Properties of Solutions

Classification of Matter Solutions are homogeneous mixtures

Solute A solute is the dissolved substance in a solution. A solvent is the dissolving medium in a solution. Solvent Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Water in salt waterWater in soda

Calculations of Solution Concentration Mass percent - the ratio of mass (in grams) of solute to mass (in grams) of solution, expressed as a percent

Calculations of Solution Concentration Mass/volume (m/v) % - the ratio of mass (in grams) of solute to volume of solution (in mL), expressed as a percent

Calculations of Solution Concentration Volume/volume (v/v) % - the ratio of volume (in mL) of solute to volume of solution (in mL), expressed as a percent

Calculations of Solution Concentration Mole fraction – the ratio of moles of solute to total moles of solution

Calculations of Solution Concentration Molarity (M) - the ratio of moles of solute to liters of solution

Calculations of Solution Concentration Normality (N) – moles of equilvalents/Liter of solution

Calculations of Solution Concentration Molality (m) – moles of solute per kilogram of solvent

“Like Dissolves Like” Fats Benzene Benzene Steroids Steroids Hexane Hexane Waxes Waxes Toluene Toluene Inorganic Salts Water Water Sugars Sugars Small alcohols Small alcohols Acetic acid Acetic acid Polar and ionic solutes dissolve best in polar solvents Nonpolar solutes dissolve best in nonpolar solvents

Heat of Solution The Heat of Solution is the amount of heat energy absorbed (endothermic) or released (exothermic) when a specific amount of solute dissolves in a solvent. Substance Heat of Solution (kJ/mol) NaOH NH 4 NO KNO HCl-74.84

Steps in Solution Formation  H 1 Expanding the solute  H 2 Expanding the solvent  H 3 Interaction of solute and solvent to form the solution Separating the solute into individual components Overcoming intermolecular forces of the solvent molecules

Enthalpy Changes in Solution The enthalpy change of the overall process depends on  H for each of these steps. Start End Start

Why do endothermic processes sometimes occur spontaneously? Some processes, like the dissolution of NH 4 NO 3 in water, are spontaneous at room temperature even though heat is absorbed, not released.

Predicting Solution Formation Solvent/Solute H1H1H1H1 H2H2H2H2 H3H3H3H3  H sol’n Outcome Polar/Polar + large - large +/- small Solutionforms Polar/Nonpolar + small + large +/- small + large No solution forms Nonpolar/Nonpolar + small +/- small +/-smallSolutionforms Nonpolar/polar + large + small +/- small + large No solution forms

Solubility Trends  The solubility of MOST solids increases with temperature.  The rate at which solids dissolve increases with increasing surface area of the solid.  The solubility of gases decreases with increases in temperature.  The solubility of gases increases with the pressure above the solution.

Enthalpy Is Only Part of the Picture Entropy is a measure of: Dispersal of energy in the system. Number of microstates (arrangements) in the system. b. has greater entropy,  is the favored state (more on this in chap 19)

Therefore… Solids tend to dissolve best when: o Heated o Stirred o Ground into small particles Gases tend to dissolve best when: o The solution is cold o Pressure is high

Saturation of Solutions  A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated.  A solution that contains less solute than a saturated solution under existing conditions is unsaturated.  A solution that contains more dissolved solute than a saturated solution under the same conditions is supersaturated.

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

Solubility Chart

Gases in Solution In general, the solubility of gases in water increases with increasing mass. Why? Larger molecules have stronger dispersion forces.

Gases in Solution 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. Increasing pressure above solution forces more gas to dissolve.

Temperature Higher temperature drives gases out of solution.  Carbonated soft drinks are more “bubbly” if stored in the refrigerator.  Warm lakes have less O 2 dissolved in them than cool lakes.

Henry’s Law The concentration of a dissolved gas in a solution is directly proportional to the pressure of the gas above the solution Applies most accurately for dilute solutions of gases that do not dissociate or react with the solvent Yes  CO 2, N 2, O 2 No  HCl, HI

Colligative Properties Colligative properties depend only on the number of solute particles present, not on the identity of the solute particles. Among colligative properties are  Vapor pressure lowering  Boiling point elevation  Melting point depression  Osmotic pressure

Vapor Pressure As solute molecules are added to a solution, the solvent becomes less volatile (=decreased vapor pressure). Solute-solvent interactions contribute to this effect.

