AP Chemistry Chapter 13 � Properties of Solutions by Mm8OFfu

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									AP Chemistry Chapter 13 – Properties of Solutions

The Solution Process
      A solution is a homogeneous mixture of solute and solvent.
      Solutions may be gases, liquids or solids.
      Each substance present is a component of the solution.
          o The solvent is the component present in the largest amount.
          o The component(s) dissolved (present in the lesser amounts) are solutes.
      Soln. process with condensed phases: rearrangement of intermolecular forces (IMF)
          o Solute-solvent interactions > solute-solute or solvent-solvent interactions
          o Example: NaCl (solute) dissolving in water (solvent)
          o H-bonds between water molecules are broken
          o Ion-dipole forces form between Na+ and water and between Cl- and water
                   Such an interaction between a solute and a solvent is called solvation
                   If water is the solvent, the interaction is called hydration

Energy Changes and Solution Formation
    Three steps involving energy in the formation of a solution:
         o Separation of solute molecules (∆H1) – breaking IMF always endothermic
         o Separation of solvent molecules (∆H2) – endothermic
         o Formation of solute-solvent interactions ((∆H3) – forming IMF always exothermic
    ∆Hsoln = ∆H1 + ∆H2 + ∆H3
         o ∆Hsoln can be endothermic or exothermic depending on whether:
         o ∆H3 > (∆H1 + ∆H2) OR ∆H3 < (∆H1 + ∆H2) See Fig. 13.4
    How to predict if a soln. will form:
         o Soln. forms if ∆Hsoln is negative
         o Soln. will not form if ∆Hsoln is too endothermic
         o Rule of thumb: “like dissolves like”
         o Polar (nonpolar) solvents dissolve polar (nonpolar) solutes
         o Examples: mixing NaCl and water, NaCl and gasoline, water and octane (C8H18)

Solution Formation, Spontaneity, and Disorder
    A spontaneous process occurs without outside intervention
    When the energy of the system decrease, the process is spontaneous
            o Spontaneous processes tend to be exothermic
    Some spontaneous processes do not involve the movement of the system to a lower energy state
       (i.e. endothermic process), but are driven by an increase in disorder (entropy)
    In most cases, soln. formation is favored by the increase in disorder that accompanies mixing
            o Example: mixture of CCl4 and C6H14 has ∆Hsoln close to zero but is less ordered than the
                two separate liquids   See Fig. 13.6
    A soln. will form unless solute-solute or solvent-solvent interactions are too strong relative to
       solute-solvent interactions

Solution Formation and Chemical Reactions
    Some solns. form by physical processes, others by chemical processes:
          o Example:      Ni + 2HCl  NiCl2 + H2 (chemical)
                          NaCl + H2O  Na+ + Cl- (physical)
Saturated Solutions and Solubility
      As a solid dissolves, a solution forms (dissolution): solute + solvent  solution
      Crystallization is the opposite of dissolution: solution  solute + solvent
      If crystallization and dissolution are in equilibrium with undissolved solute, the solution is
       saturated
           o No further increase in the amount of dissolved solute (ratedissolution = ratecrystallization)
      Solubility is the amount of a given solute that dissolves in a given amount of solvent to form a
       saturated solution
      A solution with a concentration of dissolved solute that is less than the solubility is said to be
       unsaturated
      A solution is said to be supersaturated if more solute is dissolved than in a saturated solution
           o Supersaturated solutions are unstable  crystallization

Factors Affecting Solubility
      The tendency of one substance to dissolve in another depends on: the nature of the solute/solvent,
       the temperature, and the pressure (for gases)

Solute-Solvent Interactions
    Pairs of liquids that mix in any proportions are said to be miscible
          o Example: ethanol (C2H3OH) and water are miscible liquids
          o Why? What happens as the number of C atoms on the alcohol increases? As the number of
             –OH groups increases?
    Immiscible liquids do not mix significantly
          o Example: gasoline and water
    Intermolecular forces are an important factor: “like dissolves like”
          o Substances with similar IMF tend to be soluble in one another
          o The more polar (or if ionic) the bonds are in the molecule, the better it dissolves in a polar
             solvent.
          o The less polar the bonds, the less likely it is to dissolve in a polar solvent and the more
             likely it is to dissolve in a nonpolar solvent
          o Network solids do not dissolve because the strong IMF in the solid are not reestablished in
             any solution
                                              Pract. Ex. 1

