pH, BUFFERS & ISOTONIC SOLUTIONS.pptx

B.PHARM III-SEMESTER.

1. pH , BUFFERS & ISOTONIC SOLUTIONS BALASUNDARESAN M

2. C o n t e nts Sorensen’s pH scale pH determination (electrometric and calorimetric) Buffer equation Buffer capacity Buffers in pharmaceutical and biological systems Applications of buffers Buffered isotonic solutions. 2

3. Definitions 🠶 pH is a measure of how acidic/basic water is. The range goes from - 14 , with 7 being neutral. pHs of less than 7 indicate acidity, whereas a pH of greater than 7 indicates a base . (Or) pH is really a measure of the relative amount of free hydrogen and hydroxyl ions in the water 🠶 Buffer: “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 .” 🠶 Buffer Action: The resistance of a buffer solution to a change in pH is known as buffer action . 🠶 Buffer Capacity: It is defined as the ratio of the increment of strong base (or acid) to the small change in pH brought about by this addition . The buffer capacity is expressed as the amount of strong acid or base, in gram-equivalents, that must be added to 1 liter of the solution to change its pH by one unit. 🠶 Isotonic: These are the solutions which produce the same osmotic pressure as that of the cell contents in question, without net gain or loss of water by both solutions, provided the cell membrane is impermeable to solutes. 🠶 Tonicity: Tonicity is a measure of the effective osmotic pressure gradient , as defined by the water potential of two solutions separated by a semipermeable membrane 3

4. 1. Sorensen’s pH scale 4

5. Sorensen’s pH Scale 🠶 The concept of pH scale was first introduced by Danish chemist Soren Peder Lauritz Sorensen (S.P.L. Sorensen), in the year 1909. The scale was later revised to the modern pH in the year 1924. 🠶 Sorenson defined pH as the logarithm of the reciprocal of hydrogen ion concentration (Or, it can be rearranged as) Thus, pH may be defined as negative logarithm of hydrogen ion concentration p H i s a t e r m u s e d t o s p e c i f y t h e a c i d i t y o r b a s i c i t y o f a n s o l u t i o n . A p H s c a l e h e l p s i n m e a s u r i n g h o w a c i d i c o r ba s i c a s u b s t a n c e . 5

6. Sorensen’s pH Scale…. 🠶 Based on the pH values and different concentrations of H + ions, a scale has been devised and named after Sorensen. 🠶 The scale starts with zero pH , i.e, hydrogen ion concentration is 10 . It means the solution is strongly acidic . 🠶 At the other end of the scale , pH is 14 . i.e, hydrogen ion concentration is 10 -14 . It means the solution is strongly alkaline . 🠶 The central point pH in the scale is 7.0 , because [H + ] is e q u a l t o [ O H + ] , i . e . , h y d r o ge n i o n c o n c e n t r a t i o n i s 10 - 7 🠶 Solutions with a pH less than 7 are acidic and solutions with a pH greater than 7 are basic . Pure water neutral, being neither an acid nor a base. 🠶 pH Applications: Enhancing solubility, Increasing stability, improving purity, Optimizing biological activity and storage of products. 6

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8. 2. pH determination (electrometric and calorimetric ) 8

9. a. Calorimetric method of determination of pH Principle: 🠶 Colorimetric means to measure color . 🠶 Colourimetric method is used to determine the pH of the solution upto ±0.2 units in t h e r a n g e o f p H 3 t o 11 . , ba s e d o n t h e c o l o u r c h a n g e s . 🠶 The principle involves the comparison of colour of the test solution with that of the standard solutions of definite pH values. 🠶 S e v e r a l s t a n d a r d s o l u t i o n s a r e a v a i l a b l e c o m m e r c i a l l y , w h i c h a r e p r e - m i x e d solutions of buffer and indicator. 🠶 Capillators and comparators are available commercially. 🠶 Capillators: The buffer solutions and universal indicator are mixed and placed in capillary tubes . A set of such standard solutions is known as capillators. These are used for small volumes. 🠶 Comparators: For large volumes , comparators are used. These are similar to capillators , but test tubes are used inste ad of capillary tubes. These are particularly useful for examining turbid or coloured solutions. The test sample is also mixed with the universal indicator. The colour produced is compared with the standard colour of the capillators or comparators. 9

