Solution

							Chemistry 20
Solutions (15)

Definitions:

Solution:                       a homogenous mixture of two or more compounds that are uniformly
                                distributed
Heterogenous Mixture:           observable segregation of the components of the mixture
Solvent:                        substance in a solution in greater quantity
Solute:                         substance in a solution in lesser quantity

Solutions
     Solutions can be solid, liquid, or gaseous.
           o Ex: Gaseous solution: air
           o Ex: Solid solution: alloy (metal mixture) – brass
     In liquid solutions, if water is the solvent in a liquid solution, the solution is called aqueous
     A solution that can dissolve more solute is called unsaturated
     At some point, a solution cannot dissolve any more solute and the solution is called saturated
           o At this point, any excess solute will remain undissolved – this means an equilibrium or
               balance has been reached. We represent it like this:

                        Solvent + solute  solution + solute

                 (forward reaction balanced by reverse reaction)

                 Ex. Saturated salt solution:     NaCl  Na+ + Cl-
       The limit of specific solute that will be dissolved in a given amount of solvent is known as
        solubility
            o Ex. At 0.0°C, 35.7 g of NaCl will dissolve in 100.0 g of water
       The rate at which a solution is created (when a solid is dissolved in water) can be increased by:
            o Raising the temperature of the solution
            o Stirring the solution
            o Increasing the surface area of the solute, or crushing the solute
       This happens because:
            o Dissolving takes place at the surface of the solute; increased surface area = increased
                 dissolving
            o Increased temperature = increased energy of particles (increased molecular motion)
            o Stirring brings fresh solute to the surface = increased solution rate
       There are some liquid-liquid solute/solvent combinations that form solutions in any proportion.
        There is no limit to how much of each you can add.
            o These solutions are termed perfectly miscible
                      Ex: Ethanol and water
       Substances that do not dissolve in one another are immiscible
            o Ex: Oil and water

Concentration
    Remember: Molarity = moles of solute/litres of solution
    In chemistry it is important to know the concentration of individual ions.
           o Ex:      K2SO4  2K+ + SO4-2
    So each mole of potassium sulfate produces 2 mole of potassium ions and 1 mole of sulfate ions.
    By convention, we express the concentration of a substance in mol/L by placing square brackets
      around each formula.
           o Ex: If we have 0.5 M solution of H2SO4 it looks like this:       [H2SO4] = 0.5 M
    Example Question:
           o Calculate [Cu+2] and [Cl-] if 10.0 g of CuCl2 dissolves in 500.0 mL of solution.
Factors Affecting Solubility
a) Nature of Solvent and Solute
         Solvation: interaction between solvent and solute particles
                   Ex: Hydration: surrounding an ion by water molecules (polar)
                   Ex: Dissociation: decomposing a crystal into ions in solution (KCl  K+ + Cl-)
         Solvation also occurs between polar compounds and polar solvents BUT oil (non-polar) will
            not interact with water (polar) while salt (ions) will not interact with oil (non-polar)
         On the other hand, non-polar substances readily dissolve non-polar substances
         This leads to the saying; LIKE DISSOLVES LIKE

b) Effect of Temperature
        Gases dissolve in liquids more at decreased temperature than at increased temperature
        Liquids and solids dissolve more at increased temperature
        A solution becomes supersaturated when a solution dissolves more solute than normal at an
           increased temperature
c) Effect of Pressure
        Change in pressure has practically no effect on the solubility of solids and liquids
        Gases are more soluble in liquid as the pressure is increased (ex carbonated beverages)
d) Heat of Solution
        Energy involved in the solvation process is called heat of solution
        Exothermic: heat is released as a solute dissolves
                  Solute + solvent  saturated solution + heat
        Endothermic: heat is absorbed as a solute dissolves
                  Solute + solvent + heat  saturated solution
        If heat of solution is negative, the solution is exothermic
        If heat of solution is positive, the solution is endothermic

Precipitation and Net Ionic Equations
    Net ionic equations show only the species of reactants and produces that produces a visible
        produce (precipitate)
    To write an ionic equation, you must know which compounds disassociate into ions in solution
        and which product forms a precipitate.
    Solubility can be determined from the Solubility Table
    To write ionic equations, 3 basic steps are involved:
            o Write the balanced equation showing all reactants and products
            o Show the reactant and product ions in the equation
            o Eliminate the spectator ions – those ions that remain unchanged during the reaction (they
                don’t form a precipitate) – and write the net ionic equation

Qualitative Analysis
    A process used to determine what substance is present in a certain solution
    Using reactants, you can determine what ions are present and which ions can be removed by
        precipitation
    Need to use the Solubility Table for this.