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HONORS CHEMISTRY
UNIT 6: CHEMICAL EQUILIBRIUM
CHAPTER 15: SOLUTIONS
OBJECTIVES AND NOTES - V8
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Chapter Objectives
1. Describe the properties of solutions. (15.1)
2. Describe and identify the different components of solutions. (15.1)
3. Identify the different types of solutions. (15.1)
4. Describe and give examples of alloys. (15.1)
5. Describe soluble, insoluble, miscible, and immiscible systems. (15.1)
6. Describe and identify strong electrolytes, weak electrolytes, and nonelectrolytes. (15.1)
7. Explain what is meant by the concentration of a solution. (15.2)
8. Define and solve problems involving the molarity of a solution. (15.2)
9. Define and solve problems involving the molality of a solution. (15.2)
10. Differentiate among saturated, unsaturated, and supersaturated solutions and describe methods to form such
solutions. (15.2)
11. Differentiate between dilute and concentrated solutions and describe methods to form such solutions. (15.2)
12. Describe how to prepare dilute solutions from concentrated solutions of known molarity. (15.2)
13. Explain how solutions form. (15.3)
14. Describe the role of solvation in the dissolving process. (15.3)
15. Define solubility, describe the factors that affect it, and interpret solubility curves. (15.3)
16. Use the rule that "like dissolves like" to predict the solubility of one substance in another. (15.3)
17. Describe the differences between the solubility of solids in liquids and gases in liquids. (15.3)
18. Describe Henry's law. (15.3)
19. Describe the factors that affect the rate at which solutes dissolve in solvents. (15.3)
20. Describe colligative properties of solutions. (15.4)
21. Describe why a solution has a lower vapor pressure than the pure solvent of that solution. (15.4)
22. Explain how the addition of a solute to a pure solvent causes an elevation of the boiling point and a
depression of the freezing point of the resultant solution. (15.4)
23. Describe colloids and suspensions; distinguish them from solutions. (15.4)
24. Demonstrate and be able to describe all aspects of laboratory safety rules and procedures. (Applicable every
chapter)
1
15.1 The Nature of Solutions
1. solution (homogeneous mixture): A homogeneous material that is composed of two or more physically
combined materials that are uniformly distributed throughout the mixture. The component parts do not
have to be in a fixed ratio and they do not lose their own identity.
a. Because the components are physically mixed together, they can be separated by physical methods
such as distillation or crystallization
1. Note that while filtration is a physical method of separating mixtures, filtering a solution does not
separate it into its component parts as the parts are so small as to pass through the pores in the
filter paper.
b. Most solid solutions are formed when all components are in the liquid state.
c. An aqueous solution must be clear, although it may be colored.
1. Solutions do not scatter light.
d. soln is the abbreviation for solution that is used by really cool people or even really groovy people,
but it is almost never used by chemistry nerds or geeks.
a. Properties of Solutions
1. The particles in a solution are atoms, ions, or molecules.
2. The uniform distribution occurs on a molecular level.
3. solvent: The substance that is doing the dissolving in a solution. When both components are in the same
state, such as a liquid-liquid solution, the solvent is the component which is in the greatest amount.
4. solute: The substance that is being dissolved in a solution. The solute is usually the component which is
present in the lesser amount.
5. soluble: A substance that is capable of being dissolved in the chosen solvent.
a. Example: sodium chloride is soluble in water.
6. insoluble: A substance that is not capable of being dissolved in the chosen solvent.
a. Example: calcium carbonate (limestone) is insoluble in water.
b. Note that this is not exactly correct, most insoluble substances actually do dissolve in the chosen
solvent to a very small degree. However, the amount of solute that dissolves is so small (how small?)
that for now it will be considered to be insignificant. (For a more detailed view of this type of
equilibrium system, see Chapter 17 in the Prentice-Hall textbook).
c. Note that this term is not imsoluble, unsoluble, nonsoluble, or antisoluble.
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b. Types of Solutions
1. miscible: A liquid solute that is capable of being dissolved in a chosen liquid solvent.
a. Example: ethyl alcohol is miscible in water.
2. immiscible: A liquid solute that is not capable of being dissolved in a chosen liquid solvent.
a. Example: oil is immiscible in vinegar.
b. Note that this term is not inmisicible, unmisicible, nonmisicible, or antimisicible.
