Solution is a homogeneous mixture (uniform composition)
Solute = what gets dissolved (Example: NaCl)
Solvent = what the solute dissolves in (Example: H2O)
Solubility
A solution can be :
saturated = the maximum amount of solute has dissolved in the solvent (equilibrium has been reached)
unsaturated = more solute can be dissolved
supersaturated = more than the maximum amount of solvent has dissolved
Concentration
There are multiple ways to express concentration of a solution.
These are:
Problem: What is the mole fraction of water in 100. g of 89% (by mass) ethanol, C2H5OH?
Solution:
mass of ethanol = 89g
mass of water = 100-89=11g
moles of water: 11g(1mol/18g) = 0.61 mol water
moles of ethanol: 89g(1mol/46g) = 1.9mol ethanol
Mole fraction of water = moles of water/ moles of solution
mole fraction of water = 0.61/(0.61+1.9)=0.24
Problem: A solution of NaCl in water has a concentration of 20.5% by mass.
What is the molal concentration of the solution?
Solution : Assume 100 g of solution (since we are given % only)
Solute: NaCl
Solvent: water
Mass of NaCl = 100g(.205)= 20.5g NaCl
Mass of water = 100-20.5g = 79.5 g water(1kg/1000g) = .0795kg
Moles of NaCl = 20.5g/58.44g/mol =0.351mol
Molality= moles of solute/ kg of solvent
molality = 0.351mol/.0795kg= 4.41m
Colligative Properties
Colligative properties are the physical changes that occur when solute is added to pure solvent. Colligative properties depend only on the identity of the solvent and the concentration of the solute. They do not depend on the identity of the solute.
Colligative properties are:
vapor-pressure lowering
boiling point elevation
freeezing point depression
osmotic pressure
Vapor Pressure Lowering
When a solute is added to pure solvent, the vapor pressure of solution will be lower than the vapor pressure of pure solvent.
Pa = Pa°Xa
Where P°a= vapor pressure of pure solvent, Xa = mole fraction of solvent, Pa= vapor pressure of solution
Boiling point elevation
When a solute is added to pure solvent, the boiling point of solution will be higher than the boiling point of pure solvent.
ΔTb =iKb*m, where i=number of ions for ionic compounds or 1 for molecular compounds (all nonmetals)
ΔT = change in boiling point (boiling point of solution - boiling point of pure solvent), Kb = boiling point elevation constant and m is molality (moles of solute/kg solvent)
Freezing point depression
When a solute is added to pure solvent, the freezing point of solution will be lower than the freezing point of pure solvent.
ΔTf =iKf*m, where i=number of ions for ionic compounds or 1 for molecular compounds (all nonmetals), ΔT = change in freezing point (freezing point of solution - freezing point of pure solvent), Kf = freezing point constant and m is molality (moles of solute/kg solvent).
Problem: An aqueous solution is 0.0222 m glucose. What is the freezing point of this solution?
ΔTf =iKf*m
Glucose is molecular compound and therefore i=1. We can look up the freezing point constant for water, which is 1.86C/m
ΔTf =1(1.86C/m)(0.0222m)= 0.0413C
This is the change if freezing point.
Freezing point of pure water is 0C.
0C-0.0413C = -0.0413C
Osmotic pressure is the pressure that is needed to stop osmosis (solvent flow through a semipermeable membrane to equalize the solute concentrations on both sides of the membrane).
π= iMRT where I =number of ions for ionic compounds or 1 for molecular compounds (all nonmetals), M is molarity (moles of solute/L of solution), R is the ideal gas constant (0.082 L*atm/(K*mol) and T is temperature in Kelvin.
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