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 Chapter 12
Types of Mixtures

   Ch 12 Section 1
 Homogeneous Aqueous Systems
• Soluble
  – Capable of being dissolved

   (a) Milk consists of visible particles in a non-uniform arrangement.
   (b) Salt water is an example of a homogeneous mixture. Ions and water
       molecules are in a uniform arrangement.
 Homogeneous Aqueous Systems
• Solution
  – A homogeneous mixture of two or more
    substances in a single phase
• Solvent
  – The dissolving medium in a solution
• Solute
  – The dissolved substance in a solution
  Homogeneous Aqueous Systems
Types of Solutions
• Gaseous mixtures
  – Air is a solution
• Solid solutions
  – Metal alloys
• Liquid solutions
  – Liquid dissolved in a liquid (alcohol in water)
  – Solid dissolved in a liquid (salt water)
     • Note: the composition of each solution is uniform.
(a) 24-karat gold is pure gold.
(b) 14-karat gold is an alloy of gold and silver. 14-karat gold is
   14/24, or 58.3%, gold.
  Homogeneous Aqueous Systems
Solutes: Electrolytes vs. Nonelectrolytes
• Electrolyte
  – A substance that dissolves in water to give a
    solution that conducts electric current
  – Solutions of acids, bases and salts are electrolytes
• Nonelectrolyte
  – A substance that dissolves in water to give a
    solution that does not conduct an electric current
 Homogeneous Aqueous Systems
Measuring Conductivity
• Good conductors
  – Lamp glows brightly, ammeter registers a
    substantial current
• Moderate conductors
  – Lamp is dull, ammeter registers a small current
• Nonconductors
  – Lamp does not glow, ammeter may not register a
    current at all
     Homogeneous Aqueous Systems

(a) Sodium chloride dissolves in water to produce a salt solution that conducts electric current.
    NaCl is an electrolyte.
(b) Sucrose dissolves in water to produce a sugar solution that does not conduct electricity.
    Sucrose is a non-electrolyte.
(c) Hydrogen chloride dissolves in water to produce a solution that conducts current. HCl is an
 Heterogeneous Aqueous Systems
• Suspensions
  – A mixture from which particles settle out upon
     • Ex. Jar of muddy water
        – Soil particles are much larger and more dense than water
          molecules so they settle out.
    Heterogeneous Aqueous Systems
• Colloidal Dispersions (Colloids)
     – Tiny particles suspended in some medium
     – Particles range in size from 1 to 1000 nm.
• Tyndall Effect
     – Scattering of light by particles
           • Light passes through a solution
           • Light is scattered in a colloid
•   The mixture of gelatin and water in the jar on the right is
    a colloid. The mixture of water and sodium chloride in
    the jar on the left is a true solution.
 Heterogeneous Aqueous Systems
                           Types of Colloids
                               Dispersing      Dispersed
             Examples                                       Colloid Type
                                Medium         Substance
Fog, aerosol sprays         Gas             Liquid         Aerosol
Smoke, airborne bacteria    Gas             Solid          Aerosol
Whipped cream, soap suds    Liquid          Gas            Foam
Milk, mayonnaise            Liquid          Liquid         Emulsion
Paint, clays, gelatin       Liquid          Solid          Sol
Marshmallow, polystyrene    Solid           Gas            Solid Foam
Butter, cheese              Solid           Liquid         Solid emulsion
Ruby glass                  Solid           Solid          Solid sol
The Solution Process

      Section 2
        Factors Affecting the Rate of
• Increasing the Surface Area of the Solute
  (Particle size)
   – Finely divided substances dissolve more rapidly
      • Ex. Crush the sugar cube first
• Agitating a Solution
   – Stirring or shaking brings
      solvent into contact with
     more solute particles
   – Added energy temporarily
     increases solubility
• Heating a Solvent
   – Heating always increasing the
     rate of dissolution of solids in
Solution Equilibrium - The physical state in which the opposing
   processes of dissolution and crystallization of a solute occur at
   equal rates

Saturation Levels
• Saturated solution
   – A solution that contains the maximum amount of dissolved solute
• Unsaturated solutions
   – A solution that contains less solute than a saturated solution under
     the existing conditions
• Supersaturated Solutions
   – A solution that contains more dissolved solute than a saturated
     solution contains under the same conditions
A saturated solution in a closed system is at equilibrium. The solute is
recrystallizing at the same rate that it is dissolving, even though it appears that
these is no activity in the system.
• Solubility Values
   – The solubility of a substance is the amount of that
     substance required to form a saturated solution with a
     specific amount of solvent at a specified temperature
   – The rate at which a substance dissolves does not alter the
     substances solubility

  This graph shows the range of solute masses
  that will produce an unsaturated solution.
  Once the saturation point is exceeded, the
  system will contain undissolved solute.
      Solute-Solvent Interactions
• "Like dissolves like"
  – Polar substances dissolve in polar solvents
  – Nonpolar substances dissolve in nonpolar solvents
       Solute-Solvent Interactions
• Dissolving Ionic Compounds in Aqueous Solutions
   – Electropositive hydrogen of the water molecule is attracted to
     negatively charged ions
   – Electronegative oxygen of the water molecule is attracted to positively
     charged ions
   – Hydration
       • The solution process with water as the solvent
   – Hydrates
       • Ionic substances that incorporate water molecules
         into their structure during the recrystallization
       the ” •" means that the water is loosely attached

