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

1. SUMIT TIWARI
B.PHARM (AKTU)
M.PHARM (SRK)
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2.  Negative logarithm of the hydrogen ion (H+) in a given
solution.
 A figure expressing the acidity or alkalinity of a
solution on a logarithmic scale on which 7 is neutral
 The pH is equal to −log10 c, where c is the hydrogen ion
concentration in moles per litre.
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3.  Buffer solution usually containing an acid and a base,
or a salt, that tends to maintain a constant hydrogen
ion concentration.
 Ions are atoms or molecules that have lost or gained
one or more electrons.
 An example of a common buffer is a solution of acetic
acid (CH3COOH) and sodium acetate.
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4.  In water solution, sodium acetate is completely
dissociated into sodium (Na+) and acetate (CH3COO-)
ions.
 The hydrogen ion concentration of the buffer solution
is given by the expression:
[H+] = Ka [CH3COOH/CH3COO-]
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5.  Buffer solutions with different hydrogen ion
concentrations may be prepared by varying the buffer
ratio and by choice of an acid .
 Buffer solutions commonly used include phosphoric,
citric, or boric acids and their salts.
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6.  Because acids and bases tend to promote a wide range
of chemical reactions, the maintenance of a certain
level of acidity or alkalinity in a solution through the
use of buffer solutions is essential to many chemical
and biological experiments.
 Many biochemical processes occur only at specific pH
values, which are maintained by natural buffers
present in the body.
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7.  Solution that has the same salt concentration as
cells and blood.
 Common examples of isotonic solutions are: -
 0.9% normal saline and lactated ringers.
 These fluids are useful when the patient has lost fluid
volume from blood loss, trauma, or dehydration due to
excessive nausea/vomiting or diarrhea.
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8.  Based on the values and different concentration of
ions, a scale is devised and named after Sorenson, who
had developed it.
 Sorensen's scale assigns a pH of 1 to 14.
 With 1 being the most acidic, 14 being the most basic,
and 7 being neutral (neither acidic or basic).
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9.  The pH scale works in powers of ten, so each jump in
number is a multiple of 10 in concentration.
 For example, a pH of 1 is 10 times more acidic than a
pH of 2
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10.  There are two methods for measuring pH:
a. Colorimetric methods :-
Using indicator/solutions or papers
b. Electrochemical methods :-
Using electrodes and pH meter
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11.  pH (colorimetric) method is based on the property
of acid-base indicator dyes, which produce color
depending on the pH of the sample.
 The color change can be measured as an absorbance
change spectrophotometrically.
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12.  INTENDED USE :- Reagent for photometric
determination of pH (colorimetric) in homogenous
liquid samples.
 METHOD :- Colorimetric test with pH indicator dyes
in an aqueous solution. Method is performed at 37 °C,
using 575 nm filter and 700 nm as side Wavelenght.
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13.  PRINCIPLE OF THE PROCEDURE :-
pH (colorimetric) method is based on the property of
acid-base indicator dyes, which produce color
depending on the pH of the sample.
The color change can be measured as an absorbance
change Spectrophotometrically.
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14.  The determination may be made with any pH meter
provided with a glass electrode, using instructions
from the manufacturer. Express the result to the
nearest 0.1 units. Electronic pH measurement system
consists of :-
 Measuring electrode
 Reference electrode
 Potential measuring system
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15.  Measuring electrode: - A glass electrode is made of a thin
glass membrane of special composition that could develop
potential proportional to the difference in H+ ion
concentration of liquid on either side of the membrane.
 The glass envelope has pH sensitive glass membrane at the
bottom that contains constant pH buffer solution.
 This electrode is dipped in the measuring solution so that
potential is developed at the platinum electrode which is
proportional to the pH of the measuring solution.
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16.  This potential is measured by completing the circuit
with the reference electrode.
 Reference electrode: Calomel electrode can be used as
reference electrode which has glass envelope that
contains glass tube which contains calomel (mercury
and mercurous chloride) solution along with platinum
wire dipped in it.
 This tube is surrounded by KCl solution that slowly
diffuses or leaks into process liquid through liquids
junction provided by asbestos fibre.
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17.  Due to this, the reference electrode develops constant
potential.
• Potential measuring system: The measuring and the
reference electrode together form an electrolytic cell
whose output equals the sum of the voltage produced
by the two electrodes.
• This net voltage is applied to a null balance mill volt
potentiometer in which the slide wire can be
calibrated in terms of the pH of the measuring liquid.
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18.  Since the electrode operation depends upon the
electrical resistivity of glass, change in temperature
may cause an error in pH reading.
 To compensate for changes in temperature of the
measuring solution a temperature compensation
resistance is included in the circuit which is immersed
in the solution. The resistance of this resistor changes
with temperature.
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19.  A buffered pH is a necessity of most enzymes to
function efficiently and correctly. Furthermore,
buffering is important for ensuring proper colour
concentration when using dyes.