3. alloy: A solid solution which has metallic properties; alloys are composed of only metallic elements or
mostly metallic elements.
a. The formation of the solution probably took place when all of the components were liquids.
b. Examples: brass is zinc (33%) dissolved in copper (66%), sterling silver is copper (7%) dissolved in
silver (93%), pewter is antimony (2%), bismuth (6%), and copper (7%) dissolved in tin (85%), and
steel is carbon (up to 2%) dissolved in iron.
1. There are also alloy steels in which other elements are dissolved in the steel (listen to "Allentown"
by Billy Joel).
2. If you memorize any of the substances and/or % in b. above you need to get a life.
4. aqueous solution: A solution in which the solvent is water.
Types Of Solutions Based On The Physical State Of Solutes And Solvents
(Memorize Only If You Like Lima Beans)
Solute
Solvent
Common Example
Names of Example
Gas
Gas
Oxygen in Nitrogen
Air
Liquid
Gas
Water in Air
Humidity
Solid
Gas
Hydrogen in Palladium
Hydrogen Storage
Gas
Liquid
Carbon Dioxide in Water
Soda
Liquid
Liquid
Acetic Acid in Water
Vinegar
Solid
Liquid
Salt in Water
The Oceans
Gas
Solid
Poisonous Gases in Charcoal Filter
Air Filter
Liquid
Solid
Mercury in Silver/Tin
Dental Fillings
Solid
Solid
Carbon in Iron
Steel
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5. electrolyte: A substance that conducts electricity when dissolved in water. These substances must ionize
or dissociate so as to produce ions that carry the electrical current.
a. strong electrolyte: A substance, that when placed in water, dissociates or ionizes to a large degree. A
strong electrolyte produces many ions that allow the solution to be a good conductor of electricity.
1. Examples: most soluble salts and strong acids and bases.
b. weak electrolyte: A substance, that when placed in water, dissociates or ionizes only slightly. A weak
electrolyte produces few ions so the solution is a poor conductor of electricity.
1. Examples: weak acids and bases.
6. nonelectrolyte: A substance that does not conduct electricity when dissolved in water. These substances
produce no ions when dissolved.
a. Example: sugar.
15.2 Concentration of Solutions
Note: In many concentration expressions the term mole is used. However, when working in the chemistry
laboratory we do not have measuring devices with moles as units. Therefore, in the real world, as in most
test questions, moles of a substance are usually converted to or from grams of that substance.
a. Molarity
1. M =
moles of solute
liter of solution
2. This type of concentration is referred to as a molar solution.
3. The units are moles/liter.
3. MM is an abbreviation for molar mass; the units are grams/mole.
a. MM is used when converting from # of moles of solute  grams of solute.
4. 1 liter = 1 dm3
5. 1 cubic centimeter (1 cc) = 1 centimeter cubed (1 cm3) = 1 milliliter (1 mL)
6. It is the most common form of concentration expression.
7. Remember that in this expression it is liters of solution, not liters of solvent.
8. In the chemistry laboratory, this type of solution is usually made in a volumetric flask.
a. When preparing this type of solution textbooks often state to place the necessary grams of solute in
the volumetric flask and then fill to the mark with distilled water. This is a great way to have the
solute clump to the bottom of the flask. A far better technique is to first put some distilled water into
the flask, add the solute and swirl the contents until the solute dissolves, and then fill to the mark with
distilled water.
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b. Molality
moles of solute
kg of solvent
2. This type of concentration is referred to as a molal solution.
1. m =
3. The units are moles/kg.
4. Molality is the form of concentration expression found in a number of selected colligative property
relationships such as freezing point depression (see 15.4 in the notes below).
5. Remember that in this expression it is kg of solvent, not kg of solution.
c. Saturation
1. saturated solution: A solution that has dissolved as much solute as it can under those conditions;
dissolved solute is in equilibrium with any undissolved solute.
a. A saturated solution can be dilute or concentrated.
1. dilute solution: A solution that contains a small amount of solute per volume of solution. There are
no mathematical parameters for this term.
a. A saturated aqueous solution of plumbous iodide is very dilute as plumbous iodide tends not to
dissolve in water.
2. concentrated solution: A solution that contains a large amount of solute per volume of solution.
There are no mathematical parameters for this term.
a. A saturated aqueous solution of sugar is so concentrated as to be very viscous (no, not vicious,
viscous).
2. unsaturated solution: A solution that is capable of dissolving more solute under the present conditions.
a. An unsaturated solution can be dilute or concentrated.
