Physical Pharmaceutics-I 3rd sem Nirali.pdf

A Text Book Of PHYSICAL PHARMACEUTICS - I As Per PCI Regulations SECOND YEAR B. PHARM. Semester III Dr. Ashok A. Hajare M. Pharm. Ph.D. Professor and Head, Department of Pharmaceutical Technology, Bharati Vidyapeeth College of Pharmacy, Kolhapur, Maharashtra, India N3952 Physical Pharmaceutics - I ISBN 978-93-88194-17-4 First Edition : July 2018 © : Authors The text of this publication, or any part thereof, should not be reproduced or transmitted in any form or stored in any computer storage system or device for distribution including photocopy, recording, taping or information retrieval system or reproduced on any disc, tape, perforated media or other information storage device etc., without the written permission of Authors with whom the rights are reserved. Breach of this condition is liable for legal action. Every effort has been made to avoid errors or omissions in this publication. In spite of this, errors may have crept in. Any mistake, error or discrepancy so noted and shall be brought to our notice shall be taken care of in the next edition. It is notified that neither the publisher nor the authors or seller shall be responsible for any damage or loss of action to any one, of any kind, in any manner, therefrom. Published By : Polyplate Printed By : NIRALI PRAKASHAN YOGIRAJ PRINTERS AND BINDERS Abhyudaya Pragati, 1312, Shivaji Nagar, Works: Sr. No. 10\1,Ghule Industrial Estate, Off J.M. Road, PUNE – 411005 Nanded Village Road, Tel - (020) 25512336/37/39, Fax - (020) 25511379 TAL-HAVELI, DIT-PUNE 411041. Email : [email protected] Mobile - 9850046517, 9404225254 ☞ DISTRIBUTION CENTRES PUNE Nirali Prakashan : 119, Budhwar Peth, Jogeshwari Mandir Lane, Pune 411002, Maharashtra Tel : (020) 2445 2044, 66022708, Fax : (020) 2445 1538; Mobile : 9657703145 Email : [email protected], [email protected] Nirali Prakashan : S. 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The reader must cross check all the facts and contents with original Government notification or publications. [email protected] | www.pragationline.com Also find us on www.facebook.com/niralibooks Preface It is indeed a matter of great pride for us that, the Pharmacy Council of India (PCI), New Delhi has framed Bachelor of Pharmacy (B. Pharm.) course regulations 2014. The Physical Pharmaceutics-I is a very important subject at Second year of the course. It gives us a great pleasure to present this book in the hand of readers. The book is strictly written as per syllabus framed by PCI under Section 6, 7 & 8 of Regulation 2014 and we had made an attempt to make it simple and understandable to the readers. The sequence of syllabus content is designed unit wise imparting better understanding of subject matter. In this book wherever needed full forms and abbreviations, pictorial diagrams, tabular data, examples of formulation including marketed products and manufacturers are given. Model questions are given at the end of every subunit to exercise on contents studied. We owe a great debt of gratitude to Hon. Dr. Patangraoji Kadam, Founder, Bharati Vidyapeeth Pune, for their encouragement. We are indeed very grateful to Prof. Dr. Shivajirao Kadam, Pro-Chancellor, Bharati Vidyapeeth University, Pune for his consistent and cheerful support and Dr. Vishwajit Kadam, Secretary Bharati Vidyapeeth, Pune for encouragement and motivation. We honestly extend our gratitude to Dr. H. N. More, Principal, Bharati Vidyapeeth College of Pharmacy, Kolhapur for freedom to work and timely help. We are thankful to Mrs. Snehal Patil Librarian for her co-operation during literature work. We thank Mrs. Suvarna and Mrs. Shital, Digvijay and Aarush and our parents from bottom of heart for sustained support, encouragement and for their forbearance. We are thankful to Mr. Jignesh Furia, Ilyas Shaikh, Anagha Medhekar, Manasi Pingle of Nirali Prakashan, Pune and Staff of Nirali Prakashan for bringing out nicely printed book. Dr. A. A. Hajare Syllabus UNIT I (10 Hours) Solubility of Drugs: Solubility expressions, mechanisms of solute solvent interactions. ideal solubility parameters, solvation and association, quantitative approach to the factors influencing solubility of drugs, diffusion principles in biological systems. Solubility of gas in liquids, solubility of liquids in liquids, (Binary solutions, ideal solutions) Raoult's law, real solutions, partially miscible liquids, Critical solution temperature and applications. Distribution law, its limitations and applications. UNIT II (10 Hours) State of matter and properties of a matter changes in the state of matter, latent heats, vapour pressure, sublimation critical point, eutectic mixtures, gases, aerosols - inhalers, relative humidity, liquid complexes, liquid crystals, glassy states, solid-crystalline, amorphous and polymorphism. Physicochemical properties of drug molecules: Refractive index, optical rotation, dielectric constant, dipole moment, dissociation constant, determinations and applications. UNIT III (08 Hours) Surface and Interfacial Phenomenon: Liquid interface, surface and interfacial tensions, surface free energy, measurement of surface and interfacial tensions, spreading coefficient, adsorption at liquid interfaces, surface active agents, HLD scale, solubilization, detergency, adsorption at solid interface. UNIT IV Complexation and Protein Hinding: Introduction, Classification of Complexation, Applications, Methods of analysis, Protein binding, Complexation and drug action, Crystalline structures of complexes and thermodynamic treatment of stability constants. UNIT V (07 Hours) 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. ✍✍✍ Contents 1. Solubility of Drugs 1.1 - 1.32 1.1 Introduction 1.1 1.2 Solubility Expressions 1.2 1.3 Mechanisms of Solute Solvent Interactions 1.5 1.4 Ideal Solubility Parameters 1.6 1.5 Solvation 1.7 1.6 Association 1.8 1.7 Quantitatve Approach to the Factors Influencing Solubility of Drugs 1.9 1.8 Diffusion Principles in Biological Systems 1.11 1.9 Solubility of Gas in Liquids 1.12 1.10 Solubility of Liquids in Liquids 1.14 1.10.1 Binary Solutions 1.14 1.10.2 Ideal Solutions 1.15 1.11 Raoult’s Law 1.17 1.12 Real Solutions 1.21 1.13 Partially Miscible Liquids 1.21 1.14 Critical Solution temperature and its Applications 1.22 1.15 Distribution Law 1.26 1.15.1 Limitations of Distribution Law 1.29 1.15.2 Applications of Distribution Law 1.29 Exercise 1.30 2. States and Properties of Matter and Physicochemical Properties of Drug Molecules 2.1 - 2.78 2.1 States of Matter and Properties of Matter 2.1 2.1.1 States of Matter 2.1 2.1.2 Changes in the State of Matter 2.5 2.1.3 Latent Heat 2.6 2.1.4 Vapour Pressure 2.7 2.1.5 Sublimation 2.11 2.1.6 Critical Point 2.12 2.1.7 Eutectic Mixtures 2.13 2.1.8 Gases 2.14 2.1.9 Aerosols 2.22 2.1.10 Inhalers 2.22 2.1.11 Relative Humidity 2.24 2.1.12 Liquid Complexes 2.29 2.1.13 Liquid Crystals 2.30 2.1.14 Glassy State 2.35 2.1.15 Solids 2.38 2.2 Physicochemical Properties of Drug Molecules: Determinations and Applications 2.60 2.2.1 Refractive Index 2.62 2.2.2 Optical Rotation 2.66 2.2.3 Dielectric Constant 2.71 2.2.4 Dipole Moment 2.73 2.2.5 Dissociation Constant 2.75 Exercise 2.76 3. Surface and Interfacial Phenomenon 3.1 - 3.44 3.1 Liquid Interface 3.2 3.2 Surface Tension 3.3 3.3 Interfacial Tension 3.4 3.4 Surface Free Energy 3.5 3.5 Classification of Methods 3.7 3.6 Measurement of Surface Tension 3.9 3.7 Measurement of Interfacial Tension 3.18 3.8 Spreading Coefficient 3.20 3.9 Adsorption at Liquid Interfaces 3.23 3.10 Surface Active Agents 3.24 3.11 HLB Scale 3.25 3.12 Solubilization 3.30 3.13 Detergency 3.36 3.14 Adsorption at Solid Interface 3.37 Exercise 3.43 4. Complexation and Protein Bonding 4.1 - 4.30 4.1 Introduction 4.2 4.2 Classification of Complexation 4.2 4.3 Applications of Complexation 4.14 4.4 Methods of Analysis 4.16 4.5 Protein Binding 4.24 4.6 Complexation and Drug Action 4.26 4.7 Crystalline Structures of Complexes 4.27 4.8 Thermodynamic Treatment of Stability Constants 4.28 Exercise 4.29 5. pH Buffers and Isotonic Solutions 5.1 - 5.42 5.1 Introduction 5.1 5.2 Sorensen’s pH Scale 5.2 5.3 Electrometric pH Determination 5.6 5.3.1 Colorimetric pH Determination 5.8 5.4 Applications of Buffers 5.9 5.5 Buffer Equation 5.10 5.6 Buffer Capacity 5.13 5.7 Buffers in Pharmaceuticals 5.18 5.8 Buffers in Biological Systems 5.26 5.9 Buffered Isotonic Solutions 5.28 Exercise 5.41 ✍✍✍ Unit … 1 SOLUBILITY OF DRUGS ‚ OBJECTIVES ‚ Solubility is the physical property of substances that varies with temperature and pressure as well as the nature of the solute and the solvent. In the view of making formulations bioavailable and stable the knowledge of basic concepts of solubility is must. After knowing the importance of studying and understanding the phenomenon of solubility the student should be able to: Identify the descriptive terms for solubility, their meaning, and various types of solutions. Understand the terms and concepts of miscibility. Understand the factors controlling the solubility of drugs. Understand partition coefficient and its importance in pharmaceutical systems. Overcome problems arising during preparation of pharmaceutical solutions. Calculate the partition coefficients for different types of solutes in aqueous/organic solvent systems. 