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CHAPTER 2

CHAPTER 2. Solutions By Dr. Hisham Ezzat . http://www.staff.zu.edu.eg/ezzat_hisham/browseMyFiles.asp?path=./userdownloads/physical%20chemistry%20for%20clinical%20pharmacy/. The Dissolution Process. Solutions are homogeneous mixtures of two or more substances.

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CHAPTER 2

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  1. CHAPTER 2 • Solutions By Dr. Hisham Ezzat http://www.staff.zu.edu.eg/ezzat_hisham/browseMyFiles.asp?path=./userdownloads/physical%20chemistry%20for%20clinical%20pharmacy/

  2. The Dissolution Process • Solutions are homogeneous mixtures of two or more substances. • Dissolving medium is called the solvent. • Dissolved species are called the solute. • There are three states of matter (solid, liquid, and gas) which when mixed two at a time gives nine different kinds of mixtures. • Seven of the possibilities can be homogeneous. • Two of the possibilities must be heterogeneous.

  3. The Dissolution Process Seven Homogeneous Possibilities SoluteSolventExample • Solid Liquid salt water • Liquid Liquid mixed drinks • Gas Liquid carbonated beverages • Liquid Solid dental amalgams • Solid Solid alloys • Gas Solid metal pipes • Gas Gas air Two Heterogeneous Possibilities • Solid Gas dust in air • Liquid Gas clouds, fog

  4. Ways of Expressing Concentration • Qualitative Terms: • Dilute Solution – A dilute solution has a relatively small concentration of solute. • A concentrated solution has a relatively high concentration of solute.

  5. Quantitative terms Molarity Chapter 13

  6. Molality • Molality is a concentration unit based on the number of moles of solute per kilogram of solvent. Chapter 13

  7. Molality and Mole Fraction • Weight percent (wt %) • Volume percent or percent by volume (vol %)

  8. Molality and Mole Fraction • Molarity • We must introduce two new concentration units in this chapter.

  9. Molality and Mole Fraction • Example 14-1: Calculate the molarity and the molality of an aqueous solution that is 10.0% glucose, C6H12O6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has several medical uses. 1 mol C6H12O6 = 180 g

  10. Molality and Mole Fraction • Molality is a concentration unit based on the number of moles of solute per kilogram of solvent.

  11. Molality and Mole Fraction • Example 14-1: Calculate the molality and the molarity of an aqueous solution that is 10.0% glucose, C6H12O6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has several medical uses. 1 mol C6H12O6 = 180 g You calculate the molarity!

  12. Molality and Mole Fraction • Example 14-2: Calculate the molality of a solution that contains 7.25 g of benzoic acid C6H5COOH, in 2.00 x 102 mL of benzene, C6H6. The density of benzene is 0.879 g/mL. 1 mol C6H5COOH = 122 g You do it!

  13. Molality and Mole Fraction • Mole fraction is the number of moles of one component divided by the moles of all the components of the solution • Mole fraction is literally a fraction using moles of one component as the numerator and moles of all the components as the denominator. • In a two component solution, the mole fraction of one component, A, has the symbol XA.

  14. Molality and Mole Fraction • The mole fraction of component B - XB

  15. Molality and Mole Fraction • Example 14-3: What are the mole fractions of glucose and water in a 10.0% glucose solution (Example 14-1)? You do it!

  16. Molality and Mole Fraction • Example 14-3: What are the mole fractions of glucose and water in a 10.0% glucose solution (Example 14-1)?