10. a. Calorimetric method of determination of pH. . Method: Standard buffer solutions of known pH values ranging from 3.0 to 11.0 are prepared with 1.0 pH intervals. A few drops of the universal indicator are added to each buffer solution. Different colours are produced . A few drops of universal indicator are added to the test solution so that it also p o s s e s s e s t h e c o l o u r d e p e n d i n g o n i t s p H . The colour of the unknown (test) solution is matched with the standard colours produced by buffer solution. The pH of the test solution must be same as that of the buffer, which has the same colour shade. Based on the approximate pH value obtained in step (4), the pH interval is reduced to a narrow range. For example, if the approximate pH is identified as 5, then standard buffer solutions are prepared from pH 4 to 6 with 0.2 pH intervals. Steps from (2) to (4) are repeated to obtain the actual pH of the test solution; The exact pH of the test solution is reported . 10

11. a. Calorimetric method of determination of pH. . Advantages 🠶 C o l o ur i m e t i c m e t h o d i s l e s s e x p e n s i v e . 🠶 It is useful for the study of acid -base reactions in non- aqueous solution. Disadvantages 🠶 C o l o ur i m e t r i c m e t h o d i s l e s s a c c u r a t e a n d l e s s c o n v e n i e n t . 🠶 It is particularly used when the solution is not coloured or not turbid . 🠶 S i n c e i n d i c a t o r s t h e m s e l v e s a r e a c i d s ( o r ba s e s ) , t h e i r a d d i t i o n t o u n b uf f e r e d solution (whose pH is to be determined) changes the pH of the solution. 11

12. b. Electrometric method of determination of pH Principle: 🠶 The basic principle of the electrometric pH measurement is determination of the activity of the hydrogen ion by potentiometric measurement using a standard hydrogen electrode and a reference electrode . Apparatus: 12

13. b . E l ec t r o m e t r i c m e th o d o f de t e r m i n a t i o n o f p H . . Method: 🠶 Before use, remove electrode from storage solution, rinse, and blot, dry with a soft tissue paper. 🠶 Calibrate the instrument with standard buffer solution. [Ex: pH 4.0, 7.0 and 10] 🠶 Once the instrument is calibrated remove the electrode from standard solution; rinse, blot and dry. 🠶 Dip the electrode in the sample whose pH has to be measured. 🠶 S t i r t h e s a m p l e t o e n su r e h o mo g e n e i t y . 🠶 Note down the reading (pH) from the pH meter. Advantages Sensitivity of the electrometric method is high . Hence, accurate measurements can be obtained. The solution is uncontaminated , because the addition of indicators is avoided. T h e p H r a n g e o f m e a s u r e m e n t i s l a r g e . Disadvantages Electrometric method is not suitable for viscous solutions and gels , because of poor ionic mobility . Initial cost of pH meter is high compared to the colourimetric method. 13

14. 5. buffers in pharmaceutical and biological systems 14

15. I. In biological systems a. Blood 🠶 B l oo d c o n s i s t s o f p r i m a ry a n d s e c o n d a r y b u f f e r s sys t e m s c o n t r i b u t i n g t h e p H 7 .4 . 🠶 When the pH of the blood goes below 7.0 or above 7.8, life is in danger. 🠶 The pH of the blood in diabetic coma is reported to drop as low as 6.8 . It is maintained at about 7.4 by two buffer systems. That are; Primary buffers : These are present in plasma . The plasma contains; carbonic acid/carbonate & acid/alkali sodium salt of phosphoric acid. Secondary buffers: these are present in erythrocytes which are; oxy-ha emoglobin / haemoglobin & acid / alkali potassium salts of phosphoric acid. 15

