300 Chemistry
Solutions Notes
Introduction to Solutions
Solution = homogeneous mixture
Aqueous solution = solution in which water is the dissolving medium (most of the ones we’ll deal with are this kind… although we can have solutions of any kind & involving other states of matter too)
Solute = substance that is dissolved (there is less of it)
Solvent = substance that does the dissolving (there is more of it)
Electrolyte = aqueous solution containing ions (ions conduct electricity)
Nonelectrolyte = aqueous solution that does not contain ions (does not conduct electricity)
Types of Solutions
As a solid solute dissolves in a solvent, the number of dissolved solute particles increases… then they can collide with each other and recombine… so there is always dissolving and crystallizing happening
When the rate of dissolving = rate of crystallization, we have a dynamic equilibrium… and that defines what we call a saturated solution
Saturated solution = solution in equilibrium (has undissolved solute on the bottom)
Unsaturated solution = solution where more solute can be dissolved
Supersaturated solution = unstable solution where conditions have been altered so that you are able to dissolve more than it takes to be saturated; a sudden change can cause rapid crystallization (ex: using a “seed” crystal provides a template for crystallization)… these are rare
Solubility = amount of solute needed to form a saturated solution in a given quantity of solvent (is dependent on temperature)
Factors Influencing Solutions and Solubility
• Interaction between solute and solvent particles- substances with similar intermolecular attractive forces tend to be soluble in one another… “LIKE dissolves LIKE”
Polar substances will dissolve other polar substances
Non-polar substances will dissolve other non-polar substances.
NOTE: The “Like dissolves like” is a general rule and does not always work!
• Pressure– FOR GASES DISSOLVED IN A LIQUID ONLY!
For gases as solutes, the pressure over a solvent increases the solubility of a gas in a solvent (ex: open bottle of soda, pressure goes down, bubbles go out of solution because solubility decreased)
Not an issue for liquid or solid solutes
• Temperature
For solid and liquid solutes, as temperature increases, solubility increases except for a few rare exceptions
For gas solutes, as temperature increases, solubility decreases (inverse relationship) (think about what happens to soda if left at a warm temperature… or why drinking really cold soda makes you burp!)
Solubility Curves
Notice we can look at general patterns for substances, compare them to each other, or can get specific (what is the solubility of X at a certain temperature; how many grams of a substance can be dissolved in a certain amount of water; conditions for unsaturated, saturated, and supersaturated solutions; effect of changing temperature on amount dissolved or precipitated, etc.)
If your point on the graph falls ABOVE the line, your solution is SUPERSATURATED
If your point on the graph falls BELOW the line, your solution is UNSATURATED
If your point on the graph falls ON THE LINE, your solution is SATURATED.
Solution Concentration
Most solutions are unsaturated, so we need a way to express the amount of solute and solvent they contain
Concentration = amount of solute in a solution in relation to a particular amount of solvent
Dilute solution = weak solution, contains a comparatively low amount of solute
Concentrated solution = strong solution, contains a comparatively high amount of solute
Many different ways to measure concentration (we’ll use 2: molarity and molality)
1. Molarity (M) = moles of solute / liters of solution , units are molar (M)
2. Molality (m) = moles of solute / kg of solvent , units are molal (m)
Dilution
A liquid solution, such as an inorganic acid (HCl, H2SO4, HC2H3O3, H3PO4, HNO3) can be diluted to make a solution with a new molarity. To do this, we use the following equation:
M1V1 = M2V2
Where:
M1 = the original molarity of the solution
V1 = the volume required of the original solution needed to make the new molarity solution
M2 = the NEW or desired molarity
V2 = the volume of the new molarity solution that is needed.
For example:
What volume of 12 M HCl is required to make 1 liter of a 6 M HCl solution?
M1 = 12 M
V1 = ?
M2 = 6 M
V2 = 1 liter
(12M)(x) = (6M)(1 L)
x = 0.500 L of 12 M is needed to make the 6 M solution
Colligative Properties
Properties of a solution differ from properties of a pure solvent
Colligative property = property that depends on the number of solute particles
2 colligative properties of interest to us: boiling point elevation, and freezing point depression
Boiling Point Elevation
• Boiling point elevation is related to vapor pressure lowering
• If vapor pressure is lower, an even higher increase in temperature is needed to make the solution boil (why do you add salt to water when cooking pasta?), so a boiling point elevation results
• ΔTb = kb m i = boiling point of solution – boiling point of pure solvent
• ΔTb = boiling point elevation (units: K or deg C)
• kb = molal boiling point elevation constant (depends on the solvent) (units: K/m or deg C/m)
• i = van’t hoff factor (no units)
o A solute that breaks up into “pieces” has more solute particles and has a different effect than one that does not break up
o i = the van’t hoff factor or mole number (i.e., the number of “pieces” the substance breaks up into when it dissolves)
o For a nonelectrolyte (molecular substance), i =1 since it doesn’t break apart into ions
o For an electrolyte (ionic substance), i = the number of ions the substance breaks into
o ex: Na2SO4 would have i = 3 because 3 ions result
Freezing Point Depression
Freezing point depression results because solute particles interfere with crystallization (ex: antifreeze, salting icy roads, making ice cream)
ΔTf = kf m i = freezing point of pure solvent – freezing point of solution
ΔTf = freezing point depression (units: K or deg C)
kf = molal freezing point depression constant (depends on the solvent) (units: K/m or deg C/m)
i = van’t hoff factor (no units)
IMPORTANT: You can compare solutions in terms of their concentrations and mole numbers to determine which one would experience a greater freezing point depression or boiling point elevation. BE CAREFUL when doing this… a bigger freezing point
DEPRESSION means that the freezing point will be LOWER than usual… a bigger boiling point ELEVATION means that the boiling point will be HIGHER than usual.