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					Chapter 11: Solutions

      Cory Stuart
Let’s Review the Solubility Rules
1.Most nitrate (NO31-) salts are soluble.
2.Most salts of Na+, K+, and NH4+ are soluble.
3.Most chloride salts are soluble. Notable exceptions are
  AgCl, PbCl2, and Hg2Cl2.
4.Most sulfate salts are soluble. Notable exceptions are
  BaSO4, PbSO4, and CaSO4.
5.Most hydroxide compounds are only slightly soluble.* The
  important exceptions are NaOH and KOH, Ba(OH)2 and
  Ca(OH)2 are only moderately soluble.
6.Most (S2-), (CO32-), and (PO43-) salts are only slightly
  soluble.
              The Basics
• A solution is a homogenous mixture

• A solute is the substance being dissolved

• A solvent does the dissolving
    Solution = solvent + solute
• Solution = Solute + Solvent
• Aqueous (water)
   – Tincture (alcohol)
   – Amalgam (mercury)
   – Organic
      • Polar
      • Non-polar
                 Molarity
• A term used to describe solution
  composition
• The number of mole of solute per liter of
  solution
• Symbolized by M
                 Molarity
• What is the molarity of a solution
  containing 0.32 moles of NaCl in 3.4 liters?


• molarity =    0.32 moles NaCl ÷ 3.4
           =    0.094 M NaCl
     Mass Percent Composition

1. Calculate the molar mass of the compound
2. Calculate the molar mass of the element of interest
   in the compound by multiplying the subscript by
   the molar mass of the element
3. Divide the molar mass of the element in the
   compound by the total molar mass of the
   compound
     Mass Percent Composition
• Example 1. What percentage of the mass of
  carbon dioxide (CO2) is made up by the carbon?
•   Solution: first find the mass of the total
  compound.
      C = 12.0 u x 1 atom = 12.0 u
      O = 16.0 u x 2 atoms = 32.0 u
                              --------
                              44.0 u
•
                  continued
• % of the mass of CO2 that is made up by carbon
  = 12.0 u (partial mass of carbon) ÷ 44.0 u (total
      mass CO2) x 100

• % of the mass of CO2 that is made up by carbon
  = 27.3%
•
             Mole Fraction
• Ratio of the number of moles of a given
  component tot he total number of moles of
  solution
• Symbolized by the lowercase Greek letter
  chi ( )
• Mole Fraction = moles of substance A ÷
                  total Moles of solution
Mole Fraction
• Example
  Problem #1:
             Molality


• Unlike molarity, temperature does not
            change the unit
                 Molality
• Molality of C2H5OH =
 (moles of C2H5OH ÷ kilograms of H2O)
                        =
 (2.17E-2 mole) ÷ (100.0g x (1kg ÷ 1000g))
                        =
 92.17E-2 mole ÷ 0.1000kg
                        =
 0.217 m
   Energy of Making Solutions
• The Heat of Solution is the amount of heat
  energy absorbed (endothermic) or released
  (exothermic) when a specific amount of
  solute dissolves in a solvent.
• Most easily understood if broken into
     steps.
    Steps in Solution Formation
1 Expanding the solute
     Separating the solute into individual
2 Expanding the solvent
     Separating the solute into individual components
3 Interaction of solute and solvent to form the
  solution
    Mixing Solvent and Solute
• DH3 depends on what you are mixing.
• Molecules can attract each other DH3 is
  large and negative.
• Molecules can’t attract- DH3 is small and
  negative.
• This explains the rule “Like dissolves
      Like”
        “Like Dissolves Like”
• Polar and ionic solutes dissolve best in polar
  solvents
  – fats, steroids, and waxes in dissolve best in
    benzene, hexane, and toluene
• Nonpolar solutes dissolve best in nonpolar
  solvents
  – inorganic salts and sugars dissolve best in
    water, small alcohol amounts, and acetic acid
                      Practice
• Decide whether liquid hexane or liquid methenol
  is the more appropriate solvent for the substances
  grease and potassium iodide.


~ Hexane is a non-polar solvent because it contains C-H
  bonds. Thus hexane will be better for the non-polar solute
  grease. Methanol has an O-H group that makes it polar.
  Thus it is a better solvent for the ionic solid potassium
  iodide.
                   Solubility
• Two types of solubility:
     • Fat-soluble- must be a nonpolar solvent dissolved in
       a nonpolar material, such as fat
     • water soluble- must have dipole moments
• Hydrophobic ( fat soluble) means afraid of
  water
• Hydrophilic (water soluble) means water-
  loving
         Feelin’ the pressure?
• Pressure does not affect the solubility of
  liquids or solids, only gases.

