INTERMOLECULAR ATTRACTIONS AND THE PROPERTIES OF LIQUIDS AND SOLIDS
A. Solids and liquids are virtually incompressible
1. Molecules are packed closely together
B. Intermolecular attractions (forces) -- attractive forces that exist between neighboring molecules
1. Physical properties are controlled by these forces
2. IM forces are caused by the chemical makeup of the molecules
3. Properties depend heavily on chemical makeup
4. IM forces are always much weaker than intramolecular forces (bonds)
a. Intermolecular forces determine physical properties
b. Intramolecular forces determine chemical properties
II. Intermolecular Attractions
A. Dipole-Dipole Attractions
1. net attractive force betweeen polar molecules
a. strength is 1% of a covalent bond
b. Only partial charges
c. Collisions between molecules causes the dipoles to be misaligned
B. Hydrogen Bonds
1. dipole-dipole attraction that occurs between hydrogen that is covalently bonded to a small, highly electronegative atom, and other molecules
a. Covalent bonds are very polar
b. Charges are highly concentrated at the two ends
c. 5-10 times larger than dipole-dipole attractions
C. London Forces (van der Waals forces)
1. Small IM forces found in all substances
a. Instantaneous dipole -- at any given moment the electron density of a particle can be unsymmetrical
b. Induced dipole -- dipole formed by a dipole in a neighboring atom
1. Very short lived since electrons are always moving
c. Short lived dipole-dipole attractions are called London forces
1. Exist between all molecules and ions
2. Strength of London forces depend on three factors
a. Polarizability of the electron cloud -- ease of electron cloud distortion
1. Increase electron cloud size, increase polarizability
2. Large electron cloud, stronger London forces, higher boiling points
b. Number of atoms in a molecule
1. Increase atoms, increase London forces
2. Longer molecular chain, higher boiling point, stronger IM forces
c. Molecular shape
1. More compact molecule, less London forces, less interaction with other molecules
3. London Forces and Dipole-Dipole Forces Dompared
a. London forces are often as strong as dipole-dipole forces or stronger
4. Ion-Dipole and Ion-Induced Dipole Forces of Attraction
a. Ion-dipole -- Ions interact with polar molecules
b. Ions can also induce dipoles in neutral atoms and molecules
III. Some General Properties of Liquids and Solids
A. Properties that depend primarily on tightness of packing
1. Compressibility
a. most of the space is taken up by the molecules --> nearly incompressible
2. Diffusion
a. Many more collisions --> diffusion is much slower
B. Properties that depend primarily on the strengths of intermolecular attractions
1. Retention of voume and shape
a. IM forces are great --> hold particles together
2. Surface Tension -- tendency of a liquid to seek a shpae that yields minimum surface area (sphere)
a. A molecule at the surface has greater potential energy than a molecule in the bulk of the liquid (not surrounded by other molecules)
b. By lowering surface area, you lower potential energy (more stable)
c. Strong IM forces, large surface tension
3. Wetting of a Surface by a Liquid
a. Wetting -- the spreading of a liquid across a surface to form a thin film
b. IM forces between the liquid and the surface must be about the same strength as the forces within the liquid itself
c. surfactants -- chemicals that lower surface tension of water
4. Viscosity -- resistance to flow
a. stronger IM forces, higher viscosity
b. Increase temperature, decrease viscosity (lowers IM forces)
5. Evaporation and Sublimation
a. When a molecule has enough KE, it can leave the liquid or solid and become a vapor
6. Evaporation and Cooling
a. Evaporation is a cooling effect
b. Molecueles take KE with them when they escape --> leave behind less KE (Lower temperature)
7. Factors that control the rate of evaporation
a. Surface area -- increase surface area, increase rate of evaporation
b. Temperature -- increase temperature, increase rate of evaporation
c. IM forces -- weaker IM forces, faster rate of evaporation
IV. Changes of State and Dynamic Equillibrium
A. Change of State -- a substance is transformed from one physical state to another
B. Dynamic Equilibrium -- forward rate = reverse rate
1. Rate of evaporation = rate of condensation (boiling point)
2. Rate of Sublimation = rate of condensation
3. Rate of Fusion = Rate of Melting (melting point)
4. # of molecules in each state remains constant
5. Activity continues (evaporation and condensation are always occurring)
V. Vapor Pressures of Liquids and Solids
A. Vapor Pressure -- pressure exerted by a vapor above a liquid or solid
1. Equilibrium vapor pressure of the liquid -- when dynamic equilibrium is reached
B. Factors that Affect Equilibrium Vapor Pressure
1. Increase temperature, increase rate of evaporation, increase vapor pressure
a. Vapor pressure is a function of the rate of evaporation
2. Increase IM attractions, rate of evaporation decrease, vapor pressure decrease
a. Can use vapor pressure to indicate relative IM forces
C. Factors that do not Affect the Vapor Pressure
1. Increase surface area, no change in total rate of evaporation/condensation
2. Increase amount of liquid, no change to rate of evaporation (happens at surface)
3. Increase space above liquid, change more liquid into vapor, but no change to Vapor pressure
D. Vapor Pressures of Solids
1. Equilibrium Vapor Pressure of the Solid -- when dynamic equilibrium is reached
VI. Boiling Points of Liquids
A. Boiling Point -- the temperature at which the vapor pressure of the liquid is equal to the prevailing atmospheric pressure
