1. Chapter 11: States of matter; liquids and solids
Chemistry 1062: Principles of Chemistry II
Andy Aspaas, Instructor

2. Gases, liquids, and solids
Gas: compressible fluid
Mostly empty space
Particles in constant random motion
Vapor: gas form of substance normally solid or liquid
Liquid: incompressible fluid
Tightly packed particles
Still in constant random motion
Solid: incompressible and rigid
Particles only vibrate about their fixed sites

3. Phase transitions
Melting (fusion): transition from a solid to a liquid
Freezing: transition from a liquid to a solid
Vaporization: change of a solid or liquid to a vapor
Sublimation: change of a solid directly to a vapor
Condensation: change of a gas to a liquid or a solid
Deposition: change of a gas directly to a solid

4. Vapor pressure
Vapor pressure of a liquid: partial pressure of vapor over the liquid, measured at equilibrium
Vaporization and condensation are happening continually
When the rate of condensation equals the rate of vaporization, equilibrium has been reached
Vapor pressure depends on temperature
More kinetic energy in the liquid makes vaporization occur more quickly
Volatile liquids and solids have high vapor pressure

5. Boiling point and melting point
Boiling point: temperature at which the vapor pressure of a liquid equals the atmospheric pressure
Bubbles form in the liquid when it’s heated enough that its vapor pressure equals external pressure
Once boiling begins, the temperature of the liquid stops rising
Freezing point: temperature at which a pure liquid changes to a solid
Identical to melting point: temperature at which a solid changes to a liquid

6. Heat of phase transition
Heat of fusion: amount heat required to melt a solid
H2O(s)  H2O(l); ∆Hfus = 6.01 kJ/mol
Heat of vaporization: amount of heat required to vaporize a liquid
H2O(l)  H2O(g); ∆Hvap = 40.7 kJ/mol
Heating curve: no temperature change during a phase transition

7. Phase transitions
Heat must be added for melting or vaporization
Endothermic processes
Positive value of ∆H
Heat must be removed for condensation or freezing
Exothermic processes
Negative value of ∆H

8. Clausius-Clapeyron equation
Vapor pressure of a liquid depends on temperature
Also depends on ∆Hvap for that liquid
Clausius-Clapeyron equation:
The two-point form removes the constant B from the equation

9. Phase diagrams Phase diagram for water:
Curves show exptl points at which phases are in equilibrium
Solid-liquid line (melting-point curve)
Leans left if liquid is more dense than solid (as in water)
Otherwise leans right

10. Phase diagrams Curve AC gives the vapor pressure at any temperature
Triple point, A: point at which all 3 phases exist in equilibrium
Sublimation occurs when heating at pressures below triple point
Critical point, C: point above which gas and liquid are no longer distinct
Supercritical fluid: at temperature and pressure above critical point

11. Liquid state: surface tension
Surface tension: energy required to increase the surface area of a liquid by a unit amount
Liquids will minimize their amount of surface area (since intermolecular forces pull molecules at the surface inward)
Dissolved substances reduce surface tension by interrupting intermolecular forces
Capillary rise: when a liquid is attracted to a solid capillary (like water to glass) it will rise in the capillary to minimize surface tension at the inverted meniscus

12. Viscosity
Viscosity: resistance to flow in a liquid or gas
Strong intermolecular forces increase viscosity and cause a liquid to flow more slowly (like syrup)
Motor oil viscosity in SAE units
10W/30 has SAE 10 in winter, 30 in summer (because viscosity of engine oil increases as temperature decreases)

13. Intermolecular forces
Dipole-dipole force: attractive interaction between polar molecules
Polar molecules have a dipole moment
+ and – sides of the molecule resulting from electronegativity differences
Negative side of one molecule will be attracted to the positive side of another molecule
London dispersion forces: attractive forces resulting from instantaneous induced dipoles caused by random electron groupings in an atom
High molecular weight = higher London forces
Compact molecule = lower London forces

