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Solutions(3)

VIEWS: 22 PAGES: 63

									Solutions
        Occur in all phases
 The  solvent does the dissolving, usually
  is the substance that is present in the
  greatest amount.
 The solute is dissolved.
 There are examples of all types of
  solutes dissolving all types of solvent (9
  types).
 We will focus on aqueous solutions.
    Formation of a Solution
 When   1 or more substances disperses
  uniformly throughout a solution.
 We talked about intermolecular forces,
  they can also exist between solute and
  the solvent particles in a solution.
 For example, If you were to dissolve
  solid KCl in water, it would readily
  dissolve because the attractive
  interactions between the ions and the
 Polar water molecules are strong and
  overcome the lattice energy holding KCl
  together. The force of the ion-dipole
  attraction is strong enough to pull the
  ions out of the crystalline structure.
 Solvation, when the particles of the
  solute are completely surrounded by the
  solvent. (called hydration when the
  solvent is water)
         Forming a solution
 To  form a solution, not only do we need
  to break apart the bond (NaCl), but
  between water molecules also so there
  is room for the ions to fit in.
 The enthalpy change is made of 3
  energies.

 ΔHsoln= ΔH1 + ΔH2 + ΔH3
 Energy of Making Solutions
 Heat of solution ( DHsoln ) is the energy
  change for making a solution.
 Most easily understood if broken into
  steps.
 1.Break apart solvent
 2.Break apart solute
 3. Mixing solvent and solute
     1. Break apart Solvent
 Haveto overcome attractive forces.
 DH1 >0

2. Break apart Solute.
 Haveto overcome attractive forces.
 DH2 >0
3. Mixing solvent and solute
   DH3 depends on what you are mixing.

