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Unit 8, Lesson 1 of 3
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Properties of Solutions
OBJECTIVES: After this lesson, I will understand the basic properties of solutions. In addition, I will
understand the difference between saturated, unsaturated, and supersaturated solutions.
Recall that a solution:
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Thus far when we have spoken about solutions, we have thought about a physical process, such as dissolving
salt in water, NaCl(s) + H2O(l)  NaCl(aq) +H2O(aq). We know that this is a physical process, because we
can recover the salt in its original form, simply by evaporating or boiling off the water. Solutions can also form
by chemical reactions, such as a single replacement reaction, Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g). In this
instance, the chemical form of the substance being dissolved is changed from Ni to NiCl2. If the solution is
evaporated, NiCl2•6H2O(s), not Ni(s) is recovered.
Solutions are composed of two parts:
Solute:
Solvent:
The process of dissolving a solute in a solvent is known generally as solvation. When the solvent is water, the
dissolving process is also referred to as hydration.
Ionic Compounds in Water
When ionic compounds dissolve in water, each ion separates from the solid structure and disperses throughout
the solution. This process is known as dissociation. The ions are solvated or surrounded by water molecules,
which prevents the cations and anions from recombining and allows the water molecules to freely move about
so that the ions are dispersed uniformly throughout the solution.
Electrolyte:
Molecular Compounds in Water
When a molecular compound dissolves in water, the solution usually consists of intact molecules dispersed
throughout the solution. There are, however, a few molecular substances whose aqueous solutions contain ions.
The most important of these are acids. For example, when HCl(g) dissolves in water to form hydrochloric acid,
HCl(aq), it ionizes, that is, it dissociates into H+(aq) and Cl-(aq) ions.
Nonelectrolyte:
Unit 8, Lesson 1 of 3
After considering the way in which a solute dissolves in a solvent, we can consider the degree to which a solute
dissolves in a solvent. This refers to the concept of solubility.
Solubility:
As a solid solute begins to dissolve in a solvent, the concentration of solute particles increases, and so do their
chances of colliding with the surface of the solid. Such a collision may result in the solute particles becoming
reattached to the solid. This process, which is the opposite of the solution process, is called crystallization. As
a result, two opposing processes occur in a solution in contact with undissolved solute. This situation is
represented in a chemical equation by use of a double arrow: Solute + Solvent  Solution
When the rates of dissolving and crystallizing are equal, no further net increase in the amount of solute in
solution occurs. This is called solution equilibrium.
Saturated Solution:
Unsaturated Solution:
Supersaturated Solution:
Factors Affecting Solubility
Substances have a tendency to mix due to the natural tendency of systems to move toward a more dispersed or
random, state. However, some substances will mix more easily than others due to solute-solvent interactions,
temperature, or pressure conditions.
*Solute-Solvent Interactions
Chemical bonds and intermolecular forces determine how solutes and solvents will interact. In general, when
other factors are comparable, the stronger the attractions between solute and solvent molecules, the greater the
solubility. As a result, we use the phrase, “Like Dissolves Like” to determine the relative solubility of solutes
in solvents.
Solute Type
Nonpolar
Nonpolar Solvent
Polar Solvent
Polar
Ionic
Alcohol Solubility
Alcohols are identified by the presence of OH groups on hydrocarbon chains, for example CH3CH2OH is an
alcohol called ethanol. As a result of this structure, alcohols have a polar end and a nonpolar end making them
miscible in both water and nonpolar solvents. The degree of their solubility is dependent upon the length of the
hydrocarbon chain and the number of alcohol groups.
In general: the longer the hydrocarbon chain, the greater its solubility in a hydrocarbon rather than water.
In general: the more alcohol groups, the greater its solubility in water rather than a hydrocarbon.
Unit 8, Lesson 1 of 3
Predict the relative solubility of the following alcohols in water versus C6H14.
a. CH3OH
b. CH3CH2CH2CH2OH
c. HOCH2CH2CH2CH2OH
Predict whether each of the following substances is more likely to dissolve in CCl4 or in H2O.
a. C7H16
b. Na2SO4
c. HCl
d. I2
*Pressure
Pressure changes do not appreciably affect the solubilities of solids and liquids; however pressure does affect
the solubility of gases. The solubility of a gas in any solvent is increased as the pressure over the solvent
increases. Remember your soda can!
*Temperature
As the temperature increases, most solid solutes become more soluble in water. In contrast, the solubility of
gases in water decreases with increasing temperature.
Solubility Curves
Solubility is often presented on a graph, which shows the relationship between temperature and grams of solute
that may be dissolved. Look at the solubility curve below to find out what information we can obtain from
solubility curves.
Point A:
Point B:
Point C:
Point D:
Question to consider: How do you make a saturated solution?
Unit 8, Lesson 1 of 3
Solubility curves show the number of grams
of a substance that can be dissolved in a given
number of grams of water at a range of
temperatures. Each line represents the
maximum amount of that substance that can
be dissolved at a given temperature.
1. How many grams of potassium iodide will
dissolve at 30°C?
2. At what temperature will 20 grams of
potassium chlorate dissolve in water?
3. How many grams of sodium nitrate will
dissolve at 90°C?
4. At approximately what temperature will
85 grams of potassium nitrate dissolve
5. What substance dissolves best at the freezing point of water (0°C ) ?
6. Why would more material (solute) dissolve in water (solvent) at higher temperatures?
7. Classify the following solutions as saturated, unsaturated, or supersaturated.
a. 80 grams of NaNO3 in 100 g of water at 50°C ________________.
b. 30 grams of KClO3 in 50g of water at 63°C ________________.
c. 100 grams of NaNO3 in 200 g of water at 25°C ________________.
8. Predict the products of the following reactions. (You will need to use the solubility rules from Unit 6,
Lesson 6.) For each reaction, write a complete ionic and net ionic equation.
a. Na3PO4(aq) + Ca(OH)2(aq) 
b. Al2(SO4)3(aq) + Pb(C2H3O2)2(aq) 