Properties of Solutions

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• Properties of Solutions
• Definitions and Examples
• A solution is a homogeneous mixture.
• Solutions are made up of at least two components.
• One component is the solvent - usually the
  component in greatest quantity.
• The other component(s) is(are) the solutes.
• Examples of solutions

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• Properties of Solutions
• Methods of expressing solution concentration
• Weight percent is the mass of solute per mass of
  solution times 100, where the masses are in the
  same units.

• Parts per million ppm is the mass of solute per
  mass of solution times 106

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• Properties of Solutions
• Methods of expressing solution concentration
• Molarity is the moles of solute per liter of
  solution abbreviation M
• Molarity is temperature dependent because density
  changes with temperature and the volume of
  solution containing a given number of moles
  solute will vary.

• Molality is the moles solute per kilogram
  solvent abbreviation m
• Molality does not vary with temperature because
  masses are not temperature dependant.

• Mole fraction is the moles solute per total moles
  of all components in a solution abbreviation
  X.
• Mole fraction is temperature independent.

Numerous examples dealing with calculations using
these concentrations are given in the text and
should be examined carefully.
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• Properties of Solutions
• Example 45.0 g of camphor (C10H16O) is dissolved
  in 425 mL ethanol (C2H5OH). The density of
  ethanol is 0.785 g/mL
• Calculate the molality of camphor Calculate the
  mole fraction of camphor
• Calculate the weight of camphor

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• Properties of Solutions
• Example Concentrated aqueous ammonia has a
  molarity of 14.8 and a density of 0.90 g/cm3
• What is the molality of the solution?
• What is the mole fraction NH3? What is the
  wt. NH3?

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• Properties of Solutions
• Saturated Solutions and Solubility
• A saturated solution can contain no more solute
  than already dissolved in a solution.
• Adding more solute to a saturated solution adds
  to the mass of undissolved phase.
• The solubility of a solute is the concentration
  of solute in a saturated solution.
• The solubility is often expressed in g/100 mL of
  solvent or molarity.
• When a saturated solution exists, there is a
  dynamic equilibrium between dissolved and
  undissolved solute.
• solute solvent solution
• Solutions can contain less than the amount of
  solute required to form a saturated solution,
  an unsaturated solution.
• A supersaturated solution contains more solute
  than expected based on the equilibrium
  conditions of a saturated solution.
• This is a metastable equilibrium which results
  from the lack of a nucleation site on which the
  crystal of solid solute can grow.

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• Properties of Solutions
• The solution process
• To form a solution requires overcoming
  intermolecular interactions between solute
  molecules and solvent molecules so solute
  molecules and solvent molecules can be
  homogeneously distributed throughout the bulk of
  the solution.
• The kinds of interactions to be overcome depend
  on the nature of the solutes and solvents.
• Dissolving NaCl in H2O involves breaking ion-ion
  interactions in NaCl and the hydrogen bonds in
  H2O.
• Interactions between Na and H2O and Cl- and
  water will be established to stablize the
  solution solvation occurs - hydration if H2O is
  solvent.

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Properties of Solutions Enthalpy changes
accompanying solution formation Exothermic heat
of solution Endothermic heat of solution

• DHsolnDH1DH2DH3
• DHsoln(NaOH)-44.48 kJ/mol
• DHsoln(NH4NO3)26.4 kJ/mol

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• Properties of Solutions
• Enthalpy changes accompanying solution formation
• DH3 represents the enthalpy of interaction
  between solute particles and solvent particles.
• Should DH3 be small in magnitude, the formation
  of the solution will be too endothermic to
  form.
• This occurs when non-polar solvents interact with
  ionic or polar solutes.
• Solution formation and increase in disorder
• Generally, processes that are exothermic are
  spontaneous, but not all spontaneous processes
  are exothermic.
• The formation of an aqueous solution of NH4NO3 is
  endothermic.
• One factor that influences the spontaneity of
  processes besides energy change is the degree of
  randomness created as the result of the process.
• Forming an aqueous solution of NH4NO3 requires
  breaking up the order associated with the
  hydrogen bonded water molecules and the highly
  ordered lattice of NH4NO3.
• The solution is more disordered than the unmixed
  components of the solution.
• If the ordered pure solution components have
  strong interparticle attractive interactions it
  is possible that the disorder created by forming
  a solution will not compensate for these
  energetic interactions.

