The Structure of Water
Water is a Polar Molecule. This means that
it has a positive and a negative end.
This is why water is “sticky.”
It also allows us to BEND
WATER!!!!
Why is water polar?
It has to do with something called
ELECTRONEGATIVITY. This is an atom’s
ability to attract electrons. There is a trend
of electronegativity on the periodic table.
The Solution Process
• Ionic and covalent molecules dissolve
differently
– Covalent – share e- (from the RS of the PT)
– Ionic – give and take e- (from opposite sides
of the PT)
• Dissociation – when the ions separate
from the crystal structure
– We write dissociation equations like this:
AgCl(s)  Ag+(aq) + Cl-(aq)
(Pay attention to the states!)
• Solvation – the process of surrounding the
solute ions with solvent molecules
• Hydration – when the solvation process
involes water as the solvent. The solute
ions are said to be HYDRATED.
Heat of Solution
• E is required to break attractive forces.
• Therefore, the separation of solute
particles from one another and the
separation of solvent particles from one
another are endothermic processes. THEY
USE HEAT.
– Example:
• The attraction between solute and solvent
particles during solvation is exothermic.
THEY RELEASE HEAT.
– Example:
“Heat of solution” is the next energy
change!
Energy Absorbed > Energy Released
ENDOTHERMIC (cold)
Energy Absorbed < Energy Released
EXOTHERMIC (hot)
Assignment:
1. Ex 44 1-10
2. Draw diagrams of the solution process for
a. An ionic compound
b. A covalent molecule
This assignment is due tomorrow!!!
Factors that Affect the Rate of
Solubility
1. Agitation (stirring)
2. Temperature
3. Surface Area
Terms you need to know….
Soluble – capable of
being dissolved in a
particular solvent
Miscible – liquids that
mix in all proportions
and have no max
concentration
Insoluble – cannot be
dissolved in a
particular solvent
Immiscible – two
liquids that form
separate layers
instead of dissolving
Saturated – when a solution contains the max
amount of solute it can dissolve at a given
temp
Unsaturated – when a solution doesn’t contain
the max amount of solute it can dissolve at a
given temp
Supersaturated – more than the max amount of
solute dissolved for a specific temp (made by
dissolving at high temp and cooling)
Saturated solution
• The animation represents a saturated solution:
• the blue "molecules" escape into solution from
the ordered crystal. At the same time, molecules
are coming out of solution and depositing on
the solid. Since this is a continual process and
the concentrations do not change, it is called
dynamic equilibrium.
Solubility – the concetration of a saturated
solution at a specific Temp and Pressure
Ex – the solubility of Na2SO4 at 0°C is
4.76 g/100mL H2O
Solubility curve
Saturated
Supersaturated
Unsaturated
3 Ways to Saturate a Solution
1. Add more solute to the saturation point
2. Evaporate the solvent
and……
3. Decrease the temperature to the saturation
point
SOLUBILITY
Depends on:
1. Temperature
2. Pressure
3. Chemical natures of the solute and solvent
Temperature
•In general, as the temp increases, so does
solubility (for endothermic reactions)
Example: Cold packs - The dissolving reaction
is endothermic - requires heat. Heat comes
from the surroundings, and the pack feels cold.
• If the reaction is exothermic, the
increase in heat will actually cause the
solubility to DECREASE.
• This is not very common.
Reading graph: at 38 °C the solubility of copper sulphate, CuSO4, is
28g per 100g of water.
Reading graph: at 84 °C the solubility of potassium sulphate, K2SO4, is
22g per 100g of water.
Ex Q1: How much potassium nitrate will dissolve in 20g of water at 34 °C?
At 34 °C the solubility is 52g per 100g of water, so scaling down, 52 x 20 / 100 = 10.4g
will dissolve in 20g of water.
Ex Q3: 100 mL of saturated copper sulphate solution was prepared at a temperature of
90 °C. What mass of copper sulphate crystals form if the solution was cooled to 20 °C?
Solubility of copper sulphate at 90 °C is 67g/100g water, and 21g/100g water at 20 °C. Therefore for
mass of crystals formed = 67 - 21 = 46g (for 100 cm3 of solution)
Exercise 43: 1-9
Solubility of Gases
• The solubility for gases always decreases when the
temperature increases.
Solubility of Gases: Temperature
• ↑ Temp = ↑ KE Particles will move more, break IMFs
and the gas will escape!
• Ex – pop that has been sitting open at room temp
will be flat as the CO2 has come out of solution.
Pressure and Solubility of Gases
• With a change in pressure
- liquids and solids show
no change in solubility
•
Gases increase in
solubility with an increase
in pressure.
• If the pressure is
increased, the gas
molecules are "forced"
into the solution. The
concentration of gas
molecules in the solution
have increased!
Pressure and Solubility of Gases
• Carbonated beverages provide the best example of
this phenomena. All carbonated beverages are
bottled under pressure to increase the carbon
dioxide dissolved in solution.