1.1 INTRODUCTION Solubility is the ability of one substance to fully dissolve in another substance under specified conditions. The word soluble comes from the fourteenth century, from the Latin word ‘solvere’ meaning to dissolve. The concentration of a solution is usually quoted in terms of mass of solute dissolved in a particular volume of solvent. The solubility is generally expressed in gram per litre. Therefore, solubility of a solute in a solvent at a particular temperature is the number of grams of the solute necessary to saturate 100 grams or mL of the solvent at that temperature. Most commonly encountered solutions are solids dissolved in liquids. The solid that dissolve in a liquid is the solute and the liquid in which it dissolves is solvent. A solute is the dissolved agent usually the less abundant part of solution whereas solvent is more abundant part of solution. If a solid can dissolve in a liquid, it is said to be soluble in that liquid, if not it is said to be insoluble. As we add more solids to a liquid the solution becomes more concentrated. The greater the solubility of a substance the more concentrated it is possible to make the solution. Solubility is measured after solute of interest has had sufficient contact time (however long it takes) with the solvent. There are two types of solubility: one is called intrinsic solubility and the other one is apparent solubility. Intrinsic 1.1 Physical Pharmaceutics - I Solubility of Drugs solubility is defined as the maximum concentration to which a solution can be prepared with a specific solute and solvent. It is often derived from calculation, and is a single numeric number (for example, 0.5 µg/mL) that is independent of the environmental factors. The apparent solubility is dependent on the environmental factor such as pH and ionic strength and is obtained from the experimental measurements. The rate of solubility is affected by many factors such as type of solvent, size and amount of solute particles, stirring speed and temperature. The concept of solubility is very important because it governs the preparation of solutions as dosage forms and a drug must be in solution before it can be absorbed by the body or have any biological activity. Since activity of drug depends on solubility, it is equally important to control environmental conditions which impact various types of solution. 1.2 SOLUBILITY EXPRESSIONS The solubility of a drug or other substance in a solvent can be expressed quantitatively in numerous terms viz. percent by mass, percent by volume, molality (m), molarity (M), mole fraction (x), and parts per million (ppm), etc. The particular terminology we use depends largely on the use to which we will put it. Solubility of substance is deﬁned as the amount of solute dissolved in a speciﬁc amount of solvent at specific temperature. The British Pharmacopoeia and other official chemical and pharmaceutical compendia frequently use the term parts per parts of solvent (for example, parts per million, ppm). The expressions ‘insoluble’, ‘very highly soluble’ and ‘soluble’ also can be used to express solubility of solutes but being inaccurate often not found to be helpful. For quantitative work specific concentration terms must be used. Most substances have at least some degree of solubility in water and while they may appear to be ‘insoluble’ by a qualitative test, their solubility can be measured and quoted precisely. In aqueous media at pH 10, chlorpromazine base has a solubility of 8 × 10−6 mol/dm3. It is very slightly soluble and it might be considered as ‘insoluble’ upon visual inspection due to lack of disappearance of solid. In many solutions the concentration has a maximum limit that depends on various factors, such as temperature, pressure, and the nature of the solvent. Relative concentrations of a solute/solvent system can often be expressed by the terms dilute and concentrated, or by the terms unsaturated, saturated, and supersaturated. Solutes in water are often categorized as either strong electrolytes, if completely ionized in water or weak electrolytes, if only partially ionized or non-electrolytes when non-ionized. In regard to solubility, general terms can be used when describing whether a compound is soluble or not. These terms are given in Table 1.1, and are based on the part of solvent needed to dissolve 1 part of the solute for example, testosterone is considered insoluble in water but soluble in alcohol, ether or other organic solvents. Fortunately, when injected to body, insoluble testosterone is diluted and the larger volume of body water permits testosterone to go into solution. 1.2 Physical Pharmaceutics - I Solubility of Drugs Table 1.1: General Terms of Solubility Term Parts of solvent required per part of solute Very soluble Less than 1 part Freely soluble 1 - 10 Soluble 10 - 30 Sparingly soluble 30 - 100 Slightly soluble 100 - 1000 Very insoluble 1000 - 10,000 Insoluble More than 10,000 Saturated Solution A solution in which dissolved solute is in equilibrium with the undissolved solute or solid phase is known as saturated solution. It is when no more of the solid will dissolve into the solution. When we add solute to a solvent a point is reached where no more solute dissolve under speciﬁed condition. The solution is saturated. The concentration of the solute in a saturated solution is the solubility of the solute in that solvent at that temperature. Saturation of solution also can be deﬁned as the point where the solution is in equilibrium with undissolved solute. In a saturated solution containing undissolved solid solute, the rate at which the molecules or ions leave the solid surface is equal to the rate which the solvated molecules return to the solid. KSOL H2O H O H2O 2 H2O H2O H2O H2O KPPT H2O Undissolved solute Dissolved solute Figure 1.1: Saturated solutions In Fig. 1.1, KSOL is the rate constant at which solid is solvated and KPPT is the rate constant at which the solvated molecule is returned to the solid. The solubility of substance is ratio of these rate constants at equilibrium in a given solution. At equilibrium the rate of a solute precipitating out of solution is equal to the rate in which the solute goes into solution. Unsaturated Solution: An unsaturated solution is a solution containing the dissolved solute in a concentration less than a saturated solution. If less solute is added to the solvent, then the solution is said to be unsaturated. Most pharmaceutical solutions are considered to be unsaturated. 1.3 Physical Pharmaceutics - I Solubility of Drugs Supersaturated Solution: A solution which contains more concentration of solute than saturated solution is known as supersaturated solution. It requires an increase in temperature to make it possible to dissolve more solute into solvent than is required to produce a saturated solution. This yields a supersaturated solution. These solutions can be prepared by heating the saturated solutions at higher temperatures. The solute is dissolved into the solvent at a high temperature and then the solution is slowly cooled, such solution is unstable and the addition of small amount of solute cause all of the excess dissolved solute to crystallize out of the solution. A saturated potassium chloride solution at 10oC will have 31 grams of this substance dissolved in 100 grams of water. If there are 40 grams of potassium chloride in the container, then there will be 9 grams of undissolved potassium chloride remaining in the solution. Raising the temperature of the mixture to 30oC will increase the amount of dissolved potassium chloride to 37 grams and there will be only 3 grams of solid undissolved. The entire 40 grams can be dissolved if the temperature is raised above 40oC. Cooling the hot solution (40oC) will reverse the process. When the temperature decreased to 20oC the solubility will eventually be decreased to 34 grams. There is a time delay before the extra 6 grams of dissolved potassium chloride crystallizes. This solution is “supersaturated” and is a temporary condition. The “extra” solute will come out of solution when the randomly moving solute particles can form the crystal pattern of the solid. A “seed” crystal is sometimes needed to provide the surface for solute particles to crystallize on and establish equilibrium. Concentration Units: A wide