      Hydrated copper (II) sulfate has water trapped in the
      crystal structure. Heating releases the water and
      produces the anhydrous form of the substance, which has
      the formula CuSO4.
     Solute-Solvent Interactions
• Nonpolar Solvents
  – Polar and ionic compounds are not soluble in
    nonpolar solvents
  – Fats, oils and many petroleum products are
    soluble in nonpolar solvents
  – Nonpolar solvents include CCl4 and toluene
    (methyl benzene), C6H5CH3
       Liquid Solutes and Solvents
   – Immiscible - Liquid solutes and
     solvents that are not soluble in each
       • Oil and water

Toluene and water are immiscible. The components
of this system exist in two distinct phases.
           Liquid Solutes and Solvents
     – Miscible - Liquids that dissolve freely
       in one another in any proportion
          • Benzene and carbon tetrachloride (both
          • Water and ethanol (both polar)

(a) Water and ethanol are miscible. The components of this system exist in a single phase
with a uniform arrangement. (b) Hydrogen bonding between the solute and solvent
enhances the solubility of ethanol in water.
Effects of Pressure on Solubility
– Pressure has no real effect on the solubilities of
  liquids and solids in liquid solvents
– Increasing pressure increases
   the solubility of gases in
     • The solubility of a gas in a liquid is directly
       proportional to the partial pressure of that gas on the
       surface of the liquid. – Henry’s law

(a) There are no gas bubbles in the unopened
bottle of soda because the pressure of CO2
applied during bottling keeps the carbon dioxide
gas dissolved in the liquid. (b) When the cap on
the bottle is removed, the pressure of CO2 can
escape from the liquid. The soda effervesces
when the bottle is opened and the pressure is
    • Effects of Temperature on Solubility
          – Solubility of solids (generally) increases with
          – Solubility of gases decreases with temperature

The solubility of gases in
water decreases with
increasing temperature.
Which gas has the greater
solubility at 30°--CO2 or
Solubility curves for various solid solutes generally show increasing solubility
with increases in temperature. From the graph, you can see that the solubility
of NaNO3 is affected more by temperature than is NaCl.
                       Heats of Solution
• Solvent molecules are held together by intermolecular forces (solvent-
  solvent attraction)
• In the solute, molecules are held together by intermolecular forces
  (solute-solute attraction).
• Energy is required to separate solute molecules and solvent molecules
  from their neighbors.
• The net amount of heat energy absorbed or released when a specific
  amount of solute dissolves in a solvent is the heat of solution.
    – When the heat of solutions is negative (heat is released) the sum of attractions from
      Steps 1 and 2 is greater than Step 3.
                        Heats of Solution
• A solute particle that is surrounded by solvent molecules, as shown
  in the figure, is said to be solvated.

The graph shows the changes in the heat content that occur during the formation of a solution.
How would the graph differ for a system with an endothermic heat of solution?
Concentration of Solutions

         Section 3
     Concentration of Solutions
Concentration - A measure of the amount of
  solute in a given amount of solvent or solution

         Not this type of concentration 
• The number of moles of solute in one liter of
  – To find molarity of a solution, you must know the
    molar mass of the solute
     • The units of molarity, mol/L, are usually represented by
       a scripted capital “M”.

       Nope, not this type of mole 
The preparation of a 0.5000 M solution
of CuSO4•5H2O starts with calculating
the mass of solute needed.
You have 3.50 L of solution that contains 90.0 g of sodium
chloride, NaCl. What is the molarity of that solution?

 Given: solute mass = 90.0 g NaCl
        solution volume = 3.50 L
 Unknown: molarity of NaCl solution
       g of solute  number of moles of solute  molarity

 90.0 g NaCl x    1molNaCl = 1.54 mol NaCl
                 58.44 g NaCl

  1.54 mol NaCl = 0.440 M NaCl
 3.50 L of solution
You have 0.8 L of a 0.5 M HCl solution. How many
moles of HCl does this solution contain?

Given: volume of solution = 0.8 L
       concentration of solution = 0.5 M HCl
Unknown: moles of HCl in a given volume
concentration (mol of HCl/L of soln) x volume (L of soln) = mol of HCl

   0.5 mol HCl x 0.8 L of solution = 0.4 mol HCl
1.0 L of solution
To produce 40.0 g of silver chromate, you will need at least 23.4 g of
potassium chromate in solution as a reactant. All you have on hand in the
stock room is 5 L of a 6.0 M K2CrO4 solution. What volume of the solution is
needed to give you the 23.4 g K2CrO4 needed for the reaction?
Given: volume of solution = 5 L
         concentration of solution = 6.0 M K2CrO4
         mass of solute = 23.4 g K2CrO4
         mass of product = Ag2CrO4
Unknown: volume of K2CrO4 solution in L
                        g of solute  moles of solute
            moles solute and molarity  liters of solution needed
        23.4 g K2CrO4 x 1.0 mol K2CrO4 = 0.120 mol
                        194.2 g K2CrO4
         6.0 M K2CrO4 = 0.120 mol K2CrO4
                       X L K2CrO4 soln

                   X = 0.020 L K2CrO4 soln

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