1. Solutions and Colligative Properties
Solutions are homogeneous mixtures of two or more substances in a single phase.
Solute – substance being dissolved
Solvent – dissolving medium

2. Types of solutions Gas Oxygen in nitrogen Liquid
Solute
Solvent
Example
Gas
Oxygen in nitrogen
Liquid
CO2 in H2O (carbonated beverages)
Alcohol in water, ethylene glycol in water (antifreeze), acetic acid in water (vinegar)
Solid
Hg in Ag and Sn (dental amalgam)
Sugar in water, ocean water (salt in water)
Copper in nickel (Monel™ alloy)
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3. This one fits in solution PPT

4. Heterogeneous mixtures are suspensions and colloids
Heterogeneous mixtures are suspensions and colloids. Brownian motion is the erratic movement of colloid particles. The Tyndall effect (named after John Tyndall) is when light is scattered by colloidal particles dispersed in a transparent medium. For example: car headlights in fog.

5. Comparison of solutions, colloids and suspensions
Homogeneous
Heterogeneous
Particle size: nm; can be atoms, ions, molecules
Particle size: nm, dispersed; can be aggregates of large molecules
Particle size: over 1000 nm, suspended; can be large particles or aggregates
Do not separate on standing
Particles settle out if not constantly stirred
Cannot be separated by filtration
Can be separated by filtration
Do not scatter light
Scatter light (Tyndall effect)
May scatter light, but are not transparent
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6. Classes of colloids Class of colloid Phases Example Sol
Solid dispersed in liquid
Paints, mud
Gel
Solid network extending throughout liquid
Gelatin
Liquid emulsion
Liquid dispersed in liquid
Milk, mayonnaise
Foam
Gas dispersed in liquid
Shaving cream, whipped cream
Solid aerosol
Solid dispersed in gas
Smoke, airborne particulate matter, auto exhaust
Liquid aerosol
Liquid dispersed in gas
Fog, mist, clouds, aerosol spray
Solid emulsion
Liquid dispersed in solid
Cheese, butter
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7. Solutes are classified according to the extent they dissociate
into ions in aqueous
solutions.
Strong electrolyte: substance that dissolves in water to give a solution that conducts electricity. (i.e. salt – NaCl)
Weak electrolyte: less than 50% of dissolved solute exists as ions.
(i.e. – acetic acid[vinegar])

10. Nonelectrolyte: < 0.01% exists as ions; does not conduct electricity when solute is dissolved in water. (i.e. sugar – C12H22O11)

12. The Solution Process Factors that affect the rate of dissolution:
stirring
temperature
surface area
Amount of solute already dissolved in solution

13. Solubility is the maximum quantity of solute that can be dissolved in a given quantity of solvent at a particular temperature. It is usually expressed in g of solute per 100 g of solvent at a given temperature.
The three interactions that determine the solubility are solute-solute, solute-solvent, and solvent-solvent interactions.

14. Temperature effects on solubility

15. Solution equilibrium is the physical state in which the opposing processes of dissolution and crystallization of a solute occur at equal rates.
Saturated solution: a solution that contains the maximum amount of dissolved solute.
Unsaturated solution: a solution that contains less than a saturated solution under the current conditions.
Supersaturated solution: a solution that contains more dissolved solute than a saturated solution contains under the same conditions.

17. Nature of solute and solvent (polarity, etc
Nature of solute and solvent (polarity, etc.) Solute-solvent interactions - “like dissolves like”
Temperature (for gases an increase in T decreases solubility). For the majority of solids in liquids, increasing T increases their solubility. The degree of solubility can differ greatly and in some cases, even decrease a solid’s solubility.
pressure (for solids and liquids has no effect; for gases an increase in P increases solubility)
gas + solvent  solution

18. Henry’s Law the solubility of a gas in a liquid is
Directly proportional to the partial
pressure of that gas on the surface of
the liquid.
S1 = S2
P1 P2
In carbonated beverages forcing it into solution at pressure of 5-10 atm increases CO2 solubility. The containers are then sealed. When opened, the CO2 gas escapes as the pressure returns to 1 atm. The rapid escape of a gas from a liquid in which it is dissolved is called effervescence.

19. The net amount of heat energy absorbed or released when a specific amount of solute dissolves in a solvent is the heat of solution.

