Aqueous Solutions
Tuesday March 22 and Wednesday March 23
I. Solution Types of Mixtures Table 3, pg 404
A. Homogenous - all samples are identical.
B. Particle size is less than 1nm, cannot be filtered or settle out.
C. Can be solid (metal alloys: bronze, steel, sterling silver), liquid (salt water), or gas (air)
D. Two parts to a solution
1. Solvent - does the dissolving (example: water - the universal solvent)
2. Solute - the substance being dissolved
E. Soluble - able to be dissolved.
II. Suspensions
A. Heterogeneous – all samples are not identical
B. Solute particle size is larger than 1000nm
C. Particles are large enough to be filtered out (coffee grounds)
D. Particles settle out after mixture stands undisturbed (sand in beach water)
E. Scatter visible light. Like sun through dust in the air of a room.
III. Colloids
A. Heterogeneous - all samples are not identical
B. Solute particles are of medium size, between 1nm and 1000nm.
C. Cloudy looking. Examples: glues, gelatins, paints, smoke, muddy water
D. Can NOT be filtered and will NOT settle out.
E. Tyndall Effect – transparent particles > 1nm scatter visible light in all directions.
1. Brownian Motion - chaotic movement of solute particles.
2. Like headlights in the fog
3. Solution particles are too small. Suspensions are not transparent.
4. Helps distinguish between a solution and a colloid (pg404)
IV. When two liquids try to mix....
A. Miscible - two liquids that can dissolve each other. Example: water and alcohol
1. “Like dissolves like”
2. Polar solvents dissolve polar and ionic solutes (charges attract).
Hydrogen bonds & dipole interactions pull apart charged solute ions
Exceptions: Water will not dissolve BaSO4 or CaCO3 because the ionic bonds holding the molecule together are stronger than the hydrogen bonds trying to pull them apart.
3. Nonpolar solvents dissolve nonpolar solutes.
London dispersion forces
Example: Nail polish remover dissolves Styrofoam.
Some nonpolar substances: gasoline, oil, fat, grease
B. Immiscible - two liquids that cannot dissolve each other. Example: water and oil
1. Need an emulsifying agent, like soap, to help mix immiscible solutions.
Ex: mayonnaise (oil and vinegar with egg yolk…lecithin is emulsifying agent)
2. Emulsions are a type colloid – Two liquids that normally will not mix.
3. Solvent = the greater percentage & Solute = the lower percentage
V. Electrolytes - compounds that conduct an electric current when in solution.
A. Necessary for life - conduct continuous flow of energy throughout body.
B. Strength depends on different degrees of ionization, NOT concentration of solute.
B. Strong Electrolytes - ionic compounds: charged ions separate completely when
dissolved in water (aqueous) or molten (melted).
C. Weak Electrolytes - aqueous polar compounds: only a fraction of solute exists as ions
when dissolved in water - most ions remain bound in compound
D. Nonelectrolytes - nonpolar molecules
1. organic compounds such as alcohols and sugars
2. compounds usually contain carbon.
3. glucose and glycerol, methane, grease, gasoline
VI. Rate of Dissolution - how fast a solute goes into solution.
A. Agitation - shake, mix, stir - get fresh solute in contact with solvent.
1. Only increases how fast, NOT how much solute goes into solution. “Shake it.”
B. Temperature – inc. temp, increase energy, increase force & frequency of collisions.
1. Increases how fast AND how much solute a solution can hold. “Bake it.”
C. Particle Size - the smaller the particle size, the greater the surface area of the particle exposed to the solvent. Sugar cube verses granules.
1. Only increases how fast NOT how much solute goes into solution. “Break it.”
VII. Solubility - how much solute goes into solution (Figure 15, pg 414)
A. Saturated solution - maximum solute in a given quantity of solvent at constant temp.
1. At equilibrium. Appears clear!!
2. Equilibrium - rate of dissolution (dissolving) = rate of crystallization
3. Formulas -
a. solubility = xg solute / 100g solvent (Table 4 on pg 410)
b. mass solute = solubility of solute x mass of solvent
Example: How much KCl can be dissolved in 350g of H2O at 50˚C?
B. Unsaturated solution - less than the maximum solute in a given amount of solvent at
constant temp.
1. Appears clear.
C. Supersaturated solution - more solute than it can theoretically hold at given temp.
1. Add solute when solution is hot and set aside to cool undisturbed.
2. No un-dissolved solute - appears clear!!!!
3. Crystallization of excess solute can be initiated by a single “seed” crystal.
Example: Seeding rain clouds with AgI causes water vapor to condense and drop
Questions: What could you do to make a saturated solution unsaturated?
What could you do to make an unsaturated solution saturated?
VIII. Factors that affect solubility (how much solute is able to be dissolved):
A. Temperature
1. Solid as solutes - increase temp, energy, and collisions, increases solubility
Example: Hot Tea and Sugar
2. Gas as solutes - increase temp, energy, ability to escape, decreases solubility.
Example 1: Thermal Pollution - increase temp, increases O2 escaping into
the air, increases death in lake.
Example 2: Open Soda - increase temp, carbon dioxide escapes faster, and
soda goes flat faster.
