Chapter 12 Insert picture from First page of chapter Intermolecular Forces and the Physical Properties of Liquids and Solids Copyright McGraw-Hill 2009 1
12.1 Intermolecular Forces Intermolecular forces are the attractive forces holding particles together in the condensed (liquid and solid) phases of matter Result from coulombic attractions Dependent on the magnitude of the charge Dependent on distance between charges Weaker than forces of ionic bonding Involve partial charges Copyright McGraw-Hill 2009 2
The Three Phases of Matter condensed phases Copyright McGraw-Hill 2009 3
Types of Intermolecular Forces van der Waals forces between atoms and molecules of pure substances Dipole-dipole interactions attractive forces between polar molecules Hydrogen bonding attractive force in polar molecules containing a H atom bonded to a small, highly electronegative element (N, O and F) (London) Dispersion forces attractive forces arising from instantaneous dipoles and induced dipoles Copyright McGraw-Hill 2009 4
Arrangement of Polar Molecules in a Liquid and a Solid. Copyright McGraw-Hill 2009 5
Intermolecular forces determine certain physical properties. Copyright McGraw-Hill 2009 6
Hydrogen Bonds Between HF Molecules Copyright McGraw-Hill 2009 7
Effect of Molar Mass and Hydrogen Bonding on Boiling Points Copyright McGraw-Hill 2009 8
Instantaneous and Induced Dipoles Magnitude depends on the ability to be polarized which is greater for larger molecules. Copyright McGraw-Hill 2009 9
Polarization and Molar Mass Copyright McGraw-Hill 2009 10
What kind(s) of intermolecular forces exist in CH 2 ClCH 2 COOH(l) Cl H O H C C C O H H H dispersion forces dipole-dipole interactions hydrogen bonding Copyright McGraw-Hill 2009 11
Ion-dipole Interactions Occur in mixtures of ionic and polar species Coulombic attraction between ions and polar molecules Dependent upon Size and charge of ion Dipole moment of the molecule Size of the molecule Can also be repulsive Copyright McGraw-Hill 2009 12
12.2 Properties of Liquids Surface Tension a quantitative measure of the elastic force at the surface of a liquid Manifestations Formation of a mensicus Capillary action which results from a combination of Cohesion (attractions between like molecules, cohesive forces) Adhesion (attractions between unlike molecules, adhesive forces) Copyright McGraw-Hill 2009 13
Effect of Surface Tension Copyright McGraw-Hill 2009 14
Intermolecular Forces: Surface versus Interior of a Liquid Copyright McGraw-Hill 2009 15
Cohesion and Adhesion Water Mercury Adhesion > Cohesion Cohesion > Adhesion Copyright McGraw-Hill 2009 16
Viscosity a measure of a fluid s resistance to flow Units: N. s/m 2 The higher the viscosity the greater the resistance to flow Varies inversely with temperature Stronger intermolecular forces produce higher viscosities Copyright McGraw-Hill 2009 17
Glycerol high viscocity due to Three hydrogen bonding sites Molecular shape Copyright McGraw-Hill 2009 18
Vapor Pressure of a Liquid Depends on the magnitude of intermolecular forces Temperature dependent T 1 <T 2 Copyright McGraw-Hill 2009 19
Equilibrium Vapor Pressure a dynamic state Evaporation (vaporization) liquid molecules escape into the gas phase Condensation gas molecules return to the liquid phase Copyright McGraw-Hill 2009 20
Comparison of Rates of Evaporation and Condensation at Constant Temperature Copyright McGraw-Hill 2009 21
Effect of Temperature and Intermolecular Forces on Vapor Pressure Copyright McGraw-Hill 2009 22
