Intermolecular forces: Generalizing properties

1. Intermolecular forces: Generalizing properties • Low boiling point = particles are more likely to leave liquid solution • Weaker IM forces = lower boiling point • Lower boiling point = more vapor = higher vapor pressure • High boiling point = slow evaporation • If IM forces are the same, look at formula weight. Heavier molecules have higher boiling points. • Strength of IM forces: Hydrogen bond>dipole-dipole>London dispersion

2. Polar molecules? Ions involved? Are polar molecules and ions both present? Are H atoms bonded to N,O or F atoms? Intermolecular Forces Interacting molecules or ions No Yes Yes No No Yes No Yes London Forces only Ex. Ar(l), I2(s) Dipole-Dipole Ex. H2S Hydrogen Bonding Ex. NH3, H2O Ion-dipole Forces Ex. KBr in H2O Ionic bonding Ex. NaCl

3. Water’s Properties • Hexagonal crystal shape • Molecule is polar. • Hydrogen bonding • Ice floats. • Expands during freezing until -4.0 º C. • Solid form is less dense than liquid • Surface tension: • Water “beads” on smooth surfaces. • Insects walk on water surfaces.

4. Surface tension • Force that pulls adjacent parts of a liquid surface together. • The higher the attractive forces between particles in the liquid, the higher the surface tension. • Hydrogen bonds make water have higher surface tension than most liquids. Soap Water droplet

5. Phases of matter: Comparison

6. Solids • Crystalline solids • Particles are arranged in an orderly, geometric, repeating pattern. • Examples: Emerald, diamond, calcite • Amorphous solids: (Without shape) • Particles are arranged randomly. • Examples: Glass, plastic • Network solids • Covalent bonds, usually single element arranged in orderly pattern • Examples: Diamond, graphite

7. Bonding in Solids • Molecular solids • Most are liquids or gases at room temp. • Ex. H2O, Ar • Covalent Network solids • Covalent bonds are stronger than IM forces, so substances have relatively high melting points and are harder than molecular ones. • Ex: quartz, diamond, graphite, SiO2 • Ionic solids • Ionic bonds are the strongest of all • Strength of bond depends on charge: Higher charges = higher melting point. • Crystal structures: Examples • Face-centered cubic, body-centered cubic, hexagonal close-packed structures. • Metallic Solids (metallic bonds)

8. The Crystal Lattice • 3-dimensional pattern that repeats itself over and over again. • Each ion is bonded with all oppositely charged ions that directly surround it. • NaCl forms a cube shape, called a body-centered-cubic structure. • There are 7 crystal shapes, determined by how the ions are arranged in the lattice. Mullis

9. Crystal Growth • Crystals grow by adding ions to all sides. • They grow equally in all directions from the outside. • Crystals form in 2 ways: • Solution containing a dissolved ionic compound evaporates. • An ionic solid is heated until it melts, then liquid is cooled. (Igneous rocks) Mullis

10. Energetics of Ionic Bond Formation • Recall that heat of formation of NaCl was exothermic (∆H°f = -410.9kJ/mol) • Separation of NaCl is endothermic: • (∆H = +788 kJ/mol) • The energy required to separate 1 mol of ions in an ionic lattice into gaseous ions is called lattice energy, ∆Hlattice . • Lattice energy depends on the charge on the ions and the size of the ions. Mullis

11. Lattice Energy (Ionic bonds) • Lattice energy depends on the charge on the ions and the size of the ions. • The stability of the compound comes from the attraction between ions of unlike charge. • The specific relationship is given by Coulomb’s equation: • E = kQ1Q2 Q is the charge on the particles, d is the distancedbetween their centers and k is aconstant. • As Q1 and Q2 increase, E increases and as d increases, E decreases. Mullis

12. Lattice Energies for Some Ionic Compounds

13. CrystalsImages created by Daniel Mayer or Wikimedia Commons and licensed under terms of the GNU FDL. Can examine structure using X-ray diffraction • Uses Bragg’s Law to determine distance between planes of atoms. Computer instrumentation is used to translate wave functions into photographic images. • Bragg’s Law: nλ = 2dsinΘ • n is an integer (1), λ is wavelength, d is distance between atoms, Θ angle of incidence • 3D link: http://www.le.ac.uk/eg/spg3/atomic.html cbc cfc hexagonal

14. Allotropes (different forms of same element) • Carbon (C) • Diamond • Graphite (pencil “lead”) • Charcoal • Sulfur (S) • Rhombic (puckered ring)= S8 • Phosphorous (P) • White phosphorous, P4 is most reactive, tetrahedral • Red phosphorous is more stable.

15. (1) Sulfur rings (2) graphite’s hexagonal crystal structure (3) Diamond • http://www.green-planet-solar-energy.com/the-element-sulfur.html • http://www.avogadro.co.uk/structure/chemstruc/network/g-molecular.htm • http://www.enmu.edu/services/museums/miles-mineral/images/diamond_large.jpg

16. Silicon Doping(N-type is more conductive when voltage is applied.) O:O:O:O: O:O:O:O: O:O:O:O: O:O:O:O: O:B.O:O: O:P:O:O: Silicon (4 e-) P-type N-type semiconductor hole created extra e- in lattice p = positive n= negative To customize conductive properties, add a dopant such as B (p-type), As or P (n-type) .. .. .. .. .. .. .. .. .. … .. .. .. .. .. .. .. .. .. .. .. .. .. ..

