AP Chemistry - Notes : Chapter 11 Mr. Niedenzu Providence High School
A. Properties of Solutions - homogeneous mixtures
- may be a gas in a gas, solid in solid, gas in liquid, etc.
1. Solution composition
a. Qualitative
- dilute - little solute as compared to the amount of solvent
- concentrated - a lot of solute compared to the amount of solvent
b. Quantitative
- molarity (M)- mol solute/Liter of solution
- mass percent
- mole fraction
- molality(m) - mol solute /Kg solvent
- normality (N) - equivalents/Liter of solution
- equivalent (eq) :
- for an acid-base reaction an equivalent is the mass of an acid or a base that will provide one mole of H+ or OH- ions
- for a redox reaction an equivalent is that mass of an oxidizing or reducing agent that can accept or donate one mole of electrons
2. The Energies of Solution Formation
a. Cardinal rule : "Like dissolves like."
b. Three steps to solution formation :
- Step 1 : break up of the solute into individual components (ions, atoms or molecules)
- endothermic - requires energy to overcome attractive forces of particles (ΔH1>0)
- Step 2 : separation of the solvent
- endothermic - intermolecular forces must be overcome(ΔH2>0)
- Step 3 : solute and solvent interact to form solution
- often exothermic - bonding of solute and solvent results in a more stable, lower energy arrangement(ΔH3<0)
c. enthalpy (or heat) of solution = ΔHsoln = ΔH1 + ΔH2 + ΔH3
- if ΔHsoln is negative (exothermic), energy will be released (magnitude of ΔH3 is greater than ΔH1 + ΔH2).
- if ΔHsoln is positive (endothermic), energy is absorbed (ΔH3 is not exothermic or magnitude of ΔH3 is less than ΔH1 + ΔH2).
Summary of Energies and Solution Formation :
ΔH1 / ΔH2 / ΔH3 / ΔHsoln / ResultPolar solvent, polar solute / Large / Large / Large, negative / Small / Solution forms*
Polar solvent, nonpolar solute / Small / Large / Small / Large, positive / No solution forms
Nonpolar solvent, nonpolar solute / Small / Small / Small / Small / Solution forms*
Nonpolar solvent, polar solute / Large / Small / Small / Large, positive / No solution forms
*Note "Like dissolves like," substances with similar intermolecular forces tend to be soluble in each other.
d. Entropy - a measure of the disorder of a system - along with enthalpy changes determines if a solution will form or not
- factors favoring solution formation :
- decrease in enthalpy (negative value, exothermic : ΔH3 > (ΔH1 + ΔH2))
- increase in entropy (positive value, increase in disorder)
- factors inhibiting solution formation :
- increase in enthalpy (positive value, endothermic : ΔH3 < (ΔH1 + ΔH2))
- decrease in entropy (negative value, decrease in disorder)
3. Factors Affecting Solubility
a. Molecular structure - determines polarity
e.g. - fat soluble vitamins(A,D,E and K) are nonpolar and so dissolve in nonpolar oils and fats
- water soluble (B's and C) vitamins are polar and dissolve in water
- An important implication of this is that we can store fat soluble vitamins in our fatty tissues. We ca also overdose on fat soluble vitamins since they can be stored. It would be very difficult to overdose on water soluble vitamins since they are easily excreted in urine, which also prevents their being stored.
b. Pressure effects - pressure effects the solubility of gases :
- increased pressure increases the solubility of gases - really an equilibrium effect. An increase in pressure above a solvent will concentrates the gas particles above the solvent. To reach equal concentrations in and out of the solvent more gas particles will move into the solvent to form a more concentrated solution.
- decreased pressure decreases the solubility of gases - opposite of above
- Henry's Law : States that the solubility of a gas is directly proportional to the pressure of the gas above the solution. (Assuming no chemical interaction between the gas and the solvent). (P = kC, where : P = pressure, k = solubility constant, C = concentration of the dissolved gas).
- pressure has little effect on the solubility of solids and liquids
** To remember the above think of a bottle of soda. If you open it (decreases pressure) the gases will leave (become less soluble), but the sugars (solids) will not leave(solubility unaffected).
c. Temperature Effects (for aqueous solutions)
- the solubility of gases decreases with increases temperature and vice versa (thermal pollution of bodies of water decreases oxygen content and can lead to a die off of aquatic animals).
- the solubility of solids generally increases with increase in temperature (sodium sulfate and cerium sulfate become less soluble as temperature rises).
4. The Vapor Pressure of Solutions :
- The presence of a nonvolatile solute lowers the vapor pressure of a solvent. It does so by diluting the solvent, reducing the numbers of solvent particles at the surface and thereby reducing the numbers of solvent particles escaping and creating vapor pressure. The vapor pressure of a solution is then directly proportional to the mole fraction of the solvent present as stated in Raoult's Law : Psoln = XsolventPºsolvent, and illustrated below.
Raoult's law can be used to count molecules, or determine the number of moles present in a solution, and thereby determine the molar mass of a substance. A mass of the solute is first dissolved in a given amount of solvent(number of moles known). The vapor pressure can then be measured and the mole fractions determined. The moles of solute is then calculated and molar mass determined by dividing the mass of solute by moles of solute.
