2 10.1 Intermolecular forces Inside molecules (intramolecular) the atoms are bonded to each other.Intermolecular refers to the forces between the molecules.These are what hold the molecules together in the condensed states.
3 Intermolecular forces Strongcovalent bondingionic bondingWeakDipole dipoleLondon dispersion forcesDuring phase changes the molecules stay intact.Energy used to overcome forces.
4 Dipole - Dipole 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 or ionic bonds.Weaker with greater distance.Small role in gases.
9 London Dispersion Forces Non - polar molecules also exert forces on each other.Otherwise, no 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.
11 London Dispersion Forces Weak, short lived.Lasts longer at low temperature.Eventually long enough to make liquids.More electrons, more polarizable.Bigger molecules, higher melting and boiling points.Exist in all molecules, much weaker than other forces.Also called Van der Waal’s forces.
16 10.2 LiquidsMany of the properties due to internal attraction of atoms.BeadingSurface tensionCapillary actionStronger intermolecular forces cause each of these to increase.
17 Surface Tension Molecules at the the top are only pulled inside. Molecules in the middle are attracted in all directions.Minimizes surface area.
18 Capillary Action Liquids spontaneously rise in a narrow tube. Intermolecular forces are cohesive, connecting like things.Adhesive forces connect to something else.Glass is polar.It attracts water molecules.
20 Beading If a polar substance is placed on a non-polar surface. There are cohesive,But no adhesive forces.And Visa Versa
21 Viscosity How much a liquid resists flowing. Large forces, more viscous.Large molecules can get tangled up.Cyclohexane has a lower viscosity than hexane.Because it is a circle- more compact.
22 Model for LiquidsCan’t see molecules so picture them in motion but attracted to each other.With regions arranged like solids butwith higher disorder.with fewer holes than a gas.Highly dynamic, regions changing between types.
23 Phases The phase of a substance is determined by three things. The temperature.The pressure.The strength of intermolecular forces.
24 10.3 SolidsTwo major types.Crystalline - have a regular arrangement of components in their structure. (table salt)Amorphous - those with much disorder in their structure. Said to be “frozen in place”. (glass, plastic)
25 CrystalsLattice - a three dimensional grid that describes the locations of the pieces, (atoms, ions, or molecules) in a crystalline solid.Unit Cell - The smallest repeating unit in of the lattice.Three common types.
29 10.4 Bonding Models for Metals Why do metal atoms stay together? How does their bonding effect their properties?Two Models:Electron Sea Model: A regular array of metals in a “sea” of electrons.Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.
32 The 1s, 2s, and 2p electrons are close to nucleus, so they are not able to move around. Empty Molecular OrbitalsFilled Molecular Orbitals3s2p2s1sMagnesium Atoms
33 The 3s and 3p orbitals overlap and form molecular orbitals. Empty Molecular Orbitals3pFilled Molecular Orbitals3s2p2s1sMagnesium Atoms
34 Electrons in these energy level can travel freely throughout the crystal. Empty Molecular Orbitals3plFilled Molecular Orbitals3s2p2s1sMagnesium Atoms
35 Empty Molecular Orbitals 3p l Filled Molecular Orbitals 3s This makes metals conductorsMalleable because the bonds are flexible.Empty Molecular Orbitals3plFilled Molecular Orbitals3s2p2s1sMagnesium Atoms
36 10.5 C & Si - Network Atomic Solids Composed of strong directional covalent bonds that are best viewed as a “giant molecule”.brittledo not conduct heat or electricitycarbon, silicon-basedgraphite, diamond, ceramics, glassDiamond - hardest natural substance on earth, insulates both heat and electricity.Graphite - slippery, conducts electricity.
37 Diamond - each Carbon is sp3 hybridized, connected to four other carbons. Carbon atoms are locked into tetrahedral shape.Strong bonds give the huge molecule its hardness.
38 Graphite is different Each carbon is connected to three other carbons and sp2hybridized.The molecule is flat with 120º angles in fused 6 member rings.The bonds extend above and below the plane.
39 The layers slide by each other. This bond overlap forms a huge bonding network.Electrons are free to move through out these delocalized orbitals.The layers slide by each other.
40 10.6 Molecular solids. 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.
41 Those without dipoles. Most are gases at 25ºC. The only forces are London Dispersion Forces.These depend on size of atom.Large molecules (such as I2 ) can be solids even without dipoles.
42 Those with dipolesDipole-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 freezing points.
43 Water is special Each molecule has two polar O-H bonds. Each molecule has two lone pair on its oxygen.Each oxygen can interact with 4 hydrogen atoms.HOd+
44 Water is specialHOd+This gives water an especially high melting and boiling point.
45 10.7 Ionic SolidsThe extremes in dipole dipole forces-atoms are actually held together by opposite charges.Huge melting and boiling points.Atoms are locked in lattice so they are hard and brittle.Every electron is accounted for so they are poor conductors - good insulators.
46 What type of solid? Diamond Quartz PH3 NH4NO3 H2 Ar Mg Cu KCl C6H12O6 SF2
47 10.8 Vapor PressureVaporization - change from liquid to gas at boiling point.Evaporation - change from liquid to gas below boiling point.Heat (or Enthalpy) of Vaporization (Hvap) - the energy required to vaporize 1 mol at 1 atm.