B. Pressure - increase pressure increase solubility of a gas!
1. Henry’s Law - solubility of gas is directly proportional to pressure
above liquid. Example: Sealed coke, lots of pressure, keeps gas in
(soluble). Open soda, less pressure, gas escapes (less soluble
a. Effervescence: the quick release of gas particles from a solution
b. Formula: S1 = S2
P1 P2
Example: Scuba diving bends: Nitrogen gas was absorbed by the blood at deep ocean pressure.
Question: If the solubility of the gas in water is 0.77 g/L at 350kPa, what is its solubility, in g/L at 100 kPa? (Temp. is constant at 25°C.)
IX. Water of Hydration - the water in a solid crystal
A. Hydrate - crystal compounds containing water - looks dry
Example: CuSO4-5H2O = 1molecule CuSO4 is connected to 5molecules H2O
B. Anhydrous - dry, no water in crystal, dehydrated, powdery
C. Efflorescent - the process of losing water, becoming dry.
Example: Wintergreen Certs - bite with mouth open in dark closet, see sparks.
D. Hygroscopic - compound absorbs water from air - may dissolve at RT
Example: NaOH - seal container tight to avoid ruining.
X. Molar enthalpy of dissolution - amount of energy released or gained by solute as it dissolves.
A. ΔHsol (solution) (Table 5, pg 416)
Thursday March 24 and Friday March 25
XII. Concentration - measurement of how much solute is in a given amount of solvent.
A. Diluted Solution - small amount of solute, less dense, qualitative description
B. Concentrated Solution - large amount of solute, more dense, qualitative description
C. Molarity - M - number of moles of solute dissolved in 1L of solution –
molar concentration - quantitative description.
1. Formula: Molarity (M) = number of moles of solute
number of liters of solution
2. NOTE: Liters of solution is different from liters of solvent.
Example: 1M = 1 molar solution = 1 mole solute
1 liter of solution
Example: How many moles of HCl are in a sample of 500ml of 12 M HCl?
Question: If you added more solvent (water) to the above example, how many moles
of HCl would be left?
D. Making Dilutions - moles of solute before dilution equals the moles of solute after
1. Dilution only changes the solvent amount.
2. Formulas: M (molarity) = mol (moles)
V (liters) mol = M x V M1V1 = M2V2
Example 1: What is the molarity of a solution made by dissolving 60g NaNO3 in
enough water to make 500ml solution?
Example 2: How would you prepare 100ml of 0.40M MgSO4 from stock solution of 2.0M MgSO4?
E. Molality (not molarity this time) – variable Greek letter mu = μ
1. a way to express the solute to solvent ratio, molal concentration
2. number of moles of solute dissolved in 1kg of solvent
3. Formula: molality = mole solute
Kg solvent
Ex: 2µ = 2 molal solution = 2 mole solute = 2 mole solute
1kg solvent 1000g solvent
Note: Remember the following conversion: 1g H2O = 1ml H2O and 1kg = 1000g
Example: How many grams of KI must you dissolve in 500g of water to produce
a 0.06 molal KI solution?
Example: Calculate the molality of a solution prepared by dissolving 10g of NaCl in 600g of water.
XIII. Solubility Rules page 437
Monday March 28 and Tuesday March 29
XIV. Colligative Properties of Solutions – depends on the number of particles dissolved in a
given mass of solvent - Concentration
A. Vapor Pressure Lowering
1. Solvent particles are held down by the solute particles.
2. No need to have a lot of pressure pushing down on the surface of the solvent.
3. The more solute added to solvent, the lower the vapor pressure needs to be to
prevent the solvent from evaporating.
4. The greater the number of particles, the more the vapor pressure decreases.
Example: 5mol of NaCl (2x particles when dissolved) verses 5mol of CaCl2 (3x particles when dissolved)… the same number of compound moles of CaCl2 lowers the vapor pressure more than NaCl.
B. Boiling-Point Elevation – more solute holds onto more solvent
1. solvent requires more energy to break free from the liquid phase
2. boiling point must be higher to obtain the greater quantity of energy.
3. Formula ΔTb = (kb)(molality) kb is a given constant found on pg448
Old BP + ΔTb = New BP (elevated)
C. Freezing-Point Depression – solute disrupts structured pattern of pure frozen solvent
1. makes it harder to freeze, solvent must lose more energy prior to solidifying
2. Therefore, the FP will be lower than normal.
3. Example: sprinkle salt on an icy wet sidewalk – the salt makes it more difficult
for the ice (water) to stay frozen (freeze) – the FP is lower because more energy
must be lost prior to the water being able to solidify.
4. Formula ΔTf = (kf)(molality) kf is a given constant found on pg448
Old FP + ΔTf = New FP (depressed)
D. Molar Mass – use changes in BP and FP to determine the molar mass
1. Add known mass of solute to a know mass of solvent
2. Look up the kb and kf constants
3. Measure the changes in BP and FP
4. Calculate the molar mass
Example: A solution of 7.50g of a nonvolatile compound in 22.60g of water boils at
100.78˚C at 760mmHg. What is the molar mass of the solute? Assume that the solute exists as molecules, not ions.