1. Liquids & Solids
Kinetic Molecular Theory explains liquids and solids as well as gas.

2. Figure 10.1: Schematic representations of the three states of matter.

3. What compound is in all three phases RIGHT NOW ??
WATER
So why aren’t all the other gases in the air - like CO2 or O2 or N2 in all three phases ??
Liquid O2 is –183 oC
Dry Ice is –78 oC
Liquid N2 is –196 oC

4. Condensed Phase shape Assumes ____________ ___________
PROPERTY
GAS
LIQUID
SOLID
shape
Assumes ____________ ___________
Assumes ____ __________ -flows
____________ shape
Kinetic Energy
Density
High
Order
Moderate
Compressibility
Low
Importance of Intermolecular forces
Thermal expansion
Small

5. Condensed Phase shape Assumes volume of container
PROPERTY
GAS
LIQUID
SOLID
shape
Assumes volume of container
Assumes shape of container - flows
Retains own shape
Kinetic Energy
High (a lot of movement)
Medium
Low (vibration only motion)
Density
Low
High
Order
None
Moderate
Compressibility
Importance of Intermolecular forces
Thermal expansion
Large
Small

6. What holds things together?
Intramolecular Forces
Hold particles together _________ molecule (covalent bonds).
Intermolecular Forces
Cause attractive interactions _______ __________ (molecules, ions or atoms) of a substance.
WITHIN
BETWEEN
Particles

7. Intermolecular Forces
Intermolecular forces hold solids and liquids together! There are four types:
Ion-dipole forces
Dipole-dipole forces
(London) Dispersion forces
Hydrogen bonding

8. Electronegativity
Electronegativity relationships between atoms ultimately determine the nature of bonding.
Electronegativity
A=B
A<B
A<<B
Electron “Share”
A : B
Ad+ : Bd-
A+ : B-
Type of Interaction
_____ sharing
______ sharing
transfer
Bonds formed
Elements Involved
Unequal
Equal
Complete
Non Polar
Covalent
Polar
Covalent
Ionic
Like
nonmetals
Unlike
nonmetals
Metal &
nonmetal

9. Bond Polarity
Because of a difference in electronegativity, polar covalent bonds have shift in electron density There is a permanent unequal distribution of electrons due to different electron pulling power.
H - F
H is partially (+) because it has lost some e- density
F is partially (-) because it has gained some e- density d+ d-

10. Molecular Polarity
Electronegativity differences between atoms combine with molecular shape to determine if the whole molecule will be polar. Molecules will be polar if:
They contain polar bonds
The polar bonds are asymmetrically arranged around the central atom

11. Molecular Polarity Examples: asym sym asym sym Polar nonpolar nonpolar
polar bonds
geometry
molecule C
asym
sym
asym
sym
Polar
nonpolar
nonpolar
Polar

12. Ion-Dipole Forces
Ion-dipole Forces are the electrostatic attraction between an ion and the oppositely charged end of a polar molecule.
Example:
Ions in water
Interaction _______ energy = (enthalpy) of hydration = DHhyd
releases

13. Dipole-Dipole Forces
Dipole-Dipole Forces are the electrostatic attraction between oppositely charged ends of polar molecules
Example: HCl (l) Water (l)
HCl HCl
ClH ClH

14. Polarizability
Polarizability is the tendency for the electron density of atom or molecule to be deformed or ”sloshed around”
Polarizability will increase with:
_________ mass
_____________ of electrons
electrons _____________
increased
greater number
less tightly held

15. Polarizability
Polarization of molecules/atoms results in “____________” dipoles.
Instantaneous dipole moments are due to the polarizability of a molecule!
Example: I2 iodine has no permanent dipole (non-polar)
instantaneous

16. London Dispersion Forces
Dispersion Forces are the electrostatic attraction involving __________ dipole moments in (non-polar) molecules 
The greater the polarizability of the atoms/molecules, the ________ the interaction
___ molecules are polarizable
___ molecules experience dispersion forces (even polar molecules)
instantaneous
stronger
All
All

17. Which would be affected more by London Dispersion Forces?
H2 or I2
How many electrons does each have?

18. Hydrogen Bonding
Hydrogen Bonding is an interaction that exists _______ the H-atom in a polar covalent bond with ________ and the _______ electrons on a nearby highly electronegative atom (________).
Special case of dipole-dipole interaction
Since H _____ shares all its e-, the positive nucleus is relatively exposed and is attracted to unshared electron pairs on another molecule
between
F, O, or N
lone-pair
F, O, or N
always

