1. Ch. 14/15 – Solids, Liquids and Solutions

2. Intermolecular Forces
Forces of attraction between different molecules rather than bonding forces within the same molecule.
Dipole-dipole attraction
Between particles with charged sides ml
Hydrogen bonds - Strongest
Between H’s & F, O, N on other molecules
Dispersion forces - Weakest
Caused by electrons shifting towards one end of a molecule.

3. “Water: The Magnificent Dipole”
One side of water is negatively charged because the oxygen atom keeps the shared
electrons longer than the hydrogen atoms. As a
result the oxygen side is
negatively charged and the hydrogen side of water is positively charged. O

4. It is its polarity and
hydrogen bonding that give
water its many unusual
Properties!
i.e. high boiling point, expansion upon freezing, and surface tension

5. Hydrogen bonds - Strongest Between H’s & F, O, N on other molecules
Ice
Liquid

6. Water is always trying to pull itself into a tight ball as long as there is nothing nearby that has a charge on it. Therefore, this surface is not repelling water; it’s simply not attracting it and keeping water from doing what it does naturally.
Water pulls on itself
so much that it forms a “skin.”
It’s called surface tension.

7. Wax does not repel water
We’ve heard that wax or oils repel water. But that isn’t true. Water is so attracted to other water molecules that anything between them is squeezed out of the way. O O
Oil droplet O O O

8. Forces and Phases
Substances with very little intermolecular attraction exist as gases
Substances with strong intermolecular attraction exist as liquids
Substances with very strong intermolecular (or ionic) attraction exist as solids

9. “V” is for Vocabulary!
Vapor – gaseous state of a
substance that is not normally
a gas at room temperature.
Volatile – evaporates rapidly
(due to weak intermolecular forces)
Ex. ‘thin’ liquids
Viscous – evaporates slowly
(due to strong intermolecular forces)
Ex. ‘thick’ liquids

10. Phase Differences

11. Three Phases of Matter

12. Liquids
A decrease in the average kinetic energy of gas particles causes the temperature to decrease.
As it cools, the particles tend to move more slowly, if they slow down enough, attractive forces - called van der Waal’s forces –pull them very close together so they can only slip & slide past each other.
Condensation –
Change of a gas
to a liquid
It is now in
liquid form!

13. The Nature of Liquids
The conversion of a liquid to a gas or vapor below its boiling point is called Vaporization
(occurs at the surface of a liquid)
In an open container, this process is called Evaporation
Particles near the surface with enough kinetic energy that happen to bounce in the right direction escape!

14. Microscopic view of a liquid near its surface.

15. The Nature of Liquids
Eventually the particles will lose energy and return to the liquid state, or condense.
What are the odds that they will return to the original liquid?
What if we cover the container?
So, the particles begin to evaporate, then some begin to condense. Eventually, the number of particles evaporating will equal the number condensing & the space above the liquid will be saturated with vapor

16. A dynamic equilibrium now exists where
Rate of evaporation = rate of condensation
Note that there will still be particles that evaporate and condense
But, there will be no NET change
It will not look like there are changes taking place!

17. Evaporation is a cooling process Cooling occurs because particles with the highest energy escape first
Particles left behind have lower average kinetic energies; thus the temperature decreases
Similar to removing the fastest runner from a race- the remaining runners have a lower average speed
Evaporation helps to keep our skin cooler on a hot day!

18. The Nature of Liquids
A liquid will evaporate faster when heated because the added heat increases the average kinetic energy needed to overcome the attractive forces so more particles have enough energy to ‘escape’!

19. If you were to add a
drop of water below
the tube to the left
what would happen?
It would rise to the
top & evaporate.
What would it do to
the surface of the
mercury?
(push it down!)
Vapor Pressure –
pressure exerted by
vapor!

20. (
Since different liquids evaporate at different rates, they have different vapor pressures (at the same temps)

21. ( Which liquid is the most volatile?
Yes! Diethyl Ether! It depressed the mercury the most (highest vapor pressure!)

22. Vapor Pressures of Liquids
Which is the most volatile liquid here?
Diethyl ether
It has the highest vapor pressure at any
Temp.
Which has the strongest forces of attraction?
Water

23. A liquid boils when its vapor pressure equals the external pressure, so the boiling point changes if the external pressure changes.
Bubbles form throughout the liquid, rise to the surface, and escape into the air
Normal boiling point- is when the vapor pressure of a liquid equals standard pressure. (1 atm)

24. The boiling point (bp) is the temperature at which the vapor pressure of the liquid is equal to the external pressure on the liquid VP BP
Direct relationship!

