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Solubility. Solutions. INTRODUCTION. Solutions are homogeneous mixtures of two or more pure substances. In a solution, the solute is dispersed uniformly throughout the solvent. Importance of studying the phenomenon of solubility

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  1. Solubility

  2. Solutions INTRODUCTION • Solutions are homogeneous mixtures of two or more pure substances. • In a solution, the solute is dispersed uniformly throughout the solvent.

  3. Importance of studying the phenomenon of solubility • Understanding the phenomenon of solubility helps the pharmacist to • Select the best solvent for a drug or a mixture of drugs. • Overcome problems arising during preparation of pharmaceutical solutions. • Have information about the structure and intermolecular forces of the drug.

  4. How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.

  5. If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal. A sodium ion solvated by water molecules.

  6. The amount of solute that can be dissolved in a certain volume of solvent at a certain temperature. • The solubility of a solid increases as temperature increases. • More solid can be dissolved if a greater volume of solvent is used. SOLUBILITY

  7. True solution: is a mixture of two or more components that form a homogenous molecular dispersion. The components are referred to the solute & the solvent . • Solute: is the dissolved agent (less abundant part of the solution). • Solvent: is the component in which the solute is dissolved (more abundant part of the solution).

  8. Types of Solutions • Saturated • Solvent holds as much solute as is possible at that temperature. • Dissolved solute is in dynamic equilibrium with solid solute particles.

  9. Unsaturated • Less than the maximum amount of solute for that temperature is dissolved in the solvent.

  10. Supersaturated • Solvent holds more solute than is normally possible at that temperature. • These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

  11. Solubility: in a quantitative way: it is the concentration - solute in a saturated solution at a certain temperature. • in a qualitative way: it is the spontaneous interaction of two - or more substances (solute & solvent) to form a homogeneous molecular dispersion.

  12. a) The solubility of a drug can be expressed in terms of: • Molarity: no of moles (or gram molecular weight) of solute dissolved in 1 lit of solution • Normality: it is measure of concentration equal to gram equivalent weight per lit of solution. • Molality: no of moles of solute dissolved in 1000 gram of solvent. • Mole fraction: it is unit of conc. Define to be equal to no of moles of a component divided by total no of moles of solution. Solubility expression

  13. percentage (% w/w, % w/v, % v/v). • % w/w: no of grams of solute dissolved in 100 grams of solution. • % w/v: no of grams of solute dissolved in 100 ml of solution. • % v/v: no of ml of solute dissolved in 100 ml of solution. b) The USP lists the solubility of drugs as the number of ml of solvent in which 1g of solute will dissolve. E.g. 1g of boric acid dissolves in 18 mL of water, and in 4 mL of glycerin.

  14. Solubility In Descriptive Terms

  15. -Solubility depends on chemical, electrical & structural effects that lead to mutual interactions between the solute and the solvent. • In pre-or early formulation, selection of the most suitable solvent is based on the principle of “like dissolves like”. • That is, a solute dissolves best in a solvent with similar chemical properties. i.e. • Polar solutes dissolve in polar solvents. E.g salts & sugar dissolve in water . • Non polar solutes dissolve in non polar solvents. Eg. naphtalene dissolves in benzene. Solute-Solvent interactions

  16. To explain the above rule, consider the forces of attraction between solute and solvent molecules. • If the solvent is A & the solute is B, and the forces of attraction are represented by A-A, B-B and A-B, one of the following conditions will occur: 1. If A-A >> A-B The solvent molecules will be attracted to each other & the solute will be excluded. Example: Benzene & water, where benzene molecules are unable to penetrate the closely bound water aggregates. 2. If B-B >> A-A The solvent will not be able to break the binding forces between solute molecules. Example: NaCl in benzene, where the NaCl crystal is held by strong electrovalent forces which cannot be broken by benzene. 3. If A-B >> A-A or B-B, or the three forces are equal The solute will disperse & form a solution. Example: NaCl in water.

  17. Polar solvents: • Polar solvents (water, glycols, methyl & ethyl alcohol), dissolve ionic solutes & other polar substances. • Solubility of substances in polar solvents depends on structural features: - Straight chain monohydroxy alcohols, aldehydes & ketones with » 5 C are slightly soluble in water. -Branching of the carbon chain in aliphatic alcohols increases water solubility. • Tertiary butyl alcohol » soluble than n-butyl alcohol • Polyhydroxy compounds as glycerin, tartaric acid, PEG are water soluble. Classification of solvents & their mechanism of action