20. http://img. sparknotes

21. When two liquids dissolve in each other at all proportions they are said to be miscible (i.e. alcohol and water); when they do not they are called immiscible (i.e. water and oil). You can separate immiscible liquids with a piece of equipment called a separatory funnel (shown at right

22. Fractional Distillation Apparatus
knowledgerush.com

23. ljcreate.com
Paper Chromatography

24. Chromatogram Development (KB)
Draw a line to indicate the height of the solvent front
X the pigments to find the center
Put a pencil mark to locate where the pigment goes
A B C
Calculate the Rf = B/A and C/A

25. Gas chromatography (GC) is an analytical technique for separating compounds based primarily on their volatilities. Compounds move through a GC column as gases. The compounds partition between a stationary phase, which can be either solid or liquid, and a mobile phase (gas). The differential partitioning into the stationary phase allows the compounds to be separated in time and space.

26. cee.vt.edu

27. Hydration: solution process in which the solvent is water.
Solvation: process of surrounding solute particles with solvent particles. A solute particle that is surrounded by solvent molecules is said to be solvated.
Hydration: solution process in which the solvent is water.
see picture at right:

28. Dissociation is the separation of ions that occurs when an ionic compound dissolves.
ex. NaCl (s)  Na+ (aq) + Cl- (aq)
Ionization is the process whereby ions are formed from solute molecules by the action of the solvent.
ex. HCl  H+ (aq) + Cl- (aq)
In water solutions, the H+ ion normally exists attached to the water molecules in solution to form the hydronium ion (H3O+)
ex. HCl (g) + H2O (l)  H3O+ (aq) + Cl- (aq)

29. Concentration of solutions: a measure of the amount of solute in a given amount of solvent or solution.
Mass percent =
mass of component in solution x 100
total mass of solution ppm
mass of component in solution x 106
total mass of solution

30. Molarity
number of moles of solute in one liter of solution. Note that the total volume of the solution is 1 liter, not that you add solute to 1 liter of solvent. The symbol for molarity is “M” and usually referred to as an “X molar solution.”
Molarity (M) = moles of solute
liters of solution
Dilutions of solutions M1V1 = M2V2

31. Molality = number of moles of solute per kilogram of solvent
Molality = number of moles of solute per kilogram of solvent. The symbol for molality is “m,” and is usually referred to as an “X molal solution.”
molality (m) = moles of solute kg of solvent
moles of solute
1000g of solvent

32. Precipitation Reactions
Typically when two aqueous solutions are combined and one of the products is minimally soluble a PPT is formed

33. Table 1 in your textbook page 437; Table A-12 (appendix) page 860
Table 1 in your textbook page 437;
Table A-12 (appendix) page 860

34. Na2CO3 (aq) + Ca(NO3)2 (aq) 
CaCO3 (?) + 2NaNO3 (?)
Looking at your solubility table we
see that all nitrates are soluble and
most carbonates are insoluble,
except Na, K and NH4
Thus CaCO3 (s) + 2NaNO3 (aq)

35. Net ionic equations
Include only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution.
Spectator ions are those that do NOT take part in a chemical reaction and are found in solution before and after the reaction takes place

36. Ca2+ (aq) + CO32- (aq)  CaCO3 (s)
Ionic equation:
2Na+ (aq) + CO32- (aq) + Ca2+ (aq) + 2NO3- (aq)  CaCO3 (s) + 2Na+ (aq) + 2NO3- (aq)
Net ionic equation:
Do not include spectator ions.
Ca2+ (aq) + CO32- (aq)  CaCO3 (s)

37. Colligative Properties
properties that depend on the concentration of solute particles but not on their identity.
Vapor pressure lowering. Boiling point is higher and freezing point of a solution is lower than that of a pure solvent. This is due to the presence of nonvolatile solutes. These are substances that have little tendency to become a gas under existing conditions.

39. Freezing point depression, ∆Tf, is the difference between the freezing points of the pure solvent and a solution of a non-electrolyte in that solvent, and it is directly proportional to the molal concentration of the solution.
∆Tf = Kf m
Kf = the molal freezing-point constant is the freezing-point depression of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute. The value for this constant depends on the solvent used. (See a reference table for specific values)

40. Boiling point elevation, ∆Tb, is the difference between the boiling points of pure solvent and a nonelectrolyte solution of that solvent, and it is directly proportional to the molal concentration of the solution ∆Tb = Kb m
Kb = molal boiling-point constant is the boiling-point elevation of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute. The value for this constant depends on the solvent used. (See a reference table for specific values)

41. Electrolytes and colligative properties
Keeping in mind that colligative properties depend on the number of solute particles, not their identity, we must know the extent to which electrolytes dissociate.
The higher the degree of dissociation, the more moles of ions we have in solution, thus increasing the net effect.