19. Hydrogen Bonding HF H2O NH3
H-bonds are much ______than covalent or ionic bonds.
H-bonds are ________ than other dipole-dipole or dispersion (London) forces.
H-bonds are strong enough to determine structure of _____________!
weaker
stronger
ice and proteins

20. Cool feature of water Due to arrangement and hole in middle
In the solid crystal form, the solid water (ice) is __________ than liquid water. This is NOT true for other solids.
This allows -
Less dense
Due to arrangement
and hole in middle
Ice to float
Fish to live at bottom
of lake

22. IM Force Practice -> dispersion forces (both are nonpolar)
What are the predominant forces in the following interactions?
Dissolving candle wax with kerosene
Ethyl alcohol absorbing into your blood stream
Iodine gas molecules condensing into solid iodine
-> dispersion forces (both are nonpolar)
-> dipole-dipole
-> dispersion forces (both nonpolar)

23. IM Force Practice
Detergent removing gravy from a garment, in water
Dissolving table salt in water
Dissolving sugar in water
-> dispersion forces (nonpolar end of detergent and fat)
ion-dipole (detergent and water)
-> ion-dipole
-> dipole-dipole

24. IM Force Pracitice
Which has stronger intermolecular forces?
 Dissolving table salt vs. sugar in water
 Dissolving I2 vs. O2 in water
-> table salt; ion-dipole stronger than dipole-dipole.
-> I2; both are dispersion forces, but I2 is more polarizable giving it a larger instantaneous dipole.

25. IMF and Solid/Liquid Physical Properties
Intermolecular Forces (IMF) between particles determine physical properties:
Melting point (mp)
Boiling point (bp)
Viscosity
Surface tension
Vapor pressure

26. Liquid Properties
1. high density compared to gas due to ___________ of molecules
2. Incompressibility
3. Ability to ______ - due to fluidity
compactness
diffuse

27. Physical Properties
Viscosity is the resistance of a liquid to flow (viscosity is high if molecules _______move past each other easily).
The _______ the IMF, ______ the viscosity of the liquid.
At _______, viscosity _________; particles have ________ and can ________ attractive forces (note: viscosity is also affected by molecular chain lengths, physical entanglement)
Do NOT
stronger
higher
Higher T
decreases
Examples of viscous liquids include magma, lava, molasses, honey
If the temperature increases the energy of the molecules increase, thus allowing them to slide past each other more efficiently, which is why they can overcome the attractive forces more easily.
The viscosity also depends on the length of the molecule and the way it is arranged but we won’t have time to get into all of that.
more KE
overcome

28. Liquid Properties
Surface tension - a force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the ____________________ Varies from liquid to liquid.
5. Capillary action-attraction of the surface of a _____ to the surface of a _____
smallest possible size
Surface tension is dealing with a liquid-liquid interaction
Generally things want the smallest space exposed to the outside world, so the surface area is decreased (thus reducing heat loss mostly if we think in biology terms).
Capillary action is why we have to read the meniscus in a graduated cylinder and why plants can grow tall (the wick the water up their roots and stems)
liquid
solid

29. Physical Properties
Surface Tension is the energy required to _________________ of a liquid.
For a given number of liquid molecules:
Highest surface area is to put them all in a row.
Lowest surface area is a sphere
 ______ IMF results in liquids that have _______ surface tensions.
increase surface area
It takes a lot of energy for molecules to lay flat (unless they have low IMFs, like gasoline etc)
If you put a drop of water on a surface and a drop of rubbing alcohol or acetone on the surface you will notice that water will form a bubble and the alcohol will spread out and not form a bubble…the lower IMFs will cause the alcohol to spread out. Water, with the higher IMFs will form that bubble (the hydrogen bonds rule here!).
Stronger
higher

30. Example of liquid mercury which has very high IMFs

31. Physical Properties
Cohesive forces are a measure of how well molecules _______________.
Adhesive forces are a measure of how well molecules stick to ___________ (e.g., to the molecules of a container). 
If adhesive force > cohesive force, then the molecules will maximize surface area to maximize adhesive interactions!
-this determines _________________!
stick to each other
other molecules
Cohesion is a liquid to a liquid interaction (so water being attracted to itself)
Adhesion is a liquid being attracted to a non-liquid (so water being attracted to the side of a container)
direction of meniscus