16. b . L a c r i ma l f lu id ( p H 7 . 4 , i n r a n g e o f 7 - 8 o r s l i g h t l y ) Lacrimal fluids (or tears) have been found to have a great degree of buffer capacity ,allowing dilution of 1:15 with neutral distilled water. The pH of tears is a b o u t 7 . 4 , w i t h a r a n g e o f 7 . t o 8 . . c. Urine 🠶 p H : 6 . ( r a n g e 4 . 5 – 7 . 8 ) 🠶 below normal…hydrogen ions are excreted by the kidney. 🠶 above pH 7.4…hydrogen ions are retained by action of the kidney. I. In biological systems.... 16

17. 🠶 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. a. 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 . II. In pharmaceutical systems 17

18. b. 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 rang 7 – 8 (7.4) , but it has good buffering capacity and can tolerate preparations having pH values between 3.5 – 10.5 with little discomfort. Out side this range (i.e. 3.5 below and 10.5 above), increase lacrimation (the flow of tears) may occur with other complications . 🠶 The buffering agents most commonly used in ophthalmic preparations include borate, carbonate and phosphates . c. 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 commonly used buffers in ointments and creams are citric acid / its salts & phosphoric acid / its salts. 18

19. Factors influencing pH of buffer Addition of small amt of water cause small + ve or – ve deviation bcz it alters activity coefficient and water itself behave as a weak acid or weak base. ⚫ +ve value of dilution :pH rises with dilution - v e v a l u e o f d i l u t i o n : p H f a ll s w i t h d i l u t i o n . Temperature : The pH values in the current use are based on the studies at 25 o C A s t h e t e m p . i n c r e a s e s: A c e t a t e b u f f e r s : p H i n c r e a s e Boric acid- sodium borate buffers: pH decrease 19

20. Io n i c st r e n g t h / SALT EFF E C T: 🠶 Addition of neutral salt to buffer solution changes the pH of the solution due to altered ionic strength. 🠶 Dilution of buffers also changes the pH due to altered ionic strength. 🠶 Therefore, whenever pH of buffer solution is mentioned, ionic strength should also be specified. Factors influencing pH of buffer… 20

21. Select a weak acid having a pKa near to a pH at which the buffer is to be used to ensure a max buffer capacity. Calculate the ratio of salt and weak acid required to obtain the desired pH. The buffer eqn is satisfactory for approximate calculation within the pH range of 4 to 10. C o n s i d e r t h e i n d i v i d u a l c o n c e n tr a t i o n o f t h e b uf f e r s a l t a n d a c i d n e e d e d t o o b t a i n a suitable buffer capacity. A conc of 0.05 to 0.5M is usually sufficient and buffer capacity of 0.01 to 0.1 is generally adequate. Steps to develop a new buffer solution. 21

22. 🠶 Availability of chemicals, sterility of the final solution, stabilty of the drug and buffer on aging, cost of materials, and freedom from toxicity should be considered. 🠶 E.g. a borate buffer, bcz of its toxic effects, certainly can not be used stabilize a solution to be administered orally or parenterally. 🠶 Determine the pH and buffer capacity of the completed buffered solution using a reliable pH meter. 🠶 W h e n t h e e l e c t r o l y t e c o n c i s h i g h , t h e p H c a l c u l a t e d b y u s e o f t h e b u f f e r e q n i s s o m e w h a t d i f f e r e n t f r o m t h e e x p e r i m e n t a l v a l u e . 🠶 It is to be expected when activity coefficient are not taken in to account . Steps to develop a new buffer solution... 22

23. Pharmaceutical Buffers The buffers of clark and Lubs were determined at 20 o C and re- determined at 25 o C. The mix and their ph ranges are: 1 . H C l a n d K C l , 1 . 2 t o 2 . 2. 2 . H C l a n d K H P , 2 . 2 . t o 4 .0 3 . N a O H a n d K H P , 4 . 2 t o 5 . 8 4 . N a O H a n d K H 2 P O 4 , 5 . 8 t o 8 5 . H 3 BO 3 , N a O H , a n d K C l , 8 t o 10. B e l o w p H 2 H C l a l o n e h a s c o n s i d e r a b l e b u f f e r e ff i c i e n cy a n d K C l i s n e u tr a l s a l t a n d i s a d d e d to adjust the ionic strength. 23