• Solubility increases as pressure increases
  and vice versa.
           Dissolving Gases

• Pressure effects the amount of gas that can
  dissolve in a liquid.

• The dissolved gas is at equilibrium with the
  gas above the liquid.
               Equilibrium
• The gas is at equilibrium with the dissolved
  gas in this solution.
• The equilibrium is dynamic.
• Increase the pressure and the solubility
  increases, thus disturbing the equilibrium.
• A new equilibrium is created, with more
  dissolved gas.
                 Henry’s Law
• The change in equilibrium is modeled by
  Henry’s Law:
     • P = kC
        – P is pressure, k is the constant, and C is the concentration
          of dissolved gas.

• States that the amount of a gas dissolved in
  a solution is directly proportional to the
  pressure of the gas above the solution.
               Temperature
• Increased temperature increases the rate at
  which a solid dissolves.
• Unpredictable whether it will increase the
  amount of solid that dissolves.
• Predictable to tell whether it will increase
  the amount of gas dissolved.
        Temperature and Gas
• As temperature increases, the amount of
  solubility decreases.
• As temperature increases, the gas molecules
  can move fat enough to escape its confines.
• Explains thermal pollution (the oxygen-
  depleting effect on lakes and rivers of using
  water for industrial cooling and return to its
  natural source at a higher temperature).
            Vapor Pressure
• The vapor pressure of a solvent is lowered
  by nonvolatile solute.

• The nonvolatile solute decreases the amount
  of soluble solvents per unit volume, thus
  lowering the number of solvent molecules
  at the surface.
            Vapor Pressure
• The escaping tendency of the molecules is
  lowered proportionately to the lowered
  number of solvent molecules on the surface.

• The molecules of the solvent must
  overcome the force of both the other solvent
  molecules and the solute molecules.
       Modified Raoult's Law:
   • Liquid-liquid solutions in which both
           components are volatile

• P0 is the vapor pressure of the pure solvent
• PA and PB are the partial pressures
             Raoult’s Law
• A solution that obeys Raoult’s Law is called
  an ideal solution
Colligative Properties
     Freezing Point Depression
• When a solute is dissolved in a solvent, the
  freezing point of the solution is lower than
  that of the pure substance. This is due to the
  change in vapor pressure of the solvent.
• Expressed by the equation:
     Freezing Point Depression


• Along with boiling-point elevation, this can
  be used to determine the molar masses and
  to characterize solutions.
 Freezing Point Depression and Boiling
    Point Elevation Constants, °C/m
Solvent                Kf        Kb
Acetic Acid            3.9       3.07
Benzene                5.12      2.53
Nitrobenzene           8.1       5.24
Phenol                 7.27      3.56
Water                  1.86      0.512
 Freezing Point Depression and Boiling
    Point Elevation Constants, °C/m
Solvent                Kf        Kb
Acetic Acid            3.9       3.07
Benzene                5.12      2.53
Nitrobenzene           8.1       5.24
Phenol                 7.27      3.56
Water                  1.86      0.512
           Osmotic Pressure
• Osmosis is the the flow of solvent into a
  solution through a semi-permeable
  membrane.
• Osmotic Pressure is (       ) is the pressure
  that must be applied to a solution to stop
  osmosis
• expressed by
           Osmotic Pressure
• Osmosis can be prevented by applying a
  pressure to the solution.

• The minimum pressure that stops the
  osmosis is equal to the osmotic pressure of
  the solution.
                 Dialysis
• A phenomenon in which a semi-permeable
  membrane allows transfer of both solvent
  molecules and small solute molecules and
  ions.
• Used to cleanse the kidneys
• Takes place in the walls of most plant and
  animal cells.
          Isotonic Solutions
• Solutions that have identical osmotic
  pressures
• What concentration of NaCl in water is
  needed to produce an aqueous solution
  isotonic with blood( = 7.7 atm at 25 C)?
       – The solution is 0.158 M
           Reverse Osmosis
• The process process when the external
  pressure on a solution causes a net flow of
  solvent through a semi-permeable
  membrane from the solutions to the solvent.
• The semi-permeable membrane acts as a
  “molecular filter” to remove solute
  particles. This is applicable due to
  desalination (removal of dissolved salts).
        Electrolytes and Non-
             electrolytes.
• An electrolyte is:
  – A substance whose aqueous solution conducts
    an electric current.
• A non-electrolyte is:
  – A substance whose aqueous solution does not
    conduct an electric current.
• Try to classify the following substances as
  electrolytes or non-electrolytes…
Electrolyte or not?
         1. Pure water
          2. Tap water
       3. Sugar solution
 4. Sodium chloride solution
5. Hydrochloric acid solution
    6. Lactic acid solution
   7. Ethyl alcohol solution
  8. Pure sodium chloride
               Answers...
Electrolytes              Non-electrolytes

