1. Intermolecular Forces, Liquids, and Solids

2. Intermolecular Forces, Liquids, and Solids
Some Characteristic Properties of the States of Matter
Assumes both the volume and shape of the container
Is Compressible
Diffusion within a gas occurs rapidly
Flows readily
Gas:
Liquid
Assumes the shape of the portion of the container it occupies
Does not expand to fill the container
Is virtually incompressible
Diffusion within a gas occurs slowly
Flows readily
Retains its own shape and volume
Is virtually incompressible
Diffusion within a solid occurs extremely slowly
Does not flow
Solid

3. The strength of intermolecular forces vary over a wide range,
but are generally weaker than covalent or ionic bonds
16 kJ will overcome the intermolecular attraction leading to vaporization
431 kJ of energy is required to break the covalent bond between HCl

5. the motions of electrons on its near neighbors
London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment
The instantaneous distribution at any given moment can be different from the average distribution producing an instantaneous dipole moment
Because electrons repel, the motions of electrons in one atom influence
the motions of electrons on its near neighbors

6. London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment
The ease with which the charge distribution in a molecule can be distorted by an external force is called its polarizability. The greater the polarizability of a molecule, the more easily its electron cloud can be distorted to give a momentary dipole.
In general, larger molecules tend to have greater polarizability because their electrons are farther from the nucleus
Because molecular size and molecular mass tend to parallel one another, dispersion forces tend to increase in strength with increasing molecular weight.
Halogen Boiling pt (K) Noble gas Boiling point (K)
F He
Cl Ne
Br Ar
I Kr

7. London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment
The shapes of molecules can play a role in the magnitude of dispersion forces.

8. London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment
London dispersion forces may also operate within polar molecules and contribute to the overall attractive forces between the molecules; some times more than the dipole-dipole forces
HCl: bp = K
Dipole Moment = 1.03 D
HBr: bp = K
Dipole Moment = 0.79 D

9. Dipole Dipole forces exist between neutral polar molecules, are effective only when polar molecules are very close together, and are generally weaker then ion-dipole interactions
Substance Dipole moment Bp
For two particles of equal mass and size, the strengths of the intermolecular attractions increase with increasing polarity
Propane
Dimethyl ether
Methyl chloride
Acetaldehyde
Molecules that are attracting one another spend more time near each other than those molecules which repel. Thus the overall effect is a net attraction

10. Hydrogen bonding is a special type of intermoleculer attraction that exists between the hydrogen atom in a polar bond (e.g., H F, H O, or H N) and an unshared electron pair on a nearby electronegative atom (usually an F, O, or N atom on another molecule)
Because F, N, and O are so electronegative, a bond between hydrogen and
any of these elements is quite polar, with hydrogen at the positive end.

11. Ion-Dipole Forces exist between an ion and the partial charge on the end of a polar molecule
Cl-
The magnitude of the interaction depends upon the charge of the ion, the dipole moment of the polar molecule and the distance from the center of the ion to the midpoint of the dipole

12. Metallic solids consist of metal atoms
Metallic solids consist of metal atoms. Bonding is due to valence electrons that are delocalized throughout the entire solid. We can visualize these electrons as an array of positive ions immersed in a sea of delocalized electrons.

13. Intermolecular Forces, Liquids, and Solids
Metallic bond
examples: Cu,
Fe, or alloy

14. ionic solids

15. network solids

16. Intermolecular Forces, Liquids, and Solids
Properties of Liquids: Viscosity and Surface Tension
The resistance of liquids to flow is called their viscosity. The greater the viscosity, the
more slowly the liquid flows. Viscosity decreases with increasing temperature.
Viscosity is related to the ease with which individual molecules of a liquid can move
with respect to one another. It thus depends on the attractive forces between molecules
and and their structural character.

17. Intermolecular Forces, Liquids, and Solids
Properties of Liquids: Viscosity and Surface Tension
Molecules in the interior are attracted equally in all directions, whereas those at the surface experience a net inward force
This net inward force pulls molecules from the surface into interior, thereby reducing the surface area
Surface tension is therefore defined as the energy required to increase the surface of a liquid by a unit amount
Forces that bind like molecules to one another are
called cohesive forces.
Forces that bind molecules to a surface are called
adhesive forces

19. Changes of State
Vapor Pressure: pressure exerted by evaporating liquid in the space
above that liquid once dynamic equilibrium has been established
Substances with a high vapor pressure evaporate more rapidly than thosewith low vapor pressures. Theses substances are more volatile
The vapor pressure of a substance increases as the temperature increases