24. 6. APPLICATIONS OF BUFFERS The applications remain same for pH and buffer solutions, because buffers are used for maintaining a definite pH of the solution. 24

25. a. Enhancing solubility 🠶 If pH of a solution is not adjusted properly, drug present in the solution may precipitate . This principle is applied in the manufacture of dosage forms. 🠶 For example, sodium salicylate precipitates as salicylic acid , when acidified . Therefore, optimum pH should be adjusted for maximum solubility. 🠶 Acidic drugs are more soluble in alkaline pH due to in situ formation of salt. The pH is maintained by choosing a suitable buffer. Similarly, basic drugs are more soluble in acidic solutions , because they are predominantly in i o n i c f o r m , w h i c h i s m o r e s o l u b l e i n w a t e r ( a q u e o u s m e d i a ) . 25

26. b. Increasing stability 🠶 Many drugs get hydrolysed in aqueous solutions. Adjusting the pH of the s o l u t i o n s t a b i l i z e s s u c h d ru g s. 🠶 F o r e x a m p l e , v i t a m in s a r e s t a b l e o n l y w i t h i n a n a r r o w r a n g e o f p H . S u i t a b l e b uf f e r i s s e l e c t e d f o r o p t i m um s t a b i l i t y . 26

27. c. Improving purity 🠶 Proteins are purified based on the fact that amphoteric compounds are least soluble at their isoelectric points. The isoelectric pH is maintained using suitable buffer. 🠶 For example, insulin precipitates from aqueous solutions at pH 5.0 to 6.0 . This m e t h o d i s u s e d f o r t h e p ur i f i c a t i o n o f i n su l i n . Amphoteric compound: Able to react both as a base and as an acid. The isoelectric point is the pH at which a particular molecule carries no net electrical charge. 27

28. d. Optimizing biological activity 🠶 En z y m e s h a v e m a x i m u m a c tivi t y a t d e fini t e p H v a l u e s . H e n c e b uf f e r o f d e s i r e d p H i s a d d e d t o t h e p r e pa r a t i o n . 🠶 For example, pepsin has maximum activity at pH 1.5. 28

29. e. Comforting the body 🠶 Some of the solutions when applied to tissues cause irritation , if their pH is greatly different from normal pH of the relevant body fluids . 🠶 Therefore, while formulating the solutions meant for applying to the sensitive body parts, such as eyes (irritation), blood (hemolysis ) and abraded skin surfaces (burning sensation) the pH of the preparation must match with the pH of the physiological fluids . 29

30. 7. Buffered Isotonic solutions 30

31. Int r o d u c t i o n 🠶 Isotonic buffered solution is defined as a solution which maintains the isotonicity and the pH as that of the body fluids. Isotonic solutions cause no swelling or c o n tr a c t i o n . E . g . i s o t o n i c N a C l solutions. Isotonic : These are the solutions which produce the same osmotic pressure as that of the cell contents in question, without net gain or loss of water by both solutions, provided the cell membrane is impermeable to solutes. 31

32. a. Isotonic Solutions These are the solutions which produce the same osmotic pressure as that of the cell contents in question, without net gain or loss of water by both solutions, provided the cell membrane is impermeable to solutes. E x : . 9 % w / v s a l i n e ( N a C l ) s o l u t i o n Blood cells + 0.9 % NaCl = cells retain normal size (Isotonic with blood) 32

33. b. Hypertonic Solutions T h e s e a r e d e f i n e d a s t h e s o l u t i o n s c o n t a i n i n g t h e s o lu t e i n h i g h e r c o n ce nt r a t i o n t h a n t h a t i s r e q u i r e d f o r i s o t o n i c s o l u t i o n s Ex: 2 % w/v saline ( NaCl) solution (concentration > 0.9 w/v) Blood cells + 2 % NaCl = cells shrink and become wrinkled (Hypertonic with blood) 33

34. l o w e r c. Hypotonic Solutions T h e s e a r e d e f i n e d a s t h e s o l u t i o n s c o n t a i n i n g t h e s o lu t e in c o n ce nt r a t i o n t h a n t h a t i s r e q u i r e d f o r i s o t o n i c s o l u t i o n s Ex: 0.2 % w/v saline ( NaCl) solution (concentration < 0.9 w/v) B l o o d c e ll s + . 2 % N a C l = c e ll s s w e ll s a n d b urs t l i b e r a t i n g h e mo g l o b i n (Hypotonic with blood) 34