  17. Molality and Mole Fraction • Now we can calculate the mole fractions.

  18. Chapter Goals The Dissolution Process • Spontaneity of the Dissolution Process • Dissolution of Solids in Liquids • Dissolution of Liquids in Liquids (Miscibility) • Dissolution of Gases in Liquids • Rates of Dissolution and Saturation • Effect of Temperature on Solubility • Effect of Pressure on Solubility • Molality and Mole Fraction

  19. Chapter Goals Colligative Properties of Solutions • Lowering of Vapor Pressure and Raoult’s Law • Fractional Distillation • Boiling Point Elevation • Freezing Point Depression • Determination of Molecular Weight by Freezing Point Depression or Boiling Point Elevation • Colligative Properties and Dissociation of Electrolytes • Osmotic Pressure

  20. Chapter Goals Colloids • The Tyndall Effect • The Adsorption Phenomena • Hydrophilic and Hydrophobic Colloids

  21. Spontaneity of the Dissolution Process • As an example of dissolution, let’s assume that the solvent is a liquid. • Two major factors affect dissolution of solutes • Change of energy content or enthalpy of solution, Hsolution • If Hsolution is exothermic (< 0) dissolution is favored. • If Hsolution is endothermic (> 0) dissolution is not favored.

  22. Spontaneity of the Dissolution Process • Change in disorder, or randomness, of the solution Smixing • If Smixing increases (> 0) dissolution is favored. • If Smixing decreases (< 0) dissolution is not favored. • Thus the best conditions for dissolution are: • For the solution process to be exothermic. • Hsolution < 0 • For the solution to become more disordered. • Smixing > 0

  23. Spontaneity of the Dissolution Process • Disorder in mixing a solution is very common. • Smixing is almost always > 0. • What factors affect Hsolution? • There is a competition between several different attractions. • Solute-solute attractions such as ion-ion attraction, dipole-dipole, etc. • Breaking the solute-solute attraction requires an absorption of E.

  24. Spontaneity of the Dissolution Process • Solvent-solvent attractions such as hydrogen bonding in water. • This also requires an absorption of E. • Solvent-solute attractions, solvation, releases energy. • If solvation energy is greater than the sum of the solute-solute and solvent-solvent attractions, the dissolution is exothermic, Hsolution < 0. • If solvation energy is less than the sum of the solute-solute and solvent-solvent attractions, the dissolution is endothermic, Hsolution > 0.

  25. Spontaneity of the Dissolution Process

  26. Dissolution of Solids in Liquids • The energy released (exothermic) when a mole of formula units of a solid is formed from its constituent ions (molecules or atoms for nonionic solids) in the gas phase is called the crystal lattice energy. • The crystal lattice energy is a measure of the attractive forces in a solid. • The crystal lattice energy increases as the charge density increases.

  27. Dissolution of Solids in Liquids • Dissolution is a competition between: • Solute -solute attractions • crystal lattice energy for ionic solids • Solvent-solvent attractions • H-bonding for water • Solute-solvent attractions • Solvation or hydration energy

  28. Dissolution of Solids in Liquids • Solvation is directed by the water to ion attractions as shown in these electrostatic potentials.

  29. Dissolution of Solids in Liquids • In an exothermic dissolution, energy is released when solute particles are dissolved. • This energy is called the energy of solvation or the hydration energy (if solvent is water). • Let’s look at the dissolution of CaCl2.

  30. OH2 2+ O H H OH2 H2O H O H H O H Ca Cl- H2O OH2 H H O OH2 Dissolution of Solids in Liquids

  31. Dissolution of Solids in Liquids • The energy absorbed when one mole of formula units becomes hydrated is the molar energy of hydration.

  32. Dissolution of Solids in Liquids • Hydration energy increases with increasing charge density IonRadius(Å)Charge/radiusHeat of Hydration K+ 1.33 0.75 -351 kJ/mol Ca2+ 0.99 2.02 -1650 kJ/mol Cu2+0.72 2.78 -2160 kJ/mol Al3+ 0.50 6.00 -4750 kJ/mol

  33. Dissolution of Liquids in Liquids (Miscibility) • Most polar liquids are miscible in other polar liquids. • In general, liquids obey the “like dissolves like” rule. • Polar molecules are soluble in polar solvents. • Nonpolar molecules are soluble in nonpolar solvents. • For example, methanol, CH3OH, is very soluble in water

  34. Dissolution of Liquids in Liquids (Miscibility) • Nonpolar molecules essentially “slide” in between each other. • For example, carbon tetrachloride and benzene are very miscible.

