Download Metathesis Problems (and Some Solutions) Identified Through

Chapter 12: Solutions
Chemistry 1062: Principles of Chemistry II
Andy Aspaas, Instructor
Solutions
• Solution: homogeneous mixture of two or more
substances (atoms, molecules, or ions)
• Can exist in any state of matter
Solute
Gaseous
solution
Component in
smaller amount
Liquid component in
Liquid solution smaller amount, or
solid or gas
Component in
Solid solution
smaller amount
Solvent
Component in
larger amount
Liquid component
in larger amount
Component in
larger amount
States of solutions
• Miscible fluids: fluids that dissolve with each other in all
proportions
• All nonreactive gases are generally miscible
– Air is a gaseous solution (N2, O2, CO2, etc.)
• Liquid solutions: dissolving a solid, liquid, or gas into a liquid
– Ethanol and water are miscible, and when mixed make a
liquid solution
– Brine is solid sodium chloride dissolved in water
• Solid solutions are called alloys when metals are mixed
– Dental fillings, brass, steel, etc.
Solubility and saturation
• If 40.0 g of NaCl were stirred in 100 mL of 20 °C
water, most of the salt would dissolve but some
would remain on the bottom
– Ions dissolve by leaving the surface of the crystal
and entering the liquid solution
– Some crystals may re-deposit on the crystal
• Equilibrium is reached at point where particles
dissolve at same rate as they return to the crystal
– The solution has become saturated
Solubility and saturation
• The point at which the solution becomes saturated
can be expressed as solubility (at a given
temperature)
– The solubility of NaCl in 20 °C water is 36.0 g
NaCl / 100 mL H2O
• The solution is unsaturated when not enough solid
has been added for the equilibrium to be reached
– Unsaturated solutions can support addition of
more solid to be dissolved
Supersaturation
• Supersaturated solution: solution which contains more
dissolved substance than a saturated solution does
• Sodium acetate and many other ionic compounds more
soluble in hot water than in cold water
– If a saturated solution is prepared at high temperature,
and then the temperature is slowly lowered, the solution
may become supersaturated
– Introduction of any solid to a supersaturated solution will
cause the whole solution to quickly crystallize
Factors behind solubility
• “Like dissolves like” - similar substances tend to dissolve in
each other
• The natural tendency of substances to mix through the
random motion of their particles can be overridden if one
component has strong intermolecular forces, and the other
does not
– Oil and water: water’s intermolecular forces are
maintained if nonpolar oil molecules do not interrupt the
water molecules
– Alcohol and water: similar hydrogen bonds can be
formed, so 3-carbon alcohols and smaller are miscible in
water (larger alcohols are too dissimilar to water)
Solubility of ionic compounds
• Partial charges of water molecules orient
themselves towards oppositely charged ions in
solutions (ion-dipole force)
• Hydration: water molecules surrounding ions - this
favors dissolving of ionic solids in water
• Lattice energy: energy holding together ions in a
crystal lattice - this works against dissolving
– Lattice energy increases with ion charge
– Lattice energy decreases with ionic radius
– Hydration energy increases with ionic radius
Temperature effects on solubility
• Most gases are less soluble in water at higher temperatures
(bubbles that appear when heating water)
• Most ionic solids are more soluble in water at higher
temperatures
– Some have very little change, like NaCl
– Some are less soluble in higher temperatures
• Heat of solution: heat absorbed or released when a solid is
dissolved
– Depends on combination of lattice energy and hydration
energy
– Chemical hot packs and cold packs take advantage of this
Pressure effects on solubility
• All gases become more soluble in a liquid at a given
temperature when the partial pressure fo the gas
over the solution is increased
– Le Chatelier’s principle: if an equilibrium is
disturbed by a temperature, pressure or
concentration change, the equilibrium will shift to
compensate for the change
– Increasing the partial pressure of CO2(g) over
water will cause more CO2 to dissolve
CO2(g)
CO2(aq)
• This equilibrium will shift to the right (more CO2 will