35. Measurement of Tonicity 🠶 A p a r t f r o m N a c l a n o . o f d r u g s a n d c h e m i c a l s a r e u s e d in f o r m u l a t i o ns w h i c h a l s o c o n t r i b u t e t o t o n i c i ty o f s o lu t i o n . T h e r e methods are need to measure tonicity and to adjust the tonicity . 🠶 Two methods Hemolytic method Cryoscopic method or depression of freezing point 35

36. Measurement of Tonicity…. A. Hemolytic Method 🠶 Red blood cells are suspended in various solutions and the appearance of RBCs is observed for swelling, bursting, shrinking and wrinkling of the blood cells . In hypotonic solutions, oxyhemoglobin is released, which is in direct proportion to the number of cells hemolysed . In hypertonic solutions, the cells shrink and become wrinkled or crenated (notched surface) In isotonic solutions, the cells do not change their morphology . 36

37. HYPERTONIC ISOTONIC HYPOTONIC NaCl 2% NaCl 0.9% NaCl 0.2% solute ‹ solute Inside outside sol u te = sol u te Inside outside solute › solute Inside outside SHRINKAGE EQUILIBRIUM SWELLING 37

38. Measurement of Tonicity… B. Cryoscopic method or depression of freezing point Colligative properties of solutions are helpful in determining the isotonicity values . Among them, depression of freezing point is extensively used. Water has a freezing point of O C. When substance such as Nacl are added to water, the freezing point of water decreases . Such as a solution shows same osmotic pressure as that of the blood . Hence, the functions of RBC a n d t i s s u e s d o n o t a l t e r . ∆ T f = -0.52 ºC (Freezing point depression of human blood & lacrimal fluid) T h e d e p r e ss i o n o f f r ee z i n g p o i n t ( ∆ T f ) o f b l o o d i s - 0.5 2 º C . T h e r e f o r e , t h e ∆ T f v a l u e o f t h e d r u g s o l u t i o n s h ou l d a l s o b e - 0.5 2 º C . 38

39. 🠶 Since osmotic pressure of a solution is not a readily measurable quantity , other easily measurable colligative properties such as the freezing point depression are used in the isotonicity calculations . 🠶 Isotonicity value is defined as the concentration of an aqueous Nacl solution having same colligative properties (freezing point, boiling point, vapor pressure and osmotic pressure)as the solution in question. 🠶 Class I methods: These methods involve addition of Sodium Chloride (or another substance) to lower the freezing point of soln. to - 0.52° C Cryoscopic Method Sodium Chloride Equivalent method 🠶 Class II methods: These methods involve addition of water to form an isotonic solution . Sprowls method White Vincent method Methods for adjustment of Tonicity 39

40. A. Cryoscopic Method: 🠶 Pure water has a freezing point ( T f ) of 0°C. When solutes are added to water, its freezing point is lowered. 🠶 B l oo d p l a s m a h a s a f r e e z i n g p o i n t o f − . 5 2 b e c a u s e o f ac i d s , s a l t s a n d Haemoglobin . 🠶 . 9 % s od i u m c h l o r i d e h a s t h e s a m e o s m o t i c p r e ss u r e a n d t h e s a m e f r e e z i n g point depression of -0.52 as that of blood plasma, red blood cells, and tears. 🠶 Dr u g s o l u t i on s w h i c h h a v e a f r e e z i n g p o i nt d e p r e ss i o n o f - . 5 2 a r e , t h e r e f o r e , isotonic with blood. Methods for adjustment of Tonicity… So lut io n ( 1 % w / v d r ug) ∆ T f, o C E Apomorphine Hcl -0.08 0.14 Boric acid -0.29 0.50 Calcium gluconate -0.09 0.16 Pilocarpine nitrate -0.14 0.23 Potassium chloride -0.45 0.76 Sodium chloride -0.58 1.00 Sodium sulphacetamide -0.14 0.23 W= weight in grams of the adjusting substance per 100 ml. a= Freezing point depression of 1 % solution of pure drug. b= Freezing point depression of 1 % solution of adjusting substance. 40