1. Bubbles are vapor of the liquid! not air!
B. Normal Boiling Point -- boiling point of a liquid at 1 atm
1. Increase IM forces, Increase boiling (more energy to remove molecules from liquid)
a. Down Periodic table, IM forces increaese, Higher boiling point
b. Across Periodic table, IM forces increase, higher boiling point
VII. Energy Changes during Changes of State
A. Heating Curves and Cooling Curves
1. Heating Curve -- Heat versus Temperature graphs
2. Temperature relates to KE; sloped graph, increase KE and temperature; flat graph, KE constant (heat energy must be as PE) (phase changes)
3. Cooling Curves -- Opposite of heating curves
a. Supercooling -- liquid cools below its freezing point
1. Molecules are still trying to get into orderly arrangements, KE continues to decrease
2. sudden increase in KE and then freezes at its normal freezing temperature
B. Molar Heat of Fusion, Vaporization, and Sublimation
1. Molar heat of fusion -- heat absorbed by one mole of a solid when it melts to give a liquid
2. Molar heat of vaporization -- heat absorbed by one mole of a liquid when it evaporates to give a vapor
3. Molar heat of sublimation -- heat obsorbed by one mole of a solid when it sublimes to give a vapor
4. All three are Positive (endothermic, add heat)
C. Heat of Vaporizaiton and Vapor Pressure
1. Clausius-Clapeyron equation -- used to relate vapor pressure and heat of vaporization
ln P = -DeltaHvap/RT + C
a. Similar to y = mx + b
b. ln P versus 1/T is a straight line sloping downward
D. Energy Changes and Intermolecular Attractions
1. DeltaHvap > DeltaHfus and DeltaHsub > DeltaHvap
2. Increase IM forces, increase DeltaH values; more energy needed to overcome forces
VIII. Dynamic Equilibrium and Le Chatelier's Principle
A. Le Chatelier's Principle -- when a dynamic equilibrium in a system is upset by a disturbance, the system responds in a direction that tends to counteract the disturbance, and if possible, restore equilibrium
1. heat + liquid <--> vapor
a. Increase temperature (heat), some liquid will cahnge into vapor to readjust equilibrium
b. Position of equilibrium -- relative amounts of the substances on opposite sides of the double arrows
c. position of equilibrium has shifted left in the above example
IX. Phase Diagrams
A. A graphical representation of the pressure-temperature relationships that apply to the equilibria between the phases of a substance
1. Triple Point -- Temperature and pressure at which triple equilibrium between vapor, liquid, and solid occurs
a. Most substances
1. Increase pressure, increase melting point, evaporation point, sublimation point
b. water is unique
1. Increae pressure, decrease melting point, increase evaporation point, sublimation point
2. Strong IM forces (Hydrogen bonds) cause ice to form a crystal structure that is less dense than liquid water; increase pressure, ice will melt to return to equilibrium (Le Chatelier's)
2. Critical Point -- above the critical temperature, a liquid cannot exist, regardless of pressure
a. Increase temperature, vapor density increaes, liquid density decreases
b. when densities are equal, no liiquid exists
c. Supercritical fluid -- a substance that has a temperature above its critical temperature and a density near its liquid density
d. High IM forces, high critical temperature
X. Crystalline Solids
A. Crystal Lattice and Unit Cells
1. Crystal Lattice -- repeating patterns of particles in a structure
a. Unit Cell -- Small units of the larger structure
b. Simple Cubic Unit Cell -- simplest, most symmetrical unit cell
2. Types
a. Face-centered cubic unit cell
b. Body-centered cubic unit cell
XII. Physical Properties and Crystal Types
A. Ionic Crystals
1. lattice sites occupied by oppositely charged ions
2. Hard, high melting points, brittle, strong IM forces (ion attraction), conduct electricity in solution or melted
B. Molecular Crystals
1. solids, lattice sites occupied by atoms or by molecules
2. Soft, low melting points, weak IM forces (H bonds, London forces, dipole-dipole), do not conduct electricity
C. Covalent Crystals
1. Solids, lattice positions are occupied by atoms taht are covalently bonded to other neighboring atoms
2. Sometimes called network solids
3. Very Hard, very high melting points, very strong IM forces, do not conduct electricity
D. Metallic Crystals
1. Solids, lattice sites occupied by positive ions, surrounded by free flowing electrons
2. Soft to hard, low to high melting points, conduct electricity in solid and liquid state, lusterous, high IM forces (attraction between ions and electrons)
XIII. Noncrystalline Solids
A. Amorphous solids -- no long-range repetitive internal strucutres that are found in crystals (random)
1. Glass is a good example
2. Supercooled liquids -- a term used to describe amorphous solids (kind of disorder)
3. Amorphous solids soften gradually when heated
Outline based upon:
Brady, J. E., Holum, J. R., Russell, J. W. (2000). Chemistry: The Study of Matter and Its Changes. (3rd ed.). New York: John Wiley & Sons, Inc. pp. 467-510.
