2. Liquids and solids

3.  compared to gases.  They are incompressible.  Their density doesn’t change with temperature.  These similarities are due ◦ to the molecules being close together in solids and liquids ◦ and far apart in gases  What holds them close together?

4.  Inside molecules (intramolecular) the atoms are bonded to each other.  Intramolecular forces are covalent, ionic and metallic bonds  Intermolecular refers to the forces between the molecules...these may also involve those listed above.  These are what hold the molecules together in the condensed states.

5.  Strong ◦ covalent bonding ◦ ionic bonding  Weak ◦ Dipole-dipole ◦ London dispersion forces  During phase changes the molecules stay intact.  Phase changes involve energy input or release…  Energy used to overcome I.M. forces.

6.  Remember where the polar definition came from?  Molecules line up in the presence of a electric field. The opposite ends of the dipole can attract each other so the molecules stay close together.  1% as strong as covalent bonds.  Weaker with greater distance.  Small role in gases.  (Which correction factor in the Van der Waals Equation?)

7. + - + - + - + - + - + - + - + - + - + -

8.  Especially strong dipole-dipole forces when H is attached to N, O, or F  These three because- ◦ They have high electronegativity. ◦ They are small enough to allow close approach of the dipoles.  Affects boiling point, melting point, and other colligative properties (freezing point depression, boiling point elevation and osmotic pressure).

9. CH 4 SiH 4 GeH 4 SnH 4 PH 3 NH 3 SbH 3 AsH 3 H2OH2O H2SH2S H 2 Se H 2 Te HF HI HBr HCl Boiling Points 0ºC 100 -100 200

10. ++ -- ++

11.  Ionic substances are soluble in water due to the high degree of polarity.  Ionic solubility in other solvents depends upon the solvent polarity.  Examine the solubility of an ionic compound K 2 SO 4 in water vs. isopropyl alcohol. 70% isopropyl alcohol Saturated K 2 SO 4 soln.

12.  Non-polar molecules also exert forces on each other.  Otherwise, nonpolar substances could not exist as solids or liquids.  Electrons are not evenly distributed at every instant in time.  Have an instantaneous dipole.  Induces a dipole in the atom next to it.  Induced dipole-induced dipole interaction.

13. HH HH HH HH ++ ++ HH HH ++ -- ++ 

14.  Weak, short lived.  More significant at lower temperatures.  Eventually enough to change gases to liquids.  More electrons on a particle = more polarizable.  Bigger molecules, higher melting and boiling points.  Much, much weaker than other forces when compared…but…can be strong!  Also called Van der Waal’s forces.

15.  Many of the properties due to internal attraction of atoms. ◦ Beading ◦ Surface tension ◦ Capillary action  Stronger intermolecular forces cause each of these to increase.

16.  Molecules in the middle are attracted in all directions. u Molecules at the the top are only pulled inside. u Minimizes surface area.

17.  Liquids spontaneously rise in a narrow tube.  Intermolecular forces are cohesive, connecting like things.  Adhesive forces connect to something else.  Glass is polar.  It adhesively attracts water molecules.

19.  If a polar substance is placed on a non-polar surface (think of water beading up on a waxed car). ◦ There are cohesive, ◦ But no adhesive forces.

20.  How much a liquid resists flowing.  Large forces, more viscous.  Large molecules can get tangled up, which increases viscosity (corn syrup).  Cyclohexane has a lower viscosity than hexane (examine their structures).  Because it is a circle-more compact.

21. ◦ Hydrogen bonding ◦ Polar bonding ◦ LDF ◦ Compare iodine, water and oil…

22.  The phase of a substance is determined by three things.  The temperature.  The pressure.  The strength of intermolecular forces.

23.  Two major types.  Amorphous- those with much disorder in their structure.  Crystalline- have a regular arrangement of components in their structure.

24.  Lattice- a three dimensional grid that describes the locations of the pieces in a crystalline solid.  Unit Cell-The smallest repeating unit in of the lattice.  Three common types.

28.  There are many amorphous solids. ◦ Like glass.  We tend to focus on crystalline solids.  Two Types. ◦ Ionic solids have ions at the lattice points. ◦ Molecular solids have molecules.  Sugar vs. Salt.

29.  Using diffraction patterns to identify crystal structures.  Talks about metals and the closest packing model.  It is interesting, but trivial…unless you are a crystallographer or geologist!  We need to focus on metallic bonding.  Why do metal atoms stay together.  How the bonding effects their properties.