32. Physical Properties ____ ____
____ ____
Capillary action is the result of _______________.
Paper chromatography is a separation technique based on __________ in adhesive forces.
Ad>Co
Co>Ad
meniscus
adhesive forces
Some molecules with greater cohesion (like mercury) will cause a meniscus to go the other way!
differences

33. Nonpolar liquid mercury forms a convex meniscus in a glass tube, whereas polar water forms a concave meniscus.
Example of mercury vs water and the cohesion and adhesion forces present

34. Liquid Properties Evaporation and Boiling
6.Evaporation - process by which particles escape from the surface of a _________ liquid and enter the gas state.
Boiling - is a change of a liquid to ________ of vapor.
non-boiling
VERY IMPORTANT: Evaporation is from the surface of a non-boiling liquid! Only from the surface!
Boiling is concerning a liquid where the whole body of the liquid is boiling and has bubbles (think finding Nemo here). Very important to know the difference here!
bubbles

35. Evaporation Without boiling
Escape of molecules from a liquid to a gas _________________________________
Warmer molecules have _________________ & they are able to ____________________ This leaves the ______________ molecules behind.
Without boiling
High Kinetic Energy (KE)
Escape the surface
cooler
Remember that being cold does not mean the presence of cold, it means that there is an absence of heat.

36. Evaporation crucial Sweat COOLS Water cycle for RAIN
This is _________________ in nature
_____________________________
______________________________
Sweat COOLS
Water cycle for RAIN
This is why our skin feels cooler sometimes when we sweat, because the more energetic water molecules are leaving thus leaving behind the cooler water molecules.

37. Liquid Properties removal of heat
7.     freezing or solidification - change of liquid to solid by _______________
removal of heat
Again, freezing or solidification or crystallizing is a absence of heat not an addition of cold

38. Physical Properties Melting Point/Boiling Point:
The _______ the IMF, ______ the melting point of the solid.
The _______ the IMF, ______ the boiling point of the liquid.
Stronger IMF means that _________ is required to disrupt interactions!
stronger
higher
stronger
higher
Essentially, if something has stronger IMFs, it will take more energy to break those IMFs thus the melting point will be higher, same goes with the boiling point. For example, water has strong IMFs with the hydrogen bond, it will take a lot more energy to get water to boil than gasoline which has only London dispersion forces acting on it.
more heat

39. Physical Properties
Vapor pressure is the pressure of the vapor of a substance when the liquid and vapor phases are in _______________.
If a liquid is in a closed container, some molecules have enough KE to escape the IMF and form vapor phase in container.
____ container will _____ reach equilibrium
dynamic equilibrium
Vapor pressure is sometimes hard to understand, think about it in terms of how many molecules of the liquid are in the gas phase in a manometer. If you have a hydrocarbon (gas for example) in a manometer, it will have a lot of molecules in the gas phase because it has low IMFs, thus it will have a high vapor pressure. If you compare that to water, which the high IMFs, it will have a low vapor pressure because there are not as many molecules in the gas phase because it is harder for those molecules to overcome the IMFs and enter the gas phase.
liquid
liquid
open
never

40. Physical Properties
Substance reaches dynamic equilibrium when there is __________ in particles in gas phase (even though gas and liquid particles in constant flux!)
Vapor pressure _______ w/ _________
________ T ________ # of molecules with sufficient energy to escape surface of liquid.
Liquid with high vapor pressure is ______!
no net change
increases
Increases T
Increasing
increases
Now think about if you were to heat a manometer, if you increase the temperature of the liquid, you are going to increase the energy of the molecules, thus allowing them to overcome the IMFs in the liquid, thus you will increase your vapor pressure.
volatile

41. Volatile liquids evaporate weak BETWEEN
Liquids ____________ easily
Examples:
Have _______________ attractive forces _______________ molecules
Would water be volatile?
Would CH4 be volatile?
weak
BETWEEN
We dealt with volatile liquids in the IMF lab, anything you can smell very quickly is volatile. Gas, acetone, rubbing alcohol etc are all volatile . They all have low IMFs
No – lots of attraction, H bonding
Yes, non-polar, no attraction

42. Initially only molecules Leave surface
Container at
Equilibrium
BOTH leaving
AND returning
Just closed container
Initially only molecules
Leave surface
This is a diagram explaining the start of a liquid in a closed container, it is only just beginning to have molecules leave the surface of the water and enter the gas phase, the second is explaining that the liquid and the gas are in dynamic equilibrium because they have the same number of gas and liquid molecules and they are in constant flux.