1. An aqueous solution of ammonia can be concentrated but unsaturated.
2. An aqueous solution of most soluble salts can be dilute as well as unsaturated.
3. supersaturated solution: A very unstable system in which the solution has more solute dissolved than
theoretically possible under the present conditions.
a. Supersaturated solutions are normally formed when a solution is saturated at a certain temperature and
then the temperature of the system is decreased. For most solutions, some of the solute should
crystallize out of solution as the temperature decreases. However, in some systems, the solute does
not precipitate, therefore, more solute is dissolved than should occur. Often shaking the solution, or
adding a "seed crystal" is all that is needed to cause a massive precipitation to occur.
1. Aqueous sodium acetate or sodium thiosulfate (Na2S2O3) solutions can become supersaturated
fairly easily using the process found directly above.
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4. Dilute solutions can be prepared by adding solvent, usually water, to a more concentrated solution.
a. M(initial soln) x volume(initial soln) = M(final soln) x volume(final soln)
b. While a more concentrated solution can be formed by evaporating or boiling away the solvent, in
most situations this is not a feasible process.
15.3 The Formation of Solutions
a. How a Solution Forms
1. dissolution: The process of a solute dissolving in a solvent.
a. The opposite of a substance dissolving to form a solution is a solute coming out of solution, this is
usually called crystallization or precipitation.
2. solvation: The process in which solute particles are surrounded by solvent particles when a solution is
formed.
a. hydration: Solvation when the solvent is water.
3. enthalpy of solution (heat of solution): The energy change that occurs during dissolution.
a. Separating solute particles from other solute particles is always an endothermic process as attractive
forces must be broken.
b. Solvation is always an exothermic process as attractive forces are formed.
c. The heat of solution is the sum of the above two processes.
1. When the energy released during solvation is greater than the energy needed to separate the solute
particles the process is exothermic.
a. Example: dissolving sodium hydroxide in water.
2. When the energy released during solvation is less than the energy needed to separate the solute
particles the process is endothermic.
a. Example: dissolving sodium nitrate in water (the chemical process used in "cold packs").
d. The energy change that accompanies a supersaturated solution coming out of solution is analogous to
heat of solution.
1. The energy change when a supersaturated solution of sodium acetate or sodium thiosulfate comes
out of solution is the chemical basis for many reusable "heat packs".
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b. Solubility
1. solubility: The amount of solute that can dissolve in a given amount of solvent.
a. Solubility is usually expressed in grams of solute per 100 g of solvent.
2. Solubility is affected by the nature of the solute and solvent.
a. A common phrase is "like dissolves like", that is, solutes dissolve in solvents that have similar
bonding structures.
1. Many salts and the polar molecule ethyl alcohol are soluble in the very polar solvent water.
2. Nonpolar oil-based paints dissolve in nonpolar turpentine or paint thinner.
3. Nonpolar grease is insoluble in polar water.
3. Solubility is affected by temperature.
a. "Figure 15-17" on page 517 in the Prentice-Hall textbook is a classic example of the relationship
between solubility and temperature for a number of solutes dissolved in water.
b. Note that for most solids, solubility increases as temperature increases.
1. The added energy helps to break the solute-solute attractions.
c. However, note that solubility for gases usually decrease as temperature increases.
1. The solute-solute forces are nonexistent as the solute is a gas. The added energy breaks the solutesolvent attractions and thus the gas escapes the solvent.
a. Thermal pollution is the heating of water that supports aquatic life. As the temperature of the
water rises the solubility of oxygen gas decreases and this causes fish to suffocate.
b. Soda in an opened container will go flat slower in a refrigerator than if left at normal room
temperature.
4. The solubility of gases in liquids is affected by pressure.
a. Henry's law: At a given temperature, the solubility of a gas in a liquid (S) is directly proportional to
the pressure (P) of the gas above the liquid.
b. When a container of soda is opened the sound heard is the carbon dioxide that was at a pressure
greater than one atmosphere escaping the container. Bubbles of CO2 immediately start to form in the
soda; the pressure over the liquid was decreased and the solubility of the CO2 decreased.
1. Soda bottle tops with built-in pumps work on the reverse process. By pumping air into the bottle
the "flattening" of the soda can be slowed down.
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c. Factors Affecting the Rate of Dissolving
1. Increasing surface area increases the rate of dissolving. Solvation takes place at the surface of a solid
solute. Increasing surface area increases the amount of solute exposed to the solvent so that solvation
can take place faster.
a. Granulated sugar dissolves faster than sugar cubes.
