Heterogeneous vs. Homogeneous

A homogeneous mixture of 2 or more substances in a single phase.
 Solvent: the dissolving medium
 Solute: the substance which is dissolved
 Types of Solutions: (particle size is small, 0.01-1nm)
 oxygen in nitrogen, carbon dioxide in water
 water in air, alcohol in water, mercury in silver/tin
 sugar in water, zinc in copper (brass)


Suspensions:
Particles in a solvent are so large, >1000 nm, they will settle out unless stirred.
Colloids:
Particles are between 1-1000 nm. They are small enough to be suspended throughout the solvent.
 TABLE 13.2
 Tyndall Effect: scattering of light by colloidal particles.


Electrolytes: Dissolve in water to give a conducting solution.

Ex: NaCl

Nonelectrolytes: Dissolves in water to give solution that does NOT conduct eleectricity.

Ex: C
12
H
22
O
11

1. Factors Affecting the Rate of Dissolving:

Surface Area

Agitation

Heat
2. Solubility: the amount of a substance required to form a saturated solution at a given temp.




Saturated: contains
amount of dissolved solute
Unsaturated: contains
solute than saturated solnution.
Supersaturated: contains
solute than a saturated solution under the
conditions.


Predicts whether one substance will dissolve another:

Polarity of molecule

Intermolecular forces between solute/solvent.
 Hydration: Solution process with water as solvent.
Ionic compounds are not normally soluble in nonpolar solvents
(CCl
4
). Why??


Immiscible: Liquid solute/solvent are NOT soluble in each other.
Ex: Oil/Water

Miscible: Liquid solute/solvent are soluble in each other . Ex: Oil/Gasoline

Nonpolar molecules exert no strong attractive/repulsive forces- molecules mix freely.


Very little effect on liquids or solids

Solubility of a gas is unchanged at a given pressure: gas + solvent  solution

Henry’s Law: The solubility of a gas in a liquid increases as the pressure of that gas on the surface of the liquid increases.
 Effervescence: The rapid escape of a gas from a liquid.



Inverse relationship: An increase in temperature,
decreases gas solubility. Why?

More solute molecules can escape the attraction of solvent molecules.


Difficult to predict



The net amount of heat energy absorbed or released when a specific amount of solute dissolves in a solvent.
Overhead Transparency #74

- H solution
Step 3
= Heat is released = Steps 1 and 2

+ H solution
Step 3.
= Heat is absorbed = Steps 1 and 2



A measure of the amount of solute in a given amount of solvent or solution.
Molarity = moles solute/ Liter of solution

Molality = moles solute/ kg solvent

Doesn’t change with temperature changes.


You are asked to prepare 5.00 Liters of a 2.00M solution of Sodium Acetate.

How many grams of sodium acetate would you measure out?

Given all of the necessary glassware, what steps would you take to make your solution?



Dissociation: Separation of ions when dissolved.
NaCl (s)  Na+ (aq) + Cl – (aq)
Precipitation Reactions: Compounds of very low solubility…Practically Insoluble.

Table 14. 1 shows us general guidelines to predict solubility.

Dissociation reactions are NOT written for insoluble compounds.


Includes only those compounds/ions that undergo a chemical change in a reaction in an aqueous solution.
1. Convert chemical equation  ionic equation
2. Cancel out spectator ions (ions that do not take part in a chemical reaction).
Ex: Sodium Chloride + Silver Nitrate 


The term used for the formation of ions from molecular compounds. (The creation of ions from where there are none)

Degree of ionization depends on strength of bonds between solute AND strength of attraction between solute/solvent.

Ex: HCl (g) + H
2
O (l)  H
3
O + (aq) + Cl (aq)
Hydronium
Ion


Strong Electrolyte: A compound whose dilute aqueous solution conducts electricity
because of the high presence of ions when dissolved.

Ex: HCl, NaCl

Weak Electrolytes: Solution does NOT conduct electricity well because there is a low presence of ions that are dissolved.

Ex: HF, CH
3
COOH


Properties that depend on the
of solute particles and not their identity

Vapor Pressure Lowering: The vapor pressure of water containing sugar (or any other nonvolatile solute) is less than the vapor pressure of pure water.
WHY?
Concentration of water molecules is lower at the surface of the liquid b/c of increased attraction between solute/solvent.


The boiling point of water containing a nonvolatile solute is
than the boiling point of pure water.
WHY?
Since the vapor pressure is lower, water particles need a higher KE to overcome atmospheric pressure and boil.


The freezing point of a solution is lower than the freezing point of a pure solvent.
WHY?
Solute particles get in the way of water particles trying to freeze, so water particles need to move slower to ensure the correct orientation of the lattice.

Δ t f
= K f m
Molality
Freezing point
depression: The difference between the freezing point of a pure solvent and solution
Freezing Point
Constant: The freezing point depression of the solvent in a 1-molal solution (°C/m)
Δ t b
= K b m
Boiling Point
Elevation
Boiling Point
Constant
(Table 14.2)
Molality


The external pressure required to stop osmosis.

Osmosis: The movement of solvent through a semipermeable membrane from

solute concentration.
Osmotic Pressure increases with the number of solute particles in solution .


Colligative Properties depend on the
of particles produced in solution.

1m solution of NaCl produced more than 1m solution of dextrose due to NaCl dissociation.

Table 14.3



Write your answer on a sheet of paper and hold up when asked
Which would produce the greatest change in freezing/boiling points? a. 1m solution Sucrose (C
12
H
22
O
11
) b. 1m solution Glucose (C
6
H
12
0
6
) c. 1m solution KCl d. 1m solution of CaCl
2
ANSWER: D. More ions are produced in solution.