43. Equilibrium Vapor Pressure
Here is an example of a manometer and how it can measure vapor pressure. If I have water in here it will have a low vapor pressure because it has high IMFs, thus the molecules won’t be able to enter the gas phase as easily. If I have gasoline in here it will have high vapor pressure because it has low IMFs and the molecules are able to enter the gas phase easily because there isn’t as much to overcome.
Now think about if I increase the temperature here, the molecules would be able to enter the gas phase more easily and the vapor pressure would increase.

44. Vapor Pressure
Here is another example of vapor pressure. Water has a lower vapor pressure, the second molecule has a little bit higher vapor pressure than water but it still isn't that high because there is a hydrogen bond present because of the OH group. The last molecule we can assume has the lowest IMFs because it has the highest vapor pressure.

45. Physical Properties
Boiling is the conversion of liquid to vapor within a liquid as well as its surface. Occurs when ____________________ = ________________
Boiling point is the temperature where equilibrium vapor pressure of the liquid is _______ the ambient atmospheric pressure. (i.e., equal to ____________ __________________________). 
equilibrium vapor pressure
atmospheric pressure
exactly
Boiling is important to remember that it happens throughout the body of liquid, not just the surface.
Boiling point is essentially when the molecules are able to not only overcome their IMFs (vapor pressure) they are also able to overcome the atmospheric pressure pushing down on the liquid.
pressure acting
on the surface by the atmosphere

46. Physical Properties
Example: At sea level, temp at which vapor pressure is 1 atm, the bp = 100oC for water
At this point, bubbles of vapor can form ______ the bulk of the liquid
 molecules of gas not just coming from the surface but _____________
within
Again, important to remember that boiling is from everywhere!
EVERYWHERE

47. Physical Properties
Normal boiling point refers to standard pressure, 1 atm.
If Patm , then vp must be _ for boiling to be reached (requires __ T).
If Patm , then vp must be _ for boiling to be reached (requires __ T).
Since bp occurs at lower T, takes ______ to cook things at high alt (low P)    
If I have more atmopshperes pushing down on the liquid, it will take more energy to overcome the pressure from the atmosphere, thus the boiling point will be higher. This also means that the vapor pressure will also have to increase.
Vise versa if we decrease the pressure, the boiling point will decrease
Since boiling occurs faster at higher atmospheres, it will boil at a lower temperature because it has not taken as much energy to get the liquid to boil. This is why when you are at higher elevations you have to cook food longer, because the temperature (while boiling) isn't as hot.
longer

48. Boiling Point changes due to pressure
Lots Little
This is a diagram drawn in class, get it from a lab partner or the teacher when you get back

49. A graph relating boiling point to pressure, as you can see, things with lower IMFs have much lower boiling points than things with higher IMFs

51. Kinetic Molecular Theory
Gases consist of particles that are very far apart.
If KE > IMF, substance ___________
If KE < IMF, substance is __________
Molecules of condensed phase are close together. Energy (heat) must be added to _________________; energy must be released when ____________ ________ in liquid or solid.
must be a gas
liquid or solid
separate the molecules
particles come
together

52. Heating/Cooling Curve

53. When temperature is changing, Energy is increasing or
Decreasing Kinetic Energy (KE)
When temperature is NOT changing,
Energy is used to overcome attractive forces like H-bonding or dipole-dipole

54. Phase Change
Reactions often involve changes in phases of substances which also involve energy or enthalpy changes.
The enthalpy terms associated with common phase changes include:
Type of phase change
Symbol
Heat of fusion
DHfus
Heat of vaporization
DHvap

55. Molar Heat of Fusion
Amount of heat energy _______ to melt 1 mole of solid at its melting point. Must ________ attractive forces. (solid  liquid)
Or the heat energy ________ when one mole freezes. (liquid  solid)
____Temperature change during phase change.
Temperature this occurs varies for every solid
required
overcome
released NO

56. Molar Heat of Vaporization
Amount of heat energy _______ to vaporize 1 mole of liquid at its boiling point. (liquid  gas)
Varies for every liquid. The ___________ the intermolecular attraction, the higher the energy needed to separate, the higher the temperature.
Or, amount of heat ________ to condense 1 mole of gas (gas  liquid)
____Temperature change
needed
STRONGER
released NO