range of units is commonly used to express solution concentration, and confusion often arises in the inter-conversion of one set of units to another. Wherever possible throughout this book we have used the SI system of units. Although this is the currently recommended system of units in Great Britain, other more traditional systems are still widely used and these are also used in latter sections. Weight Concentration: Concentration is often expressed as a weight of solute in a unit volume of solution; for example, g/dm3, or % w/v, which is the number of grams of solute in 100 cm3 of solution. This is not an exact method when working at a range of temperatures, since the volume of the solution is temperature dependent and hence the weight concentration also changes with temperature. Whenever a hydrated compound is used, it is important to use the correct state of hydration in the calculation of weight concentration. Thus, 10% w/v CaCl2 (anhydrous) is approximately equivalent to 20% w/v CaCl2·6H2O and consequently the use of the vague statement ‘10% calcium chloride’ could result in gross error. The SI unit of weight concentration is kg/m3 which is numerically equal to g/dm3. Molarity and Molality: Molarity and molality are similar-sounding terms and must not be confused. The molarity of a solution is the number of moles (gram molecular weight) of solute in 1 litre (1 dm3 or 1.4 Physical Pharmaceutics - I Solubility of Drugs 1000 mL) of solution. The molality is the number of moles of solute in 1 kg of solvent. Molality has the unit, mol/kg, which is an accepted SI unit. Molarity may be converted to SI units using the relationship 1 mol/L = 1 mol/dm3 = 1M= 1000 mol/m3. Interconversion between molarity and molality requires knowledge of the density of the solution. Of the two units, molality is preferable for a precise expression of concentration because it does not depend on the solution temperature as does molarity; also, the molality of a component in a solution remains unaltered by the addition of a second solute, whereas the molarity of this component decreases because the total volume of solution increases upon the addition of the second solute. Milliequivalents: The unit milliequivalent (mEq) is commonly used clinically in expressing the concentration of an ion in solution. The term ‘equivalent’, or gram equivalent weight, is analogous to the mole or gram molecular weight. When monovalent ions are considered, these two terms are identical. A 1 molar solution of sodium bicarbonate, NaHCO3, contains 1 molar 1 Eq of Na+ and 1 mol or 1 Eq of HCO3 per litre (dm3) of solution. With multivalent ions, attention must be paid to the valency of each ion; for example, 10% w/v CaCl2·2H2O contains 6.8 mmol or 13.6 mEq of Ca2 in 10 cm3. The Pharmaceutical Codex gives a table of milliequivalents for various ions and also a simple formula for the calculation of milli equivalents per litre. In analytical chemistry a solution which contains 1 Eq/dm3 is referred to as a normal solution. Unfortunately the term ‘normal’ is also used to mean physiologically normal with reference to saline solution. In this usage, a physiologically normal saline solution contains 0.9 g NaCl in 100 cm3 aqueous solutions and not 1 equivalent (58.44 g) per litre. 1.3 MECHANISMS OF SOLUTE SOLVENT INTERACTIONS A solute dissolves in a solvent when it forms favourable interactions with the solvent. This dissolving process all depends upon the free energy changes of both solute and solvent. The free energy of solvation is a combination of several factors. The process can be considered in three stages: (i) A solute (drug) molecule is ‘removed’ from its crystal. + Drug crystal Drug crystal Drug molecule Figure 1.2 (a) : Removal of solute molecule The solute must separate out from the bulk solute. This is enthalpically unfavourable as solute-solute interactions are breaking but is entropically favourable. 1.5 Physical Pharmaceutics - I Solubility of Drugs (ii) A cavity for the drug molecule is created in the solvent. Cavity Solvent Cavity in solvent Figure 1.2 (b) : Creation of cavity A cavity must be created in the solvent. The creation of the cavity will be entropically and enthalpically unfavourable as the ordered structure of the solvent decreases and there are fewer solvent-solvent interactions. (iii) The solute (drug) molecule is inserted into this cavity. + Solvent with cavity Drug molecule Drug in solvent Figure 1.2 (c) : Insertion of solute The solute must occupy the cavity created in the solvent. Placing the solute molecule in the solvent cavity requires a number of solute–solvent contacts; the larger the solute molecule, the more contacts are created. If the surface area of the solute molecule is A, and the solute–solvent interface increases by γ12 A, where γ12 is the interfacial tension between the solvent1 and the solute2 then it leads to favourable solute-solvent interactions. This is entropically favourable as the mixture is more disordered than when the solute and solvent are not mixed. Dissolution often occurs when the solute-solvent interactions are similar to the solvent- solvent interactions, signified by the term ‘Like dissolves Like’. Hence, polar solutes dissolve in polar solvents, whereas non-polar solutes dissolve in non-polar solvents. Dissimilar nature of solute and solvent makes solute insoluble in the solvent. Substances dissolve when solvent- solute attraction is greater than solvent-solvent attraction and solute-solute attraction. 1.4 IDEAL SOLUBILITY PARAMETERS Regular solution theory characterises non-polar solvents in terms of solubility parameter, δ1, which is defined as ∆U1/2 ∆H − RT1/2 δ1 =   = V  … (1.1) V   1.6 Physical Pharmaceutics - I Solubility of Drugs Where, ∆U is the molar energy and ∆H is the molar heat of vapourization of the solvent. The ∆H is determined by calorimetry at temperatures below the boiling point at constant volume and V is the molar volume of the solvent. The solubility parameter is thus a measure of the intermolecular forces within the solvent and gives us information on the ability of the liquid to act as a solvent. The ratio ∆U/V is the liquid’s cohesive energy density, a measure of the attraction of a molecule from its own liquid, which is the energy required to remove it from the liquid and is equal to the energy of vapourization per unit volume. As cavities have to be formed in a solvent by separating other solvent molecules to accommodate solute molecules the solubility parameter δ1 enables predictions of solubility to be made in a semi- quantitative manner, especially in relation to the solubility parameter of the solute, δ2. By itself the solubility parameter can explain the behaviour of only a relatively small group of solvents – those with little or no polarity and those unable to participate in hydrogen bonding interactions. The difference between the solubility parameters expressed as (δ1-δ2) will give an indication of solubility relationships. For solid solutes a hypothetical value of δ2 can be calculated from (U/V) 1/2, where U is the lattice energy of the crystal. In a study of the solubility of ion pairs in organic solvents it has been found that the logarithm of the solubility (log S) correlates well with (δ1/δ2)2. 1.5 SOLVATION The process of solvation is sometimes called dissolution. Solvation is a kinetic process and is quantified by its rate. It is the attraction and association of molecules of a solvent with molecules or ions of a solute. When a solute is soluble in a certain solvent, the solute's molecules or ions spreads out and became surrounded by solvent molecules. A complex formed of molecule or ion of solute in a solvent is known as a solvation complex. Solvation is the process of rearranging solvent and solute molecules into solvation complexes to distribute solute molecules evenly within the solvent. Solvation process is affected by hydrogen bonding and van der Waals forces (which consist of dipole-dipole, dipole-induced dipole, and induced dipole-induced dipole interactions). Which of these forces are at play depends on the molecular structure and properties of the solvent and solute. Insoluble solute molecules interact with other solute molecules rather than break apart and become solvated by the solvent, for example, solvation of functional groups on a surface of ion-exchange resin. In fact solvation is an interaction of a solute with the solvent, which leads to stabilization of the solute species in the solution. Solvation of a solute by water is called hydration. Solvation is, in concept, distinct from solubility. Solubility quantifies the dynamic equilibrium state achieved when the rate of dissolution equals the rate of precipitation. The consideration of the units makes the distinction clearer. The typical unit for dissolution rate is mol/sec. The units for solubility express a concentration as mass per volume (mg/mL), molarity (mol/L) etc. The similarity between solvent and solute determines how well a solute can be solvated by a solvent. 