48 Vaporization is an endothermic process - it requires heat. Energy is required to overcome intermolecular forces.Responsible for cool earth.Why we sweat.
49 Calculating vapor pressure The vapor pressure of water at 25°C is 23.8 torr, and the Hvap of water at 25°C is 43.9 kJ/mol. Calculate the vapor pressure of water of water at 50. °C
52 Practice #79The enthalpy of vaporization of mercury is 59.1 kJ/mol. The normal boiling point of mercury is 357°C. What is the vapor pressure of mercury at 25°C?
53 Practice #79The enthalpy of vaporization of mercury is 59.1 kJ/mol. The normal boiling point of mercury is 357°C. What is the vapor pressure of mercury at 25°C? Ans: 2.56 x 10-3 torr
54 Condensation Change from gas to liquid. Achieves a dynamic equilibrium with vaporization in a closed system.What is a closed system?A closed system means matter can’t go in or out.What the heck is a “dynamic equilibrium?”
55 Dynamic equilibriumWhen first sealed the molecules gradually escape the surface of the liquid.
56 Dynamic equilibriumWhen 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.
57 Dynamic equilibriumWhen first sealed the molecules gradually escape the surface of the liquidAs 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 Dynamic equilibrium Rate of Vaporization = Rate of Condensation Molecules are constantly changing phase “Dynamic”The total amount of liquid and vapor remains constant “Equilibrium”
59 Vapor pressure The pressure above the liquid at equilibrium. Liquids with high vapor pressures evaporate easily. They are called volatile.Increases with increasing temperature.Decreases with increasing intermolecular forces.Bigger molecules (bigger LDF)More polar molecules (dipole-dipole)
60 Dish of HgVacuumPatm=760 torrA barometer will hold a column of mercury 760 mm high at one atm.
61 Dish of HgVacuumPatm=760 torrA 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.
62 WaterA 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.Patm=760 torrDish of Hg
63 Dish of Hg The mercury is pushed down by the vapor pressure. Water VaporThe mercury is pushed down by the vapor pressure.Patm = PHg + PvapPatm - PHg = Pvap= 24 torr736 mm HgDish of Hg
64 Temperature Effect Energy needed to overcome intermolecular forces # of moleculesKinetic energy
65 Energy needed to overcome intermolecular forces At higher temperature more molecules have enough energy - higher vapor pressure.Energy needed to overcome intermolecular forcesEnergy needed to overcome intermolecular forcesT1T1# of moleculesT2Kinetic energy
66 Changes of stateThe 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 Hfus
67 Water and Steam Water and Ice Heating Curve for WaterSteamWater and SteamWaterWater and IceIce
68 Heat of Vaporization Heat of Fusion Heating Curve for WaterSlope is Heat CapacityHeat ofVaporizationHeat ofFusion
69 Melting PointMelting 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 Water Vapor Vapor Solid Water Liquid Water If the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium.Solid WaterLiquid WaterWater Vapor Vapor
72 While the molecules of vapor condense to a liquid. Solid WaterLiquid WaterWater Vapor Vapor
73 This can only happen if the temperature is above the freezing point since solid is turning to liquid.Solid WaterLiquid WaterWater Vapor Vapor
74 Water Vapor Vapor Solid Water Liquid Water If the vapor pressure of the liquid is higher than that of the solid, the liquid will release molecules to achieve equilibrium.Solid WaterLiquid WaterWater Vapor Vapor
75 While the molecules condense to a solid. Solid WaterLiquid WaterWater Vapor Vapor
76 The temperature must be below the freezing point since the liquid is turning to a solid. Solid WaterLiquid WaterWater Vapor Vapor
77 Water Vapor Vapor Solid Water Liquid Water If the vapor pressure of the solid and liquid are equal, the solid and liquid are vaporizing and condensing at the same rate. The Melting point.Solid WaterLiquid WaterWater Vapor Vapor
78 Boiling PointReached when the vapor pressure equals the pressure of the surrounding atmosphere.Normal boiling point is the boiling point at 1 atm pressure.Super heating - Rapid heating above the boiling point allows liquid state to exist above normal boiling point.Supercooling - Rapid cooling below the freezing point allows liquid state to exist below the normal freezing point.
81 10.9 Phase DiagramsA plot of temperature versus pressure for a closed system, with lines to indicate where there is a phase change.Critical temperature: temperature above which the vapor can not be liquefied.Critical pressure: pressure required to liquefy at the critical temperature.Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).
87 Like most substances, bromine exists in one of the three typical phases. Br2 has a normal melting point of -7.2°C and a normal boiling point of 59°C. The triple point for Br2 is -7.3°C and 40 torr, and the critical point is 320°C and 100 atm. Using this information, sketch a phase diagram for bromine indicating the points described above.Based on your phase diagram, order the three phases from least dense to most dense.What is the stable phase of Br2 at room temperature and 1 atm?Under what temperature conditions can liquid bromine never exist?What phase changes occur as the temperature of a sample of bromine at 0.10 atm is increased from -50°C to 200°C?