Chapter 13 - the Properties of Solutions and Colloids

What is the difference between a solution and colloid?

• solution - homogeneous mixture, individual single phase

• colloid - heterogeneous mixture, multiple phases

Define “solubility”

the maximum amount of solute that dissolves in a fixed quantity of solvent

describe the phrase “like dissolves like”

substances w/ similar intermolecular forces dissolve in each other

describe the intermolecular force: ion-dipole

• polar

• between charged ion and molecular dipole

• 40-600

describe the intermolecular force: hydrogen-bond

• polar bond w/ a hydrogen

• high electronegativity of N, O, F

• 10-40

desribe the intermolecular force: dipole- dipole

dipoles attracted to each other

• 5-25

describe the intermolecular force: dipole-induced dipole

• polar

• permanent dipole & temporary dipole

• 2-10

describe the intermolecular force: ion-induced dipole

• polar

• charged ion & temporary dipole

• 3-15

describe the intermolecular force: dispersion

nonpolar & nonpolar

describe the enthalpy of hydration

• the energy released when an ion is surrounded by water molecules

• depends on the charge density!

(1) describe the solute-solute energy change in solution formation

• endothermic, takes in energy to break lattice

• solute (aggregated) + heat → solute (separated)

• enthalpy of hydration solute > 0

(2) describe the solvent-solvent energy change in solution formation

• endothermic

• solvent (aggregated) + heat → solvent (separated)

• enthalpy of hydration solvent > 0

(3) describe the solute-solvent energy change in solution formation

• exothermic; extra energy in formation

• solute (separated) + solvent (separated) → solution + heat

• enthalpy of hydration mix< 0

HEAT OF SOLUTION:

what is the formula for the enthalpy of hydration solution?

enthalpy hyd. solution = (1) enthalpy hyd. solute !POSITIVE! + (2) enthalpy hyd. solvent !POSITIVE! + (3) hyd. mix !NEGATIVE!

→ endothermic (+) when heat enthalpy of solution > 0

→ exothermic (-) when heat enthalpy of solution < 0

Define/describe solvation

a solute particle surrounded by solvent particles

• the hydration of an ion is always exothermic

• change in heat enthalpy of solvation = heat enthalpy solvent + heat enthalpy mix

  • (for water) → change in heat enthalpy of solution = heat enthalpy of solute + heat enthalpy hydration

Describe enthalpy hydration & its trend on the periodic table

• depends on charge density

• charge/vol.

• charge density & enthalpy hydr. decrease down a group

• depends across period (metals = CD increases; nonmetals = CD decreases)

what would an increase in charge density do to the enthalpy of hydration?

more exothermic (negative)!

what would a decrease in charge density do to the enthalpy of hydration?

more endothermic (positive)!

the higher the charge density & smaller its radius…

STRONGER ATTRACTION!

• the closer the ion is to the oppositely charged pole

ppm?

10^6

• the # of units of mass of solute per million units of solution

ppb?

10^9

• the # of units of mass of solute per billion units of solution

what are the factors that affect solubility?

• temperature: @ higher temp.. solids more & gases less

• pressure: @ higher pressure.. gases more

what is an unsaturated solution?

contains less than the equilibrium concentration of solute WATER IS HUNGRY!!

what is a saturated solution?

equilibrium: maximum amount of dissolved solute in the presence of undissolved solute

what is a supersatuated solution?

contains more than the equilibrium concentration of solute, unstable!

henry’s law?

• Sgas = kH x Pgas

  • solubility of a gas is directly proportional to its partial pressure

molarity f(M):

mol/vol.

molality (m):

mol (solute)/mass,kg (solvent)

parts by mass:

mass (solute)/mass (solution)

parts by volume:

vol.(solute)/vol.(solvent)

mole fraction (X):

mol(solute)/(mol, solute + mol, solvent)

conversion for ppm

mass(solute)/mass(solvent) x 10^6

conversion for ppb

mass(solute)/mass(solvent) x 10^9

what are colligative properties?

properties that depend on the NUMBER of solute particles

difference between electrolytes and non-electrolytes?

• electrolytes → dissociate into ions in water

• non-electrolytes → don’t dissociate

what are the colligative properties of electrolyte solutions?

• vapor pressure lowering ∆𝑃 = 𝑖 𝑋𝑠𝑜𝑙𝑢𝑡𝑒 × 𝑃𝑠𝑜𝑙𝑣𝑒𝑛𝑡

• freezing point depression ∆Tf = i(Kbm)

• boiling point elevation ∆Tb = i(Kbm)

• osmotic pressure π = MRT

i= Van’t Hoff’s factor; the # of particles a formula breaks up into

Describe the colligative property: Vapor pressure lowering

• ∆𝑃 = 𝑖 𝑋𝑠𝑜𝑙𝑢𝑡𝑒 × 𝑃𝑠𝑜𝑙𝑣𝑒𝑛𝑡

• vapor pressure of solution < vapor pressure of pure solvent

• Raoult’s law → the vapor pressure of the solvent ABOVE the solution is proportional to the mole fraction of the solvent present: P_solvent = X_solvent * P°_solvent

  • ONLY DILUTE SOLUTIONS

• the vapor pressure LOWERING is proportional to the mole fraction of the solute present: ∆P_solvent = X_solvent * P°_solvent

Describe the colligative property: Freezing point depression

• ∆Tf = i(Kbm)

• solution freezes @ a lower temperature than the pure solvent

• directly proportional to the MOLALITY of the solution

• ex: industry salts (..NaCl)

Describe the colligative property: Boiling point elevation

• ∆Tb = i(Kbm)

• solution boils @ a higher temperature than the pure solvent (! result of vapor pressure lowering !)

• directly proportional to the MOLALITY of the solution

Describe the colligative property: Osmotic pressure

• the pressure applied to prevent the net flow of solvent

• π = MRT

  • M = molarity

  • R= 0.0821 (atm x L)/(mol x K)

  • T = kelvin temp.

osmosis → net flow of solvent to the more concentration solution

the higher the ***general solubility (molarity/molality, concentration)*** the..

• ___ the osmotic pressure

• ____ the boiling point

• ____ the freezing point

• ____ the vapor pressure

• higher

• higher

• lower

• lower

how many particles of this formula break into?

NH4Br

i = 2

which one has the highest freezing point & vapor pressure?

1. 0.1m NaNO3

2. 0.1m glucose

3. 0.1m CaCl2

2. 0.1m glucose

whiich one has the highest boiling point & osmotic pressure?

1. .1m NaNO3

2. 0.1m glucose

3. 0.1m CaCl2

3. 0.1m CaCl2

which solution has the lowest freezing point?

1. 0.1m CaCl2

2. 0.1m (NH4)2SO4

3. 0.1m Ca3(PO4)2

4. 0.1m sucrose

3. Ca3(PO4)2