1. 2Na(s) + 2H2O  2NaOH(aq) + H2(g)
Solutions
Solution - A homogeneous mixture of substances that consists of a solvent and one or more solutes.
Solvent - Generally the most abundant substance in a solution.
Solute - The impurity in a solution and less abundant than the solvent.
Dissolution
Dissolution - The formation of a solution by the addition of a solute to a solvent. This may occur with or without reaction.
Solvation - Dissolution without the occurrence of a reaction.
2Na(s) + 2H2O  2NaOH(aq) + H2(g)
H2O
NaCl(s)  NaCl(aq)

2. Ease of Dissolution Two factors that affect the ease of dissolution:
1) The change in energy (enthalpy)
2) The change in order (entropy)
Dissolution is favored by exothermic processes that increase disorder
Interactions that favor dissolution
1) weak solute-solute intermolecular interactions
2) weak solvent-solvent intermolecular interactions
3) strong solvent-solute intermolecular interactions

3. (page 544)
Enthalpy of solution - Hsol - The heat involved in the solvation of a solute.
Generally dissolution accompanies an increase in disorder so, endothermic dissolutions can occur spontaneous depending on the magnitude of Hsol.

4. Solvation of solids
Crystal lattice energy (solute-solute) - The energy change involved in the formation of one mole of formula units in a crystalline state from it’s particles in the gaseous state. (always exothermic)
Na+(g) + Cl-(g) NaCl(s) + heat
Hsol = (heat of solvation) - (crystal lattice energy)
Hydration - Solvation where the solvent is water.
The larger the ion the more waters hydrate it (4-9 waters, 6 average)

5. Dissolution of liquids
miscibility - the ability of one liquid to dissolve in another.
Polar liquids tend to dissolve in polar liquids
Non-polar liquids tend to dissolve in non-polar liquids
The dissolution of liquids is generally exothermic

6. Dissolution of gases in liquids
Gases that are capable of hydrogen bonding or ionize are generally soluble in water (HF, HCl, HBr ...)
Some non-polar gases are slightly soluble in water
O2 due to dispersion forces between O2 and H2O
CO2 due to reaction CO2(g) + H2O  H2CO3
H2CO3  HCO3- + H+
HCO3-  CO H+
An increase in pressure increases
the solubility of a gas

7. Effect of temp. on solubility
Exothermic: reactant  product + heat
Endothermic: reactant + heat  product
negative Hsol (exothermic) solubility decreases with increasing temp.
positive Hsol (endothermic) solubility increases with increasing temp.
gas dissolutions in liquid are exothermic
soda goes flat at room temperature faster than in the refrigerator
(Don’t put your fish tank near a window... )

8. m = mol of solute/kg of solvent (mol/kg)
Molality
molality - moles of solute per 1 kg of solvent.
m = mol of solute/kg of solvent (mol/kg)
Why molality instead of molarity?
Molality is used when there is a change in temperature involved because it is temperature invarient.
Molarity (moles/L) is dependent on the volume which changes with a change in temperature.
What is the molality of a solution prepared from 50 g of sucrose (C12H22O11) and 117g of water?
m = (50g sucrose * 1mol/342g sucrose)/0.117 kg H2O = 1.25 m

9. Mole Fraction Mole Fraction of A - XA = moles of A/total moles
dimensionless quantity
What is the mole fraction of CH3OH and H2O in a solution composed of 128 g CH3OH and 108 g H2O?
mol CH3OH = 128gCH3OH * (1 mol/32.0gCH3OH) = 4.00 mol CH3OH
mol H2O = 108 g H2O * (1 mol/18.0 g H2O) = 6.00 mol H2O
XCH3OH= 4.00 mol/(4.00 mol mol) = 0.400
XH2O = 6.00 mol/(4.00 mol mol) = 0.600
The mole fraction of all species in a solution add up to 1