20. Changes of State
Whenever a change of state involves going to a less ordered state, energy
must be supplied in order to overcome intermolecular forces
The heat needed for vaporization is called the heat of vaporization
(e.g kJ/mol for water)
ENDOTHERMIC
EXOTHERMIC
Melting is also called fusion. The enthalpy of change associated with
melting is called heat of fusion. (e.g kJ/mol for water)

21. Changes of State
Enthalpy and Temperature Changes Accompanying Heating

22. Ht = .940 kJ + 6.01 kJ + 7.52 kJ + 40.67 kJ + .830 kJ = 55.97 kJ
Calculate the enthalpy change associated with converting 1.00 mol of ice at -25°C to water vapor at 125 °C and 1 atm. The heat capacities per gram (specific heats) of ice, water, and steam are 2.09 J/g-°C , 4.18 J/g-°C , and 1.84 J/g-°C, respectively. The heat of fusion of ice is 6.01 kJ/mol, and the heat of vaporization of water is kJ/mol
Heat required to bring the
ice from -25 °C to 0 °C :
18.0 g.
2.09 J
1.00 mol H2O
25 °C
= 940 J
1.00 mol H2O
g -°C
Heat of fusion:
6.01 kJ
1.00 mol H2O
= 6.01 kJ
mol H2O
Heat required to bring liquid
water from 0 °C to 100 °C :
18.0 g.
4.18 J
1.00 mol H2O
100 °C
=7520 J
1.00 mol H2O
g -°C
Heat of vaporization
40.67 kJ
1.00 mol H2O
= kJ
mol H2O
Heat required to raise the
temperature of the vapor
to 125 °C
18.0 g.
1.84 J
1.00 mol H2O
25°C
=830 J
1.00 mol H2O
g -°C
Ht = .940 kJ kJ kJ kJ kJ = kJ

23. Vapor Pressure and Boiling Point: A liquid boils when its vapor
Changes of State
Vapor Pressure and Boiling Point: A liquid boils when its vapor
pressure equals the external pressure acting on the surface of the liquid
The higher the outside pressure, the higher the boiling point

24. A phase diagram is a graphical way to summarize the conditions under which equilibria exist between the difference states of matter.
Line A-D represents the
change in melting point
of the solid under
increased pressure
Line A-B is the vapor
curve of the liquid
The line A-C represents the variation in
vapor pressure during sublimation.

25. A phase diagram is a graphical way to summarize the conditions under which equilibria exist between the difference states of matter.
Note that the solid phase is the stable phase under condition of high pressure and low temperature
Note that the gas phase is the stable phase under condition of low pressure and high temperature

26. Why are the fusion-freezing lines different for H2O and CO2?
Why does CO2 sublime under normal conditions?
How is freeze-dry food prepared?

27. AS THE TEMPERATURE IS INCREASED : At point 6-A, the H2O exists entirely as a solid. When the temperature reaches point 4-B, the solid begins to melt, and an equilibrium condition occurs between the solid and the liquid. At a yet high temperature, point 7-C, the solid has been converted entirely to a liquid. When point 8-D is encountered, vapor forms, and a liquid-vapor equilibrium is achieved. Upon further heating, to point 9-E, the H2O is converted entirely to the vapor phase. 8 7 5 3 2 1

28. A crystalline solid is a solid whose atoms, ions, or molecules are ordered in well-defined arrangements. Amorphous solids have no structure.
A unit cell is the repeating unit that makes up the crystalline solid. The array of repeating points is called the crystal lattice

29. Covalent Network Solids are Held together in large networks or
chains by covalent bonds. These solids are much harder and have
higher melting points than molecular solids
Molecular solids consist of atoms or molecules held together by dipole-dipole or London dispersion forces or hydrogen bonds

30. of the bonds depend upon the charges of the ions. Melting point
Ionic solids consist of ions held together by ionic bonds. The strength
of the bonds depend upon the charges of the ions. Melting point
increases as bond strength increases
NaCl
The structure of the solid is dependent upon the charges and
relative sizes of the ions.
The coordination number increases as the ratio of the cation radius to the anion radius increases.

31. Atomic solids – (london) eg. He, Ar etc.
atoms exist at lattice pts
molecular solids (London, dipole, H-Bond) CH4, NH3
molecules exist at lattice pts
Ionic solids (electrostatic attractions) NaCl
cations & anions exist at lattice pts
Metallic solids- (Cations held by a sea of electrons) Cu, Fe
cations exist at lattice pts
Network solids – (3 dimensional covalent bonds) diamond (C), quartz (SiO2)