1.7 Physical Pharmaceutics - I Solubility of Drugs 1.6 ASSOCIATION Association or ion association is a chemical reaction wherein ions of opposite electrical charge come together in solution to form a distinct chemical entity. Ion associates are classified, according to the number of ions that associate with each other, as ion pairs, ion triplets etc. Ion pairs are also classified according to the nature of the interaction as contact, solvent-shared or solvent-separated. The most important factor that determines the extent of ion association is the dielectric constant of the solvent. Ion associates have been characterized by means of vibrational spectroscopy. Ion pairs are formed when a cation and anion come together: An+ + Bm− AB(n−m)+ There are three distinct types of ion pairs depending on the extent of solvation of the two ions: Fully solveted Solvent shared and Solvent separted Contact Figure 1.3: Schematic of types of ion pair In the above schematic representation, the circles represent spheres. The sizes are arbitrary and not necessarily similar as shown in Fig. 1.3, the cation is coloured dark and the anion is coloured grey. The area surrounding ions represents solvent molecules in a primary solvation shell; secondary solvation is ignored. When both ions have a complete primary solvation sphere, the ion pair may be termed fully solvated. When there is about one solvent molecule between cation and anion, the ion pair may be termed solvent-shared. Lastly, when the ions are in contact with each other, the ion pair is termed a contact ion pair. In contact ion pair the ions retain most of their solvation shell and the nature of this solvation shell is generally not known. In aqueous solution and in other donor solvents, metal cations are surrounded by between 4 and 9 solvent molecules in the primary solvation shell, but the nature of solvation of anions is mostly unknown. Another name for a solvent-shared ion pair is an outer-sphere complex. Usage of outer- sphere complex is common in co-ordination chemistry and denotes a complex between a solvated metal cation and an anion. Similarly, a contact ion pair may be termed an inner- sphere complex. The major difference between these three types is the closeness with which the ions approach each other: The order of closeness is prevented as Fully solvated > Solvent-shared > Contact. With fully solvated and solvent-shared ion pairs the interaction is primarily electrostatic, but in a contact ion pair some covalent character in the bond between cation and anion is also present. An ion triplet may be formed from one cation and two anions or from one anion and two cations. Higher aggregates, such as a tetramer (AB)4, may be formed. Ternary ion associates involve the association of three species. Another type, named intrusion ion pair, has also been characterized. 1.8 Physical Pharmaceutics - I Solubility of Drugs 1.7 QUANTITATVE APPROACH TO THE FACTORS INFLUENCING SOLUBILITY OF DRUGS The solubility of most solid solutes is significantly affected by temperature. When some solid dissolves in a liquid a change in the physical state of the solid analogues (melting) takes place. Heat is required to break the bonds holding the molecules in the solid together. At the same time, heat is given off during the formation of new solute-solvent bonds. The typical solubility data for some common inorganic compounds at respective temperatures is given in Table 1.2. Table 1.2: Solubility of Common Inorganic Compounds in g/100 mL of Water Substance 0°°C 10°°C 20°°C 30°°C 40°°C 50°°C Potassiumiodide 127.5 136 144 152 160 168 Potassium chloride 27.6 31.0 34.0 37.0 40.0 42.6 Sodium chloride 35.7 35.8 36.0 36.3 36.6 37.0 Sodium bicarbonate 6.9 8.15 9.6 11.1 12.7 14.45 Sodium hydroxide − − 109 119 145 174 Epsom salts, magnesium − 23.6 26.2 29 31.3 − sulfate heptahydrate These values are the amount of solute that will dissolve and form a saturated solution at the temperatures listed. The solubility can be increased if the temperature is increased. The solubility of solute usually increases with increasing temperature but there are exceptions such as Ce2(SO4)3 as shown in Fig. 1.4. 100 90 NO 3 Na Solubility (g/100 g of water) 80 70 O7 l2 Cr C 2 60 Ca )2 K 2 3 ( NO 50 Pb KC l 40 NaCl 30 lO 3 20 KC 10 Ca(SO4)3 0 10 20 30 40 50 60 70 80 90 100 o Temperature ( C) Figure 1.4: Solubility of common inorganic compounds 1.9 Physical Pharmaceutics - I Solubility of Drugs Generally, increase in temperature increases solubility of solids in solvent. Although in many cases solubility increases with the rise in temperature and decreases with the fall of temperature, it is not necessary in all cases. It means there are exceptions that solubility decreases with increase in temperature. CASE I: Increase in Solubility with Temperature In endothermic processes solubility increases with the increase in temperature and vice versa. For example, solubility of potassium nitrate increases with the increase in temperature. If the heat given off in the dissolving reaction is less than the heat required to break apart the solid, the net dissolving reaction is endothermic (energy required). Therefore, the heat is drawn from the surroundings. The addition of more heat facilitates the dissolving reaction by providing energy to break bonds in the solid. This is the most common situation where an increase in temperature produces an increase in solubility for solids. CASE II: Decrease in Solubility with Temperature In exothermic processes solubility decrease with the increase in temperature. For example, solubility of calcium oxide decreases with the increase in temperature. Gases are more soluble in cold solvent than in hot solvent. If the heat given off in the dissolving process is greater than the heat required to break apart the solid, the net dissolving reaction is exothermic (energy given off). The addition of more heat (increases temperature) inhibits the dissolving reaction since excess heat is already being produced by the reaction. This situation where an increase in temperature produces a decrease in solubility is not very common, for example, calcium hydroxide is more soluble at cold temperatures than at warm. When we dissolve a substance we must separate the intermolecular forces which surround the molecules. Separation of molecules requires a certain amount of energy which, in this case, can be provided in the terms of heat. There is also the possibility that compound will form a bond with the solvent resulting in energy release. However, care must be taken while supplying heat that may destroy a drug or cause other changes in the solution. For example, sucrose solution when we heat in presence of acid results in formation of invert sugar. The energy is supplied in the form of heat, providing a cooling effect. On the other hand, there is possibility of interaction between solute and solvent with formation of dipole-dipole type bond and this interaction will tend to give off heat. Based on which of these interactions are greater, we can get increase or decrease in temperature. A good example is mixture of chloroform and acetone. There exists a strong interaction between acetone and chloroform molecules. The heat produced by solute-solvent interaction is so much higher than the heat necessary to separate the molecules of acetone and chloroform, that the excess heat can be detected as rise in temperature of the liquid. Solubility Curves: Solids are usually more soluble at higher temperatures; more salt will dissolve in warm water than in an equal amount of cold water. A graph showing the solubility of different solids as a function of temperature are very useful in chemical analysis. A curve drawn between solubility and temperature is called solubility curve. It indicates the effect of 1.10 Physical Pharmaceutics - I Solubility of Drugs temperature on solubility of substances. Substances such as calcium acetate and calcium chromate show decreased solubility with increase in temperature while sodium nitrate and lead nitrate show increase in solubility with increase in temperature. The solubility curve of sodium chloride shows very minute rise with increase of temperature. There are two types of solubility curves as shown in Fig. 1.5. Continuous Solubility Curve: Solubility curve of substance such as calcium salts of fatty acids, potassium chlorates, lead nitrate and sodium chloride are continuous solubility curves. They show no sharp break in the curves anywhere. The solubility curve of hydrated calcium sulphate shows a rise and then fall but it remains continuous at maximum point. Discontinuous Solubility Curve: The solubility curve which shows sudden change in direction is called as discontinuous solubility curve. For example, sodium sulphate, calcium chloride, ammonium nitrate etc. At the break a new solid phase appears and another solubility curve of that new phase starts. The break in a solubility curve shows with sharp point where two different curves meet each other. 