1. Condensed States of Matter: Liquids and Solids Chapter 14

2. Condensed States
Liquids and solids = condensed states because they have significantly higher densities than gases

3. only slightly compressible
Gas
Liquid
Solid
highly compressible
only slightly compressible
low density
high density
fills container completely
does not expand to fill container - has definite volume
rigidly retains its volume
assumes shape of container
retains own shape
rapid diffusion
slow diffusion
extremely slow diffusion - only occurs at surfaces
high expansion on heating
low expansion on heating

4. KMT of Matter
According to the kinetic molecular theory, the state of a substance at room temperature depends on the strength of the attractions between its particles.

5. Forces
Intramolecular forces are between the atoms within a molecule = bonds
Intermolecular forces are between molecules

6. Intramolecular Forces
Covalent Bonds – between nonmetals; sharing of electrons
Ionic Bonds – between metals and nonmetals; transfer of electrons to form ions
Metallic Bonds – between metals; sea of electrons

7. Intermolecular Forces
Dipole-dipole attractions
a. Hydrogen bonding
Ion-dipole attractions
London dispersion forces

8. Dipole-Dipole Attractions
Attractions due to permanent dipoles in polar molecules
Remember, a dipole is created when positive and negative charges are separated by some distance.
Only 1% as strong as covalent or ionic bonds and even weaker if distance between charges increases.

9. Hydrogen Bonds
Unusually strong dipole-dipole attractions that occur among molecules in which hydrogen is bonded to a highly electronegative atom, such as F, N, or O

10. Ion-Dipole Attractions
Polar molecules surround ions based on their attraction for the charge on the ion

11. London Dispersion Forces
Explain the attraction that exists between non-polar molecules
Even in these molecules the electrons are not uniformly distributed at every second
Temporary dipolar arrangement of charge creates an instantaneous dipole.

12. London Dispersion Forces
Instantaneous dipoles can induce similar dipoles in neighboring atoms

13. Intermolecular Attraction
Kinetic Energy
Intermolecular Attraction
State at Room Temp.
low
high
solids
medium
liquids
gases
Kinetic
Kinetic energy determines if particles will overcome the intermolecular forces keeping them together. Thus, higher attractions mean higher boiling points, because a higher kinetic energy will be needed to overcome the attraction.

14. Properties of Solids
Intermolecular forces keep the particles of a solid packed together tightly
- Particles are highly ordered with fixed positions
- Only particle movement = vibrations
Liquids have similar distance between particles

15. Bonding in Solids
2 basic types of solids based on the nature of the particles that make them up:
Crystalline
Atomic
Ionic
Covalent-network
Amorphous

16. Crystalline Solids
Solid in which the representative particles are in a highly ordered, repeating pattern called a crystal
Can be studied as unit cells, small representative units that repeat throughout the structure

17. Bonding in Crystalline Solids
The physical properties of solids, such as hardness, electrical conductibility, and melting point, depend on the kind of particles that make up the solid and on the strength of the attractive forces between them. (Includes ionic, covalent, and atomic substances)

18. Covalent-Network Solids
A type of crystalline solid in which strong covalent bonds forma a network extending throughout the solid
Have very high melting points due to strong covalent bonds

19. Amorphous Solids
Solids in which the arrangement of the representative particles lacks a regular, repeating pattern
Also known as “supercooled liquids”
- Liquids cooled until viscosity becomes so high that no flow can occur

20. Amorphous Solids
Get softer when heated instead of reaching an abrupt melting point like crystalline solids
Examples: glass, rubber, some plastics

21. Liquids
The physical properties of liquids are determined by the nature and strength of the intermolecular forces present between their molecules

22. Properties of Liquids
Viscosity – resistance to motion that exists between molecules of a liquid when they move past each other
Increased intermolecular forces yield increased viscosity
Glycerine has lots of hydrogen bonds
Decreased temperatures yield increased viscosity

23. Properties of Liquids
Surface tension – imbalance of attractive forces at the surface of a liquid that cause the surface to behave as if it had a film across it
Explains “beading” of liquids
As with viscosity, increased intermolecular forces or decreased temperatures yield increased surface tension

24. Phase Changes Changes of state are physical changes, not chemical.
Intermolecular forces break or form, not intramolecular
Phase changes occur when energy enters or leaves a compound.
Energy in: solid g liquid g gas
Energy out: gas g liquid g solid

25. Heating/Cooling Curves
Show the changes in temperature as energy as heat is added to (or removed from) a substance and it changes states.

26. Heating/Cooling Curve

27. Phase Changes
During phase changes, heat is still transferred, but no temperature change takes place.

28. Phase Diagrams
Shows the state of matter of a substance at increasing pressures and temperatures.
Can be used to determine the state of matter of a substance at a particular pressure and temperature.

29. Phase Diagram of Water

30. Energy Requirements
Molar heat of fusion: the energy required to melt one mole of a solid
Molar heat of vaporization: the energy required to vaporize one mole of a liquid
Molar heats of vaporization are higher because more intermolecular forces will have to be overcome to go from a liquid to a gas than a solid to a liquid.

31. Vaporization
Not all of the particles in a liquid have the same kinetic energy.
Temperature is a measurement of the average KE.
The particles of a liquid must be moving at a sufficient speed to overcome the intermolecular forces of the liquid to escape as a gas.
Thus, evaporation is endothermic: it requires energy to occur.

32. Vapor Pressure
Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid phase at a certain temperature.
For a liquid in a closed container, vaporization and condensation will occur at an equal rate once the equilibrium vapor pressure is reached.

33. Vapor Pressure and Boiling Point
When water begins to boil, bubbles of H2O gas are created as some of the molecules escape the intermolecular bonds holding the liquid together.
The bubbles will expand and rise to the surface only if the pressure exerted by the water vapor in the bubble is greater than the atmospheric pressure.

34. Vapor Pressure and Boiling Point
Thus, the vapor pressure of the water must be equal to atmospheric pressure before boiling can occur.
At temperatures below 100oC, the vapor pressure of water is below 1 atm, so the atmospheric pressure prevents boiling from occurring.

35. Elevation and Boiling Point
What happens to atmospheric pressure as you increase your elevation above sea level?
What would you expect to happen to the boiling point of water at high elevations?