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Chapter 11 Properties of solutions. 11.1 – Solution composition. Non-specific ways to indicate solution composition: Dilute = low concentration Concentrated = high concentration Specific ways to indicate solution composition: Molarity Mass percent Mole fraction Molality Normality.
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11.1 – Solution composition • Non-specific ways to indicate solution composition: • Dilute = low concentration • Concentrated = high concentration • Specific ways to indicate solution composition: • Molarity • Mass percent • Mole fraction • Molality • Normality
11.1 – Solution composition • Example, car batteries contain 3.75M sulfuric acid that has a density of 1.230g/ml. Calculate the mass percent, molality, molarity, and normality of the sulfuric acid.
11.2 – The energies of solution formation • Dissolving a solute in a solvent involves three processes, each with its own enthalpy
11.2 – The energies of solution formation • Enthalpy (heat) of solution (ΔHsoln): • Typically, • ΔH1 and ΔH2 are endothermic • ΔH3 is exothermic
11.2 – The energies of solution formation • Depending on the magnitudes of ΔH1, ΔH2, and ΔH3, the ΔHsoln can be either endothermic or exothermic • If ΔHsoln >> 0, then no solution forms • If ΔHsoln is small (or negative), a solution will form
11.2 – The energies of solution formation • Cardinal rule: like dissolves like • If you have a polar solute, it will dissolve in a polar solvent • If you have a non-polar solute, it will dissolve in a non-polar solvent.
11.2 – The energies of solution formation • Example,
11.3 – Factors affecting solubility • Pressure • The amount of a gas dissolved in a solution is directly proportional to the pressure of the gas above the solution, according to Henry’s Law: • C is the concentration of gas dissolved in solution • k is a constant, specific to a particular solution • P is the pressure above the solution
11.3 – Factors affecting solubility • Pressure
11.3 – Factors affecting solubility • Pressure Example, a soft drink at 25oC contains CO2 gas at a pressure of 5.0atm over the liquid. The partial pressure of CO2 in the atmosphere is 4.0x10-4atm. Calculate the equilibrium concentrations of CO2 in the soft drink before and after it is opened. The k for CO2 in solution is 3.1x10-2mol/Latm
11.3 – Factors affecting solubility • Temperature • At higher temperatures: • The dissolving of a solid occurs more rapidly • The amount of solid that can be dissolved generally increases, but may decrease • The only way to determine is through experimentation
11.3 – Factors affecting solubility • Temperature • At higher temperatures: • The solubility of a gas in water will decrease (an important concept in thermal pollution)
11.4 – Vapor pressure of solutions • A container filled with a solvent will have a number of solvent molecules escaping into the air above it, exerting a pressure (vapor pressure) on the walls of the container.
11.4 – Vapor pressure of solutions • If a non-volatile solute is mixed in, it will lower the vapor pressure
11.4 – Vapor pressure of solutions • The vapor pressure is lowered according to Raoult’s Law: • Psoln is the observed vapor pressure of the solution • xsolvent is the mole fraction of solvent • Posolvent is the vapor pressure of the pure solvent • For a solution of ½ solute and ½ solvent molecules, the vapor pressure of the solution is half that of the pure solvent
11.4 – Vapor pressure of solutions Example. Glycerin, C3H8O3 is a nonvolatile liquid. What is the vapor pressure of a solution made by adding 164g of glycerin to 338ml of H2O at 39.8oC? The vapor pressure of water at 39.8oC is 54.74 torr and its density is 0.992g/ml
11.4 – Vapor pressure of solutions • The propellant in an aerosol spray can helps to disperse and mix the components of the aerosol can. When the valve is pressed, the dip tube opens to the atmosphere. Because the propellant is at a higher pressure than the atmosphere, it escapes through the tube, bringing the active ingredient with it.
11.5 – Boiling-point elevation and Freezing-point depression • Colligative properties: • Boiling-point elevation • Freezing-point depression • Osmotic pressure • They are properties of solution that do not depend on the identity of the solute dissolved in them, only the number of dissolved particles
11.5 – Boiling-point elevation and Freezing-point depression • van’t Hoff factor (i)
11.5 – Boiling-point elevation and Freezing-point depression • Molality (m) = moles/kg solvent • Used instead of molarity because it does not depend on the temperature
11.5 – Boiling-point elevation and Freezing-point depression • Boiling-point elevation • Generally, boiling point is defined as the temperature at which vapor pressure is equal to atmospheric pressure • We saw earlier that adding solutes lowers the vapor pressure of a solution. If vapor pressure is lowered, boiling point is increased (elevated)
11.5 – Boiling-point elevation and Freezing-point depression • Boiling-point elevation • Given by the equation: • ΔT is the change in boiling point of the solvent • kb is a constant, depending on the solvent • i is the van’t Hoff factor of the solute • m is the molality of the solution
11.5 – Boiling-point elevation and Freezing-point depression • Boiling-point elevation Example,
11.5 – Boiling-point elevation and Freezing-point depression • Freezing-point depression • The transition from liquid to solid is usually a highly ordered process where the individual molecules place themselves in a particular orientation. Example, water to ice:
11.5 – Boiling-point elevation and Freezing-point depression • Freezing-point depression • Freezing-point depression is given by: • ΔT is the change in freezing point of the solvent • kb is a constant, depending on the solvent • i is the van’t Hoff factor of the solute • m is the molality of the solution
11.5 – Boiling-point elevation and Freezing-point depression • Freezing-point depression • Example, What mass of ethylene glycol (MW = 62.1g/mol) must be added to 10.0L of water to produce a solution in the car’s radiator that freezes at -23.3oC? Assume the density of water is exactly 1g/ml
11.6 – Osmotic Pressure • Osmosis – the movement of water across a semi-permeable membrane (only allows water to pass) Osmotic pressure is defined as the minimum pressure required to prevent osmosis from occurring.
11.6 – Osmotic Pressure • Given by the equation: • Π is the osmotic pressure in atm • M is molarity • R is the ideal gas constant • T is the temperature in kelvins
11.6 – Osmotic Pressure • Example,
11.6 – Osmotic Pressure • Hypertonic = higher solute concentration • Hypotonic = lower solute concentration • Isotonic = equal solute concentration
11.6 – Osmotic Pressure • Reverse osmosis – occurs when you apply a pressure greater than the osmotic pressure to a solution, causing water molecules to pass “backwards” through the semi-permeable membrane
11.8 - Colloids • A suspension of tiny particles in some medium is called a colloidal dispersion, or a colloid. • The suspended particles are 1 – 1000nm in size. • Shining a light through a colloid, produces a scattering of light. If it was a true solution, it would not scatter light.
11.8 - Colloids • If the particles are <1nm in size, solution • If the particles are between 1nm and 1μm, colloid. • If the particles are >1μm, suspension
11.8 - Colloids • The scattering of light is referred to as the Tyndall effect. • The destruction of a colloid is called coagulation. This can be done by adding heat or an electrolyte
