IB Chemistry Notes Chapter 11: Properties of Solutions

Solutions  ______– a homogeneous mixture of pure substances (occur in all phases, but we will focus on aqueous solutions)  The is the medium in which the are dissolved. o (The solvent is usually the most abundant substance.)

Concentration of Solution - refers to the amount of solute dissolved in a solution.

MOLARITY (M) =

MOLALITY (m) =

MASS PERCENT (%) =

MOLE FRACTION () =

NORMALITY (N) =

Energy of Making Solutions

 Heat of solution ( Hsoln ) is the for making a solution.  Most easily understood if broken into steps. o Break apart solvent . Have to overcome attractive forces. o Break apart solute . Have to overcome attractive forces. o Mixing solvent and solute

. H3 depends on what you are mixing. . Molecules can attract each other – . Molecules can’t attract - o This explains the rule

Size of H3 determines whether a solution will form  Types of Solvent and solutes

 If Hsoln is small and positive, a solution will still form because of entropy.  There are many more ways for them to become mixed than there is for them to stay separate. Solution Formation – Factors Favoring Spontaneity  Processes in which the energy content of the system decreases (exothermic) tend to occur spontaneously.  Processes in which the disorder (entropy) of the system increases tend to occur spontaneously.

Structure and Solubility IB Chemistry Notes Chapter 11: Properties of Solutions  Water soluble molecules must have dipole moments -

 To be soluble in non polar solvents the molecules must be .

Pressure  Changing the pressure doesn’t affect the amount of solid or liquid that dissolves o They are incompressible.  Pressure does affect solubility of gases.

Dissolving Gases  Pressure affects the amount of that can dissolve in a liquid.  The dissolved gas is at equilibrium with the gas above the liquid.  If you increase the pressure the gas molecules dissolve faster. o The equilibrium is disturbed. o The system reaches a new equilibrium with more gas dissolved.  Henry’s Law:

Temperature Effects  Increased temperature increases the rate at which a solid dissolves.  We can’t predict whether it will increase the amount of solid that dissolves.  We must read it from a graph of experimental data.  Gases are predictable  As temperature increases, solubility decreases.  Gas molecules can move fast enough to escape.  Thermal pollution.

Vapor Pressure of Solutions  A nonvolatile solvent lowers the vapor pressure of the solution.  The molecules of the solvent must overcome the force of both the other solvent molecules and the solute molecules.

Raoult’s Law:

 Applies only to an ideal solution where the solute doesn’t contribute to the vapor pressure.

To determine whether a sol’n is IDEAL…  Liquid-liquid solutions where both are volatile.  Modify Raoult’s Law to:

Ptotal =

• Ptotal = vapor pressure of mixture 0 • PA = vapor pressure of pure A

 If this equation works then the solution is ideal. IB Chemistry Notes Chapter 11: Properties of Solutions  Solvent and solute are alike.

Colligative Properties of Solutions = physical properties of solutions that depend on the # of particles dissolved, not the kind of particle.  Lowering vapor pressure  Raising boiling point  Lowering freezing point  Generating an osmotic pressure

Boiling Point Elevation: a solution that contains a nonvolatile solute has a higher boiling pt than the pure solvent; the boiling pt elevation is proportional to the # of moles of solute dissolved in a given mass of solvent.

where: Tb = elevation of boiling pt m = molality of solute

kb = the molal boiling pt elevation constant for a particular solvent

kb for water = 0.52 °C/m

Freezing/Melting Point Depression: the freezing point of a solution is always lower than that of the pure solvent.

where: Tf = lowering of freezing point m = molality of solute

kf = the freezing pt depression constant

kf for water = 1.86 °C/m

3 3 Ex: An antifreeze solution is prepared containing 50.0 cm of ethylene glycol, C2H6O2, (d = 1.12 g/cm ), in 50.0 g water. Calculate the freezing point of this 50-50 mixture. Would this antifreeze protect a car in Chicago on a day when the temperature gets as low as –10° F? (-10 °F = -23.3° C)

Electrolytes and Colligative Properties  Colligative properties depend on the # of particles present in solution.  Because ionic solutes dissociate into ions, they have a greater effect on freezing pt and boiling pt than molecular solids of the same molal conc. o For example, the freezing pt of water is lowered by 1.86°C with the addition of any

molecular solute at a concentration of 1 m, such as C6H12O6, or any other covalent compound o However, a 1 m NaCl solution contains 2 molal conc. of IONS. Thus, the freezing pt depression for NaCl is 3.72°C…double that of a molecular solute. IB Chemistry Notes Chapter 11: Properties of Solutions

The relationships are given by the following equations:

Tf/b = f.p. depression/elevation of b.p. m = molality of solute

kf/b = b.p. elevation/f.p depression constant i = # particles formed from the dissociation of each formula unit of the solute (van’t Hoff factor)

Ex: What is the freezing pt of: a) a 1.15 m sodium chloride solution?

Ex: What is the freezing pt of: b) a 1.15 m calcium chloride solution?

Ex: What is the freezing pt of: c) a 1.15 m calcium phosphate solution?

Osmotic Pressure: Experiments show that dependence of the osmotic pressure on solution concentration is expressed by the eqn:

Where,  = osmotic pressure (atm) M = molarity (mol/L) R = gas law constant = 0.08206 T = temp (K)