Chapter 13 Properties of Solutions

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Chapter 13Properties of Solutions
CHEMISTRY The Central Science 9th Edition
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The Solution

• A solution is a homogeneous mixture of solute
  (present in smallest amount) and solvent (present
  in largest amount).
• Solution may be gas, liquid, or solids
• Each substance present is a component of the
  solution.
• Solvent is the substances used to dissolve the
  solute.
• In the process of making solutions with condensed
  phases, intermolecular forces become rearranged.
• Intermolecular forces hold the solute particle
  and the solvent that surrounds it together.
• Solutions form when the attractive forces between
  solute and solvent can overcome the attractive
  forces with in the solute or solvent particles.

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Types of Solutions
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Formation of a Solution

• Consider NaCl (solute) dissolving in water
  (solvent)
• The water H-bonds have to be interrupted,
• NaCl dissociates into Na and Cl-,
• Ion-dipole forces form Na ?-OH2 and Cl-
  ?H2O.
• Such interaction between solute and solvent are
  called solvation.
• If water is the solvent, we say the ions are
  hydrated.

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Energy Changes in Solution Formation

• There are three energy steps in forming a
  solution
• Separation of solute molecules (?H1),
• Separation of solvent molecules (?H2),
  andformation of solute-solvent interactions
  (?H3).
• We define the enthalpy change in the solution
  process as
• ?Hsoln ?H1 ?H2 ?H3.
• ?Hsoln can either be positive or negative
  depending on the intermolecular forces.

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Enthalpic Contributions

• Breaking attractive intermolecular forces is
  always endothermic.
• Forming attractive intermolecular forces is
  always exothermic.

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Enthalpy Properties of a Solution

• To determine whether ?Hsoln is positive or
  negative, we consider the strengths of all
  solute-solute and solute-solvent interactions
• ?H1 and ?H2 are both positive.
• ?H3 is always negative.
• Solutions will not form under certain conditions.
• like substances dissolve like substance.

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The Solution Process

• Rule polar solvents dissolve polar solutes.
  Non-polar solvents dissolve non-polar solutes.
  Why?
• If ?Hsoln is too endothermic a solution will not
  form.
• NaCl in gasoline the ion-dipole forces are weak
  because gasoline is non-polar. Therefore, the
  ion-dipole forces do not compensate for the
  separation of ions.
• Water in octane water has strong H-bonds. There
  are no attractive forces between water and octane
  to compensate for the H-bonds.

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The Solution Process

• Solution can either be endothermic or exothermic
• For example
• NaOH added to water has ?Hsoln -44.48 kJ/mol.
• NH4NO3 added to water has ?Hsoln 26.4 kJ/mol.
• Endothermic meaning heat has been gained by the
  system.
• This is represented by a ?Hsoln gt 0.
• Exothermic meaning heat has been released (or
  lost) from the system.
• This is represented by a ?Hsoln lt 0.

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The Spontaneity of Solutions (Energy Decrease)

• A spontaneous process occurs without outside
  intervention.
• When energy of the system decreases (e.g.
  dropping a book and allowing it to fall to a
  lower potential energy), the process is
  spontaneous.
• Some spontaneous processes do not involve the
  system moving to a lower energy state (e.g. an
  endothermic reaction).

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The Spontaneity of Solutions (Increasing Disorder)

• If the process leads to a greater state of
  disorder, then the process is spontaneous.
• Example a mixture of CCl4 and C6H14 is less
  ordered than the two separate liquids.
  Therefore, they spontaneously mix even though
  ?Hsoln is very close to zero.
• There are solutions that form by physical
  processes and those by chemical processes.

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Ways of Expressing Concentration

• Mass Percentage, ppm, and ppb
• All methods involve quantifying amount of solute
  per amount of solvent (or solution).
• Generally amounts or measures are masses, moles
  or liters.
• Qualitatively solutions are dilute or
  concentrated.
• Definitions

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Ways of Expressing Concentration

• Mass Percentage, ppm, and ppb
• Parts per million (ppm) can be expressed as 1 mg
  of solute per kilogram of solution.
• If the density of the solution is 1g/mL, then 1
  ppm 1 mg solute per liter of solution.
• Parts per billion (ppb) are 1 ?g of solute per
  kilogram of solution.

