small,
which only maintains the pH drug by the addition of atmospheric
CO2 or ions of glass. So far the stability of drugs, same additional
buffer is used known as pharmaceutical Buffers. Some familiar
examples of pharmaceutical buffers are as follows:
1. Gifford Buffer: [H3BO3 + NaCO3– H2O] 5 9
→
This buffer solution is formed by mixing the solution of boric acid
with solution of Monohydrated sodium carbonate and it is used in
pH b/w 5 9.
→
2. Sorensen 14 Buffer: [Na3PO4] 6 8 This buffer solution is
→
obtained by mixing salts of sodium phosphate (Na3PO4). It is
used in the pH range from 6 to 8. NaCl is also added in order
to make the buffer isotonic with body fluid.

3. Palitzsch Buffer:
[H3BO3 + Na3BO3] + NaCl; 7 9 This buffer is obtained by
→
mixing the solutions of boric acid & sodium borate. In this
buffer, NaCl is added in order to make the solution
isotonic with body fluid. It is used for ophthalmic
solutions in the pH range of 7 to 9.
4. Buffers of Clark and Lubs:
a. HCl and KCl pH; 1.2 – 2.2 2.
b. HCl and K-Hydrogen Phthalate pH; 2.2 – 4.0 .
c. NaOH and K-Hydrogen Phthalate pH; 4.2 – 5.8 .
d. NaOH and KH2PO4 pH; 5.8 – 8.0 .
e. H3BO3, NaOH and KCl pH; 8.0 – 10.0

ionization The topic is pic by physical pharmacy book.pptx

More Related Content

What's hot

Similar to ionization The topic is pic by physical pharmacy book.pptx

Recently uploaded