Chapter 13: Properties of Solutions

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Chapter 13 Properties of Solutions

• Sam White
• Pd. 2

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Introduction

• A solution is any homogenous mixture, which means
  the components are uniformly intermingled on a
  molecular level
• The Solvent is the most abundant component. It
  does the dissolving.
• The Solute are any of the other components. They
  are the ones being dissolved

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Formation of Solutions

• With the exception of gas solutions, solutions
  form when the attractive forces between solute
  and solvent are comparable or greater than the
  intermolecular forces in either component

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Formation of Solutions

• Example Salt Water- Attractive forces between
  Na or Cl- and the polar water molecules overcome
  the lattice energy or solid NaCl
• Once separated, the Na and Cl- are surrounded by
  water. This interaction is known in all solutions
  as solvation
• When the solvent is water, this interaction is
  known as hydration

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Energy Change in Solution Formation

• In order to form a solution, the solvent must
  form space to house the solute and the solute
  must be dissolved, both of which take energy

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Enthalpy of Solution

• Overall Enthalpy Change
• DHsolutionDH1DH2DH3
• Example with Salt Water
• DH1 accounts for the separation of NaCl to Na
  and Cl-
• DH2 accounts for the separation of solvent
  molecules to accommodate the solute
• DH3 accounts for the attractive interactions
  between solute and solvent

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Overall Enthalpy Change
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Saturated Solutions

• As concentration of a solid solute increases, so
  does its chance of of colliding with the surface
  of the solid and becoming reattached to the solid
• This is called crystallization
• Solute Solvent Solution

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Saturated Solutions

• When the rates of crystallization and dissolving
  become equal, no increase of solute in solution
  will occur
• When a solution will not dissolve any more
  solute, it is a saturated solution
• When a solution that can still dissolve solute
  into it is an unsaturated solution

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Supersaturation

• Under suitable conditions, it is sometimes
  possible to form a solution with more solute than
  that needed for a saturated solution
• These solutions are supersaturated

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Supersaturation

• Supersaturation usually occurs because many
  solutes are more soluble at one temperature than
  another
• Example Sodium acetate, NaC2H3O2, will dissolve
  in water more readily at higher temperatures.
  When a saturated solution is made at higher
  temperatures then slowly cooled, all of the
  solute may remain dissolved even though the
  solubility decreases

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Factors Affecting Solubility

• The stronger the intermolecular attractive forces
  between solute and solvent, the greater the
  solubility
• As a result of favorable dipole-dipole
  attractions, polar liquids tend to dissolve more
  readily in polar solvents
• Water is not only polar, but has hydrogen bonds,
  making solutes that have hydrogen bonds able to
  dissolve in water as well

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Factors Affecting Solubility

• Pairs of liquids that mix in all proportions are
  miscible
• Liquids that do not dissolve significantly in one
  another are immiscible

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Hydrocarbons vs. Alcohols

• Many hydrocarbons are immiscible in water because
  they are nonpolar molecules
• Alcohols have an OH group, which are both polar
  and have hydrogen bonds, making them more readily
  soluble in water
• As the carbon chain become larger, the effect of
  the OH group becomes smaller, meaning that larger
  alcohol chains begin to become less soluble

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Pressure Effects

• Pressure only affects the solubility of gas in a
  solvent
• As pressure increases, solubility of the gas
  increases

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Henrys Law

• Cg kPg
• Cg is the solubility of the gas in solution
  (usually expressed in molarity)
• Pg is the partial pressure of the gas over
  solution
• k is the Henrys Law Constant, which is unique
  for all solute-solvent pairs as well as the
  temperature

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Temperature Effects

• As temperature increases, the solubility of solid
  solutes (such as salts) normally increases
• In contrast, as temperature increases, the
  solubility of gaseous solutes normally decreases

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Solubility Charts
Gas Solubility Curve
Solids Solubility Curve
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Ways of Expressing Concentration

• Mass percentage, ppm
• Mole Fraction
• Molarity
• Molality

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Mass Percentage and ppm

• Mass of component (mc / mt) x 100
• mc mass of component in solution
• mt total mass of solution
• ppm of component (mc / mt) x 106
• mc and mt are the same for ppm

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Mole Fraction

• Mole Fraction of Component (molc /
  molt)
• molc moles of component
• molt total moles of all components

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Molarity

• Molarity (mols / Ls)
• mols moles solute
• Ls liters solution

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Molality

• Molality (mols / kgs)
• mols moles solute
• kgs kilograms solvent

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Colligative Properties

• Colligative properties depend on the quanity of
  solute, not the type of solute
• The colligative properties are
• Vapor-Pressure Reduction
• Boiling-Point Elevation
• Freezing-Point Depression
• Osmotic Pressure

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Vapor-Pressure Reduction

• As the amount of solute increases, the vapor
  pressure of solution decreases
• This relationship can be expessed through
  Raoults Law
• PA XAPoA
• PA Partial pressure exerted by solvent
• XA Mole fraction of solvent
• PoA Vapor pressure of pure solvent

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Boiling-Point Elevation

• As amount of solute increase, boiling point
  increases
• This relationship can be expressed as
  DTb dKbm
• DTb total boiling point elevation
• d dissociation factor of the solute
• Kb molal boiling point elevation constant of
  the solvent
• m molality of solution

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Freezing-Point Depression

• As amount of solute increases, freezing point
  decreases
• This relationship can be expressed as
  DTf dKfm
• DTf total freezing point depression
• d dissociation factor of the solute
• Kf molal freezing point depression constant of
  the solvent
• m molality of solution

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Osmotic Pressure

• As amount of solute increases, osmotic pressure
  increases
• This relationship can be expressed as
  p (n / V)RT MRT
• p osmotic pressure
• n number of moles solute
• V volume of solution
• R ideal gas constant
• T temperature of solution
• M molarity of solution