  18. Polar solvents acts as a solvent according to the following mechanisms: a) Dielectric constant: due to their high dielectric constant, polar solvents reduce the force of attraction between oppositely charged ions in crystals. Example: water possessing a high dielectric constant (> = 80) can dissolve NaCl, while chloroform (> = 5) & benzene (> = 2) cannot. Ionic compounds are practically insoluble in these 2 solvents. b) Hydrogen bond formation: Water dissolves phenols, alcohols and other oxygen & nitrogen containing compounds that can form hydrogen bonds with water. c) Solvation through dipole interaction: Polar solvents are capable of solvating molecules & ions through dipole interaction forces. The solute must be polar to compete for the bonds of the already associated solvent molecules. • Example: Ion-dipole interaction between sodium salt of oleic acid & water

  19. Non polar solvents • Non polar solvents are unable to reduce the attraction between the ions due to their low dielectric constants. • They are unable to form hydrogen bonds with non electrolytes. • Non polar solvents can dissolve non polar solute with similar internal pressure through induce internal dipole interaction the solute molecules are kept in solution by weak van der Waals londan type forces • Example: solutions of oils & fats in carbon tetrachloride or benzene.

  20. Semipolar solvents • Semipolar solvents, such as ketones can induce a certain degree of polarity in non polar solvent molecules. • They can act as intermediate solvents to bring about miscibility of polar & non polar liquids. • Example: acetone increases solubility of ether in water.


  22. Nature of the solute and solvent • Pressure • Temperature • Electrolytes and Non-electrolytes • Chemical reaction Solubility Of Gases In Liquids

  23. MAIN FACTORS THAT AFFECT SOLUBILITY: • Nature of the solute and solvent – The amount of solute that dissolves depends on what type of solute it is. While only 1 gram of lead (II) chloride can be dissolved in 100 grams of water at room temperature, 200 grams of zinc chloride can be dissolved. This means that a greater amount of zinc chloride can be dissolved in the same amount of water than lead II chloride. • Temperature -- Generally, an increase in the temperature of the solution increases the solubility of a solid solute. For example, a greater amount of sugar will dissolve in warm water than in cold water. A few solid solutes, however, are less soluble in warmer solutions. For all gases, solubility decreases as the temperature of the solution rises. An example of this is Soda. The solubility of the carbon dioxide gas decreases when a soda is warm, making the soda flat.

  24. Pressure -- For solid and liquid solutes, changes in pressure have practically no effect on solubility. For gaseous solutes, an increase in pressure increases solubility and a decrease in pressure decreases solubility. Example: When the cap on a bottle of soda pop is removed, pressure is released, and the gaseous solute bubbles out of solution. This escape of a gas from solution is called effervescence.

  25. The influence of solvent on the solubility of drugs Weak electrolytes can behave like strong electrolytes or like nonelectrolytes in solution. When the solution is of such a pH that the drug is entirely in the ionic form, it behaves as a solution of a strong electrolyte, and solubility does not constitute a problem. However, when the pH is adjusted to a value at which un-ionized molecules are produced in sufficient concentration to exceed the solubility of this form, precipitation occurs. Frequently, a solute is more soluble in a mixture of solvents than in one solvent alone This phenomenon is known as cosolvency, and the solvents that, in combination, increase the solubility of the solute are called cosolvents.

  26. Combined effect of pH and solvents The solubility as a function of pH, temperature, and solvent composition was examined to determine the pKa of the salt from the solubility profile at various temperatures and in several solvent systems, for example the effect of alcohol on the solubility of phenobarbital. The results showed that the pKa of Phenobarbital, 7.41, is raised to 7.92 in a hydroalcoholic solution containing 30% by volume of alcohol. Also the solubility, of un-ionized phenobarbital is increased from 0.12 g/100 mL (0.005 M) in water to 0.64% (0.28 M) in a 30% alcoholic solution.

  27. Types of solutions Solutions of pharmaceutical importance include: • gases in liquids • liquids in liquids • solids in liquids

  28. Gases in liquid Increasing pressure above solution forces more gas to dissolve. • The solubility of liquids and solids does not change appreciably with pressure. • But, the solubility of a gas in a liquid is directly proportional to its pressure.

  29. Solubility of gases in liquids • Examples of pharmaceutical solutions of gases include: HCl, ammonia water & effervescent preparations containing CO2 maintained in solution under pressure. • The solubility of a gas in a liquid is the concentration of dissolved gas when it is in equilibrium with some of the pure gas above the solution. • The solubility depends on the pressure, temperature, presence of salts & chemical reactions that sometimes the gas undergoes with the solvent

  30. Effect of pressure According to Henry’s law: C2= p In a very dilute solution at constant temperature, the concentration (C2) of dissolved gas is proportional to the partial pressure (p) of the gas above the solution at Equilibrium. (the partial pressure = total pressure above the solution minus the vapor pressure of the solvent) Caution: When the pressure above the solution is released (decreases), the solubility of the gas decreases, and the gas may escape from the container with violence. This phenomenon occurs in effervescent solutions when the stopper of the container is removed.