17. Glass: SiO2 • High melting point (~ 1700°C) but may vary depending on the particular structure . Very strong Si-O covalent bonds arranged in a continuous lattice have to be broken throughout the structure before melting occurs. • Is also very hard due to the need to break the very strong covalent bonds. • Doesn't conduct electricity because there aren't any delocalized electrons. All the electrons are held tightly between the atoms, and aren't free to move. • Insoluble in water and organic solvents because there are no possible attractions which could occur between solvent molecules and the silicon or oxygen atoms which could overcome the covalent bonds in the giant structure. (Water is a simple covalent structure: Each molecule is made of 3 atoms so its melting point is low compared to the giant lattice structure of solids like SiO2. )

18. Allotrope: Two or more forms of the same element that have distinctly different physical or chemical properties. • Fullerenes include C60, buckminsterfullerene, a hollow sphere resembling a soccer ball. • Graphite is a black solid that feels soft and “greasy” to the touch. Planar sheets of molecules can slip by one another easily. It is used as a lubricant and leaves black marks if rubbed on a lighter-colored surface. It conducts electricity. Selling price: < $0.01/gram. • Diamond is one of the hardest substances known (Mohs hardness = 10). Its hardness is due to rigid networks of tetrahedrons, carbon atoms covalently bound. It does not conduct electricity. Selling price = $50.00 - $20,000.00/gram.

19. Material modification • Pencil “lead” is softened by adding clay to graphite. • Gold jewelry is strengthened by adding copper or other metal. 14 karat gold means that 14/24’s of the material is Au. (The relative proportion of gold originated with a medieval coin called a mark; a mark weighed 24 karats.) • Ceramics: Developed from conventional clay (Si, O, Al) and the addition of other minerals to improve strength, melting point and brittleness. Ceramics can often get much hotter before they melt than metals. • Plastics: Synthetic polymers primarily from carbon. Disadvantage is that most are made from nonrenewable petroleum resources.

20. Changing states • Equilibrium: When there is no net change in a system. • Dynamic equilibrium: • When a vapor is in equilibrium with its liquid as one molecule leaves the liquid to become a vapor, another molecule leaves the vapor to become a liquid. In other words, an equal number of molecules will be found moving in both directions.

21. Boiling Point • Vapor pressure: Pressure exerted by a vapor = Pressure of the liquid at given temperature • Liquid boils when its vapor pressure equals pressure of the atmosphere. • Boiling is the conversion of a liquid to vapor within the liquid as well as at its surface. • Boiling point is the temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure. • Volatile liquids are liquids that evaporate readily.

22. Boiling Point, cont. • High elevation: Low atmospheric pressure • Low atmospheric pressure = lower boiling point • High pressure in pressure cooker = increased boiling point, faster cooking • If pressure above liquid increases, the liquid temperature rises until it matches the new pressure and boils again.

23. Separation by Distillation • Distillation is the separation of liquid substances according to their different boiling points. • As a liquid mixture is heated, the substance with the lower boiling point will vaporize first. • Distillate: Condensed liquid substance

24. Kinetic Energy and Equilibrium Vapor Pressure In the beginning: # particles condensing to liquid phase = # particles evaporating to gas phase • Increase temp Increase kinetic energy • Now, more molecules have enough energy to leave the liquid. • More vapor molecules = higher vapor pressure • Equilibrium will soon be established, but at a higher vapor pressure.

25. Heat of Vaporization • Amount of heat necessary to boil (or condense) 1.00 mole of a substance at its boiling point • ***1) 1.00 mole of a substance***2) There is no temperature change • The molar heat of vaporization (ΔHvap ) for water is 40.7 kJ/mol. It comes from a table. • q = ΔHvap (mass / molar mass) (q = ΔHvapn) • q is total amount of heat involved.

26. Heat of Vaporization • Vapor pressure increases nonlinearly for liquids. Mathematically, the relationship is ln(Pvap) = -ΔHvap (1/T) +C R Where C = constant characteristic of a given liquid. m = slope = -ΔHvap and x = 1/T and b = intercept = C R

27. Heat of Fusion • aka standard enthalpy change of fusion • Amount of thermal energy absorbed or lost for 1 gram of a substance to change states from a solid to a liquid or vice versa. • Temperature at which it occurs is called the melting point. • Temperature falls if thermal energy is removed from a liquid or solid • At the transition point between solid and liquid (melting point), EXTRA energy is required to go from liquid to solid and increase order. For molecules to maintain the order of a solid, extra heat must be withdrawn. • In the other direction, to create the disorder from the solid crystal to liquid, extra heat must be added. • The molar heat of fusion for water is 6.02 kJ/mol. • q = ΔHvap (mass / molar mass)

28. Phase DiagramA phase diagram is a graph of pressure vs. temperature that shows the conditions under which phases of matter exist. Critical temp (Tc): Above this, the substance cannot exist in the liquid state.

29. Phase Diagrams: Density • Negative liquid/solid slope shows density of solid is LESS than liquid (like H2O). See previous slide. • Most substances will have a positive slope of this line since most solids are more dense than the liquid: http://wine1.sb.fsu.edu/chm1045/notes/Forces/Phase/Forces06.htm

30. Triple point, TP - All three phases can exist in equilibrium at this temperature and pressure. (The solid-liquid line and the liquid-vapor line meet.) Normal boiling point, Tb - The temperature at which the vapor pressure of a liquid is equal to standard atmospheric pressure. (Standard atmospheric pressure line crosses the liquid-vapor line.) Normal melting point, Tm - The temperature at which the vapor pressure of the solid and the vapor pressure of the liquid are equal. (Standard atmospheric pressure line crosses the solid-liquid line.) Critical temperature, Tc - The temperature above which no amount of pressure will liquefy a vapor. (The liquid-vapor line becomes vertical.) Four major "points" on a phase diagram