Aqueous Solutions by Y7xT503

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									                                    Aqueous Solutions
I. Solution                                                    Types of Mixtures Table 3, pg 404
        A. Homogenous - all samples are identical.
        B. Particle size is less than 1nm, cannot be filtered or settle out.
        C. Can be solid (metal alloys: bronze, steel, sterling silver), liquid (salt water), or gas (air)
        D. Two parts to a solution
                1. Solvent - does the dissolving (example: water - the universal solvent)
                2. Solute - the substance being dissolved
        E. Soluble - able to be dissolved.
II. Suspensions
        A. Heterogeneous – all samples are not identical
        B. Solute particle size is larger than 1000nm
        C. Particles are large enough to be filtered out (coffee grounds)
        D. Particles settle out after mixture stands undisturbed (sand in beach water)
        E. Scatter visible light. Like sun through dust in the air of a room.
III. Colloids
        A. Heterogeneous - all samples are not identical
        B. Solute particles are of medium size, between 1nm and 1000nm.
        C. Cloudy looking. Examples: glues, gelatins, paints, smoke, muddy water
        D. Can NOT be filtered and will NOT settle out.
        E. Tyndall Effect – transparent particles > 1nm scatter visible light in all directions.
                1. Brownian Motion - chaotic movement of solute particles.
                2. Like headlights in the fog
                3. Solution particles are too small. Suspensions are not transparent.
                4. Helps distinguish between a solution and a colloid (pg404)
IV. When two liquids try to mix....
        A. Miscible - two liquids that can dissolve each other. Example: water and alcohol
                1. “Like dissolves like”
                2. Polar solvents dissolve polar and ionic solutes (charges attract).
                    Hydrogen bonds & dipole interactions pull apart charged solute ions
                    Exceptions: Water will not dissolve BaSO4 or CaCO3 because the ionic bonds
                    holding the molecule together are stronger than the hydrogen bonds trying to
                    pull them apart.
                3. Nonpolar solvents dissolve nonpolar solutes.
                    London dispersion forces
                    Example: Nail polish remover dissolves Styrofoam.
                    Some nonpolar substances: gasoline, oil, fat, grease
          B. Immiscible - two liquids that cannot dissolve each other. Example: water and oil
                1. Need an emulsifying agent, like soap, to help mix immiscible solutions.
                Ex: mayonnaise (oil and vinegar with egg yolk…lecithin is emulsifying agent)
                2. Emulsions are a type colloid – Two liquids that normally will not mix.
                3. Solvent = the greater percentage & Solute = the lower percentage
V. Electrolytes - compounds that conduct an electric current when in solution.
        A. Necessary for life - conduct continuous flow of energy throughout body.
        B. Strength depends on different degrees of ionization, NOT concentration of solute.
        B. Strong Electrolytes - ionic compounds: charged ions separate completely when
                dissolved in water (aqueous) or molten (melted).
        C. Weak Electrolytes - aqueous polar compounds: only a fraction of solute exists as ions
                when dissolved in water - most ions remain bound in compound
        D. Nonelectrolytes - nonpolar molecules
                1. organic compounds such as alcohols and sugars
                2. compounds usually contain carbon.
                3. glucose and glycerol, methane, grease, gasoline
VI. Rate of Dissolution - how fast a solute goes into solution.
        A. Agitation - shake, mix, stir - get fresh solute in contact with solvent.
                1. Only increases how fast, NOT how much solute goes into solution. “Shake it.”
         B. Temperature – inc. temp, increase energy, increase force & frequency of collisions.
                1. Increases how fast AND how much solute a solution can hold. “Bake it.”
         C. Particle Size - the smaller the particle size, the greater the surface area of the particle
        exposed to the solvent. Sugar cube verses granules.
                1. Only increases how fast NOT how much solute goes into solution. “Break it.”
VII. Solubility - how much solute goes into solution (Figure 15, pg 414)
        A. Saturated solution - maximum solute in a given quantity of solvent at constant temp.
                1. At equilibrium. Appears clear!!
                2. Equilibrium - rate of dissolution (dissolving) = rate of crystallization
                3. Formulas -
                        a. solubility = xg solute / 100g solvent (Table 4 on pg 410)
                        b. mass solute = solubility of solute x mass of solvent
                 Example: How much KCl can be dissolved in 350g of H2O at 50˚C?
        B. Unsaturated solution - less than the maximum solute in a given amount of solvent at
                constant temp.
             1. Appears clear.
        C. Supersaturated solution - more solute than it can theoretically hold at given temp.
             1. Add solute when solution is hot and set aside to cool undisturbed.
             2. No un-dissolved solute - appears clear!!!!
             3. Crystallization of excess solute can be initiated by a single “seed” crystal.
                Example: Seeding rain clouds with AgI causes water vapor to condense and drop
                Questions: What could you do to make a saturated solution unsaturated?
                             What could you do to make an unsaturated solution saturated?
VIII. Factors that affect solubility (how much solute is able to be dissolved):
        A. Temperature
                1. Solid as solutes - increase temp, energy, and collisions, increases solubility
                        Example: Hot Tea and Sugar
                2. Gas as solutes - increase temp, energy, ability to escape, decreases solubility.
                        Example 1: Thermal Pollution - increase temp, increases O2 escaping into
                                        the air, increases death in lake.
                        Example 2: Open Soda - increase temp, carbon dioxide escapes faster, and
                                        soda goes flat faster.
        B. Pressure - increase pressure increase solubility of a gas!
                1. Henry’s Law - solubility of gas is directly proportional to pressure
                        above liquid. Example: Sealed coke, lots of pressure, keeps gas in
                        (soluble). Open soda, less pressure, gas escapes (less soluble
                      a. Effervescence: the quick release of gas particles from a solution
                      b. Formula:        S1 = S2
                                          P1 P2
                          Example: Scuba diving bends: Nitrogen gas was absorbed by the blood
                          at deep ocean pressure.
                          Question: If the solubility of the gas in water is 0.77 g/L at 350kPa,
                          what is its solubility, in g/L at 100 kPa? (Temp. is constant at 25C.)
IX. Water of Hydration - the water in a solid crystal
      A. Hydrate - crystal compounds containing water - looks dry
              Example: CuSO4-5H2O = 1molecule CuSO4 is connected to 5molecules H2O
      B. Anhydrous - dry, no water in crystal, dehydrated, powdery
      C. Efflorescent - the process of losing water, becoming dry.
              Example: Wintergreen Certs - bite with mouth open in dark closet, see sparks.
      D. Hygroscopic - compound absorbs water from air - may dissolve at RT
              Example: NaOH - seal container tight to avoid ruining.
X. Molar enthalpy of dissolution - amount of energy released or gained by solute as it dissolves.
      A. ΔHsol (solution)      (Table 5, pg 416)

								
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