120 160 3 4 O NH N 100 140 Solubility (g/100g of water) Solubility (g/100g of water) NaNo3 KNO3 80 120 CaCl2.2H2O 60 100 KCl CaCl2.4H2O 40 80 CaCl2.6H2O NaCl Na2SO4 20 40 KClO3 Na2SO4.10H2O 0 0 0 20 40 60 80 100 120 0 20 40 60 80 100 120 o o Temperature ( C) Temperature ( C) (a) (b) Figure 1.5: Solubility Curves (A) Continuous and (B) Discontinuous 1.8 DIFFUSION PRINCIPLES IN BIOLOGICAL SYSTEMS Matter moves by diffusion along energy gradients from areas of high concentration to areas of lower concentration. The rate of diffusion depends on temperature, size of the particles, and the size of the concentration gradient. In biology, the selectively permeable cell membrane creates two special forms of diffusion namely: osmosis for the diffusion of water, and dialysis for the diffusion of solutes. 1.11 Physical Pharmaceutics - I Solubility of Drugs Diffusion is one principle method of movement of substances within cells, as well as for essential small molecules to cross the cell membrane. Cell membranes act as barriers to most, but not all, molecules. A cell membrane that could allow some materials to pass while prevent the movement of other molecules is a major step in the development of the cell. The cell membrane functions as a semi-permeable barrier, allowing a very few molecules across it while holding majority of chemicals inside the cell. Cell membranes separate the inner cellular environment from the outer cellular (or external) environment. Most of the molecules move from higher to lower concentration, although there will be some molecules that move from low to high. The overall movement is thus from high to low concentration. If there is no energy input into the system, the molecules reaches a state of equilibrium and gets uniformly distributed throughout the system. A cell membrane is composed of phospholipids and proteins. Absorption of drugs across the stomach lining/mucosa and the blood/brain barrier are two representative examples of transport phenomenon. Skin is another great example of a membrane for the entry of drugs. The transport of drug molecules through a non-porous membrane occurs by diffusion. Transport through porous cell membranes occurs by diffusion and convection. The rate of diffusion is expressed by equation (1.2). dM (C1 − C2) = DSK … (12) dt h Where, M is amount of drug dissolved, t is time, D is diffusion coefficient of the drug, S is surface area of membrane, K is oil/water partition coefficient, h is thickness of the liquid film, C1 is the concentration of drug at donor side of membrane and C2 is the concentration of drug at receptor side and C1 – C2 is concentration gradient. However, C1 and C2 are not measured since these are values varies within the membrane. Typically, the gradient is measured as Cd − Cr, representing the partition at each phase, namely Ko/w = C1/Cd and Ko/w = C2/Cr. The rate of drug transport into diffusional system is predominantly dependent upon the magnitude of the concentration gradient considering the other parameters constant. Water, carbon dioxide, and oxygen are among the few simple molecules that can cross the cell membrane by diffusion (or a type of diffusion known as osmosis). Gas exchange in lungs operates by diffusion process. All cells because of cellular metabolic processes produce carbon dioxide. Since the source is inside the cell, the concentration gradient is constantly being replenished/re-elevated; leading to net flow of CO2 out of the cell. Metabolic processes in animals and plants usually require oxygen, which is in lower concentration inside the cell, have the net flow of oxygen into the cell through diffusion. 1.9 SOLUBILITY OF GAS IN LIQUIDS Solubility of gas in liquids is the concentration of dissolved gas in the liquid when it is in equilibrium with the pure gas above the solution. The example of gas in liquid includes effervescent preparations containing dissolved carbon dioxide, ammonia water and hydrochloride gas. Aerosol products containing nitrogen or carbon dioxide as propellant are also considered to be solution of gases in liquids. 1.12 Physical Pharmaceutics - I Solubility of Drugs Factors Affecting Solubility of Gas in Liquids: The solubility of gas in liquids depends on pressure, temperature, salt present, chemical reaction and micellar solubilization. Pressure: Liquids and solids exhibit practically no change of solubility with changes in pressure. When considering solubility of gases in liquids, the pressure of the gas in contact with the liquid is important. At higher gas pressure, more gas is dissolved in liquids, Fig 1.6. For example, the soda bottle is packed at high pressure of carbon dioxide before sealing. When the cap of bottle is opened, the pressure above the liquid is reduced to 1 atm and the soda fizzes. This fizzing is just carbon dioxide that was dissolved in soda, is getting released. Therefore, if lower is the pressure less carbon dioxide is soluble. Low pressure High pressure Less few gas More gas Low pressure equilibrium, Double the pressure molecules molecules Low concentration equilibrium, soluble soluble Double the concentration Figure 1.6: Solubility of Gases at Different Pressures The effect of pressure on the solubility of gas is given Henry’s law which states that in dilute solution the mass of gas which dissolves in each volume of liquid solvent at constant temperature is directly proportional to partial pressure of gas. Mathematically it is expressed as Sg = KHPg … (1.3) Where, Sg is solubility of gas, expressed as mol/L; KH is Henry law constant which is different for each solute-solvent system and Pg is partial pressure of the gas in mmHg. The amount of undissolved gas above the solution is obtained by subtracting the vapour pressure of the pure liquid from the total pressure of the solution. Example 1.2: The solubility of a pure gas in water at 25 °C and at 1 atm pressure is 1.5 × 10−3 mol/L. What will be the concentration of the gas at same temperature at 0.5 atm? Solution: Given that: Pressure = 1 atm = 101.3 kPa Concentration = 1.5 × 10−3 mol/L Solubility (Sg) = ? Sg = KHPg 1.5 × 10−3 = KH × 101.3 KH = 1.519 × 10−5 1.13 Physical Pharmaceutics - I Solubility of Drugs Now, at P = 0.5 atm = 0.5 × 101.3 kPa Sg = KHPg = 1.519 × 10−5 × 0.5 × 101.3 = 7.693 × 10-4 mole/L The concentration of gas at 25°C and at 0.5 atm pressure will be 7.693 × 10−4 mole/L. 1.10 SOLUBILITY OF LIQUIDS IN LIQUIDS 1.10.1 Binary Solutions It is very common for two or more liquids to be mixed together to make a solution. Therefore, we need to know what liquids can be mixed together without precipitation. Examples of pharmaceutical solutions of liquid dissolved in liquids are hydroalcoholic solutions, aromatic waters, spirits, elixirs, lotions, sprays and some medicated oils that contain mixture of two or more miscible oils. When two or more liquids mixed together they can be completely miscible, partially miscible or practically immiscible. Completely miscible liquids mix uniformly in all proportions and hence do not get separated. Partially miscible liquids form two immiscible liquid layers, each of which is saturated solution of one liquid in the other. Such liquid pairs are called as conjugated liquid pairs. The mutual solubility of partially miscible liquids, being temperature specific, is affected by changes in temperature. For binary phase systems, such as phenol-water system, the mutual solubility of two conjugate liquid phase increases with increase in temperature called as conjugate temperature, where as above this temperature they are soluble in any proportions. Other examples of partial miscibility include conjugate liquid pair of nicotine and water, ether and water, and triethnolamine and water. Immiscibility refers to those systems which do not mix with each other at all such as water and liquid paraffin or water and oil. The dielectric constant of a substance also affects the solubility of substance, Fig. 1.7. Dielectric constant 0 20 40 60 80 90 Solubility (mg/ml) 75 B 60 A 45 30 15 9.5 13.0 16.5 20.0 23.5 Solubility parameter Figure 1.7: Effect of Dielectric Constant on Solubility 1.14 Physical Pharmaceutics - I Solubility of Drugs It is known fact that the polarity of solvent is dependent on the dielectric constant. Also, remember that LIKE DISSOLVES LIKE. The influence of a foreign substance on a liquid-liquid system is like the idea of three component system in the phase rule. Ternary systems are produced by addition of third component to a pair of partially miscible liquids to produce a solution. If added component is soluble in only one of the two components or if its solubility in the two liquids is markedly different, the mutual solubility of the liquid pair is decreased. If added solute is roughly soluble in both the liquids approximately to the same extent, then the mutual solubility of the liquid pair is increased. This is called blending. An example of this is when succinic acid is added to the phenol-water mixture. The succinic acid is soluble or completely miscible in each phenol and water therefore it causes a blending of the liquids making the mixture one phase. 