pH, Buffers and Isotonic Solutions
Learning Objectives
• At the end of this lecture, student will be able to
– Explain the concept and importance of pH and pH scale
– Describe the importance of pH scale
– Explain the concept and composition of buffers and buffer
systems
– Describe the applications of pH and buffers
– Discuss the buffer equation for weak acid and weak base
– Describe the application of buffer equation
– Explain the buffer equation for weak acid or base and its
salt
– Discuss the application of buffer equation
– Describe pharmaceutical buffer system and its importance
– Describe the method of preparation of pharmaceutical
buffers
– Explain the various biological buffer system and its
importance
– Explain the buffer action and its mechanism
– Describe buffer capacity and maximum buffer capacity
– Discuss the relationship between buffer capacity and pH on tissue irritation
– Explain the concept of isotonic solutions and its
importance in the physiological systems
– Explain the concept of paratonic solutions and its
importance in the pharmaceutical formulations
– Describe the different methods of adjustment of tonicity
pH and Sorensen’s pH Scale
• In thermodynamic terms, pH is defined as negative
logarithm of activity of hydronium ions
• Sorensen defined pH as the logarithm of the reciprocal of
the hydrogen ion concentration
• Mathematically pH is expressed as:
pH = log.1/ [H3O+]………..(1)
Equation (1) may be rearranged as
pH = log1 – log[H3O+]…………(2)
Since log 1 is zero, equation (2) can be written as
pH = -log[H3O+]
• pH may be defined as negative logarithm of hydrogen ion
concentration
• Sorensen established the term pH, to represent hydrogen
ion potential
• Term p is used to express the negative logarithm.
• Concentration of H3O+ is expressed as molarity, moles/liter etc.
• Solutions are stated as weakly acidic or strongly alkaline
and the extent of acidity of a solution may be explained by Sorensen’s scale
Sorensen’s pH Scale
• Based on the pH values and different concentrations of
H3O+ ions, a scale is devised and named after Sorensen, who developed it
• The scale starts with a zero pH, i.e., hydrogen ion
concentration is 1, which means the solution is strongly acidic
• At the other end of the scale, pH is 14, i.e., hydrogen
ion concentration is 10-14 – strongly alkaline
• The central point pH is the scale is 7.0, because [H3O+]
is equal to [OH–], i.e., hydrogen ion concentration is 10 -7
• pH = 7 means neutral
• The region with pH values below 7.0 is designated as
acidic and above PH 7.0 is designated as basic (or alkaline)
Applications of pH
• Enhancing solubility
• Increasing stability
• Improving purity
• Optimizing biological activity
• Comforting the body
• Storage of products
Definition and Applications of Buffers
• Buffers are compounds or mixtures of compounds that, by
their presence in solution, resist changes in pH upon the addition of small quantities of acid or alkali
• The resistance to change in pH is known as buffer action
• Different characteristic properties of buffers are
– They have a definite pH value
– pH value of buffers does not alter either on keeping for
long periods or on dilution
– The pH value of the buffer is very slightly altered by the
addition of small quantities of acids or alkalis
•The applications of buffers are-Enhancing solubility
-Increasing stability
-Improving purity
-Optimizing biological activity
-Comforting the body
Buffer Systems- Composition and Examples
• A combination of two or more compounds is used in the
preparation of buffer solutions as described below-Weak acid and its conjugate base, i.e., the salt of weak acid with a strong base. Example, a solution containing acetic acid and sodium acetate
– Weak base and its conjugate acid, i.e., the salt of weak
base with a strong acid. Example, a solution containing ammonium hydroxide and ammonium chloride
– Two salts can act as an acid-base pair. Example, a
solution of monobasic potassium phosphate (KH2PO4) and
dibasic potassium phosphate (K2HPO4)
– Amphoteric electrolytes act as buffer systems. Example is
the solution of glycine
– Solutions of strong acids and solutions of strong bases
exhibit buffer action by virtue of relatively high concentration of hydronium ions and hydroxyl ions. For example, hydrochloric acid buffers cover the range of 1.2 to 2.2, which include potassium chloride
• The solutions of drugs themselves manifest buffer action,
however their buffer capacities are low. A few examples are:
– Ephedrine in acidic media forms a salt of ephedrine
hydrochloride, which acts as buffer system similar to weak base and its salt with strong acid
– When salicylic acid is stored in a soft glass bottle,
sodium ions in the container react with salicylic acid and forms sodium salicylate. The solution behaves as a buffer similar to a weak acid and its salt with strong base
Applications of Buffer Equation
• For the preparation of a specified pH solution
• To calculate the pH of an unknown solution
• To predict the drug absorption based on the ionized and
unionized fraction of the drug molecules
• Determination of pKa based on the pH of the solutions
• Prediction of solubility based on pH of different
pharmaceutical solutions
• Selection of a suitable salt form