Raoult's law can also be used to determine the number of particles a solute dissociates into. If one mole of a substance is put into solution and the vapor pressure is lowered by an amount that indicates three moles of solute then the solute dissociates into three particles.(e.g. MgCl2(s) + H2O(l) à Mg2+(aq) + 2 Cl-(aq))
- Nonideal solutions - liquid - liquid solutions where both parts of the solution are volatile
- modified Raoult's law : Psoln = PA + PB = XAPºA + XBPºB
- a liquid-liquid solution that obeys Raoult's law is considered an ideal solution
- in an ideal solution the solvent-solvent, solute-solute and solvent-solute interactions have to be very similar
- in a solution where the solute-solvent bonds are greater than solute-solute or solvent-solvent bonds :
- energy will be released (ΔHsoln is negative, exothermic formation of solution)
- there will be a negative deviation from Raoult's law (lower vapor pressure than expected)
- in a solution where the solute-solvent bonds are weaker than the solute-solute or solvent-solvent bonds :
- energy will be absorbed (ΔHsoln is positive, endothermic formation of solution)
- there will be a positive deviation from Raoult's law (higher vapor pressure than expected)
Interactive forces between solute(A) and solvent (B) / ΔHsoln / ΔT for Solution Formation / Deviation from Raoult's Law / ExampleAóA, BóB =AóB / Zero / Zero / None (=ideal solution) / Benzene and toluene (similar in size and polarity)
AóA, BóB <AóB / Negative (exothermic) / Positive / Negative / Acetone and water
AóA, BóB >AóB / Positive (endothermic) / Negative / Positive / Ethanol (polar) and hexane (nonpolar)
5. Boiling-Point Elevation and Freezing-Point Depression
a. Colligative properties - properties that depend on the number of solute particles present, not the type of particles present
- include
- boiling point elevation
- freezing point depression
- osmosis
b. Boiling-point elevation - a nonvolatile solute raises the boiling point of a solvent proportionally to the amount of solute present
- ΔTb = Kbmsolute
where : Kb = molal boiling point constant - specific for each liquid
msolute - molality of solute
c. Freezing-point depression - a nonvolatile solute lowers the freezing point of a liquid
- ΔTf = Kfmsolute
where : Kf = molal boiling point constant - specific for each liquid
msolute - molality of solute
6. Osmotic Pressure
a. Osmosis - the flow of a solvent through a semipermeable membrane (net movement of solvent is always toward the side with the greater concentration of solute)
b. Osmotic pressure - the pressure created by the movement of the solvent through the membrane
- equal to the minimum external force needed to stop the flow of the solvent
- can be used to indicated the number of particles a solute dissociates into
- can be used to determine the molar mass of a substance - very useful for this since a small amount of solute creates a large pressure
- π = MRT where : π is osmotic pressure
M is molarity
R is the gas law constant(.08206 L · atm/K · mol)
T is the Kelvin temperature
c. Dialysis - phenomenon in which small solute particles as well as solvent particles pass through a semipermeable membrane
d. isotonic solutions - solutions with equal osmotic pressures
e. hypertonic solution - a solution with a higher solute content and therefore a lower osmotic pressure than a reference solution (often cytoplasm)
- hypertonic solutions can cause crenation (shriveling due to water loss) in red blood cells
f. hypotonic solutions - a solution with a lower solute content and therefore a higher osmotic pressure than a reference solution (often cytoplasm)
- hypotonic solutions can cause hemolysis (bursting due to too much water present) in red blood cells
g. Reverse osmosis - the forcing of water through a semipermeable membrane against the flow of osmosis by the application of an external pressure greater than the osmotic pressure
- used in the desalination of seawater to obtain freshwater
7. Colligative Properties of Electrolyte Solutions
a. van't Hoff factor - factor equating moles of solute with moles of ions present
- e.g. The van't Hoff factor for magnesium nitrate ( Mg(NO3)2)would be three since three moles of ions would be formed in solution from one mole of magnesium nitrate.
- the equation for changes in the boiling or freezing-point of a solution containing an electrolyte becomes :
ΔT = imK
Where : i is the van't Hoff factor
m = molality of solute
K = freezing-point or boiling-point constant
- the equation for osmotic pressure becomes :
π = iMRT
b. Ion Pairing - in any electrolyte solution there is some ion pairing - the joining of oppositely charged ions due to electrostatic attraction. This will reduce the expected number of effective solute particles causing colligative effects. The greater the charge on an ion the greater its tendency to pair in solution.
8. Colloids
a. characteristics of colloids
- permanent mixtures (will not settle out over time)
- have a particle size intermediate to solutions and suspensions
- exhibit the Tyndall effect
- permanence in mixture is due to the electrostatic repulsion of the colloidal particles (they all have the same charge)
- structure of colloidal particles - surrounded by layers of like charged particles :
b. Coagulation - the destruction of a colloid by destroying electrostatic repulsion by :
- heating - speeds up particles and collisions remove outer layers of ions
- addition of an electrolyte- removes layers of ions
c. Examples of colloids :
Example / Dispersing Medium / Dispersed Substance / Colloid TypeFog, aerosol sprays / Gas / Liquid / Aerosol
Smoke / Gas / Solid / Aerosol
Whipped cream, soap suds / Liquid / Gas / Foam
Milk, mayonnaise / Liquid / Liquid / Emulsion
Paint / Liquid / Solid / Sol
Butter, cheese / Solid / Liquid / Solid emulsion