Pressure Effects
    The solubility of a gas in a liquid is a function of the pressure of the gas over the solution
    Solubilities of solids and liquids are not greatly affected by pressure
    With a higher pressure, more molecules of gas are close to the surface of the solution, increasing
      the probability of a gas molecule striking the surface an entering the solution
    The higher the pressure, the greater the solubility See Fig. 13.14
    Henry’s Law states that the solubility of a gas is directly proportional to the partial pressure of the
      gas above the solution
          o Mathematically:          Cg = kPg         Cg = solubility of the gas
                                                      Pg = partial pressure
                                                      k = Henry’s law constant (differs for each solute
                                                             solvent pair and with temperature)
          o Example: carbonated beverages
                                              Pract. Ex. 2
Temperature Effects
   Solubility of solids in liquids generally increases with an increase in temperature
   Gases are less soluble in liquids at higher temperatures
         o Why?
   Environmental application: thermal pollution

Ways of Expressing Concentration
      All methods involve quantifying the amount of solute per amount of solvent (or solution)
      Concentration can be expressed qualitatively or quantitatively (dilute vs. concentrated)
      Quantitative expressions use mass, moles or liters of solute, solvent or solution

Mass Percentage, ppm, and ppb
   Mass percentage is the number of grams of solute per 100 grams of solution
             mass % = mass of component in soln X 100
                              total mass of soln
   Parts per million (ppm) is the number of grams of solute per one million grams of solution
             ppm = mass of component in soln X 106
                        total mass of soln
   Parts per billion (ppb) is the number of grams of solute per one billion grams of solution
             ppb = mass of component in soln X 109
                        total mass of soln
   If soln is aqueous (density 1g/mL):
         o ppm is also 1 mg solute/kg soln or 1mg solute/L soln
         o ppb is also 1 μg solute/kg soln or 1 μg/L soln
                                             Pract. Ex. 3

Mole Fraction, Molarity, Molality
   mole fraction of a component, X =           moles of component     .
                                          total moles of all components
           o mole fraction has no units and ranges from 0 to 1

      molarity , M = moles of solute
                     liters of solution
          o molarity will change with temp. (as soln volume increases or decreases)

      molality, m = moles of solute
                     kg of solvent
          o converting between M and m requires density
                                       Pract. Ex. 4-6

Colligative Properties
      Colligative properties are physical properties of solutions that depend on the number of solute
       particles
      Four colligative properties to consider: vapor pressure lowering, boiling point elevation, freezing
       point depression, osmotic pressure
Lowering the Vapor Pressure
   Vapor pressure of a soln is always lower than that of a pure solvent
   Nonvolatile solutes reduce the ability of the surface solvent molecules to escape the liquid
          o Therefore, the vapor pressure is lowered
   Amount vapor pressure is lowered depends on the amount of solute
   Raoult’s law quantifies the extent to which a nonvolatile solute lowers the vapor pressure of the
      solvent
          o Defining the variables: PA = v.p. of solution (with solute) due to vap. of solvent A
                                    P°A = v.p. of pure solvent A (without solute)
                                    XA = mole fraction of solvent A
                                    XB = mole fraction of solute
          o Formula used depends on info. given in problem:
                 If given info about solvent and asked only for v.p. of solution, use
                           PA = XAP°A
                 If asked for v.p. reduction (i.e. how much was v.p. lowered by adding a particular
                   solute), use
                           ∆PA = P°A – PA (gives difference between pure solvent and soln v.p.)
                 If given info about the solvent
                           ∆PA = P°A - XAP°A         substituted from PA = XAP°A
                           ∆PA = P°A(1-XA)           factoring out P°A

                     If given info about the solute
                              Recall: XA + XB = 1       so, 1-XA = XB
                              So:     ∆PA = P°A(XB)
                                      ∆PA = XB P°A
                                             Pract. Ex. 7

      An ideal solution is one that obeys Raoult’s law
      Real solutions show approximately ideal behavior when:
          o The solute concentration is low
          o The solute and solvent have similarly sized molecules
          o The solute and solvent have similar types of IMF
      Raoult’s law breaks down when the solvent-solvent and solute-solute IMF are much greater or
       weaker than solute-solvent IMF
      Can have ideal solution with two or more volatile components contributing to v.p.
          o Example:          volatile components A1 and A2 contribute to v.p. as follows
                              PA = XAP°A             PA1 = XA1P°A1