Pharmaceutical Buffer System
• Pharmaceutical buffer systems are important in the
formulation of ophthalmic and parenteral drug delivery systems
• Two stock solutions suggested by Gifford, one containing
boric acid and the other monohydrated sodium carbonate, which when mixed in various proportions, yield buffer solutions with pH values from 5 to 9
• Sorensen proposed a mixture of the salts of sodium
phosphate for buffer solutions of pH values 6 to 8
• A buffer system suggested by Palitzsch and modified by
Hind and Goyan consists of boric acid, sodium borate and sufficient sodium chloride to make the mixtures isotonic
• The Clark-Lubs mixtures and their corresponding pH ranges are as follows
– HCl and KCl, pH 1.2 to 2.2
– HCl and potassium hydrogen phthalate, pH 4.2 to 5.8
• The Clark-Lubs mixtures and their corresponding pH ranges
-NaOH and potassium hydrogen phthalate, pH 2.2 to 4.0
– NaOH and KH2PO4, pH 5.8 to 8.0
-H3BO3, NaOH, and KCl, pH 8.0 to 10.0
Preparation of Pharmaceutical Buffer Solutions
• The sequence of steps involved in the preparation of
buffers is as follows:
– A weak acid should be selected, which is having pKa
value approximately equal to the desired pH of the solution. This ensuresbmaximum buffer capacity
– From the buffer equation, the ratio of the salt and acid
needed for obtaining a suitable buffer capacity should be calculated. For a pHbrange from 4 to 10, the buffer equation is satisfactory
-The individual concentrations of the buffer salt and acid
(or base), should be determined for a desired buffer capacity. A concentration of 0.05 to 0.5 M is sufficient
– The buffer capacity of 0.01 to 0.1 is generally adequate
– The ingredients should be dissolved in carbon dioxide free
water and allowed to remain for some time to establish equilibrium condition
-The pH of the solution should be verified by a suitable
means, pH meter or pH indicator paper
• The procedure remains same for the preparation of basic
buffers
Physiological (Biological) Buffer System
• Three important biological buffer systems are-
1. Blood
– The pH of the blood is maintained at a pH of 7.4 by the
primary buffers in the plasma and secondary buffers in the erythrocytes
• The plasma contains carbonic acid/bicarbonate and
acid/alkali sodium salts of phosphoric acid as buffer
• In the erythrocytes the two buffer systems are
haemoglobin/oxyhaemoglobin and acid/alkali potassium salts of phosphoric acid
Physiological (Biological) Buffer System
• The buffer equation for the carbonic acid/bicarbonate
buffer of the blood is:
pH = 6.1 + log [HCO3–]/[H2CO3]
Where [H2CO3] represents thebconcentration of CO2 present as H2CO3 dissolved in the blood.
2. Lacrimal Fluid or Tears
• The pH of tears is about 7.4 with a range of 7 to8 or
slightly higher
• Lacrimal fluids have been found to have a great degree of
buffer capacity, allowing dilution of 1:15 with neutral distilled water
3. Urine
• The urine of a normal adult has a pH averaging about 6
units
• It may be as low as 4.5 or as high as 7.8
• When the pH of the urine is below normal values, hydrogen
ions are excreted by the kidneys
• When the urine is above pH 7.4, hydrogen ions are retained
by action of the kidneys in order to return the pH to its normal range of values
Buffer Capacity
• Buffer capacity (also known as buffer efficiency, buffer
index or buffer value) is defined as the ratio of the increment of strong base (or acid) to the small change in pH brought about by this addition
• In other words, the magnitude of the resistance of a
buffer to pH change is referred to as the buffer capacity
• Buffer capacity, β, is mathematically expressed as β=ΔB/ΔpH
Where, Δ is a finite change, and ΔB is the small increment
in gram equivalents/liter of strong base added to the buffer solution to produce a change of ΔpH
• Van Slyke’s equation for buffer capacity is represented
as:
Where C is the total buffer concentration
• Maximum buffer capacity can be given by βmax=0.576C
Influence of Buffer Capacity and pH On Tissue Irritation
• If the buffer capacity is kept low, then the pH of the
solutions for introduction into the eye may vary from 4.5 to 11.5 without marked pain or damage
• pH range of non-irritation cannot be established, but it
depends upon the buffer capacity of the buffer employed in the formulations
• Tissue irritation, due to large pH differences between the
solution being administered and the physiologic environment in which it is used, will be minimal when:
-the lower is the buffer capacity of the solution
-the smaller is the volume used for a given concentration
-the larger the volume and buffer capacity of the physiologic fluid
Buffered Isotonic Solutions
• The solutions which have the same salt concentration and
the same osmotic pressure as the red blood cell contents is said to be isotonic with blood
• Examples of isotonic solutions are- 0.9% w/v sodium
chloride solution, 5%w/v dextrose solution and 2% w/v boric acid solution
• Buffered isotonic solution is defined as a solution which
maintains the isotonicity and the pH as that of the body fluids
• Hypertonic solutions are those solutions containing the
solute in higher concentration than that is required for isotonic solutions. Examples- 2% w/v sodium chloride solution, 10% w/v dextrose solution etc.