25. Normal bp of water = 100 oC
However, in Denver = 95 oC, since Denver is 1600 m above sea level and average atmospheric pressure is about 85.3 kPa (Recipe adjustments?)
In pressure cookers, which reduce cooking time, water boils above 100 oC due to the increased pressure

26. Boiling Point of Water at Various Locations
Feet above sea level
Patm (kPa)
Boiling Point (C)
Top of Mt. Everest, Tibet
29,028 32 70
Top of Mt. Denali, Alaska
20,320
45.3 79
Top of Mt. Whitney, California
14,494
57.3 85
Top of Mt. Washington, N.H.
6,293
78.6 93
Boulder, Colorado
5,430
81.3 94
New York City, New York 10
101.3
100
Death Valley, California
-282
102.6
100.3

27. Vapor Pressures of Liquids
Normal bp when
crossing here
At any pt. on a
curve line,
liquid is boiling

28. SOLIDS
If you cool a liquid, the particles lose kinetic energy and slow down.
If they slow down enough, extra forces of attraction pull them in so close together that they can only vibrate in place.
Freezing – change of a liquid to a solid.

29. Types of Solids Molecular solids Metallic solids Ionic solids
Covalent network solids (diamonds)

30. Crystals or Crystalline Solids
Particles of crystals are arranged in repeating geometric patterns
NaCl

31. Representation of Components in a Crystalline Solid
Crystal Lattice:
The orderly, regular, 3-dimensional arrangement of particles (atoms, ions, etc.) in a crystal.

32. Unit Cell
The smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice
Ex. A cubic lattice system has 3 types of unit cells

33. Table salt crystals are shaped like cubes.

34. Diamond, a form of carbon, is also a crystalline solid.
the crystals are shaped something like pyramids.

35. Non-crystalline solids
Many solids do not form crystals- Amorphous
Their molecules do not arrange into repeating patterns
often because they are too large.
No definite melting point
Examples:
Glass - also called a super-cooled liquid
many plastics, soot, asphalt, butter

36. PHASE CHANGES
PHASE CHANGES – change is physical state (melting, freezing, boiling, condensing, sublimation, deposition)
BOTH PHASES present during a phase change
Temperature remains constant during a phase change.
Sublimation – change of a solid directly to a gas (dry ice, iodine, snow)
Deposition – change of a gas directly to a solid.

37. Temperature (C°)
60°
20° 0°
-20°
Heat (kilojoules)

38. Heating and cooling curve for water heated at a constant rates.
A-B = Solid ice, temperature is increasing.
Particles gain kinetic energy, vibration of particles increases.
Ice

39. B-C = Solid starts to change state from solid to liquid
B-C = Solid starts to change state from solid to liquid. Temperature remains constant as energy is used to break inter-molecular bonds.
H2O (s)  H2O (l)
0ºC

40. C-D = temperature starts to rise once all the solid has melted
C-D = temperature starts to rise once all the solid has melted. Particles gain kinetic energy.
Liquid water

41. D-E = Liquid starts to vaporize, turning from liquid to gas
D-E = Liquid starts to vaporize, turning from liquid to gas. The temperature remains constant as energy is used to break inter-molecular forces.
H2O ()  H2O (g)
100ºC

42. E-F = temperature starts to rise once all liquid is vaporized
E-F = temperature starts to rise once all liquid is vaporized. Gas particles gain kinetic energy.
steam

43. The heating/cooling curve for water

44. Water phase changes constant Temperature remains __________
during a phase change.
constant
Water phase changes
KE not changing during
phase changes
Boiling
condensation
Melting
freezing
Kinetic energy increasing on slopes

45. Phase Diagram
Represents phases as a function of temperature and pressure.

46. Classification of Matter
Recall
Classification of Matter
Matter
Pure
Substances
Elements
Compounds
Mixtures
Homogeneous
Heterogeneous
Also called solutions

47. Solute Solvent Aqueous Solutions – water solutions
The part of a solution that gets dissolved – the part in lesser quantity
Salt in salt water
Sugar in soda drinks
Carbon dioxide in soda drinks
Solvent
The part of a solution that does the dissolving – the part in greater quantity
Water in salt water
Water in soda

48. Solutions occur in all 3 phases!
Solute
Solvent
Example
Gas
Air
Liquid
Water Vapor in Air
Solid
Mothballs
Seltzer Water
Antifreeze in radiator, cocktail
Salt water
Whipped cream
Fillings: Hg in Ag
Alloys: Brass, etc.

49. non-polar molecules are soluble in non-polar solvents
Two substances with similar intermolecular forces are likely to be soluble in each other.
non-polar molecules are soluble in non-polar solvents
Ex. Grease in gasoline
Ionic compounds & polar molecules are soluble in polar solvents
Ex. Ethanol in water; salt in water
“like dissolves like”
12.2

50. When a salt dissolves in water, the positive
ions are attracted to the slightly negative
ends of the water molecules and the negative ions are attracted to the slightly positive ends of the water molecules. They ‘dissociate’ (separate)
The ions are more strongly attracted to each
other. but they become surrounded by water molecules and can’t get back together!
Solvation – where solvent molecules
surround solute particles.

51. Dissolving of solid sodium chloride.

52. Water: the Universal Solvent
Cl-
Na+ O O O

54. SOLUBILITY A measure of how much a gas, liquid, or solid will dissolve
in a solvent.

55. Factors Effecting Solubility
The solubility of MOST solids increases with an increase in temperature.
The rate at which solids dissolve increases with increasing surface area of the solid.
(crush them & stir them!)
The solubility of gases increases with decreases in temperature.
The solubility of gases increases with an increase of pressure above the solution.