                             PT = PA + PA1 = XAP°A + XA1P°A1

Boiling-Point Elevation
    A nonvolatile solute lowers the vapor pressure of a solution See Fig. 13.22
    At the normal boiling point of the pure liquid the solution has a vapor pressure less than 1 atm
    Therefore, a higher temperature is required to reach a v.p. of 1 atm for the solution (∆Tb)
    The molal boiling-point elevation constant, Kb, expresses how much ∆Tb changes with molality
          o Kb is solvent-specific
                                   ∆Tb = Kbm
Freezing-Point Depression
    When a solution freezes, crystals of almost pure solvent are formed first
    Solute molecules are usually not soluble in the solid phase of the solvent
    Therefore, the triple point occurs at a lower temperature because of the lower vapor pressure for
       the solution See Fig. 13.22
    The melting-point (freezing-point) curve is a vertical line from the triple point
    Therefore, the solution freezes at a lower temperature (∆Tf) than the pure solvent
    The molal freezing-point depression constant, Kf, expresses how much ∆Tf changes with molality
           o Kf is solvent-specific
                                    ∆Tf = Kfm

                                            Pract. Ex. 8-9

Osmosis
   Semipermeable membranes permit passage of specific solution components
        o They often permit passage of water but not larger molecules or ions
        o Examples: cell membranes, cellophane
   Osmosis is the net movement of a solvent from an area of low solute concentration to an area of
     high solute concentration
        o Can also think of it as movement of solvent from an area of high solvent concentration to
            area of low solvent concentration
        o Movement occurs to establish an equilibrium solution concentration
   Consider a U-shaped tube with two liquids separated by a semipermeable membrane
        o Draw Fig. 13.23
        o One arm contains a dilute solution
        o One contains a concentrated solution made with the same solvent
        o If only the solvent can move through the membrane, the rate of movement of solvent from
            the dilute soln to the concentrated soln is faster than movement in the opposite direction
        o As solvent moves across the membrane, the fluid levels in the arms become uneven
        o Eventually the pressure difference due to the difference in height of liquid in the arms
            stops osmosis
   Osmotic pressure, π, is the pressure required to prevent osmosis
        o Osmotic pressure obeys a law similar in form to the ideal-gas law
                    πV = nRT             R = 0.0821 L∙atm/mol∙K
                    π = nRT = MRT        because n/V is molarity, M
                        V
        o Two solutions are said to be isotonic if they have the same osmotic pressure
        o Hypotonic solutions (more dilute) have a lower π, relative to a more concentrated solution
        o Hypertonic solutions (more conc) have a higher π, relative to a more dilute solution
   Everyday examples of osmosis – see book


Determination of Molar Mass Using Colligative Properties
    Any of the four colligative properties can be used to find molar mass by manipulation of variables.

                                            Pract. Ex. 11-12
Colligative Properties of Electrolyte Solutions
    Recall: Colligative properties depend on the total conc. of solute particles(either molecules or ions)
       in soln
    Ions of electrolytes dissociate in soln, thus increasing number/conc. of solute particles
    In a 0.100 m NaCl solution the total concentration of particles is actually 0.200 m, assuming
       complete dissociation
           o The expected freezing-point depression of 0.100 m NaCl is (0.200 m) x (1.86 °C/m) =
               0.372 °C
           o The actual freezing-point depression is found to be 0.348 °C
           o Ions in soln form ion-pairs in which oppositely charged ions associate with each other for a
               short time
           o This reduces the number (and consequently the concentration) of independent particles in
               soln
    The van’t Hoff factor (i) is a measure of the extent of dissociation of electrolytes in soln
           o It is determined by using the ratio of the actual value of a colligative property to the
               calculated value for a nonelectrolyte
                              i=             ∆Tf (measured)          .
                                  ∆Tf (calculated for nonelectrolyte)
           o See Table 13.5
           o The ideal (limiting value) of i is determined by # of ions per formula unit
           o Dilution affects the value of i for electrolytes
           o The more dilute the soln, the more closely i approaches ideal value because ion pairing
               decreases with dilution
           o Also, the lower the charges of the ions, the less pairing; again, i is closer to ideal value
                                             Pract. Ex. #59, p. 504

Read the Colloids Section
      Colloids (colloidial dispersions) are suspensions in which the suspended particles are larger than
       molecules but too small to separate out of the suspension due to gravity
      Colloidial particles scatter light (Tyndall effect)
      Hydrophilic colloids are water-loving
      Hydrophobic colloids are water-hating
      Many biological applications – read text

								
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