Chapter 11. Solutions and Their Properties

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Chapter 11. Solutions and Their Properties Powered By Docstoc
					                     Chapter 11. Solutions and Their Properties
                 Adapted from McMurry and Fay, Chemistry. 3rd Ed. 2001

I. Types of Solutions

       A. Homogeneous Mixtures

       Homogeneous mixtures can be classified as solutions or colloids. Solutions contain
particles with diameters in the range of nanometers––about the size of a molecule.
Solutions are typically transparent, and do not separate on standing. Colloids contain
larger particles, and are often opaque to light, and also do not separate on standing. (A
general property of a true homogeneous mixtures is that it does not separate on
standing.)

       B. Solution Types

       There are seven different types of solutions:

            Kind of Solution                    Example
            Gas in gas                          Air
            Gas in liquid                       Carbonated water
            Gas in solid                        H2 in palladium metal
            Liquid in liquid                    Gasoline
            Liquid in solid                     Dental amalgam (Hg in Ag)
            Solid in liquid                     Seawater
            Solid in solid                      Metal alloys

In any solution, the component present in the lesser amount is known as the solute, and
the component present in the greater amount is known as the solvent.

II. Energy Changes and the Solution Process

       A. ‘Like Dissolves Like’

       There are three types of interactions in a solution: solvent-solvent, solvent-
solute, and solute-solute. For solvent-solvent and––to a lesser extent––solute-solute
interactions, we can employ the ideas learned earlier when we studied intermolecular
interactions. Solvent-solute interactions are, on the other hand, much more complex.
Therefore it can be difficult to determine if a specific solute will be soluble in a specific
solvent. We can, however, present and defend a rule of thumb on this matter: ‘Like
Dissolves Like.’

        This ‘rule’ means that if the main solvent-solvent intermolecular interaction is
similar to the main solute-solute intermolecular interaction, the solute will be soluble in
the solvent. For example, NaCl(s) dissolves in water, since sodium chloride is an ionic
solid, and water has very polar covalent bonds (which leads to ion-dipole interactions).
Similarly, a non-polar solute like wax (made up of long-chain hydrocarbons) dissolves
in a non-polar solvent like benzene (C6H6) because both are dominated by dispersion
forces in their intermolecular interactions. On the other hand, oil will not dissolve in
water because they have different intermolecular interactions.

       Notice that the phrase ‘like dissolves like’ doesn’t explain anything… it simply
summarizes the observation that a solvent with a particular type of intermolecular force
will tend to dissolve a solvent with the same type of intermolecular force.

       B. Enthalpy of Solution

       As we have seen, nearly all chemical and physical properties have an associated
heat change. If the pressure is constant, this heat flow is called enthalpy. Specifically,
the enthalpy (or heat) of solution is the heat change that occurs when a solute dissolves
in a solvent. For example, we might write the general process of an ionic solid
dissolving to form an aqueous solution as

                         CA(s) + H2 O(l) " # CA(aq) $H = $ soln H
                                         "

where CA(aq) could also be written C+n(aq) + A–n(aq). For some solutions, ∆solnH is
positive (endothermic), and for others it is negative (exothermic). For example, CaCl2
(∆xolnH = –81.3 ! mol–1) and MgSO4 (∆solnH = –91.2 kJ mol–1) are used in hot packs, since
                kJ
they get hot when dissolved, and NH4NO3 (∆solnH = +25.7 kJ mol–1) is used for cold
packs, since it gets cold when dissolved.

      The value of the heat of solution arises from the various types of interactions
mentioned earlier:

      1. Solvent-solvent (∆H > 0). Energy is required to separate the solvent molecules
from each other to create room for the solute molecules.

        2. Solute-solute (∆H > 0). Energy is required to separate the solute molecules
from each other so they can spread out in solution. For an ionic solid solute, this energy
is the lattice energy. (Ch. 6)

       3. Solvent-solute (∆H < 0). Energy is released when solute molecules (especially
ions) are stabilized by the surrounding solvent molecules. In general, this energy
increases as the charge on the ion increases. Essentially, the energy of the system goes
down when new intermolecular interactions arise.

       The sum of these three interactions is the heat of solution, ∆solnH. Depending on
the relative magnitude of each individual ∆H above, ∆solnH can be positive or negative.

III. Units of Concentration

       A. Definitions

      Since there are seven different kinds of solutions (see I.B.), it is convenient to
have a few different units of concentration. Some we have seen before:
              1. Molarity (M)

       As we have seen, molarity is the moles of solute divided by liters of solution:

                                          mol of solute   mol
                             Molarity =                 "
                                          L of solution    L

You have seen the advantages of using this unit in your calculations––it’s relatively
easy to work with. On the downside, the molarity depends on the temperature of the
solution, since the total volume changes with the temperature. So, a solution with a
                    !
given concentration might be 0.013 M at one temperature, but 0.010 at a higher
temperature, even though the actual solution concentration hasn’t changed. Also, the
exact amount of solvent in a given volume can’t be determined unless the density of the
solution is known, since the volume is measured in L of solution.

              2. Mole Fraction (X)

       This is analogous to our discussion of mole fractions in relation to gases:

                                         mol of component
                     Mole Fraction =                          " unitless
                                        total mol of solution

Mole fractions are typically used for gas solutions, since one can show that mole
fractions of ideal gases are directly related to the partial pressures of a mixture of gases;
i.e. pA = XApT. (This is Dalton’s Law of Partial Pressures).
             !

              3. Mass Percent (mass %)

      The mass percent of any component of a solution is the mass of that component
divided by the total mass of the solution (times 100%)
                                                                                (11-1)
                                        Mass of component
                      Mass percent =                         " 100%
                                      Total mass of solution

(Numbered equations are important ones that you should understand and remember.
They won’t be given on exams.) Mass percent can be thought of as “parts per
hundred”; 1% is one part per hundred. (This is where the name comes from, right?
             !
per-cent... per-hundred. Think of century as 100 years.) We can also have parts per
million (ppm) and parts per billion (ppb):

                                      Mass of component
                            ppm =                           " 106
                                     Total mass of solution

                                      Mass of component
                             ppb =                          " 109
                                     Total mass of solution
                  !
     The advantage to using mass percent is that it is independent of temperature.
Conversely, it can be difficult to deal with liquid volumes and gases.
                  !
             4. Molality (m)

      Molality is defined as the number of moles of solute per kilogram of solvent:
                                                                                      (11-2)
                                            Moles of solute
                          Molality (m) =
                                           Mass of solvent (kg)

Note the differences between this and molarity. Molality is also temperature
independent, but again, one needs to know the density to determine volumes.
                   !
        Practice converting between each of these units of concentration. (Examples 11.2
- 11.10 in M&F are a good ones to try.)