56. Therefore… Solids tend to dissolve best when: Heated Stirred
Ground into smaller particles
Gases tend to dissolve best when:
The solution is cooler
Pressure is higher

57. When you see bubbles, it’s a heterogeneous mixture not a solution!
CO2 in and out of water
Mixture
Solution
When you see bubbles, it’s a heterogeneous mixture not a solution!

58. Salt (NaCl) is very soluble in water. 350 g/liter
Salt (NaCl) is very soluble in water g/liter. However, if water evaporates, there will be too much salt for the water to hold in solution. The salt begins to form crystals.
A lake near Death Valley is supersaturated with salt causing the salt to crystallize out.

59. Once a year the company who owns the lakes lets visitors into the area to collect salt crystals.

60. These are some of the salt crystals collected.

61. Solubility
maximum mass of solute that will dissolve in 100 g of solvent at a given temperature
varies with temp
based on a saturated solution

62. Temperature and Solubility
Solid solubility and temperature
solubility increases with increasing temperature
solubility decreases with increasing temperature
12.4

63. Solubility Chart

64. A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated.
A solution that contains less solute than it can is unsaturated.
A solution that contains more dissolved solute than it should is supersaturated.

65. Solubility UNSATURATED SOLUTION more solute dissolves
no more solute dissolves
SUPERSATURATED SOLUTION
becomes unstable, crystals form
concentration

66. Solubility Chart

67. Colligative Properties
Colligative Properties are those that depend on the concentration of particles in a solution, not upon the identity of the solute.
Boiling Point Elevation
Freezing Point Depression
Vapor Pressure Depression
More particles in solution – bigger effect
on the coll. property!
Higher bp, lower fp, etc.

68. Change in Boiling Point
Common Applications of Boiling Point Elevation

69. Change in Freezing Point
Common Applications of Freezing Point Depression
Ethylene glycol – deadly to small animals
Propylene glycol

70. Concentration

71. Concentrations of Solutions
The amount of solute in a solution is given by its concentration.
Concentration can be expressed qualitatively or quantitatively –
Dilute – contains small amount of solute
Concentrated – contains large amount of solute
(words)
(numbers)

72. Molarity = ( M ) moles solute liters of solution
We will use something called Molarity!
Molarity ( M ) =
moles solute
liters of solution
*Volume of solution MUST be in Liters!
Ex. What is the molarity of a solution made
by dissolving 3.5 moles of NaCl in a 4.3L solution?
3.5 moles NaCl
= 0.81 M NaCl(aq)
4.3 L solution
Molarity can be labeled “M” or “mol/L”
Handy in conversion factors!

73. PROBLEM: Dissolve 5.00 g of NiCl2 in enough water to make 250 mL of solution. Calculate the Molarity.
Step 1: Calculate moles of NiCl2
Step 2: convert mL to L
250 mL L
Step 3: Calculate Molarity
[NiCl2] = M

74. moles = M•V What mass of oxalic acid, H2C2O4, is
required to make 250. mL of a M
solution?
*Rearrange formula
to solve for moles.
moles = M•V
Step 1: Change mL to L.
250 mL x 1L/1000mL = L
Step 2: Calculate.
Moles = ( mol) (0.250 L) = moles
Step 3: Convert moles to grams.
( mol)(90.00 g/mol) = g oxalic acid
1 L

75. Learning Check
How many grams of NaOH are required to prepare 400. mL of 3.0 M NaOH solution?
1) 12 g
2) 48 g
3) g

76. Preparing Solutions M1V1 = M2V2
Weigh out a specific mass of solid solute and dissolve in a given quantity of solvent.
Dilute a concentrated solution to give one that is less concentrated.
M1V1 = M2V2

77. How would you make 250 mL of
3.0 M HCl from a 6 M stock solution of HCl?
M1V1 = M2V2
(250 mL) (3.0 M) = (6 M) (X)
X = 125 mL
“How would you make…” would require some words in the answer!
So,… Add enough water to 125 mL of 6 M HCl to bring the total volume to 250 mL.

78. Miscible – where liquids dissolve in each other.
Ex. Ethanol in water

79. Suspensions and Colloids
Suspensions and colloids are NOT solutions.
Suspensions: The particles are so large that they settle out of the solvent if not constantly stirred.
Colloids: The particles are intermediate in size between those of a suspension and those of a solution.

80. The Tyndall Effect
Colloids scatter light, making a beam visible. Solutions do not scatter light.
Which glass contains a colloid?
colloid
solution

81. Types of Colloids Examples Dispersing Medium Dispersed Substance
Colloid Type
Fog, aerosol sprays
Gas
Liquid
Aerosol
Smoke, airborne bacteria
Solid
Whipped cream, soap suds
Foam
Milk, mayonnaise
Emulsion
Paint, clays, gelatin
Sol
Marshmallow, Styrofoam
Solid foam
Butter, cheese
Solid emulsion
Ruby glass
Solid sol