2. Stirring increases the rate of dissolving. Stirring speeds up the collisions between unsolvated solute and
solvent particles.
a. Stirring hot chocolate mix dissolves it much faster than simply pouring the powder into the liquid.
3. Increasing temperature increases the rate of dissolving. Increasing temperature increases the speed of
the particles and thus speeds up the collisions between unsolvated solute and solvent particles.
a. Who would use hot chocolate mix in cold milk or water? It would take forever to dissolve and
probably taste lousy anyway.
15.4 Colligative Properties
1. colligative properties: Properties of solutions that are affected by the concentration of the solute
particles, not the bonding structure of the solute or solvent.
a. van't Hoff factor (i): A measure to which a solute dissociates or ionizes. For our purposes we will
assume all electrolytes dissociate or ionize completely (a gross approximation in some cases, but, it
will suffice for now). Based upon our approximation, the van't Hoff factor for several substances is
found below.
1. sugar and other nonionizing substances: i = 1
2. NaCl: i = 2
3. MgF2: i = 3
4. Ca(OH)2: i = 3
5. (NH4)3PO4: i = 4
b. For all problems involving colligative properties i must be considered.
a. Vapor Pressure Reduction
1. Increasing the concentration of a nonvolatile solute decreases the vapor pressure of the solution.
b. Boiling Point Elevation
1. Adding a nonvolatile solute to a solvent lowers the vapor pressure of the solvent and thus increases the
boiling temperature.
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c. Freezing Point Depression
1. Adding a nonvolatile solute to a solvent lowers the vapor pressure of the solvent and thus decreases the
melting temperature.
d. Colloids and Suspensions
1. colloid (colloidal dispersion/colloidal suspension): A mixture in which the disbursed particles are larger
than in a solution, but smaller than a suspension.
a. It is between a homogeneous mixture (solution) and a heterogeneous mixture (suspension).
b. The disbursed particles are stable and do not settle.
1. The particle size is 1 - 100 nm (Don’t memorize these numbers).
c. It demonstrates the Tyndall effect.
1. Tyndall effect: The scattering of visible light when light is shown through a colloid or suspension.
This scattering allows the beam of light to be visible from the side.
a. Since air is a solution, a beam of light is not visible from the side. However, if the light is
shown through dust scattered in the air, or through a cloudy liquid mixture, the beam of light
can be seen from the side; this is the Tyndall effect.
d. Examples: see below, starting with #3.
2. suspension: A mixture, that when left undisturbed will have the disbursed particles slowly settle to the
bottom of the container.
a. It is a heterogeneous mixture.
b. The particle size is greater than 100 nm (Don’t memorize this number).
c. It demonstrates the Tyndall effect.
d. It can be separated by filtration.
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Comparison Of Solutions, Colloids, And Suspensions
System
Property
Solution
Colloid
Suspension
Particle Type
Ions, Atoms,
Small Molecules
Large Molecules
or Particles
Large Particles
or Aggregates
Particle Size
(Don’t Memorize)
0.1 - 1 nm
1 - 100 nm
≥ 100 nm
Effect of Light
No Scattering
Tyndall Effect
Possible Tyndall
Effect, But
Not Transparent
Effect of Gravity
Stable, Does Not
Settle on Standing
Stable, Does Not
Settle on Standing
Unstable, Sediment
Forms
on Standing
Uniformity
Homogeneous
Heterogeneous
Heterogeneous
Colligative Properties
Affect Properties
No Effect
No Effect
Types Of Colloids Based On The Physical State
Of The Dispersed And Dispersing Mediums
(Memorize This Table and We Will Make You Listen to Barry Manilow on an 8-Track)
Dispersed
Phase
Dispersing
Medium
Type
Examples
Gas
Liquid
foam
whipped cream, shaving cream
Gas
Solid
solid foam
Styrofoam, rubber, marshmallow
Liquid
Gas
aerosol
(liquid aerosol)
fog, clouds, aerosols
Liquid
Liquid
emulsion
(liquid emulsion)
milk, mayonnaise (egg acts as the emulsifying
agent, keeping the oil dispersed in water- see #
5 below)
Liquid
Solid
solid emulsion
butter, cheese
Solid
Gas
solid aerosol
household dust in the air, smoke, smog
Solid
Liquid
sol1
water based paints, puddings, smog
1
gel: A special type of sol that is semirigid; the dispersed phase forms a network-like structure that
encloses the dispersing medium. Examples: jelly, human body cells.
Solid
Solid
solid sol
porcelain, pearls, opals
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