14. Predicting boiling points, vapor pressure, and viscosity
Polar molecules have stronger intermolecular forces than nonpolar molecules of similar MW (dipole-dipole)
Larger molecules have stronger intermolecular forces than smaller molecules of similar polarity (London)
Strong intermolecular forces
Low vapor pressure (less chance for molecule to break free)
High boiling point (high temperature required to raise vapor pressure to atmospheric pressure)
High viscosity (molecules cling strongly together)

15. Hydrogen bonds
Hydrogen bonds: much stronger than dipole-dipole or London forces
Interaction between H and Y in –X–H·····Y–
Where X and Y are F, O, or N (electronegative atoms)
H has a strong + when covalently bonded to an electronegative atom
Explains why CH3OH boils at 65 °C and CH3F boils at –78 °C (both have similar MW and polarities)

16. Classification of solids
Molecular solid: solid that consists of atoms or molecules held together by intermolecular forces
H2O(s), CO2(s)
Metallic solid: any solid metal contains a regular arrangement of positive metal nuclei surrounded by a sea of delocalized electrons
Ionic solid: consists of cations and anions held together by electrostatic attractions of opposite charges (ionic bonds)
Covalent network solid: consists of atoms held together in large networks or chains by covalent bonds

17. Melting point of solids
Molecular solid: low melting point (below 300 °C)
Only weak intermolecular forces to overcome
Ionic and covalent network solids: high melting point (800 °C or above)
Ionic and covalent bonds require much more energy to be broken
Larger charges make even higher mp’s in ionic solids
Metals: variable (Hg: –39 °C; W: 3410 °C)

18. Hardness and conductivity of solids
Molecular solids: generally soft and brittle
Ionic solids: generally hard and brittle
3-dimensional covalent network solids are tremendously hard
Metals are malleable since the atoms can move past each other
Metals (and molten ionic solids) conduct electricity because electrons can move freely

19. Crystalline solids
Crystalline solids have an ordered 3-dimensional structure of particles
Amorphous solids have their particles locked into random positions
Crystal lattice: 3-dimensional geometric arrangement of particles in a crystal
Unit cell: smallest boxlike unit from which a crystal lattice can be constructed
Unit cells are stacked in 3 dimensions

20. Cubic unit cells
Simple cubic unit cell: atoms only at the corners of a cubic unit cell
Only 1/8 of a corner atom is contained in a unit cell
Body-centered cubic: atom at the center of a unit cell as well as the corners
Face-centered cubic: atoms in the center of each of the faces as well as the corners
Only 1/2 of a face atom is contained in a unit cell

21. Cubic unit cells

22. Molecular solids; closest packing
In simple molecular solids (such as individual noble gas atoms), nondirectional London forces govern the packing of the atoms together to form a solid
Spheres pack together most closely in a hexagonal honey-comb type arrangement on one layer
A second layer can be nested in half of the crevices left by the first layer
A third layer can be nested in the crevices in the second layer, but there are two possible positions for this layer

23. Close packing arrangements
Hexagonal close packing (hcp): when 3rd layer is identical to 1st (in x holes left by 2nd layer)
Cubic close packing (ccp): when 3rd layer is different (in y holes left by 2nd layer)

24. Cubic close-packing
Cubic close-packing has a face-centered cubic lattice structure (tilted on its side)

25. Coordination number
Coordination number: number of nearest-neighbor adjacent atoms any one atom can have
Close packing structures: CN = 12
hcp, ccp/face-centered cubic
Body-centered cubic: CN = 8
Square cubic: CN = 4

26. Ionic solids
3 common structures: CsCl structure, NaCl structure, and zinc blende (cubic ZnS) structure
CsCl structure
1 of each per unit cell
When ions are similar in size
Expanded simple cubic

27. NaCl structure
Expanded face-centered cubic structure
When cation and anion are moderately different in size

28. Zinc blende structure
When cation and anion are significantly different in size
Anion is expanded body-centered cubic
Cation is in 4 opposite tetrahedral quadrants completely inside the unit cell

29. Unit cell calculations
Length of side of a unit cell
Volume of a unit cell
Mass of a unit cell
Mass of individual atoms
Molar mass
Forward or backwards!