41. B. S o d i u m C h l o r i d e E q ui v a l e nt Me th o d : 🠶 Sodium Chloride Equivalent ( E ) of a drug is the Amount of NaCl that is equivalent to(i.e., has the same osmotic effect (same f.p.d) as) 1 gm of 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 🠶 The Nacl equivalents of a number of drugs and other ingredients are given in table . In the absence of the data, the E value of a new drug can be calculated from the below equation. Methods for adjustment of Tonicity…. Where, M = Molecular weight, gram/Mole L iso =Freezing point depression of the drug solution for showing isotonicity T o m a k e a s o l u t i o n o f a pa r t i c u l a r d ru g i s o t o n i c w i t h b l o o d p l a s m a , t h e s o d i um c h o r i d e equivalent value (E) of that drug is noted from the reference table or calculated Amount of NaCI required = 0.9% - {% of solution x E) 41

42. Find the amount of sodium chloride needed to make a - solution of 0.5% of KCI isotonic with blood plasma . Sodium chloride equivalent value (E) of KCI is 0.76 . 🠶 Given solution (not isotonic) = 0.5% KCI 🠶 E value of KCI = 0.76 So, by applying formula, 🠶 Amount of NaCI required = 0.9 - (% of drug x E) 🠶 = . 9 - ( . 5 x . 7 6 ) 🠶 = . 9 - . 38 = . 52 g m 🠶 H e n c e , . 5 2 g m o f N a C I m u s t b e a d d e d i n 0.5% KCI solution to make it isotonic . 42

43. Methods for adjustment of Tonicity….. C. White Vincent Method: 🠶 1 s t A dd i t i o n o f H 2 O t o d r u g t o m a k e i t i s o t o n i c 🠶 2 nd addition of isotonic vehicle to bring solution to final volume 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 = g i v e n w e i g h t o f d ru g i n g r a m s E = N a C I e q u i v a l e n t v a l u e o f d ru g 111 . 1 = c o n s t a nt 43

44. 44

45. D. Sprowls Method: 🠶 1 s t A dd i t i o n o f H 2 O t o d r u g t o m a k e i t i s o t o n i c 🠶 2 nd addition of isotonic vehicle to bring solution to final volume The Sprowls method, a modified method of the White–Vincent method, calculates the isotonic volume by using fixed mass of the material . 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 = 0.3 x E x 111.1 Where, V = volume of water needed to make isotonic solution 0… 3 = w e i g h t o f d ru g i n g r a m s (Constant) E = N a C I e q u i v a l e n t v a l u e o f d ru g 111 . 1 = c o n s t a nt 45

46. 3 . BUFFER E Q UAT I O N (Henderson – Hasselbalch equation) 46

47. For Acid Buffers: T h e p H o f ac i d b u f f e r c a n b e c a l c u l a t e d d i s s oc i a t i o n c o n s t a n t , K a o f t h e w e a k ac i d f r o m t h e a n d t h e concentrations of the acid and salt used. 🠶 T h e d i s s oc i a t i o n e x p r e s s i o n o f t h e w e a k ac i d c a n b e represented as: 🠶 HA ↔ H+ + A- 🠶 K a = [ H + ] [ A - ] / [ H A ] 🠶 O r 🠶 [H+] = Ka [HA] / [A-] ------------- (1) 47

48. 🠶 A weak acid is only slightly dissociated, and its dissociation is further depressed by the addition of the salt (XA) 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. 48

49. 🠶 [H + ] = K a . [Acid] / [Salt] --------- (2) 🠶 Taking log on both sides, we get: 🠶 log[H + ] = logK a + log [Acid] / [Salt] 🠶 multiplying both sides by –ve sign, 🠶 -log[H + ] = -logK a - log [Acid] / [Salt] 🠶 As -log[H + ] = pH & -logK a = pk a 🠶 pH = pk a - 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. 49