1.10.2 Ideal Solutions Dilute solutions consists of negligible amount of solute compared to pure solvents. These solutions are referred as ideal solutions. An ideal solution is one in which there is no change in the properties of the components other than dilution when they are mixed to form the solution. No heat is evolved or absorbed during the solution formation. The final volume of real solution is an additive property of the individual component. In another way it can be stated as a solution which shows no shrinkage or expansion when components are mixed to form solution. Ideal solutions are formed by mixing different substances having similar properties and therefore there is complete uniformity of attractive intermolecular forces. For example, when equal amounts of methanol and ethanol are mixed together, the final volume of the solution is the sum of the volumes of the methanol and ethanol. Solutions used in pharmacy consist of wide variety of solutes and solution. The basis of solubility and solution theory is based on ideal solution. In ideal solution there is a complete absence of attractive or repulsive forces and therefore the solvent does not affect solubility. The solubility in this case depends on temperature, the melting point of solute and the molar heat of fusion (∆Hf). In ideal solution heat of solution is equal to ∆Hf. Therefore solubility in an ideal solution can be expressed by, i ∆Hf To − T − log X2 =   2.303R  ToT  … (1.4) i Where, X2 is the ideal solubility in terms of mole fraction, R is gas constant; T is the temperature of solution and To is the temperature (Kelvin) of solute. The equation (1.4) can be used to calculate molar heat of fusion by plotting the log solubility versus reciprocal of absolute temperature which results in a slope of − ∆Hf/2.303R. Unfortunately most of the solutions are non-ideal (real) because there may be interaction between solute and solvent. In these solutions mixing of solute and solvent can release or absorb heat into or from surroundings, respectively. While describing non-ideal solution, activity of solute must be considered. Activity of solute is defined as concentration of solute multiplied by the activity coefficient (X2). The activity coefficient is proportional to the volume of solute and to 1.15 Physical Pharmaceutics - I Solubility of Drugs the fraction of the total volume occupied by the solvent. On substitution these values in equation (1.4) we get;  ∆Hf  To − T − log X2 =     + log (µ2) … (1.5) 2.303R  To  As activity approaches unity, the solution becomes more ideal. For example, as a solution become more dilute the activity increases and the solution becomes ideal. The log of activity coefficient (log X2) is the term that considers the work of solubilization, volume of solute and the volume of solvent. The work of solubilization includes the intermolecular forces of attraction removing molecule from the solid and integrating into the solvent. One more term solubility parameter (γ2) which is a measure of cohesive forces between like molecules is considered for solubility. It is expressed by following equation. 2 V2φ1 − log γ2 = (ρ1 − ρ2) … (1.6) 2.303 (∆Hv − RT)1/2 Ρ =   … (1.7)  V1  Where, ∆Hv is heat of vapourization of solute, V1 is volume/mole of solute as a liquid, V2 is the molar volume of solute and φ12 is the volume fraction of solvent, T is temperature (Kelvin) and R is gas constant. Example 10.2: The molar heat of fusion and melting point of benzoic acid is 4139 cal/mole and 122°C, respectively. Calculate ideal mole fraction solubility of benzoic acid at 25oC. Given: Gas constant = 8.134 J/K mole. Solution: Given that: To = 122 °C = 273 + 122 = 395 K T = 25 °C = 273 + 25 = 298 K R = 8.134 J/K.mole ∆Hf = 4139 cal/mole = 4139 × 4.184 = 17317.58 J/mole i  ∆Hf  (To – T) − log X2 =    (T T)  (2.303R)  o   17317.58  (395 – 298) =   (395 × 298) (2.303 × 8.314)   17317.58  97  =      19.1471  117710 = 0.7453 i X2 = antilog (− 0.7453) = 0.1798 The ideal mole fraction solubility of benzoic acid is 0.1798. 1.16 Physical Pharmaceutics - I Solubility of Drugs 1.11 RAOULT’S LAW In an ideal solution volume changes are negligible. Dilute solutions show colligative properties. These properties are the factors that determine how properties of a bulk solution change depending upon the concentration of the solute in it. Colligative properties are properties of a solution that depend mainly on the relative numbers of particles of solvent and solute molecules and not on the chemical properties of the molecules themselves. These can almost be referred as statistical properties because they can be understood solely based on relative number of different particles in a solution. There are four types of colligative properties namely: 1. Lowering of vapour pressure. 2. Elevation of boiling point. 3. Depression of freezing point. 4. Osmotic pressure.. Colligative properties of non-electrolyte solutions are regular. The values of colligative properties are approximately equal for equimolar concentration of drugs. It is possible to determine the number of solute particles present in the solution by measuring these properties and comparing them with the corresponding properties of the pure solvent. If mass of solute present in known, the number average molecular weight can be calculated by dividing the mass of solute by number of particles present to obtain the average mass of particles. Osmotic pressure is the most important colligative property since it is related with physiological compatibility of parentral, ophthalmic and nasal solution. It is difficult and inconvenient to measure osmotic pressure and therefore other colligative properties are determined and related to osmotic pressure. In the following section equations for colligative properties of ideal solution are derived and are validated for this type of solutions. These equations can be applied to real solutions with respect to limit of small concentrations. While using these equations for real (non-ideal) solutions it requires correction to be made to these ideal equations because in real solutions there exist intermolecular interactions. Lowering of Vapour Pressure: Lowering of vapour pressure is the simplest of the colligative properties and easiest to understand based on physical model. The pressure brought by vapour in equilibrium with its liquid at constant temperature is known as vapour pressure. It increases with temperature. The vapour pressure of solvent is due to its escaping tendency. Temperature at which the vapour pressure of the liquid is equal to the atmospheric pressure is called as normal boiling point. The vapour pressure of pure liquid solvent depends upon the rate of escape of molecule from the surface known as escaping tendency. Solvents with greater escaping tendencies have greater vapour pressure. The added solute is generally non-volatile which does not contribute directly to the vapour pressure of the solution. The solute interferes and prevents solvent molecules from 1.17 Physical Pharmaceutics - I Solubility of Drugs escaping into the atmosphere. Therefore, the vapour pressure of solution is lower than that of pure solvent. The lowering of vapour pressure is proportional to the number of solute particles or ions. The effect of non-volatile solute on the vapour pressure may be determined in dilute solutions by applying Raoult’s law. It states that in an ideal solution the partial vapour pressure of each volatile constituent is equal to the vapour pressures of pure constituent at that temperature multiplied by its mole fraction in the solution. In equation form for two volatile constituent A and B, it can be expressed as ° PA = PA XA … (1.8) ° PB = PB XB … (1.9) ° ° where, PA and PB are partial vapour pressures, PA and PB are vapour pressures of pure constituents and XA and XB are mole fractions of the constituent A and B, respectively. The total vapour pressure of solution is sum of partial vapour pressure of each volatile constituent. Therefore, P = PA + PB … (1.10) Pure B Tota l va po ur p ress ure of s Vapour pressure (mm Hg) Pa olut r tia ion lp re ss Pure