Solutions and their
properties
Part I:
What’s a solution: The process of dissolving
Factors affecting solubility
Ways of expressing concentration
Part II:
Colligative properties
Learning objectives
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Describe a solution
Describe factors influencing solubility
Predict solubility based on molecular properties
Describe effect of temperature on solubility
Predict solubility of gases using Henry’s law
Perform concentration calculations using
molarity, molality, mole fraction and percent by
mass and volume
Mixtures are a fact of life
• We learn about matter by studying pure
substances
• In nature, most things are mixtures –
nothing is “pure”:
– Air
– Water
– Soil
– Living systems
Solution or colloid?
• Both contain a solid dispersed in a liquid
• In a solution, the particle size is on the molecular
scale (< 1 nm)
• In a colloid, the particles are larger (2 – 500 nm)
• Colloids scatter light – milk, fog
Why do solutions form?
• Things tend to get mixed up (entropy)
• Without any interactions between molecules, gases always mix
• Intermolecular forces complicate matters
– May improve mixing if forces in mixture are stronger
– May oppose mixing if forces in pure substances are stroner
• Three intermolecular forces:
– Solvent – solvent
– Solute – solute
– Solvent – solute
• Like dissolves like:
– Polar solvents dissolve polar solutes
– Nonpolar solvents dissolve nonpolar solutes
Review of intermolecular forces
Name of force
Origin
Strength
Ion-dipole
Between ions and molecules
Quite strong (10 – 50 kJ/mol)
Dipole-dipole
Between permanent dipoles
Weak (3 – 4 kJ/mol)
Hydrogen bonds
Polar bonds with H and (O,N)
Quite strong (10 – 40 kJ/mol)
London dispersion forces
Fluctuating dipoles in non-polar
bonds
Weak (1 – 10 kJ/mol)
Ion - dipole
• Characteristic of interactions in solutions
of ionic compounds in polar solvents
– Negative ion with the positive dipole end
– Positive ion with the negative dipole end
Dipole - dipole
• Important attractive force in polar
substances
• Strength of the order of 3 – 4 kJ/mol
(compared with 200 – 400 kJ/mol for
covalent bonds)
Dipole force manifested in
boiling points:
• Nonpolar substances have much lower boiling
points
– Acetone (polar) 56ºC butane (nonpolar) -0.5ºC
• Boiling point increases with dipole strength
London calling
• Even molecules with no net dipole moment attract each other.
• Electrons are not static but mobile:
– Fluctuation creates dipole in one molecule which induces dipole in
another molecule
• Effect increases with atomic number – as atom becomes more
polarizable
– Boiling increases with molar mass
• For small molecules, dispersion forces are weaker than other intermolecular forces. For large molecules this is not true. Large
Solute – solvent interactions
• The stronger the interactions between
solute and solvent, the greater the
solubility
– Dispersion forces, dipolar interactions and
liquids
– Ion-dipole interactions and ionic compounds
in water
Dissolving an ionic compound
• Attractive forces between solute and
solvent compensate attractive forces
between solute – solute and solvent solvent
Solution on the atomic scale
• The ions held tightly in the crystal lattice
become solvated by water molecules
• Ion-ion interactions are replaced by iondipole interactions in the solution
Profits and losses
• Energy costs
– Solvent – solvent interactions
– Solute – solute interactions (lattice energy)
• High lattice energy → low solubility
• Energy gain
– Solvent – solute interactions
• Small, more highly charged ions have stronger interactions
Solutions, disorder and entropy
• Substances dissolve with either gain or loss of
enthalpy
• Dissolving usually results in an increase in
entropy
• Exothermic processes tend to occur
spontaneously
• Processes that result in increased entropy tend
to occur spontaneously
Strong interactions complicate
predictions about solubility
• Weak interactions:
– Gases mix completely over all compositions
• Moderate interactions:
– Liquids exhibit limits on range of miscibility
• Strong interactions:
– Solubility of ionic compounds varies
– Some ionic compounds are completely
insoluble
Super-saturate me
• Saturated solution is in equilibrium with
undissolved solute and cannot contain any
more solute
Supersaturation: a state of the solution
where it contains more solute than allowed
by saturation. Non-equilibrium condition
• Crystallization and rainfall depend on
supersaturation
Temperature and solubility
• Solubility is amount of solute present in solution
at saturation point
• Solubility varies widely from compound to
compound
• Solubility varies with T
– Most substances increase with T (KNO3) although the
dependence can vary
– Some are unaffected by T (NaCl)
– A few decrease with T (Na2SO4)
• Results can be explained by sign of ΔH solution
and Le Chatelier’s principle (later)
Gases always decrease
solubility with temperature
• Solubility of gases always decreases as
temperature increases
Solubility and pressure
• For solutions of solids in liquids and liquids in
liquids, pressure has almost no effect.
• For gases in liquids:
• Henry’s Law operates
Solubility = k P
• Solubility of gas depends on partial pressure
above solution
• k is function of gas for given solvent
Henry’s law and equilibrium
• Amount of gas in solution is in equilibrium with
gas above solution
• When P is increased, density of molecules
above solution increases
• More molecules enter solution to restore
equilibrium
Concentration: ways of
expressing it
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Molarity
Mole fraction
Mass percent/volume percent
Molality
Molarity (M)
Molarity (M) =
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Moles solute (mol)
Vol solution (L)
Stoichiometry calculations are easy
Amounts of solution required are volumetric
Concentration varies with T
Amount of solvent requires knowledge of
density
Mole Fraction (X)
Mole fraction (X) =
Moles solute (mol)
Moles solution (mol)
• Dimensionless
• Independent of temperature
• Used for gas mixtures and solid solutions
Mass percent (mass %)
Mass percent =
Mass solute (g)
x 100%
Mass solution (g)
• For very dilute solutions:
Parts per million (ppm) = Mass solute (g)
Mass solution (g)
x 10
• Independent of temperature
• Need to work with masses rather than volumes
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Molality (m)
Molality (m) =
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Moles solute (mol)
Mass solvent (kg)
Used in calculations of colligative properties
Independent of temperature
Need to know mass of solution
Need to know density to convert to molarity