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• Properties of Solutions
• Factors affecting solubility
• Solute-solvent interactions
• The solubility of gases in liquid solvents
  depends on the strength of London dispersion
  forces between gaseous solute and liquid solvent.
• Solubilities of gases in H2O at 20 oC and 1 atm
  gas pressure

• Polar liquids dissolve readily in polar solvents
  because of dipole-dipole attractive interactions
  between solute and solvent.
• Many pairs of liquids are mutually soluble in all
  ratios they are miscible.
• Acetone is miscible with
  water, whereas 2-pentanone
• dissolves to
  the extent of 4.7 g/100 g water at 20 oC and is
  therefore not miscible with water.

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• Properties of Solutions
• Factors affecting solubility
• Solute-solvent interactions
• Hydrogen bonding is important in determining the
  solubility of solutes in water.
• The short chain alcohols are miscible with water
  but longer chain alcohols are not. For short
  chain alcohols, the hydrogen bonding between
  alcohol molecules is similar in strength to
  that in water
• Nonpolar liquids dissolve readily in nonpolar
  solvents but not polar solvents.
• The dipole-dipole attractive forces in the polar
  substance cannot be overcome by interactions
  between the nonpolar solvent and the polar
  solute.

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• Properties of Solutions
• Factors affecting solubility
• Pressure Effects and the solubility of gases
• Henrys law the solubility of a gas in a liquid
  is proportional to the pressure of the gas above
  the liquid.
• SgkPg where Sg is the solubility of
  the gas in the liquid
• Pg is the pressure of the gas above
  the liquid
• k is the Henrys law constant that
  depends on the liquid-gas pair and
  temperature.
• Example if the pressure of CO2 above the water
  in carbonated water is 4.0 atm and the Henrys
  law constant for CO2 in water at 25 oC is 3.1 x
  10-3 mol/L-atm, what is the concentration of
  CO2 in water under these conditions?

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• Properties of Solutions
• Factors affecting solubility
• Temperature Effects
• Gases decrease their solubilities with increasing
  temperature
• The solution process for gases is exothermic
• Gas liquid solvent saturated
  solution heat energy
• DHsolutionlt0
• If the temperature is increased, heat energy is
  added to the solution
• The result is to shift the equilibrium to the
  left removing gas from the solution and reducing
  the gas concentration in the solution
• LeChâteliers principle states if a stress is
  applied to a system at equilibrium, the system
  will respond to relieve the stress.
• Most solids increase their solubilities in
  liquids as temperature increases because the
  solution phenomenon is endothermic. See Fig.
  14.9, p. 655.
• solid solvent heat solution

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• Properties of Solutions
• Colligative Properties are properties of
  solutions that depend on the number of
  particles of solvent or solute present in a
  solution. Colligative properties thus depend on
  the concentration of solvent or solute present
  in a solution.
• The vapor pressure of a solution containing a
  nonvolatile solute is less than the vapor
  pressure of the pure solvent,
• The vapor pressure is a measure of the escaping
  tendency of liquid molecules.
• The fraction of solvent molecules on the surface
  of a solution is lowered by the dissolved solute
  molecules. The escaping tendency is proportional
  to the mole fraction of volatile molecules at
  the surface of the liquid.
• Raoults law gives the relationship between the
  vapor pressure of a volatile liquid containing
  a nonvolatile solute and the mole fraction of
  volatile liquid.
• PAXAPAo where PA is the vapor pressure of the
  solution
• XA is the mole fraction of solvent
• Pao is the vapor pressure of pure
  solvent

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• Properties of Solutions
• Colligative Properties
• Vapor Pressure
• Example The vapor pressure of pure water at 110
  oC is 1070 torr. A solution of ethylene glycol
  and water has a vapor pressure of 1.00 atm at 110
  oC. What is the mole fraction of ethylene glycol
  in the solution?

• Vapor pressure of a mixture of volatile liquids
• Raoults law and Daltons law of partial
  pressures can be combined. For 2 volatile
  liquids A and B
• PAXAPAo and PBXBPBo (Note XA XB 1)
• PtotalPAPB XAPAo XBPBo (conditions for an
  ideal solution)
• Example Pbenzeneo75 torr and Ptolueneo22 torr
  at 20 oC
• if X0.500, Pbenzene(0.500)(75 torr)37.5
  torr and Ptol(0.500)(22)11.0 torr
• Ptotal(37.511.0)torr48.5 torr
  Xbenz(vapor)37.5/48.5 0.773