• When the bottle is opened, the pressure above the
solution decreases. As a result, the solution
effervesces and some of the carbon dioxide bubbles
off.
• Quiz: Champagne continues to ferment in the bottle.
The fermentation produces CO2. Why is the cork
wired on a bottle of champagne?
• Answer: As more CO2 is formed , the pressure of the
gas increase.The wire is to prevent the cork from
blowing off.
• Quiz: If a diver had the "bends",
describe how this can be treated.
• Answer: Decompression chambers are
used to keep a high pressure and
gradually lower the pressure.
• Quiz: The amount of dissolved oxygen in a
mountain lake at
10,000 ft and 50oF is __?_ than the amount of
dissolved oxygen in a lake near sea level at
50oF.
• Answer: Less at higher altitude because less
pressure.
• A Coke at room temperature will have __?_
carbon dioxide in the gas space above the
liquid than an ice cold bottle.
• Answer: More gas, because the warm coke
can hold less of the gas in solution.
BONUS ASSIGNMENT
• THERMAL POLLUTION
– 3 page research paper on thermal pollution.
– Include references
– If you plaigerise, you get ZERO.
Freezing Point Depression
• This happens when the freezing point is
lowered by adding a solute to a solvent.
• In order for a liquid to freeze, it must
achieve a very ordered state that results in
the formation of a crystal.
• If there are impurities (solute) in the liquid,
the liquid is less ordered – the sol’n is now
more difficult to freeze and a lower temp is
required.
Which decreases FP more?
• NaCl
• CaCl2
NaCl  Na+ + ClCaCl2  Ca 2+ + 2ClWhich has more particles????
Boiling Point Elevation
Review – VP – the pressure at which a liquid
is in equilibrium with its vapour.
Boiling – the temp at which the vapour
pressure equals the pressure above the
liquid
• If we add solute to a solvent, the vp of the
sol’n is lowered. This happens because:
– At the surface of the sol’n (where evaporation
occurs) there are fewer solvent particles due
to the presence of solute particles – lowers vp
– The solute particles absorb energy and will
reduce the energy available to evaporate the
solvent – lowers vp
– Energy is required to overcome the IMFs
between the solute and solvent particles –
lowers vp
• If VP is lowered, the temp must be raised
to have the VP = the P above the solution
• Example
– H2O boils at 100°C
– H2O and NaCl boils at > 100°C
You need to remember:
• Adding a solute to a solvent will:
– LOWER THE FREEZING POINT
– RAISE THE BOILING POINT
• Make sure you’re able to tell me why!!!!
LAB TOMORROW!!!
• You need to get into a group
• You need to bring:
– 2 small ziplock bags
– 2 large ziplock bags
– 1 cup of sugar
– 500 mL of milk
– toppings
Ex 49: 1-8
Like Dissolves Like
• Remember – “like dissolves like”
• This means that polar molecules dissovle
polar molecules
• Non-polar molecules will dissolve in nonpolar molecules
Units of Concentrations
amount of solute per amount of solvent or solution
Percent (by mass) =
Molarity (M) =
g solute
g solution
x 100 =
g solute x 100
g solute + g solvent
moles of solute
volume in liters of solution
moles = M x VL
Examples
What is the percent of KCl if 15 g KCl are
placed in 75 g water?
%KCl = 15g x 100/(15 g + 75 g) = 17%
What is the molarity of the KCl if 90 mL of
solution are formed?
mole KCl = 15 g x (1 mole/74.5 g) = 0.20 mole
molarity = 0.20 mole/0.090L = 2.2 M KCl
Examples:
Example 1: What is the concentration when 5.2 moles of
hydrosulfuric acid are dissolved in 500 mL of water?
Step one:
Convert volume to liters, mass to moles.
500 mL = 0.500 L
Step two:
Calculate concentration.
C = 5.2 mol/0.500 L
= 10mol/L
• Example 2:
What is the volume when 9.0 moles are present in
5.6 mol/L hydrochloric acid?
• Example 3:
How many moles are present in 450 mL of 1.5
mol/L calcium hydroxide?
• Example 4:
What is the concentration of 5.6 g of magnesium
hydroxide dissolved in 550 mL?
• Example 5:
What is the volume of a 0.100 mol/L solution that
contains 5.0 g of sodium chloride?
How many Tums tablets, each 500 mg CaCO3,
would it take to neutralize a quart of vinegar,
0.83 M acetic acid (CH3COOH)?
2CH3COOH(aq) + CaCO3(s)  Ca(CH3COO)2(aq) + H2O + CO2(g)
a quart
moles acetic acid = 0.83 moles/L x 0.95 L = 0.79 moles AA
the mole ratio
mole CaCO3 = 0.79 moles AA x (1 mole CaCO3/2 moles AA)
= 0.39 moles CaCO3
molar mass
mass CaCO3 = 0.39 moles x 100 g/mole = 39 g CaCO3
number of tablets = 39 g x (1 tablet/0.500g) = 79 tablets