  35. Dissolution of Gases in Liquids • Polar gases are more soluble in water than nonpolar gases. • This is the “like dissolves like” rule in action. • Polar gases can hydrogen bond with water • Some polar gases enhance their solubility by reacting with water.

  36. A few nonpolar gases are soluble in water because they react with water. Because gases have very weak solute-solute interactions, gases dissolve in liquids in exothermic processes. Dissolution of Gases in Liquids

  37. Finely divided solids dissolve more rapidly than large crystals. Compare the dissolution of granulated sugar and sugar cubes in cold water. The reason is simple, look at a single cube of NaCl. The enormous increase in surface area helps the solid to dissolve faster. Breaks many smaller crystals up NaCl Rates of Dissolution and Saturation

  38. Rates of Dissolution and Saturation • Saturated solutions have established an equilbrium between dissolved and undissolved solutes • Examples of saturated solutions include: • Air that has 100% humidity. • Some solids dissolved in liquids.

  39. Symbolically this equilibrium is written as: In an equilibrium reaction, the forward rate of reaction is equal to the reverse rate of reaction. Rates of Dissolution and Saturation

  40. Rates of Dissolution and Saturation • Supersaturated solutions have higher-than-saturated concentrations of dissolved solutes.

  41. According to LeChatelier’s Principle when stress is applied to a system at equilibrium, the system responds in a way that best relieves the stress. Since saturated solutions are at equilibrium, LeChatelier’s principle applies to them. Possible stresses to chemical systems include: Heating or cooling the system. Changing the pressure of the system. Changing the concentrations of reactants or products. Effect of Temperature on Solubility

  42. What will be the effect of heating or cooling the water in which we wish to dissolve a solid? It depends on whether the dissolution is exo- or endothermic. For an exothermic dissolution, heat can be considered as a product. Warming the water will decrease solubility and cooling the water will increase the solubility. Predict the effect on an endothermic dissolution like this one. Effect of Temperature on Solubility

  43. Effect of Temperature on Solubility • For ionic solids that dissolve endothermically dissolution is enhanced by heating. • For ionic solids that dissolve exothermically dissolution is enhanced by cooling. • Be sure you understand these trends.

  44. Effect of Pressure on Solubility • Pressure changes have little or no effect on solubility of liquids and solids in liquids. • Liquids and solids are not compressible. • Pressure changes have large effects on the solubility of gases in liquids. • Sudden pressure change is why carbonated drinks fizz when opened. • It is also the cause of several scuba diving related problems including the “bends”.

  45. Effect of Pressure on Solubility • The effect of pressure on the solubility of gases in liquids is described by Henry’s Law.

  46. Colligative Properties of Solutions • Colligative properties are properties of solutions that depend solely on the number of particles dissolved in the solution. • Colligative properties do not depend on the kinds of particles dissolved. • Colligative properties are a physical property of solutions.

  47. Colligative Properties of Solutions • There are four common types of colligative properties: • Vapor pressure lowering • Freezing point depression • Boiling point elevation • Osmotic pressure • Vapor pressure lowering is the key to all four of the colligative properties.

  48. Lowering of Vapor Pressure and Raoult’s Law • Addition of a nonvolatile solute to a solution lowers the vapor pressure of the solution. • The effect is simply due to fewer solvent molecules at the solution’s surface. • The solute molecules occupy some of the spaces that would normally be occupied by solvent. • Raoult’s Law models this effect in ideal solutions.

  49. Lowering of Vapor Pressure and Raoult’s Law • Derivation of Raoult’s Law.

  50. Lowering of Vapor Pressure and Raoult’s Law • Lowering of vapor pressure, Psolvent, is defined as:

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