30. 1s 2s 2p 3s 3p Filled Molecular Orbitals Empty Molecular Orbitals Magnesium Atoms

31. Filled Molecular Orbitals Empty Molecular Orbitals The 1s, 2s, and 2p electrons are close to nucleus, so they are not able to move around. 1s 2s 2p 3s 3p Magnesium Atoms

32. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms The 3s and 3p orbitals overlap and form molecular orbitals.

33. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms Electrons in these energy levels can travel freely throughout the crystal.

34. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms This makes metals thermal and electrical conductors… and malleable because the bonds are flexible.

35.  There are three types of solid carbon molecules (called allotropes)  Amorphous- coal…uninteresting.  Diamond- hardest natural substance on earth; electrical and thermal insulator.  Graphite- slippery, conducts electricity.  How the atoms in these network solids are connected explains these phenomena.

36.  Carbon atoms are locked into tetrahedral shape.  Strong  bonds give the huge molecule its hardness.

37. The space between orbitals make it impossible for electrons to move around Empty MOsFilled MOs E

38.  Each carbon is connected to three other carbons and sp 2 hybridized.  The molecule is flat with 120º angles in fused 6 member rings.  The  bonds extend above and below the plane.

39.  Electrons are free to move throughout these delocalized orbitals.  The layers slide by each other.  However, he bonding within the layer is very strong  CARBON FIBER!!!

40. Applications of Carbon Fiber

41.  Molecules occupy the corners of the lattices.  Different molecules have different forces between them.  These forces depend on the size of the molecule.  They also depend on the strength and nature of dipole moments.  Molecular solids with dipoles are typically in a condensed state of matter at room temp.

42.  Dipole-dipole forces are generally stronger than L.D.F.  Hydrogen bonding is stronger than Dipole- dipole forces.  No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds.  Stronger forces lead to higher melting and boiling points.  What about vapor pressure? Examples?

43.  Most are gases at 25ºC.  The only forces are London Dispersion Forces.  These depend on the size of the particle.  Large molecules (such as I 2 ) can be solids even without dipoles. Examine the phases of the halogens!

44.  Each molecule has two polar O-H bonds. H O H   --

45.  Each molecule has two lone pairs of electrons on its oxygen. H O H  

46.  Each molecule has two polar O-H bonds.  Each molecule has two lone pairs on its oxygen.  Each oxygen can interact with 4 hydrogen atoms.  UNIQUE! H O H  

47.  This gives water an especially high melting and boiling point. H O H   H O H   H O H  

50.  The extremes in dipole-dipole forces…atoms are actually held together by opposite charges.  VERY HIGH melting and boiling points.  Atoms are locked in lattice, making the substance hard and brittle.  Every electron is accounted for so they are poor conductors-good insulators.  Molten ionic substances will conduct (free moving ions or e - = conductors)

51.  Vaporization - change from liquid to gas at boiling point.  Evaporation - change from liquid to gas below boiling point  Heat (or Enthalpy) of Vaporization (  H vap )- the energy required to vaporize 1 mol at 1 atm.

52.  Vaporization is an endothermic process - it requires heat.  Energy is required to overcome intermolecular forces.  Responsible for cool earth.  Why we sweat (boys) or perspire (girls)!

53.  Change from gas to liquid.  Always exothermic!  Achieves a dynamic equilibrium with vaporization in a closed system.  A closed system means matter can’t go in or out.  What the heck is a “dynamic equilibrium?”

54.  When first sealed the molecules gradually escape the surface of the liquid.

55.  As the molecules build up above the liquid some condense back to a liquid.

56.  When first sealed the molecules gradually escape the surface of the liquid.  As the molecules build up above the liquid some condense back to a liquid.  As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense.

57.  When first sealed the molecules gradually escape the surface of the liquid  As the molecules build up above the liquid some condense back to a liquid.  As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense.  Equilibrium is reached when

58. Rate of Vaporization = Rate of Condensation  Molecules are constantly changing phase “Dynamic”  The total amount of liquid and vapor remains constant “Equilibrium”

59.  The pressure above the liquid at equilibrium.  Liquids with high vapor pressures evaporate easily. They are called volatile.  Decreases with increasing intermolecular forces. ◦ Bigger molecules (bigger LDF) ◦ More polar molecules (dipole-dipole)

60.  Increases with increasing temperature.  Easily measured in a barometer.  Diagramed in the next series of slides…

61. Dish of Hg Vacuum P atm = 760 torr A barometer will hold a column of mercury 760 mm high at one atm

62. Dish of Hg Vacuum P atm = 760 torr A barometer will hold a column of mercury 760 mm high at one atm. If we inject a volatile liquid in the barometer it will rise to the top of the mercury.