A ur e of Vapour pressure B fA s ure o s pre tial Par Mole fraction of A Mole fraction of B Figure 1.8: Partial Vapour Pressures of Volatile Constituents A and B and the Total Vapour Pressure of their Solution at Different Mole Fraction The partial vapour pressure of A and B and the total vapour pressure of solution is shown in Fig. 1.8. There are two ways to explain Raoult’s law. First, the simple visual way and the second one is a more sophisticated way based upon entropy. To describe using a simple way, consider that equilibrium is set-up where the number of molecules of solvent breaking and escaping away from the surface and some of them are sticking on to the surface again as shown in Fig. 1.9. An added solute molecule to the solvent replaces some of the solvent molecules present at the surface causing reduction in surface area. 1.18 Physical Pharmaceutics - I Solubility of Drugs Solvent molecule Solute leaving from surface molecule Pure solvent Solution Figure 1.9: Lowering of Vapour Pressure on Addition of Non-volatile Solute A certain fraction of the solvent molecules has enough energy to escape from the surface. If these molecules are decreased as added solute replaces some of them causing reduction in the number of molecules escaping from the surface. The net result of this reduction in number is that the vapour pressure of the solvent is reduced. The composition of the solution in terms of mole fraction can be expressed as XA + XB = 1 … (1.11) ∴ XA = 1 − XB … (1.12) Substituting equation (1.12) in equation (1.8) gives ° PA = PA (1 − XB) … (1.13) Simplifying equation (1.13) we get ° (PA − PA) XB = ° … (1.14) PA Substituting terms for mole fraction in equation (1.14) gives ° (PA − PA) nB ° = … (1.15) (nA + nB) PA where, nA and nB are number of moles of solute and solvent. Above equations (1.14) and (1.15) shows that relative lowering of vapour pressure of the solution is equal to the mole fraction of the solute. The mole fraction and vapour pressure in equation (1.14) and (1.15) has no units because these are relative expressions. Hence any units consistent with the system can be used. Deviations from Raoult’s Law: In real solutions, there is no complete uniformity of intermolecular attractive forces. There are many such liquid pairs that show greater cohesive forces than the attractive forces and greater attractive forces than the cohesive forces. It can be observed even when liquids are completely miscible in all proportions. Such mixtures of liquid pairs are real or non-ideal 1.19 Physical Pharmaceutics - I Solubility of Drugs solutions. They do not adhere to the Raoult’s law over the entire range of concentrations and are represented as deviations. This behaviour shown by liquid mixtures are called as positive deviation, Fig. 1.10 (a) and negative deviation, Fig. 1.10 (b). Total vapour pressure of solution Total vapour pressure of solution Pure B Pure B Vapour pressure (mm Hg) Vapour pressure (mm Hg) Pure A Pure A Vapour pressure Vapour pressure Mole fraction of A Mole fraction of A Mole fraction of B Mole fraction of B Figure 1.10: Deviations from Raoult’s Law Limitations of Raoult’s Law: Raoult’s law work only for ideal solutions over entire range of concentrations. An ideal solution obeys Raoult’s law. While applying this law to real solutions it has following limitations. Real Solutions: In real solution, the concentration of solute is high and thus intermolecular forces between solute-solute and solute-solvent are predominant that slows down the escaping of solvent molecules from the surface. This causes deviation from Raoult’s law because it is applicable only to dilute solutions where the forces between solute and solvent are exactly same as those between solvent-solvent molecules. Nature of the Solute: Raoult’s law is applicable only for solutes which are non-volatile in nature. Volatile solutes can contribute for vapour pressure above the solution which may cause the deviation from Raoult’s law. Raoult’s law does not apply if the added solute associates or dissociates in solvent. If association takes place the number of particles or molecules decreases causing reduction in lowering of vapour pressure. On the contrary, if solute gets dissociated more number of particles or ions are formed. For example, when 1 mole of solid sodium chloride is added to water it dissociates to produce two moles of ions as Na+ and Cl–. − + − Na+ Cl(solid) → Na(aq) + Cl(aq) If 0.1 mole of sodium chloride is added to water its dissociation takes place to form 0.2 moles of particles in solution. Thus, it increases lowering of vapour pressure of solution. 1.20 Physical Pharmaceutics - I Solubility of Drugs 1.12 REAL SOLUTIONS Real solutions show change in the total volume of the solution upon mixing its different components together. Also, there is absorption or evolution of heat during mixing and solution formation. For example, at room temperature when 100 mL of sulfuric acid is mixed with 100 mL of water, the total volume of solution becomes 180 mL rather than 200 mL. During mixing of acid and water considerable heat is evolved causing reduction in total volume of the solution. 1.13 PARTIALLY MISCIBLE LIQUIDS Although three types of liquid/liquid systems are commonly encountered liquid- liquid systems are mainly divided into two categories depending on the solubility of one substance in the other. The categories are complete miscibility and partial miscibility. Miscibility is the common solubilities of the components in liquid-liquid systems. Partial miscibility is when the substances only mix partially. When mixed, there are two layers formed each layer containing some of both liquids. Of these two mixed layers, each layer contains some of both the liquids for example, phenol and water. Some liquids are practically immiscible (for example, water and mercury), whilst others (for example, water and ethyl alcohol or acetone) mix with one another in all proportions. The mutual solubility or miscibility of two liquids is a function of temperature and composition. When two liquids (liquid A and liquid B) are partially soluble in each other, two liquid phases can be observed. At equilibrium, each phase contains liquid A and liquid B in amounts that reflect their mutual solubility. Some systems are totally miscible (i.e. they form a one-phase liquid) at high temperatures, but separate into two liquid phases at lower temperatures. These systems have an upper consolute temperature, TUCT, in a plot of temperature versus mole fraction. Other systems are totally miscible at low temperatures but separate into two phases at higher temperatures giving rise to a lower consolute temperature, TLCT. Oil and water don’t mix. Pouring 10 mL of olive oil into 10 mL of water results in two distinct layers, clearly separated by a curved meniscus. Each layer has the same volume and essentially the same composition as the original liquids. Because very little mixing occurs apparently, the liquids are called “immiscible”. For example, pouring grain alcohol into the water results in a single liquid phase. No meniscus forms between the alcohol and the water, and the two liquids are considered “miscible”. Nearly any pair of liquids is miscible if only a trace amount of one of the liquids is present. Many liquid mixtures fall between these two extremes. Two liquids are “partially miscible” if shaking equal volumes of the liquids together results in a meniscus visible between two layers of liquid, but the volumes of the layers are not identical to the volumes of the liquids originally mixed. For example, shaking water with certain organic acids results in two clearly separate layers, but each layer contains water and acid (with one layer mostly water and the other, rich in acid.) Liquids tend to be immiscible when attractions between like molecules 1.21 Physical Pharmaceutics - I Solubility of Drugs are much stronger than attractions between mixed pairs. Many examples are known, however, in which the liquids are partially miscible with one another. If, for example, water be added to ether or if ether be added to water and the mixture shaken, solution will form up to a certain point; beyond this point further addition of water on the one hand, or of ether on the other, will result in the formation of two liquid layers, one consisting of a saturated solution of water in ether and the other a saturated solution of ether in water. Two such mutually saturated solutions in equilibrium at a temperature are called conjugate solutions. A conjugate system has two partially miscible liquids in contact with each other. The proportionate quantities of these liquids are responsible for their existence as two liquids in contact with. Under this condition a saturated solution of one liquid in other or vice-versa is formed. The miscibility