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Ways of Expressing Concentration

• Mass Percentage, ppm, and ppb
• Mole Fraction, Molarity, and Molality
• Recall mass can be converted to moles using the
  molar mass.

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Ways of Expressing Concentration

• Mole Fraction, Molarity, and Molality
• We define
• Converting between molarity (M) and molality (m)
  requires density.

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Class Guided Practice Problems

• 1) A solution containing equal masses of
  glycerol, C3H8O3, and water has a density of
  1.10g/mL. Calculate the (a) Molarity, (b) Mole
  Fraction of glycerol, (c) Molality of the
  solution.
• 2) Calculate the percent by mass of CaCl2 in a
  solution containing 5.2g CaCl2 in 450g of water.
• 3) Calculate the ppm, by mass, of CaCl2 in a
  solution containing 0.149 moles of CaCl2 in 443g
  of water.

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Class Guided Practice Problems

• 1) Calculate the molarity when 0.020 moles of
  glycerol, C3H8O3, is dissolved in 50 g of water
  at room temperature. Water has a density of
  1.00g/mL.

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Class Practice Problem

• 1)A 0.500 L hydrochloric acid (HCl) solution,
  which has a density of 1.10 g/mL, contains 36
  HCl by mass. Calculate the mole fraction of HCl.
• 2) A solution containing equal masses of NaCl and
  water has a density of 1.10g/mL. Calculate the
  molality of the solution.
• 3) Calculate the ppm, by mass, of CaCl2 in a
  solution containing 0.149 moles of CaCl2 in 443g
  of water.

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Saturated Solutions and Solubility

• Dissolve solute solvent ? solution.
• Crystallization solution ? solute solvent.
• Saturation crystallization and dissolution are
  in equilibrium.
• Solubility amount of solute required to form a
  saturated solution.
• Example, only 35.7g of NaCl will dissolve at 0 oC
  in 100 mL of H2O.
• Dissolving less solute than needed to saturate is
  called an unsaturated solution.
• Supersaturated a solution formed when more
  solute is dissolved than in a saturated solution.

Dynamic equilibrium
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Factors Affecting Solubility

• Solute-Solvent Interaction
• Polar liquids tend to dissolve in polar solvents.
• Miscible liquids mix in any proportions.
• Immiscible liquids do not mix.
• Intermolecular forces are important water and
  ethanol are miscible because the broken hydrogen
  bonds in both pure liquids are re-established in
  the mixture.
• The number of carbon atoms in a chain affect
  solubility the more C atoms the less soluble in
  water.

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Factors Affecting Solubility Cont.

• Solute-Solvent Interaction
• The number of -OH groups within a molecule
  increases solubility in water.
• Remember, as a generalization like dissolves
  like.
• The more polar bonds contained in the molecule,
  the better it dissolves in a polar solvent.
• The less polar the molecule the less it dissolves
  in a polar solvent and the better is dissolves in
  a non-polar solvent.
• The magnitude of ?H3 must be comparable in
  magnitude to ?H1 ?H2 before the solute will
  dissolve in the solvent.

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Factors Affecting Solubility Cont.
Solute-Solvent Interaction
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Fat and Water Soluble Vitamins
Solute-Solvent Interaction
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Pressure Effects on Solubility

• Solubility of a gas in a liquid is directly
  related to the pressure of the gas.

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Solubility of a Gas

• Pressure Effects
• The higher the pressure, the more molecules of
  gas are close to the solvent and the greater the
  chance of a gas molecule striking the surface and
  entering the solution.
• Therefore, the higher the pressure, the greater
  the solubility.
• The lower the pressure, the fewer molecules of
  gas are close to the solvent and the lower the
  solubility.
• If Sg is the solubility of a gas, k is a
  constant, and Pg is the partial pressure of a
  gas, then Henrys Law gives

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Compression of CO2

• Pressure Effects
• Carbonated beverages are bottled with a partial
  pressure of CO2 gt 1 atm.
• As the bottle is opened, the partial pressure of
  CO2 decreases and the solubility of CO2
  decreases.
• Therefore, bubbles of CO2 escape from solution.

