1. Intermolecular Forces
Ch 11
Liquids &
Intermolecular Forces

2. Remember the Kinetic Theory of Molecular Motion
Remember the Kinetic Theory of Molecular Motion. Solids have fixed volume and shape. Liquids have a fixed volume with changeable shape. Gases spread to fill any shape and volume.
Lets look at the properties of solid and liquids as well as the intermolecular forces.

3. Intermolecular forces
Van der Waals forces are any forces
between molecules.
The more forces between molecules, the higher the melting point and boiling point.
Dipole – dipole attraction
Hydrogen bonding
London dispersion forces

4. Dipole – Dipole forces
Molecules with dipole moments can attract other molecules with dipole moments. (Polar molecules) Positive end and negative sides line up.
These are permanent forces.

5. Ionic solids exhibit electrostatic forces of attraction because of the positive and negative ions.

6. N O
O O
No dipole
DEN = 0.0
bp = oC
Polar
DEN = 0.5
bp = oC
No dipole
DEN = 0.0
bp = oC
Three molecules that have similar arrangement. N2 and O2 which exhibit no dipole moment, will evaporate at lower temperatures than NO which has a small dipole moment.
If the molecule has an electrical dipole moment, the molecules will be attracted to each other.

7. Boiling points are simply a function of the molecular mass when the molecules are non-polar.

8. An increase in the dipole moment, causes an increase in the boiling point. The molecules set up a “network” of attractions between positive and negative ends. Note that all the molecules listed an similar molecular mass, but the boiling point is a function of the dipole moment.

9. Hydrogen bonding: A bond formed when a H atom bonded to F, O, or N is attracted to the F, O, or N in a neighboring molecule. Hydrogen bonding is not as strong as a covalent bond, but is very important.

10. Some of the molecules exhibiting hydrogen bonding which are very important.
Water
DNA
Proteins
One hydrogen bond is very weak. But 100’s to 1000’s of hydrogen bonds become very strong.

11. Because of hydrogen bonding, water exhibits :
Adhesion : attraction to other objects
Cohesion: attraction to itself
Adhesion: water adhering to the xylem in a plant stem, results in capillary action
Cohesion: water sticking to water, force that allows water to stick together in capillary action
High surface tension: water resisting an increase in the surface area
High Surface Tension: hydrogen bonding at the surface,
High Specific Heat: energy required to raise the temperature of 1
gram of water by 1oC
Viscosity: resistance to flow of a liquid

12. Comparison of boiling points of covalent hydrides
Comparison of boiling points of covalent hydrides. Note that the boiling points exhibited within a family increase as molecular mass increases. However, H2O, because of the hydrogen bonding between molecules, boils at 100oC.

13. Comparison of the boiling points Ethane CH3CH3 bp = -88
Comparison of the boiling points Ethane CH3CH bp = oC ethane without hydrogen bonding Ethanol CH3CH2OH bp = 78oC ethanol with hydrogen bonding Bond energy Covalent bond = 150 KJ/mole Hydrogen bond = 25 KJ/mole

14. Hydrogen bonding in DNA The complementary nitrogen bases on the two strands hydrogen bond together.

15. The secondary structures
of proteins include beta (b) pleated sheets and alpha (a) helix. Both structures are held in their shapes by hydrogen bonding.

16. Induced Dipole – London Dispersion forces
An atom at one instant of time may have a concentration of e- in one region of the atom. An instantaneous dipole moment is formed.

17. The ability to induce a dipole is directly dependent on: 1
The ability to induce a dipole is directly dependent on: 1. an increased number of e- 2. an increasing molar mass 3. long linear molecules
Ethane CH3CH m b.p oC
Propane CH3CH2CH m b.p oC
Hexane C6H m b.p. 68oC
Decane C10H m b.p. 174oC
An induced dipole causes an increase in the boiling point of a molecule.

18. Classification of solids
Crystalline solids have a highly regular repeating arrangement of components
Amorphous solids have a great amount of disorder in their structure.

20. Metallic solids exhibit a sea of e- which can flow over the entire solid. This explains the ability of metals to conduct electrical current.

21. Phase Changes Gas Liquid Solid Sublimation Condensation Deposition
Melting or Liquefaction
Freeze or Solidification
Condensation
Evaporation or Vaporization

23. All substances go through phase changes
All substances go through phase changes. A Phase Change Chart can be constructed for any substance.

24. When changing the temperature of a substance, the same formula is used.
H = m c DT
The specific heat (c) is different for each substance, and phase of the substance.

25. H = Hf m H = Hv m Melting Evaporating (Vaporizing)
The plateau’s represent phase changes.
Melting
H = Hf m
Evaporating (Vaporizing)
H = Hv m
Hf and Hv are unique values for each substance.
For water,
Hf = 344.9j/goC
Hv =2260.4j/goC

26. Triple Point: all three phase can exit
Critical Point: point where vapor and liquid are indistinguishable.

28. 760 torr (1 atm) is shows the normal evaporation temperature..
The vapor pressure (point when the liquid turns into a gas) is a function of the temperature.
The line at
760 torr (1 atm) is shows the normal evaporation temperature..

29. Liquid Crystals Viscous – milky state between a liquid & solid
See page 468
Viscous – milky state between a liquid & solid
Used as a pressure – temperature sensor in liquid crystal displays (LCD)
Liquid – molecules are random
Nematic crystals align in liquid with the long axis pointing the same way
Smectic A – crystal in liquid axis align with layers
Smectic C – crystals in liquid axis align in layers with an incline
Cholesteric – crystals align in layers with the axis rotating each layer