Tap Water (weak)          Pure Water
NaCl Solution             Sugar Solution
HCl                       Ethanol Solution
Lactate Solution (weak)   Pure NaCl
         van 't Hoff Factor, i
• Electrolytes may have two, three or more
  times the effect on boiling point, freezing
  point, and osmotic pressure, depending on
  its dissociation.
• Expressed by:
  moles of particles in solution ÷ moles of
  solute dissolved
                     Colloids
• A suspension of particles in some medium.
     • Also called a colloidal dispersion.
• Classified by the states of the dispersed
  phase and the dispersing medium (see table
  11.7 on page 550)
• Colloids are stabilized due to electrostatic
  repulsion
                  Coagulation
• The destruction of a colloid usually done by
  heating or adding an electrolyte.
  – Heating increases the velocities eventually causing the
    particles to aggregate, and then settles,
  – Adding an electrolyte neutralizes the absorbed ion
    layers.
• An example would be to remove soot from
  smoke
» the Cottrell Precipitator
            Practice Problems:
1~ A solution is made using 6.9g of NaHCO3 per
  100g of water. What is the weight of percentage of
  the solute in this solution?

2~ Predict whether each of the following substances
  are more likely to dissolve in carbon tetrachloride,
  CCl4, a non-polar solvent, or in water: C7H16,
  NaHCO3, HCL, and I2.
            Practice Problems:
3~ What is the concentration of all of the species in a
  0.1M solution of Ba(NO3)2?

4~ Calculate the freezing point and the boiling point
  of a solution of 100g of ethylene glycol (C2H6O2)
  in 900g of H2O.
           Practice Problems:
5~The average osmotic pressure of blood is 7.7 atm
  at 25 °C. What concentration of glucose, C6H12O6,
  will be isotonic with blood?

6~ How many grams of benzene acid, C6H5OOH,
  must be dissolved in 50.0 mL of benzene, C6H6 (d
  = 0.879 g/mL), to produce 0.150 m C6H5COOH?
           Practice Problems:
7~ The vapor pressure of pure water at 20.0 °C is
  17.5 mmHg. What is the vapor pressure at 20.0 °C
  above a solution that has 0.250 mol C12H22O11
  (sucrose) and 75.0g CO(NH2)2 (urea) dissolved
  per kilogram of water?

8~ What is the freezing point of an aqueous sucrose
  solution that has 25g C12H22O11 per 100g H2O?
           Practice Problems:
9~ A 375 mL sample of hexane vapor in equilibrium
  with liquid hexane, C6H14 (d = 0.6548 g/mL), at
  25 °C is dissolved in 50.0 mL of liquid
  cyclohexane, C6H12 (d = 0.7739 g/mL; vp = 97.58
  torr), at 25 °C. Calculate the total vapor pressure
  above the solution at 25 °C.
                  Solutions:
1~ 6.5%

2~ Both C7 and I2 are non-polar. We would then
  assume that they could be more soluble in CCl4
  than in H2O. On the other hand, NaHCO3 is
  isotonic and HCl is polar covalent. Water could be
  a better solvent than CCL4 for these two
  substances.
                  Solutions:
3~ Ba(NO3)2 is a salt and expected to be a strong
  electrolyte. It ionizes into Ba2+ and the polyatomic
  nitrate ion, NO3-. A .01 M solution of Ba(NO3)2 is
  .01 M Ba2+ and 0.2 M in NO3- (since 1 mol
  Ba(NO3)2 supplies 1 mol Ba2+ and 2 mol NO3-).

4~ Freezing Point = -3.33 °C
   Boiling Point = 100.39 °C
                    Solutions:
5~ M = 0.031

6~ 0.805g C6H5COOH

7~ PH2O = 17.0 mmHg

8~ -1.36 °C

9~ Ptotal = 98.0 torr
      Recommended Site to visit:
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• http://www.sciencegeek.net/APchemistry/Powerpoints/Chapter11/Cha
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• http://chemistry.about.com/library/weekly/aa030503a.htm
• plaza.ufl.edu/cesarc/Study%20Guide%202.htm
• chemdept.chem.ncsu.edu/.../fht/sld004.htm
• http://www.chem.ox.ac.uk/vrchemistry/labintro/newdefault.html
• http://www.infoplease.com/chemistry/simlab/
• http://www.chemcollective.org/vlab/vlab.php
• http://www.mmlab.ru/projects/index_en.shtml

				
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