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Temperature Effects on Solubility

• Experience tells us that sugar dissolves better
  in warm water than cold.
• As temperature increases, solubility of solids
  generally increases.
• Sometimes, solubility decreases as temperature
  increases (e.g. Ce2(SO4)3, Cerous Sulfate).
• See Table 13.15

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Temperature Solubility Table 13.15
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Colligative Properties

• Colligative properties depend on quantity of
  solute molecules. (E.g. freezing point
  depression and melting point elevation.)
• Lowering Vapor Pressure
• Non-volatile solutes reduce the ability of the
  surface solvent molecules to escape the liquid.
• Therefore, vapor pressure is lowered.
• The amount of vapor pressure lowering depends on
  the amount of solute present.

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Colligative Properties
Lowering Vapor Pressure
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Raoults Law

• Raoults Law PA is the vapor pressure with
  solute, PA? is the vapor pressure without
  solvent, and ?A is the mole fraction of A, then
• Recall Daltons Law
• Raoults law breaks down when the solvent-solvent
  and solute-solute intermolecular forces are
  greater than solute-solvent intermolecular
  forces.
• An ideal gas is a gas that obeys Raoults law

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Phase Diagram Pure Solvent and Nonvolatile Solute
Solution

• The triple point - critical point curve is
  lowered.

• ?Tb is directly related to the number of solute
  molecules present.

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Boiling-Point Elevation

• At the normal boiling point of pure liquid, the
  vapor pressure the solution will be lt 1atm .
  Therefore, a higher temperature is required to
  attain a vapor pressure of 1 atm for the solution
  (?Tb).
• Molal boiling-point-elevation constant, Kb,
  expresses how much ?Tb changes with molality, m
• Kb is dependent only on the solvent used in the
  making of the solution.

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Freezing Point Depression

• Lower vapor pressure also affects the freezing
  point of the solution.
• The freezing point is the temperature at which
  the first crystal forms.
• When a solution freezes, almost pure solvent is
  formed first.
• Therefore, the sublimation curve for the pure
  solvent is the same as for the solution.
• Therefore, the triple point occurs at a lower
  temperature because of the lower vapor pressure
  for the solution.

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Freezing Point Depression Cont.

• The change in freezing point can be defined by
• Decrease in freezing point (?Tf) is directly
  proportional to molality (Kf is the molal
  freezing-point-depression constant)
• Lowering of freezing points by added solute
  explains the use of antifreeze in cars and the
  use of calcium chloride to melt ice.

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Osmosis

• Osmosis the movement of a solvent from low
  solute concentration to high solute
  concentration.
• Semipermeable membrane permits passage of some
  components of a solution. Example cell
  membranes
• There is movement in both directions across a
  semipermeable membrane.
• As solvent moves across the membrane, the fluid
  levels becomes uneven.

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Osmotic Pressure

• Osmotic pressure, ?, is the pressure required to
  stop osmosis
• Isotonic solutions two solutions with the same ?
  separated by a semipermeable membrane.
• Hypotonic solutions a solution of lower ? with
  respects to the more concentrated solution.
• The osmotic process is spontaneous.

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Colloids

• Colloids are suspensions in which the suspended
  particles are larger than molecules but too small
  to drop out of the suspension due to gravity.
• Tyndall effect ability of a Colloid to scatter
  light. The beam of light can be seen through the
  colloid.
• Particle size 10 to 2000 Å.
• There are several types of colloid
• aerosol (gas liquid or solid, e.g. fog and
  smoke),
• foam (liquid gas, e.g. whipped cream),
• emulsion (liquid liquid, e.g. milk),
• sol (liquid solid, e.g. paint),

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Hydrophilic and Hydrophobic Colloids

• Types of colloids dispersed in water
• Water loving colloids hydrophilic.
• Water hating colloids hydrophobic.
• Molecules arrange themselves so that hydrophobic
  portions are oriented towards each other.
• If a large hydrophobic macromolecule (giant
  molecule) needs to exist in water (e.g. in a
  cell), hydrophobic molecules embed themselves
  into the macromolecule leaving the hydrophilic
  ends to interact with water.

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Colloids Removal of Colloidal Particles

• Colloid particles are too small to be separated
  by physical means (e.g. filtration).
• Colloid particles are coagulated (enlarged) until
  they can be removed by filtration.
• Methods of coagulation
• heating (colloid particles move and are attracted
  to each other when they collide)
• adding an electrolyte (neutralize the surface
  charges on the colloid particles).

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End of Chapter 13Properties of Solutions