Raoult’s Law The presence of a nonvolatile solute lowers the vapor pressure of the solvent. P solution = Observed Vapor pressure of the solution P 0 solvent = Vapor pressure of the pure solvent  solvent = Mole fraction of the solvent

Liquid-liquid solutions in which both components are volatile Modified Raoult's Law: P 0 is the vapor pressure of the pure solvent P A and P B are the partial pressures

Colligative Properties of Electrolytes Because these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) show greater changes than those of nonelectrolytes. e.g. NaCl dissociates to form 2 ion particles; its limiting van’t Hoff factor is 2.

The van’t Hoff Factor, i Electrolytes may have two, three or more times the effect on boiling point, freezing point, and osmotic pressure, depending on its dissociation.

Dissociation Equations and the Determination of i NaCl(s)  AgNO 3 (s)  MgCl 2 (s)  Na 2 SO 4 (s)  AlCl 3 (s)  Na + (aq) + Cl - (aq) Ag + (aq) + NO 3 - (aq) Mg 2+ (aq) + 2 Cl - (aq) 2 Na + (aq) + SO 4 2- (aq) Al 3+ (aq) + 3 Cl - (aq) i = 2 i = 3 i = 4

van’t Hoff Factor One mole of NaCl in water does not really give rise to two moles of ions.

van’t Hoff Factor Some Na + and Cl − reassociate as hydrated ion pairs, so the true concentration of particles is somewhat less than two times the concentration of NaCl.

The van’t Hoff Factor Reassociation is more likely at higher concentration. Therefore, the number of particles present is concentration dependent.

Boiling Point Elevation and Freezing Point Depression Solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

Boiling Point Elevation and Freezing Point Depression In both equations,  T does not depend on what the solute is, but only on how many particles are dissolved.  T b = K b  i  m  T f = K f  i  m

Boiling Point Elevation Each mole of solute particles raises the boiling point of 1 kilogram of water by 0.51 degrees Celsius. K b = 0.51  C  kilogram/mol m = molality of the solution van’t Hoff i = van’t Hoff factor

Boiling Point Elevation The change in boiling point is proportional to the molality of the solution:  T b = K b  i  m where K b is the molal boiling point elevation constant, a property of the solvent.  T b is added to the normal boiling point of the solvent.

Freezing Point Depression Each mole of solute particles lowers the freezing point of 1 kilogram of water by 1.86 degrees Celsius. K f = 1.86  C  kilogram/mol m = molality of the solution van’t Hoff i = van’t Hoff factor

Freezing Point Depression The change in freezing point can be found similarly:  T f = K f  i  m Here K f is the molal freezing point depression constant of the solvent.  T f is subtracted from the normal freezing point of the solvent.

Freezing Point Depression and Boiling Point Elevation Constants,  C/ m SolventKfKf KbKb Acetic acid Benzene Nitrobenzene Phenol Water

Osmotic Pressure The minimum pressure that stops the osmosis is equal to the osmotic pressure of the solution

Osmotic Pressure The pressure required to stop osmosis, known as osmotic pressure, , is nVnV  = i ( ) RT = iMRT where M is the molarity of the solution If the osmotic pressure is the same on both sides of a membrane (i.e., the concentrations are the same), the solutions are isotonic.

ii

Molar Mass from Colligative Properties We can use the effects of a colligative property such as osmotic pressure to determine the molar mass of a compound. K

Suspensions and Colloids Suspensions and colloids are NOT solutions. Suspensions: The particles are so large that they settle out of the solvent if not constantly stirred. Colloids: The particles intermediate in size between those of a suspension and those of a solution.

Types of Colloids ExamplesDispersingMediumDispersedSubstance Colloid Type Fog, aerosol sprays GasLiquidAerosol Smoke, airborn germs GasSolidAerosol Whipped cream, soap suds LiquidGasFoam Milk, mayonnaise LiquidLiquidEmulsion Paint, clays, gelatin LiquidSolidSol Marshmallow, Styrofoam SolidGas Solid Foam Butter, cheese SolidLiquid Solid Emulsion Ruby glass SolidSolid Solid sol