50. 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 - K b = [B + ] [OH - ] / [BOH] OR [OH - ] = K b [BOH] / [B + ] --- -- - ------ - ( 1 ) 50

51. 🠶 A w e a k ba s e i s o n l y s l i g h t ly d i s s o c i a t e d , a n d i ts d i s s o c i a t i o n 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+ b y s a l t c o n c e nt r a t i o n . 51

52. [OH - ] = K b . [Base] / [Salt] --------- (2) Taking log on both sides, we get: log[OH - ] = logK b + log [Base] / [Salt] multiplying both sides by –ve sign, -log[OH - ] = -logK b - log [Base] / [Salt] As -log[OH - ] = pOH & -logK b = pk b pOH = pk b – log [Base] / [Salt] Or pOH = pk b + log[Salt] / [Base] ---------- (3) 52

53. 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 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. 53

54. Applications Henderson–Hasselbalch equation. 🠶 Applications: For definite pH solution, it is essential to add salt and acid (or base) to water in a desired ratio. This ratio is determined by Henderson–Hasselbalch equation. S i n c e s a lt a n d a c i d a r e a d d e d i n p r e p a r a t i o n o f b uf f e r s o l u t i o n s , t h e i r concentrations are known. Hence using this data, the resultant pH of a solution can be calculated using Henderson–Hasselbalch equation. 3. The p K a of various drugs can be determined from pH of solutions. A sui t a b l e s a lt f o r m i n g s u b s ta n c e c a n b e s e l e c t e d ba s e d o n H e n d e r s o n – Hasselbalch equation. T h e s o l u b i li t y o f a s u b s t a n c e a t a n y p H c a n b e p r e d i c t e d p r o v i d e d intrin s i c solubility (S i ) and p K a are known 54

55. BUFFER C AP A C ITY 55

56. BUFFER C AP A C ITY 🠶 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.” 🠶 The magnitude of the resistance of a buffer to pH changes is referred to as the buffer capacity, β. 🠶 Buffer capacity is also referred as buffer index, buffer value, buffer efficiency or buffer coefficient. 🠶 The buffe r capacit y represented by ‘β ’ may also b e 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 o r ΔB / ΔpH 🠶 🠶 Where, ΔA or ΔB represents the small increment (in gram equivalents / litre of strong acid or base added) to the buffer to bring about a pH change of ΔpH . 56

57. 🠶 According to the above equation, a solution has a buffer capacity of 1 when one litre 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. 57

58. Types of buffers: Generally buffers are of two types; 🠶 Acidic buffers 🠶 Basic buffers 58

59. Acidic Buffers: An acidic buffer is a combination of weak acid a n d i t s s a l t w i t h a s t r o n g b a s e . i.e. Weak acid & salt with strong base ( c o n j u g a t e b a s e ) . EXAMPLES: 🠶 C H 3 C OO H / C H 3 C OO N a 🠶 H 2 C O 3 / N a H C O 3 🠶 H 3 P O 4 / N a H 2 P O 4 🠶 HCOOH / HCOONa 59

60. Basic Buffers: A b a s i c b u f f e r i s a c o m b i n a t i o n o f w e a k b a s e a n d i ts salt with a strong acid. i.e. Weak base & salt with strong acid (conjugate acid). E X A MP L ES: 🠶 NH 4 O H / NH 4 C l 🠶 NH 3 / NH 4 C l 🠶 N H 3 / ( NH 4 ) 2 C O 3 60

61. Buffer action 🠶 The resistance of a buffer solution to a change in p H i s k n o w n a s b u f f e r a c t i o n . 61

62. 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 comes in contact with air, it will absorb CO2 and becomes acidic. 🠶 On the other hand, if solution is stored in a glass bottle, alkaline impurities f r om t h e g l a ss m a y a l t e r i t s p H . 🠶 Due to these reasons, pharmaceutical solutions are buffered as the buffer solutions are capable of maintaining pH at some fairly constant value when even small amounts of acid or base are added . 62

63. 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 (KH 2 PO 4 ) Dibasic potassium phosphate (K 2 HPO 4 ). 63

64. THANK YOU 64