MR. SURRETTE VAN NUYS HIGH SCHOOL
CHAPTER 10: SOLUTIONS AND MOLARITY
CLASS NOTES
SOLUTIONS
Chemical solutions are homogeneous mixtures of solute and solvent. Solutes are substances that are dissolved in solutions. Solvents are substances present in greater amounts that dissolve the solutes. Solutions can exist in the gas, solid, and liquid phases.
SOLID SOLUTIONS
Solid solutions consist of solids dissolved in other solids. Examples of solid solutions include “white gold” (silver dissolved in gold), brass, and bronze.
GAS SOLUTIONS
Gas solutions consist of various gases mixed together. Air is a typical example of a gas solution.
LIQUID SOLUTIONS
All liquid solutions contain liquid solvents, but may contain liquid, solid, or gas solutes. For example, rubbing alcohol is a liquid-liquid solution, cola is a gas-liquid solution, and sea water is a solid-gas-liquid solution.
AQUEOUS SOLUTIONS
Normally the word “solution” refers to aqueous solutions. Aqueous solutions consist of solutes dissolved in water.
HOW SOLUTES DISSOLVE
Solutes dissolve into solvents by separating into ions or molecules. These ions or molecules from the solvent are invisible to the naked eye. They result from random molecular collisions with the solvent.
NON-POLAR VERSUS POLAR SOLVENTS
Atoms can share electrons evenly or unevenly. Atoms that share electrons evenly form non-polar molecules. Atoms that share electrons unevenly form polar molecules. Solvents can contain non-polar or polar molecules. Hexane, ethyl alcohol, and toluene are non-polar solvents. Water, acetone, and methanol are polar solvents.
LIKE DISSOLVES LIKE
In general, when it comes to solutions, “like dissolves like,” so polar solvents tend to dissolve polar solutes and non-polar solvents tend to dissolve non-polar solutes.
SOLVATION
Solvation is the attraction and association of molecules of a solvent with molecules or ions of a solute. In general, solvation is a competition between the forces that hold the solute together and the solvent. If the forces that hold the solute together are stronger, the solute will not dissolve. If the forces that hold the solute together are weaker than the solvent, then the solute will dissolve.
SOLVATION EXAMPLE
Stirring a teaspoon of sodium chloride (table salt) into a glass of water is a typical example of solvation.
NaCl in Solvation Reaction
SOLVATION EXAMPLE
In this case, the forces that hold the sodium and chlorine ions together are weaker than the water solvent. The NaCl completely dissolves as the hydrogen ends of the water molecules surround all the chlorine anions and the oxygen ends of the water molecules surround all the sodium cations.
SOLVATION RATE
Temperature, pressure, and surface area all affect solvation. In general, the following factors increase the speed of solvation:
1. Raising the temperature
2. Increasing the pressure
3. Increasing the surface area (for example, using a mortar and pestle to grind solute into powder)
SOLUBILITY AND SATURATION
Solubility measures how many grams of solute can dissolve in a solvent. With a few exceptions (like gas-liquid solutions), solubility increases with increasing temperature. Saturation is the maximum amount of solute that can dissolve in stable solution. The amount of saturation also tends to increase with increasing temperature.
GAS-LIQUID SOLUBILITY
Gas-liquid solutions are exceptions to the solubility and saturation temperature trends. As the temperature of gas-liquid solutions increases, the dissolved gases within them form additional bubbles that escape the surface of the solution. This means that the solubility and saturation of gas-liquid solutions decreases with increasing temperature.
UNSATURATED SOLUTIONS
As the name suggests, unsaturated solutions contain less than the maximum amount of solute. This means more solute can be added to them.
SUPERSATURATED SOLUTIONS
Supersaturated solutions contain unstable amounts of solute. Supersaturated solutions often form after rapid cooling. “Rock candy” is a good example. Rock candy is made by adding sugar to boiling water and chilling the solution to room temperature. Tapping the side of the rock candy container or adding more solute causes the excess sugar to solidify and precipitate out of solution.