 A buffer solution is required for calibrating
equipment. It is especially required for pH meters that
may be in the Miscalibrated in the absence of a buffer.
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20.  Buffer solutions whose preparation takes place from
acetic acid, citric acid, ammonia can have pH values as
high as 10 or as low as 2.
 This allows buffer solutions to be worked with very
strong bases or acids.
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21.  The Henderson-Hasselbalch equation provides a
relationship between the pH of acids (in aqueous
solutions) and their pKa (acid dissociation constant).
 The pH of a buffer solution can be estimated with the
help of this equation when the concentration of the
acid and its conjugate base, or the base and the
corresponding conjugate acid, are known.
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22.  Equation of Henderson-Hasselbalch
 The Henderson-Hasselbalch equation can be written
as:
pH = pKa + log10 ([A–]/[HA])
 Where [A–] denotes the molar concentration of the
conjugate base (of the acid) and [HA] denotes the
molar concentration of the weak acid.
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23.  Buffer capacity (β) is defined as the moles of an acid
or base necessary to change the pH of a solution
by 1 , divided by the pH change and the volume of
buffer in liters.
 It is a unit-less number.
 A buffer resists changes in pH due to the addition of an
acid or base though consumption of the buffer
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24.  As long as the buffer has not been completely reacted,
the pH will not change drastically.
 The pH change will increase (or decrease) more
drastically as the buffer is depleted: it becomes less
resistant to change.
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25.  Biological buffers.
 Blood is maintained at a pH of about 7.4. The plasma
contains carbonic acid/bicarbonate and acid/alkali sodium
salts of phosphoric acid as buffers.
 Plasma proteins, which behave as acids in blood, can
combine with bases and so act as buffers.
 In the erythrocytes, the two buffer systems consist of
hemoglobin/Oxy-hemoglobin and acid/alkali potassium
salts of phosphoric acid.
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26.  Lacrimal fluid or tears, have been found to have a
great degree of buffer capacity, allowing a dilution of
1:15 with neutral distilled water.
 The pH of tears is about 7.4, with a range of 7 to 8 or
slightly higher. It is generally thought that eye drops
within a pH range of 4 to 10 will not harm the cornea.
 However, discomfort and a flow of tears will occur
below pH 6.6 and above pH 9.0.
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27.  Urine The 24-hr urine collection of a normal adult has
a pH averaging about 6.0 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.
 Conversely, 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.
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28.  Buffer solutions are used frequently in pharmaceutical
practice, particularly in the formulation of ophthalmic
solutions.
 Many buffers are available today. One of the most
common biological buffers is phosphate buffered
saline (PBS).
 Phosphate buffered saline contains sodium chloride
(NaCl) and dibasic sodium phosphate (Na2PO4).
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29.  It may also contain potassium chloride (KCl),
monobasic potassium phosphate (KH2PO4), calcium
chloride (CaCl2), and magnesium sulfate (MgSO4).
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30.  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 the 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.
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31.  Isotonic sodium chloride is a familiar pharmaceutical
example of such a preparation.
 The need to achieve isotonic conditions with solutions
to be applied to delicate membranes is dramatically
illustrated by mixing a small quantity of blood with
aqueous sodium chloride solutions of varying tonicity.
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32.  If a small quantity of blood is mixed with a solution
containing 0.9 g of NaCl per 100 mL, the cells retain
their normal size.
 The solution has essentially the same salt
concentration and hence the same osmotic pressure as
the red blood cell contents
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33.  If the red blood cells are suspended in a 2.0% NaCl
solution, the water within the cells passes through the
cell membrane in an attempt to dilute the surrounding
salt solution.
 This outward passage of water causes the cells to
shrink and become wrinkled or-crenated.
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34.  If the blood is mixed with 0.2% NaCl solution or with
distilled water, water enters the blood cells, causing
them to swell and finally burst, with the liberation of
hemoglobin.
 The salt solution in this instance is said to be with
respect to the blood cell contents.
 Finally, This phenomenon is known as hemolysis.
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35.  The red blood cell membrane is not impermeable to
all drugs; that is, it is not a perfect semi-permeable
membrane.
 Thus, it will permit the passage of not only water
molecules but also solutes such as urea, ammonium
chloride, alcohol, and boric acid.
 These solutes are regarded as solvent and they do not
exert an osmotic pressure on the membrane
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36.  Os-molality and os-molarity:- Are colligative
properties that measure the concentration of the
solutes independently of their ability to cross a cell
membrane.
 Tonicity:- is the concentration of only the solutes that
cannot cross the membrane since these solutes exert
an osmotic pressure on that membrane.
 Tonicity is not the difference between the two
osmolarities on opposing sides of the membrane.
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37.  1) Hemolytic method.
 2) Measurement of the slight temperature differences.
 3) Calculating Tonicity.
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38.  Class I Methods :-
a) Cryoscopic Method
b)Sodium Chloride Equivalent (E) Method
 Class II Methods
a)White–Vincent Method
b)The Sprowls Method
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39. THANK YOU
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