of such solution mixture can be increased by increasing temperature. For example, phenol – water, nicotine – water, triethanolamine – water etc. Phenol-water solution is characterized by increasing mutual solubility with rise of temperature. Thus, when phenol is added to water at the ordinary temperature, a homogeneous liquid is produced. When the concentration of the phenol in the solution has risen to about 8 %, the addition of more phenol results in the formation of a second liquid phase, which may be regarded as a solution of water in phenol. If now the temperature is raised, the second liquid phase will disappear and more phenol must be added to produce a separation of the liquid into two layers. By increasing the amount of phenol in this way and observing the temperature at which the two layers disappear, the so-called solubility curve of phenol in water may be determined. In a similar manner, the solubility curve of water in liquid phenol may be obtained, and it is found that the solubility also increases with rise of temperature. Since with rise of temperature the concentration of water in the phenol layer and of phenol in the water layer increases. The compositions of the two conjugate solutions become more and more nearly the same and at a certain temperature the two solutions become identical in composition. The temperature at which the two layers become identical in composition and are in fact one layer is known as the critical solution temperature or the consolute temperature of the system. Above this temperature, the two liquids are miscible in all proportions. If the resulting mixture is represented by a point in the area enclosed by the solubility curve, separation into two layers will take place, whereas if the total composition of the mixture and the temperature is expressed by a point lying outside the solubility curve a clear homogeneous solution will result. 1.14 CRITICAL SOLUTION TEMPERATURE AND ITS APPLICATIONS A phase diagram is a plot describing conditions of temperature and pressure under which two or more physical states coexist in dynamic equilibrium. It means phase diagram is a graphical representation of chemical equilibrium. This diagram is also called as Pressure – Temperature graph. 1.22 Physical Pharmaceutics - I Solubility of Drugs 218 atm B Pressure (mmHg) Liquid Solid 1 atm. Vapour O 4.58 atm. Triple point A C 0 100 0.0098 Temperature (oC) Figure 1.11: Phase Diagram of Water System In phase diagram of water there are three lines or curves that separate the area of each phase. Adjacent to each line there exist a different single phase of water. At any point on line there exist equilibrium between two phases shown by area i.e. solid/liquid, liquid/vapour and solid/vapour. The line OA, OB and OC represents equilibrium between liquid and vapour, solid and liquid and solid and vapour phases, respectively. The line OA represents vapourization curve and OC represent sublimation curve. For example, above line OA the liquid-water exist and below it water vapour exists. The liquid – vapour equilibrium curve has a top limit labeled as C in the phase diagram. This is known as critical point. Water has a critical point of 374°C. The temperature and pressure corresponding to this point is known as the critical temperature and critical pressure, respectively. The solid – liquid equilibrium line (m. p. line) slopes backwards (negative slope) rather than forward (positive slope). It means in case of water; the melting point gets lower at higher pressures. At solid – liquid equilibrium the ice is less dense than liquid water formed as it melts, and the water formed occupies a smaller space. At this equilibrium if pressure is increased the equilibrium move to reduce the pressure again. That means it moves to the side with smaller volume. To make the liquid water freeze again at this high pressure, we need to reduce the temperature. Higher pressure means lower melting point. The transition temperature (TUCT) of a system helps to determine percent purity of substances. The change in TUCT is proportional to the concentration of substance added. For example, in phenol-water system addition of sodium chloride or potassium chloride changes its TUCT depending upon concentration of these substances. If known different concentration solutions of sodium chloride are prepared and added separately to phenol-water mixtures having composition say 50:50, then TUCT of the system is determined by plotting a phase diagram by taking concentrations of sodium chloride on x-axis and UCT on y-axis. An unknown solution of sodium chloride is then added to phenol – water (50:50) system and 1.23 Physical Pharmaceutics - I Solubility of Drugs again TUCT is determined. It is plotted on curve to obtain its concentration by extrapolating on x-axis. The TUCT is mostly used as criterion to test the purity of substances that form conjugate system with some other liquid. Phenol USP is a necrotic agent having freezing point 17ºC. Thus, at room temperature it exists in solid crystalline form. The corrosive characteristic and solid nature of phenol makes it difficult to handle. The Liqueﬁed Phenol BP contains 80% w/w of phenol in water. The presence of other substance or impurities solidiﬁes phenol approximately at about 10°C. The miscibility curve of phenol-water system suggest that 76% w/w of phenol should be used in the preparation. At this concentration freezing point of phenol is 3.5°C. Such preparations remain in liquid form that can be handled easily. In India, we have wide variety of climatic conditions with diverse temperatures ranging from 10 – 40°C during different seasons. Hence, a preparation which is in dry powder state in winter or rainy season would become pasty during summer. The TUCT can also be used to determine percent compositions of each component in unknown mixtures. The temperature below which when system containing partially miscible liquids exist only as a single phase is known as lower consolute temperature (TLCT). For example, triethanolamine (TEA) - water system has TLCT of about 18.5°C at 13% w/w of TEA. The temperature – concentration plot of this system is shown in Fig. 1.12. Above 18.5°C mixture of these liquids forms two layers. The left upward curve shows decrease in miscibility of TEA in water whereas right upward curve shows decrease in miscibility of water in TEA with increase in temperature of system, respectively. At 50% by weight of TEA in water at 18.5°C forms single phase. This temperature is called TLCT of TEA – water system. The region outside the curve shows mutual solubility of TEA and water in each other. Other examples of liquid pairs that shows TLCT are dimethylamine – water (43°C, 13% w/w weight of dimethylamine), 1-methyl piperidine – water (48°C, 5% w/w of piperidine), polyethylene glycol – water, paraldehyde – normal saline, water – Tween 80, etc. Temperature ( C) o Two liquid phases 18.5oC LCT 100% water Composition (%w/w) 100% Trithanolamine Figure 1.12: Phase Diagram of Triethanolamine – Water System 1.24 Physical Pharmaceutics - I Solubility of Drugs Single liquid phase UCT C 66.8oC Total liquid phases F E G T (55oC) Temperature ( C) o Tie line p/w + w/p A B p/w w/p 0 11 34.5 63 100 Concentration of phenol (%weight) Figure 1.13: Phase Diagram of Phenol – Water System Phenol and water are partially miscible liquids at room temperature. In this system, addition of small amount of phenol to water or water to phenol signiﬁcantly changes relative volumes of two layers but not their compositions. If temperature is increased by keeping composition constant the mutual solubility of both the liquids increases and at a specific temperature they become completely miscible and two layers becomes one. Thus, at a speciﬁc temperature the composition of both the components are ﬁxed and both the liquids are miscible in all proportions with each other. The temperature at which two partially miscible liquids are in the state of one phase is known as critical solution temperature (CST) or upper consolute temperature (TUCT). This behaviour of critical solution temperature is shown by phenol-water system as represented in Fig. 1.13. At any temperature (say T°C) the points F and C represents the composition of two layers in equilibrium with each other. The two solutions A (phenol in water) and B (water in phenol) are in equilibrium at a temperature is known as conjugate solution temperature. At this temperature two solutions of different concentrations exist in equilibrium with each other. The line in phase diagram of phenol-water joining the points F and E is called as tie line. It is deﬁned as the line which connects the compositions of the two layers in equilibrium on the phase diagrams of the system. At poin

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