ELECTROLYTES
Ionic solids like sodium chloride form electrolytes when dissolved in liquids. Electrolytes are charged particles (cations and anions) found in liquid solutions.
CONCENTRATED AND DILUTE SOLUTIONS
Solutions contain various amounts of solutes. Some solutions contain high percentages of solute and others contain low percentages. Concentrated solutions have a high solute percentage and dilute solutions have a small solute percentage.
MASS OF SOLUTIONS
The mass found in solutions is measured in grams. To compute the mass of solutions, add the gram weights of both the solute and the solvent.
MASS PERCENTAGE
The ratio of solute to solution can be measured as a mass percentage:
% Mass = (mass of solute / mass of solution) x 100
DENSITY OF WATER
Because aqueous solutions are common, it is important to know the density of water. The density of water is one gram per milliliter at 20 oC (typical laboratory conditions). This is sometimes written:
r(WATER) = 1g / mL
Example 1. Calculate the % mass of 11.3 grams sucrose (table sugar) dissolved in 412.1 mL water.
1A.
(1) 412.1 mL water = 412.1 g water
(2) mass of solvent = 412.1 g
(3) mass of solute = 11.3 g
(4) mass of solution = mass of solute + mass of solvent
(5) mass of solution = 11.3 g + 412.1 g
(6) mass of solution = 423.4 grams
(7) % mass = (mass solute / mass solution) x 100
(8) % mass = (11.3 g / 423.4 g) x 100
(9) % mass = 2.67%
Example 2. Soda is 11.5% sucrose. What volume of soda in mL contains 85.2 g sucrose? (Assume the density of soda = r(WATER).
2A.
(1) 11.5 % sucrose = 11.5 g sucrose (convert percentages into grams)
(2) Volume = 85.2 g sucrose (100 mL soda / 11.5 g sucrose)
(3) Volume = 740.9 mL
Example 3. Soda is 11.5% sucrose. How many grams of sucrose are found in 355 mL of soda?
3A.
(1) 11.5% sucrose = 11.5 g sucrose
(2) Mass = 355 mL (11.5 g sucrose / 100 mL)
(3) Mass = 40.8 g
CONCENTRATION
The concentration of a solution measures the amount of solute per unit solvent. Concentration is usually measured in moles per liter.
MOLARITY
Molarity (M) measures how many moles of solute are dissolved per liter of solution:
M = (moles of solute / liter of solution)
Example 4. What is the molarity of a solution that contains 15.5 grams NaCl dissolved in 1.5 liters water?
4A.
Convert grams NaCl to moles:
(1) NaCl = 15.5 g (1 mol NaCl / 58.44 g)
(2) NaCl = 0.265 mol
Compute molarity of solution:
(3) M = (0.265 mol NaCl / 1.5 L)
(4) M = 0.18
Example 5. How many liters of 0.114 M NaOH solution contain 1.24 moles NaOH?
5A.
(1) Volume = 1.24 mol NaOH (1 liter / 0.114 mol)
(2) Volume = 10.9 L
DILUTIONS
Chemical solutions are often bottled in high concentrations. These concentrated solutions often have to be diluted.
MOLARITY DILUTIONS
Molarity (M) and volume (mL) are used to determine dilution ratios:
MIVI = MFVF
(MI = molarity of initial solution
VI = volume of initial solution
MF = molarity of final solution
VF = volume of final solution)
Example 6. How much 6.0 M NaNO3 is needed to make 0.585 L of 1.2 M NaNO3 solution?
6A.
Given: MI = 6.0 M, MF = 1.2 M, VF = 0.585 L
Find: VI
(1) MIVI = MF VF
(2) VI = MFVF/ MI
(3) VI = (1.2 M)(0.585 L) / (6.0 M)
(4) VI = 0.117 L
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CHEMISTRY