• When red blood cells are suspended in a hypertonic
solution, the water within the cells passes out through the cell membranes in an attempt to dilute the surrounding salt solution. This outward passage of water causes the cells to shrink and becomes wrinkled or crenated (crenulation)
• Hypotonic solutions are those solutions containing the
solute in lower concentration than that is required for isotonic solutions. Examples- 0.2% w/v sodium chloride solution, 3% w/v dextrose solution etc.
• When red blood cells are suspended in a 0.2% w/v solution
of sodium chloride, water enters the blood cells causing them to swell and finally burst with the liberation of haemoglobin. This process is known as haemolysis
Methods of Adjusting Tonicity
•Class I methods- sodium chloride or some other substance is added to the solution of the drug to lower the freezing point of the solution to -0.520C, thus make the solution isotonic with body fluids
•Two methods under class I are: Cryoscopic method and Sodium chloride equivalent method
• Class II methods- water is added to the drug in sufficient
amount to form an isotonic solution. The preparation is then brought to its final volume with an isotonic or a buffered isotonic solution
• The two methods under class II are White-Vincent method
and Sprowls method
• Cryoscopic method- the method involves the depression of
freezing point by adding sodium chloride
• Sodium chloride equivalent method- the tonicic equivalent
or sodium chloride equivalent of a drug is the amount of sodium chloride that is equivalent to 1g of the drug
• White- Vincent method • Sprowls method
Summary
• pH- Negative logarithmic value of hydronium ion concentration
• pH scale- Devised by Sorensen, helps in determination of acidity and basicity of a chemical substance
• Buffers- These are mixture of compounds that resist the change in pH
• A combination of two or more compounds are used in the
preparation of buffer solutions
• The combination in buffer solutions are a weak acid and
its conjugate base or a weak base and its conjugate acid
• The typical example of drug as buffer systems are-
ephedrine / ephedrine hydrochloride and salicylic acid and sodium salicylate
• Buffer equation – Explained by Henderson-Hassel balchand is given by:
• Buffer equations are used for the preparation and
stabilization of different pharmaceutical preparation
• Pharmaceutical buffer system-Different buffer systems
proposed by various scientist, can be prepared in lab scale
• Typical examples of pharmaceuticalbuffer systems are: NaOH and potassium hydrogen phthalate, pH 2.2 to 4.0
– NaOH and KH2PO4, pH 5.8 to 8.0
– H3BO3, NaOH, and KCl, pH 8.0 to 10.0
• Method of preparation of pharmaceutical buffer- Different step by step methods are followed by choosing a suitable weak acid or a weak base for the preparation of a buffer system
• Biological buffer systems – Different biological buffer systems are:
– Blood
– Lachrymal fluid
– Urine
• Buffers- These are mixture of compounds that resist the change in pH
• Buffer action and its mechanism- The action by which buffers resist the change in pH. It acts by common ion effect
• Buffer capacity- The intensity of buffer action is buffer capacity
• Buffer capacity, pH and tissue irritation- Low buffer capacity and physiological pH will reduce tissue irritation
• Isotonic solutions- Solutions having the same tonicity as that of the body fluids
• Paratonic solutions- Solutions having more tonicity (hypertonic or hypotonic) than that of the body fluids
• Methods of adjustment of tonicity- Different methods are cryscopic method, sodium chloride equivalent method, white- Vincent and sprowls method
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