dissolve) to compensate for the pressure increase
Henry’s law
• Henry’s law: solubility of a gas is directly
proportional to the partial pressure of the gas above
the solution
S = k HP
where S is solubility (mass of solute per unit
volume of solvent),
kH is Henry’s law constant (for gas and liquid at a
given temperature),
P is partial pressure of the gas
Colligative properties and concentration
• Colligitave properties depend on concentration of solute
molecules or ions in solution, but not the chemical
identity of the solute
• Molarity, M = (moles solute)/(liters solution)
• Mass percentage: [(mass solute)/(mass solution)]*100%
• Molality, m = (moles solute)/(kilograms solvent)
• Mole fraction, XA = (moles A)/(total moles solution)
Vapor pressure of a solution
• Vapor pressure lowering: colligative property
– Vapor pressure of pure solvent minus vapor
pressure of solution
• Raoult’s law: PA = P°AXA
– if solute must be nonvolatile nonelectrolyte
– PA = partial pressure of solvent
– P°A = vapor pressure of pure solvent
– XA = mole fraction of solvent in solution
• Or, ∆P = P°AXB where XB = mole fraction of solute
Distillation
• An ideal solution follows Raoult’s law for all mole
fractions 0–1
– When two components are chemically similar,
their intermolecular forces are similar
• Total vapor pressure, P = P°AXA + P°BXB
• Vapor over an ideal solution is richer in the more
volatile component
• Fractional distillation condenses and re-vaporizes
the solution many times to exploit this
Boiling point elevation
• Addition of a nonvolatile solute reduces the solvent’s vapor
pressure
• Normal boiling point: temperature at which vapor pressure =
1 atm
• So, addition of solute requires a higher temperature in order
for vapor pressure to reach 1 atm
• Boiling point elevation, ∆Tb = Kbcm
Kb = bp elevation constant, depends only on solvent
cm = molal concentration of solution
Freezing point depression
• Nonvolatile solutes will lower the freezing point of a
solvent in a similar way to bp elevation
• ∆Tf = Kfcm
• Antifreeze both lowers the freezing point and raises
the boiling point of the coolant
• Molecular weight of a solute can be determined by
measuring its freezing point depression
Osmosis
• Semipermeable membrane: allows solvent
molecules to pass but large solute molecules cannot
• Osmosis: flow of solvent through a semipermeable
membrane to equalize solute concentrations on both
sides of the membrane
• π = MRT (M = molar conc., R = gas constant, T =
absolute temperature)
• Reverse osmosis: apply greater pressure to more
concentrated solution and force pure solvent
through membrane
Colligative properties of ionic solutions
• Effective concentration of ionic solutions is greater than
molecular solutions even at the same molarity or molality
– Ionic compounds dissociate into individual ions
• i = number of ions resulting from solvation of one formula unit
• Multiply i in any colligative formula if the solute is ionic
∆Tf = iKfcm
∆P = iP°AXB
∆Tb = iKbcm
π = iMRT
• i is only accurate in dilute solutions
Colloid formation
• Colloid: dispersion of particles throughout another
substance or solution
– Differs from a solution in that its dispersed
particles are more than one molecule in size (but
still too small to see with the naked eye)
• Tyndall effect: colloids scatter light, while solutions
do not
Types of colloids
• Aerosol: liquid or solid particles dispersed throughout a gas
– Fog, smoke
• Emulsion: liquid droplets dispersed throughout another liquid
– Particles of butterfat dispersed through homogenized milk
• Sol: solid particles dispersed in a liquid
• Hydrophilic colloid: when there is a strong attraction between
particles and water
– Gelatin
• Hydrophobic colloid: no attraction between particles and
water
Coagulation and association
• Coagulation: particles of a colloid are made to aggregate and
separate from solvent
– Milk curdles when its colloidal particles no longer have the
same charge (they’re no longer repelled from each other)
• Association colloid: formed when colloidal particles have both
hydrophobic and hydrophilic portions
– Soaps have long hydrocarbon chain (hydrophobic) and
charged functionality (hydrophilic)
– The hydrophobic portions gather inwards to form spheres
with a hydrophilic outside (micelles)