Ask a chemist, they always have

 Mixture: several pure substances mixed together in an indefinite ratio

Homogeneous

Heterogeneous
 Solution: a homogenous mixture that form when one or more substances dissolve into another.
 Suspensions: cloudy mixtures that form when two or more substances mix but do not dissolve.
 Emulsions: suspension of 2 liquids

 Solution: a homogenous mixture
 Solute: thing that dissolves
 Solvent: thing that does the dissolving
(found in the largest amounts)

If the solvent is water, then it is called an aqueous solution

 Example: iced tea


Solute sugar tea
Solvent water

 Ions form, separate (dissociate) and move throughout the solution

The forces that hold the ions together are overcome by the ions ’ attractions to polar water.
• Ion- dipole interaction
 Because ions are present, ionic solutions can conduct a current

Current is just movement of electrons

Solvation animation
Animation with Audio

Figure 15.2: Polar water molecules interacting with positive and negative ions of a salt.

 “ Like dissolves like ”
 Typically, hydrogen bonding occurs between the substance being dissolved and the polar water molecules
 Example:

Sugar in water

Ethanol in water

Figure 15.3: The polar water molecule interacts strongly with the polar O —H bond in ethanol

Get interactions between water molecules the polar regions on the sugar (the Os) , and some hydrogen bonding at the -OH groups

(things that don ’ t dissolve or mix)
 Anything nonpolar will not mix well with anything polar

Examples:
• Oil spill
• Salad dressing
 Can mix when shaken
(LDF) and then may separate out (other forces)

 There is a limit to the amount of a substance dissolved

Saturated : the solution holds as much solute as possible at that temperature .

Unsaturated : solution has not reached the limit

 Can you have too much? YES!


Supersaturated: have as much solute dissolved as possible, then cooled and all the solute stays dissolved.
In other words…the solution contains more dissolved solid than a saturated solution created at the same temperature.

These can begin crystallization of the solute at the slightest change

 Any chemical change (including solvation) requires a change in energy


Energy removed from or added to the reactants from the surroundings
NaOH
(s)
 Na +
(aq)
+ OH -
(aq)
ΔH= -44.5 kJ/mol
(that ’ s 44.5kJ released, so exothermic, per mole of NaOH)
• Because you are breaking the ionic bond, energy must be either released when breaking the bond, or consumed when making the new ions


ALL changes in formula indicate a change in energy.
However, sometimes the energy change is so small, you can ’ t tell that a change has occurred

 “ Strong ” coffee has more coffee dissolved in a given amount (say 1 pot) than
“ weak ” coffee.


Strong coffee = concentrated
Weak coffee = dilute
 Concentration: the amount of solute in a given amount of solvent (or solution).

M
 Most common way to express concentration
 Molarity is the number of moles of solute dissolved in each liter of solution
 Formula

M
= moles of solute liters of solution
 Dependent on temperature
 The higher the molarity the stronger the concentration

M
 Another way to calculate concentration
 Formula

M
= moles solute .
kilograms of solvent
 Not dependent on temperature
 The higher the molality the stronger the concentration

 In the winter, why do we throw salt when it snows?
 Why does Emeril add salt to boiling water when cooking pasta?


By adding salt (or other solutes) to water, the temperature of freezing drops  it freezes at a lower temperature
• Because H bonding is disturbed
• Dependent on how much solute is added

 Antifreeze protects cars from freezing and overheating. Calculate the freezing point depression of a solution of 100. g of ethylene glycol (C
K f water
2
H
6
O
=
1.86 o C/ m
2
) antifreeze in 0.500 kg of water.
 Formula:
T f
= K f m i


 K f
:
Molal Freezing Point depression constant ( o C/ m ) i= Pieces that the material dissociates into (for ionic compounds only)
(Keep I at 1 (one) for covalent compounds)

Freezing Point Depression and Boiling Point Elevation
Solvent Formula
Water H
2
O
Acetic acid HC
2
H
3
O
2
Benzene C
6
H
6
Camphor
Carbon disulfide CS
2
Cyclohexane C
6
H
12
Ethanol
C
10
H
16
O
C
2
H
5
OH
Melting
Point
( ° C)
0.000
16.60
5.455
179.5
...
6.55
...
Boiling
Point
( ° C)
100.000
118.5
80.2
...
46.3
80.74
78.3
K f
( ° C/m)
1.858
3.59
5.065
40
...
20.0
...
K b
( ° C/m
)
0.521
3.08
2.61
...
2.40
2.79
1.07


By adding salt (or other compounds) to water, the temperature of boiling goes up  it boils at a higher temperature
• Interrupts H bonding
• Need more vapor molecules and greater pressure to get bubbles to form
• Takes more time to get vapors to add to bubbles
• The molecules that do get into the bubbles need more energy

Dependent on how much solute is added

 Water with salt added boils at a higher temperature than pure water. By how much will the boiling point change if 100.g of salt is added to 500. g of water?
K b water
=
0.52 o C/ m
 Formula:
T b
= K b m i


K b
:
Molal Boiling Point elevation constant ( o C/ m ) i= = Pieces that the material dissociates into (for ionic compounds only)
 (Keep I at 1 (one) for covalent compounds)

Freezing Point Depression and Boiling
Point Elevation
Solvent Formula
Melting
Point
( ° C)
0.000
Boiling
Point
( ° C)
100.000
K f
( ° C/ m) (
1.858
K b
( ° C
/m)
0.521
Water H
2
O
Acetic acid HC
2
H
3
O
2
Benzene C
6
H
6
Camphor
Carbon disulfide CS
2
Cyclohexane C
6
H
12
Ethanol
C
10
H
16
O
C
2
H
5
OH
16.60
5.455
179.5
...
6.55
...
118.5
80.2
...
46.3
80.74
78.3
3.59
5.065
40
...
20.0
...
1.07
3.08
2.61
...
2.40
2.79

Figure 15.10: Pure water.

Figure 15.9: A bubble in the interior of liquid water surrounded by solute particles and water molecules.

Figure 15.10: Solution (contains solute).

 Vapor pressure changes as IMFs change
 For the same reasons boiling point is disturbed
 What would evaporate faster:

Salt water

Distilled water

WHY?

 Colligative properties interactive