57. Phase Change
DH’s for opposite changes will have the same mag., but negative sign. The phase transitions of water:
endo phase change
Symbol
exo phase change
Melting
+6.02 kJ/mol
DHfus
Freezing
Vaporization
+40.8 kJ/mol
DHvap
Condensation
-6.02 kJ/mol
-40.8 kJ/mol

58. Heating Curve Each segment of the curve has an enthalpy change (DH).
The total DH is sum of each segment.
For DT segments, DH = Csp x mass x DT
For phase change segments (DT = 0)
DH =  DHvap x n
n = # moles or
mass in grams
DH =  DHfus x n

59. Energy change with grams for water
Solid  Liquid (freezing or melting):
Given g H2O x Joules/gram
Heating or cooling liquid:
Given g H2O x J/g oC x  T oC
Liquid  Gas (boiling or condensing):
Given g H2O x J/g
Heating ice or heating steam
Given g H2O x 2 J/g oC x  T oC
 T = change in temperature

60. Energy Change Tables
Solid  Liquid (freezing or melting) g H2O 335 J 1 g Heating or cooling liquid g H J  T oC 1 g oC Liquid  Gas (boiling or condensing): g H2O 2260 J 1 g Heating ice or heating steam g H J  T oC

61. Heating/Cooling Curve
Plot change in T of a substance as a function of heat energy added:
Energy is _________ _______when a substance changes state T is _______ during phase change!
absorbed or released
constant

62. Heating Curve Practice
What is the DH for the process in which 2.00 g water at 27oC is converted to ice at –3oC?

63. Heating Curve Practice
What is the DH for the process in which 2.00 g water at 27oC is converted to ice at –3oC?

64. How a Wall of Water works
Added / absorbed
Heat is _____________________ to boil
Heat is _____________________ to freeze
Released

65. Phase Diagrams Plot of P vs. T, for phases:
Triple point – all phases (s, l, g,) are in equilibrium

66. Critical point - critical T & P to a substance where it cannot exist as a liquid above the critical temperature

67. Critical temperature = 373.99 oC Critical pressure = 217.75 atm
For water
Critical temperature = oC
Critical pressure = atm

68. Structure and Properties of Solids
Solids can be divided into two categories:
Amorphous
Crystalline
Ionic crystals
Metallic crystals
Molecular crystals
Network crytals

69. Solids
Amorphous solid: particles have ___ orderly structure (IMF vary in strength throughout structure)
Example: ____
Crystalline solid: atoms, ions, or molecules ordered in __________ arrangements (causes crystals to have regular shapes)
Example: ______________________ NO
glass
well-defined
NaCl, ice, diamond, all metals

70. Crystalline Solids
Ionic Solids: ions are held together by ionic (electrostatic) interactions;
bond strength ________ with ion size, and _______ with ion charge!
ionic bonds strong
high melting point, boiling point
low vapor pressure
Like a magnet
decreases
increase

71. Crystalline Solids
Metallic Solids: all atoms present are metals (inc. hydrogen at very low T)
Electrons in valence shells shared by many atoms, called __________electrons.
*____ electrical conductivity (mobile e-)
*____ thermal conductivity (tightly packed identical atoms)
*malleable and ductile
delocalized
good
good

72. Crystalline Solids strong
Molecular Solids: have ______ intramolecular bonds (covalent) and _______ intermolecular bonds, like H-bonds.
Example: sugar, water
soft substances
___ melting and boiling points
_____ vapor pressures
weaker
low
higher

73. Crystalline Solids
Network Solids: are single giant molecules with an endless number of covalent bonds among atoms. All bonds are _____________.
Example: diamond, graphite, sulfur
can be very hard (diamond)
tend toward high melting, boiling
low vapor pressure
equally strong

74. Crystalline Solids In summary: Ionic NaCl ions ionic: Metallic Cu(s)
crystalline solid type
Example
Structural Unit
Forces Between Units
Typical Properties
Ionic
NaCl
ions
ionic:
(+) and (-)
hard, brittle, high mp, poor conductor
water soluble
Metallic
Cu(s)
atoms
Metal atoms with delocalized e-
malleable, ductile, luster, conductive
Molecular
sugar
molecules
dispersion forces + other intermolecular forces (depends on type of molecule)
low to moderate mp
soft, poor conductor
Network
graphite
covalently bonded atoms
covalent
bonds
wide range of mp, poor conductor (some exceptions)