 If the molecules are attracted to each
  other DH3 is large and negative (ion dipole).
 If the molecules are not attracted to each
  other, then DH3 is small and negative (non
  polar solute- polar solvent).
   Exothermic/ Endothermic
 When   calculating the enthalpy of
  solution, the sum can be + or -,
  depending on substances.
 (-) exothermic, the solution would get
  warmer
 (+) endothermic, the solution would get
  colder
     Sizeof DH3 determines whether a
     solution will form   Solute and
                              Solvent
                                        DH3
                             DH2
E                  Solvent              Solution
n
e
r                 DH1                   DH3
      Reactants
g
y
                                        Solution
        Ice packs, heat packs
 (-) heat of solution, solution gets warm
 (+) heat of solution, the solution gets
  cool
            Spontaneity
 Exothermic   processes are usually
  spontaneous
 A solution will not form if it is too
  endothermic.
 NaCl would not dissolve in gasoline,
  gasoline is non-polar, and the weak
  forces you would get between the non-
  polar gas and NaCl would not be enough
  energy to break apart the bond in NaCl
           Like dissolves Like
 That’s where this saying came from.
 In order to understand it, you must
  understand the relative strength of
  bonds.
 When non-polar substances are mixed,
  they are all London Dispersion Forces,
  so there is no net energy change,
  however this is also a spontaneous
  process?
 There must be something besides
 energy then that accounts for whether
 a reaction will proceed spontaneously or
 not. Two things: Enthalpy and Entropy
       Types of Solvent and
             solutes
 If DHsoln is small and positive, a solution
  will still form because of entropy.
 There are many more ways for them to
  become mixed than there is for them to
  stay separate.
            Spontaneity
 Two  ways a reaction will proceed
  spontaneously:
 1) Exothermic processes are usually
  spontaneous, the change tends to lower
  the energy of the system
 2) Processes in which the disorder of
  the system increases also tend to occur
  spontaneously.
               In sum
 The formation of solutions is favored by
 an increase in disorder. A solution will
 form unless the solute-solute or
 solvent-solvent interactions are too
 strong compared to the solute-solvent
 interactions.
 Physical / Chemical Solutions
 Becareful to distinguish whether the
 formation of a solution is chemical or
 physical. If you can recover the original
 substance, than it is a physical change
 (dissolving in water).
    Structure and Solubility
 Water  soluble molecules must have
  dipole moments -polar bonds.
 To be soluble in non polar solvents the
  molecules must be non polar.
       Saturated Solutions
 When    a solution is in contact with an
  undissolved solute, two opposing
  reactions are occurring, dissolving, and
  crystallization.
 When the rate is equal, equilibrium is
  established, there is no further
  increase in the amount of solute that
  will dissolve.
 The  system is said to be saturated
 Adding more solute to the system will
  not result in an increase in the
  concentration of the solution, the solute
  particles will remain undissolved
 The amount of solute needed to form a
  saturated solution in a given quantity of
  solvent is known as the solubility of that
  solvent.
 Unsaturated   solutions have the capacity
  to hold more solute
 Supersaturated solutions are sometimes
  possible to form. When dissolve the
  substance in hot solvent and cool it, all
  of the solute may remain dissolved even
  though the solubility has decreased with
  the decrease in temp.
    Supersaturated solutions
 Extremely  unstable
 Adding a seed crystal will result in the
  excess solute crystallizing out of the
  solution.
  Factors affecting solubility
 Natural  tendency to move towards
  disorder
 Attraction between solute and solvent,
  the stronger the attraction between
  solute and solvent, the greater the
  solubility of the solute in that
  substance.
       Miscible/ immiscible
 Liquids that mix in all proportions are
  said to be miscible
 Liquids that do not mix are immiscible
 (Hydrocarbons do not mix in water)
    Solubility of alcohol in water
 Alcohols have the OH end which is polar
 Solubility of alcohols decrease with increasing
  mass because the chain is becoming more like
  a hydrocarbon
 However if the number of OH groups
  increases, the solubility will increase
 Network solids are not soluble in either polar
  or non polar solvents because of the strong
  bonding forces within the solid
         Pressure effects
 Changing  the pressure doesn’t effect
  the amount of solid or liquid that
  dissolves
 They are incompressible.
 It does effect gases, the solubility of a
  gas in any solvent is increased as the
  pressure over the gas is increased.
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.
 The  gas is at
  equilibrium with the
  dissolved gas in this
  solution.
 The equilibrium is
  dynamic.
 If you increase the
  pressure the gas
  molecules dissolve
  faster.
 The equilibrium is
  disturbed.
 The system reaches a
  new equilibrium with
  more gas dissolved.
 Henry’s Law.
P= kC
Pressure = constant x
         Concentration
                  of gas
                  Try
 Calculate the concentration of CO2 in a
 soft drink that is bottled with a partial
 pressure of 2.0 atm over the liquid at
 25°C. The Henry’s law constant for CO2
 in water at this temperature is
 3.1 x 10-2 mol/L-atm.
      Temperature Effects
 Increased   temperature increases the
  rate at which a solid dissolves.
 Usually it will increase the amount of
  solid that dissolves.
20   40   60   80   10
                    0
       Gases are predictable
 As temperature
  increases, solubility
  decreases.
 Gas molecules can
  move fast enough to
  escape.
              Concentration
 Dilute-   small concentration of solute

 Concentrate-large    concentration of
 solid
         1. Quantitatively
 Mass  percentage = mass solute x 100
                      mass solution
Very dilute in ppm = mass solute x 106
                     mass solution
A solution of 1 ppm = 1g/ million g

Really dilute ppb= mass solute x 109
                   mass solution
           2. Mole Fraction
          Moles component
    total moles of all components

Useful when dealing with gases
             3. Molarity

 Molarity   = moles of solute
                liters of solution
             4. Molality
 Molality =     moles of solute
         Kilograms of solvent