63. Dish of Hg P atm = 760 torr A barometer will hold a column of mercury 760 mm high at one atm. If we inject a volatile liquid in the barometer it will rise to the top of the mercury. There it will vaporize and push the column of mercury down. Water

64. Dish of Hg 736 mm Hg Water Vapor  The mercury is pushed down by the vapor pressure.  P atm = P Hg + P vap  P atm - P Hg = P vap  760 - 736 = 24 torr

65. Kinetic energy # of molecules T1T1 Energy needed to overcome intermolecular forces

66. Kinetic energy # of molecules T1T1 Energy needed to overcome intermolecular forces T1T1 T2T2  At higher temperature more molecules have enough energy - higher vapor pressure. Energy needed to overcome intermolecular forces

67.  The graph of temperature versus heat applied is called a heating curve.  The temperature a solid turns to a liquid is the melting point.  The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion  H fus

68. Heating Curve for Water Ice Water and Ice Water Water and Steam Steam

69. Heating Curve for Water Heat of Fusion Heat of Vaporization Slope is Heat Capacity

70.  Melting point is determined by the vapor pressure of the solid and the liquid.  At the melting point the vapor pressure of the solid = vapor pressure of the liquid

71. Solid Water Liquid Water Water Vapor Vapor

72. Solid Water Liquid Water Water Vapor Vapor  If the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium.

73. Solid Water Liquid Water Water Vapor Vapor  While the molecules of condense to a liquid.

74.  This can only happen if the temperature is above the freezing point since solid is turning to liquid. Solid Water Liquid Water Water Vapor Vapor

75.  If the vapor pressure of the liquid is higher than that of the solid, the liquid will release molecules to achieve equilibrium. Solid Water Liquid Water Water Vapor Vapor

76. Solid Water Liquid Water Water Vapor Vapor  While the molecules condense to a solid.

77.  The temperature must be below the freezing point since the liquid is turning to a solid. Solid Water Liquid Water Water Vapor Vapor

78.  If the vapor pressure of the solid and liquid are equal, the solid and liquid are vaporizing and condensing at the same rate…called The Melting point ! Solid Water Liquid Water Water Vapor Vapor

79.  Reached when the vapor pressure equals the external pressure.  Normal boiling point is the boiling point at 1 atm pressure.  Super heating - Heating above the boiling point.  Supercooling - Cooling below the freezing point.

80.  A plot of temperature versus pressure for a closed system, with lines to indicate where there is a phase change.

81.  Water’s triple point diagram is truly an anomaly.  Unfortunately, it is the most widely know substances and familiar to us!  Carbon dioxide’s triple point is much more like that of a “normal” pure substance.  Make sure you are familiar with this fact!

82. Temperature Solid Liquid Gas 1 Atm A A B B C C D D D Pressure D

83. Solid Liquid Gas Triple Point Critical Point Temperature Pressure

84. Solid Liquid Gas  This is the phase diagram for water.  The density of liquid water is higher than solid water. Temperature Pressure

85. Solid Liquid Gas 1 Atm  This is the phase diagram for CO 2  The solid is more dense than the liquid  The solid sublimes at 1 atm. Temperature Pressure

86. For carbon, the space between orbitals makes it impossible for electrons to move around…thus a great insulator! (not graphite) Empty MOsFilled MOs E Consider Carbon (in comparison to silicon)

87. In silicon, the gap between orbitals is smaller…making it possible for electrons to cross the gap. Empty MOs Filled MOs E In fact…like other semiconductors, the conductivity of Si increases with increased temperatures! (unlike metals)

88.  If you replace some of the Si atoms with some As atoms…you create a n-type semiconductor  As has one more e- than Si…which becomes available for the conduction band.  n-type means a negative type (extra neg. e-)  p-type occurs when you replace some Si atoms with B which has one less valence e-  This creates an electron hole (positive area)