  31. Effect of temperature As the temperature increases the solubility of gases decreases, owing to the great tendency of the gas to expand. • Pharmaceutical application:  The pharmacist should be cautious in opening containers of gaseous solutions in warm climates.  A container filled with a gaseous solution or a liquid with high vapor pressure,such as ethyl nitrite, should be immersed in ice or cold water, before opening the container, to reduce the temperature and pressure of the gas.

  32. Effect of Salting out Adding electrolytes (NaCl) & sometimes non electrolytes (sucrose) to gaseous solutions (eg. carbonated solutions) induces liberation of gases from the solutions. Why? Due to the attraction of the salt ions or the highly polar electrolyte for the water molecules and reduction of the aqueous environment adjacent to the gas molecules.

  33. HENRY'S LAW • ‘ In a dilute solution, the mass of a gas which dissolves in a given volume of a liquid at a constant temperature is directly proportional to the partial pressure of the gas’. c = σ p

  34. Henry’s Law & Soft Drinks • Soft drinks contain “carbonated water” – water with dissolved carbon dioxide gas. • The drinks are bottled with a CO2 pressure greater than 1 atm. • When the bottle is opened, the pressure of CO2 decreases and the solubility of CO2 also decreases, according to Henry’s Law. • Therefore, bubbles of CO2 escape from solution.

  35. Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

  36. Temperature • The opposite is true of gases. Higher temp drives gases out of solution. • Carbonated soft drinks are more “bubbly” if stored in the refrigerator.

  37. Dissolving a salt... • A salt is an ionic compound - usually a metal cation bonded to a non-metal anion. • The dissolving of a salt is an example of equilibrium. • The cations and anions are attracted to each other in the salt. • They are also attracted to the water molecules. • The water molecules will start to pull out some of the ions from the salt crystal.

  38. At first, the only process occurring is the dissolving of the salt - the dissociation of the salt into its ions. • However, soon the ions floating in the water begin to collide with the salt crystal and are “pulled back in” to the salt. (precipitation) • Eventually the rate of dissociation is equal to the rate of precipitation. • The solution is now “saturated”.

  39. Na+ and Cl -ions surrounded by water molecules NaCl Crystal In a saturated solution, there is no change in amount of solid precipitate at the bottom of the beaker. Concentration of the solution is constant. The rate at which the salt is dissolving into solution equals the rate of precipitation. Solubility Equilibrium: Dissociation = Precipitation Dissolving NaCl in water

  40. Co-solvency • Solubilisation • pH modifications • Complexation • Hydrotrophy • Chemical modification of the drug TECHNIQUES OF SOLUBILITY ENHANCEMENT

  41. COMPLEXATION • The total solubility of a complex forming drug is • equal to the inherent solubility of the uncomplexed • drug plus the concentration of the complexed drug • in solution. • The total solubility can be expressed as: ST = [D] + X [DXCY]

  42. HYDROTROPHY: • Increase in the solubility of a drug in water • owing to the presence of large amount of • additives. • E.g.,The increased solubility of caffeine in • presence of sodium salicylate. • Similarly a 30% solution of sodium benzoate • readily dissolves chlorocresol.

  43. CHEMICAL MODIFICATION OF THE DRUG • The method is based on preparing the water • soluble derivatives of poorly soluble drugs. • Corticosteroids. • Betamethasone


  45. The Hildebrand solubility parameter is the square root of the cohesive energy density. HILDEBRAND SOLUBILITY PARAMETER δ =√((∆Hv – RT)/Vm)

  46. Predictions of phase equilibrium. • Predictions on solubility and swelling of polymers by solvents. • Applies only to associated solutions. • Materials with similar solubility parameters will be able to interact with each other, resulting in solvation, miscibility or swelling USES AND LIMITATIONS

  47. Developed by Charles Hansen as a way of predicting if one material will dissolve in another and form a solution. Based on the idea that ‘like dissolves like’where one molecule is defined as being 'like' another if it bonds to itself in a similar way. HANSEN SOLUBILITY PARAMETER

  48. (Ra)2=4(δd2- δd1)2 +(δp2- δp1)2 +(δh2- δh1)2 δd -The energy from dispersion bonds & molecules. δp- The energy from dipolar intermolecular force & molecules. δh- The energy from hydrogen bonds & molecules. RED=Ra/RO RED-Relative Energy Difference.

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