 Molarity  changes with temperature,
  molality does not.
 Usually use molality when a solution is
  to be used over a range of
  temperatures.
                 Problems
 Calculate the mass percentage of Na2SO4 in a
  solution containing 14.7 g of Na2SO4 dissolved
  in 345 g of water.
 Calculate the mole fraction of methanol
  (CH3OH) when 7.5 grams of it is dissolved in
  245 g of water.
 Calculate the molarity of a solution when 10.5
  g of KCL is dissolved in 250 ml of solution.
 Calculate the molality of a solution when 13 g
  of benzene (C6H6) is dissolved in 17 g of CCl4.
        Colligative properties
 Add  ethylene glycol to car radiators as
  antifreeze
 Add salt to water to lower the freezing
  point
 Also raises the boiling point.
 Reduces the vapor pressure
 Reduces osmotic pressure
       Colligative Properties
            properties depend only on
 Colligative
 the number - not the kind of solute
 particles present
          Vapor Pressure
 Vapor pressure is the pressure exerted
  by a gas above the liquid in a closed
  system at equilibrium
 A non-volatile substance has no
  appreciable vapor pressure
 A volatile substance would
     Vapor Pressure of Solutions
A nonvolatile solute lowers the vapor
 pressure of the solution.
             Raoult’s Law:
Psoln    = csolvent x Psolvent
 Vapor  pressure of the solution =
     mole fraction of solvent x      vapor
  pressure of the pure solvent
 Doesn’t depend on the type of particle,
  only the concentration. Also, only works
  with non-volatile solutes otherwise they
  would be contributing to the pressure.
  (assume non-electrolytes)
             Real Solutions
 Raoult’s Law- for ideal solutions
 A solution only approaches ideal
  conditions when the solute
  concentration is low, and when the
  solute and solvent have similar molecular
  sizes and intermolecular forces.
     Boiling point Elevation
 Because   a non-volatile solute lowers
  the vapor pressure it raises the boiling
  point.
 The equation is: DT = Kbmsolute
 DT is the change in the boiling point
 Kb is a constant determined by the
  solvent.
 msolute is the molality of the solute
   Freezing point Depression
 Because   a non-volatile solute lowers the
  vapor pressure of the solution it lowers
  the freezing point.
 The equation is: DT = Kfmsolute
 DT is the change in the freezing point
 Kf is a constant determined by the
  solvent
 msolute is the molality of the solute
1
atm
      Vapor Pressure
      of pure water




                       Vapor Pressure
                       of solution
1
atm




      Freezing and
      boiling points
      of water
1
atm




      Freezing and boiling
      points of solution
1
atm




      DTf   DTb
    Electrolytes in solution
 Since colligative properties only
  depend on the number of molecules,
  Ionic compounds should have a bigger
  effect.
 When they dissolve they dissociate.
 Individual Na and Cl ions fall apart.
 1 mole of NaCl makes 2 moles of ions.
 1mole Al(NO3)3 makes 4 moles ions.
 Electrolyteshave a bigger impact on
 omelting and freezing points per mole
 because they make more pieces.
                  Try
 Automotive  antifreeze consists of
  ethylene glycol C2H6O2, a non-volatile
  nonelectrolyte. Calculate the boiling
  point and freezing point of a 25. mass
  percent solution of ethylene glycol in
  water.
 Kb= .52°C/m
 Kf= 1.86°C/m
             Try again!
     the following solutions in order of
 List
 their expected freezing points: 0.05 m
 CaCl2, 0.15 m NaCl, 0.10 m HCl, 0.05
 HC2H3O2, 0.10 m C12H22O11
               Osmosis
 Occurs through a semi permeable
 membrane. A membrane that usually
 allows the migration of small water
 molecules and not larger solute
 molecules.
 The net movement of solvent is always
 toward the solution with a higher solute
 concentration.
        Osmotic Pressure
 The pressure required to prevent
  osmosis (Fig 13.23).
 Osmotic Pressure (π)
 Π= (n/V) RT=MRT


 M=molarity
 R= gas laws constant
 T= temp (K)
 Isotonic-  when two solutions of identical
  osmotic pressure are separated by a
  semi-permeable membrane, no osmosis
  will occur
 Hypotonic- when one solution has lower
  osmotic pressure compared to the other
 The more concentrated one would be
  hypertonich
 People who eat salty foods retain water
 Cucumbers shrivel up in a salt water
  solution to make pickles
 Osmosis is spontaneous.


 Active   transport-opposite
                Try 2!
 The average osmotic pressure of blood
 is 7.7 atm at 25°C. What concentration
 of glucose will be isotonic with blood?



 What  is the osmotic pressure at 20°C
 of a 0.0020 M sucrose C12H22O11
 solution?
               Colloids
 Somewhere   between a solution and a
  heterogeneous mixture
 Does not separate upon standing
 Diameter of 10-2000 A
 Scatter light- Tyndall Effect
 When colloids are dispersed in water
  they can be hydrophyllic or hydrophobic

								
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