Clausius-Clapeyron Equation linear relation between temperature and vapor pressure where R = 8.314 J/K. mol At two temperatures, T 1 and T 2 : Copyright McGraw-Hill 2009 23
An unknown compound exhibits a vapor pressure of 255 mmhg at 25.5 o C and 434 mmhg at 48.8 o C. What is DH vap of this substance? Copyright McGraw-Hill 2009 24
Copyright McGraw-Hill 2009 25 K 298.65 273.15 C 25.5 1 o T K 321.95 273.15 C 48.8 2 o T D K 298.65 1 K 321.95 1 mol 8.314J/K mmhg 434 mmhg 255 vap H ln 3 3 3.3484x10 3.1061x10 mol 8.314J/K 0.53178 D H vap D 1 2 vap 2 1 1 1 T T R H P P ln
0.53178 D H vap 8.314J/K mol 2.423x10 4 1 K ( 0.53178)(8.314J/K 2.423x10 4 1 K mol) DH vap 1.82 x 10 4 J/mol DH 18.2 kj/mol DH vap vap Copyright McGraw-Hill 2009 26
12.3 Crystal Structure Crystalline solid possesses rigid and long-range order Lattice structure arrangement of particles in a crystalline solid Depends on nature of particles Depends on size of particles Stability depends on type of force between particles (ionic or covalent bonds and/or intermolecular forces) Copyright McGraw-Hill 2009 27
unit cell basic repeating structural unit of a crystalline solid lattice point each atom ion or molecule in a unit cell single unit cell 3-D array of unit cells Copyright McGraw-Hill 2009 28
Seven Types of Unit Cells Copyright McGraw-Hill 2009 29
Coordination number number of atoms (particles) surrounding an atom in a crystal lattice Indicates how tightly atoms pack Larger coordination numbers indicate tighter packing Copyright McGraw-Hill 2009 30
Types of cubic unit cells simple or primitive (scc) body-centered (bcc) face-centered (fcc) Copyright McGraw-Hill 2009 31
Alternate Perspective of bcc Arrangement Copyright McGraw-Hill 2009 32
Sharing of Atoms by Adjacent Unit Cells corner atom edge atom face atom Copyright McGraw-Hill 2009 33
Allocation of atoms among unit cells corner atoms 1 / 8 atom within unit cell face atoms 1 / 2 atom within unit cell body atoms 1 atom within unit cell Number of atoms per unit cell scc: 1 atom bcc: 2 atoms fcc: 4 atoms Copyright McGraw-Hill 2009 34
Closest Packing most efficient way to arrange atoms in a crystal hexagonal closest packed (ABA) cubic closest packed (ABC) Copyright McGraw-Hill 2009 35
Closest Packing and Cubic Unit Cells hexagonal cubic fcc Copyright McGraw-Hill 2009 36
Geometric Relationships Copyright McGraw-Hill 2009 37
When silver crystallizes, it forms facecentered cubic cells. The unit cell edge o length is 4.087A. Calculate the density of silver. Copyright McGraw-Hill 2009 38
Mass of unit cell m atoms 107.9 amu 1g unit cell atom 6.02210 4 22 23 amu 7.16710 g Volume of unit cell a V 4.087 a 3 Density of unit cell d m V 1m A 110 10 100 cm 1m A 4.08710 8 cm 8 3 23 3 4.087 10 cm 6.827 10 cm 7.16710 6.82710 22 23 g/unit cell 10.5 g/cm 3 cm /unit cell Copyright McGraw-Hill 2009 39 3
X-ray diffraction utilizes the scattering of X-rays and the resulting scattering patterns to deduce arrangement of particles Copyright McGraw-Hill 2009 40
Bragg equation Copyright McGraw-Hill 2009 41
12.4 Types of Crystals Ionic Crystals Composed of anions and cations Held together by coulombic forces Anions generally are bigger than cations Size and relative number of each ion determines the crystal structure Copyright McGraw-Hill 2009 42
Unit cell of NaCl as Defined by Cl or Na + Copyright McGraw-Hill 2009 43
Examples of Ionic Crystal Lattices CsCl ZnS CaF 2 Copyright McGraw-Hill 2009 44
How many of each ion are contained within a unit cell of CaF 2? Copyright McGraw-Hill 2009 45
Ca 2+ F Ca 2+ 8 corner ions x 1 / 8 = 1 ion 6 face ions x ½ = 3 ions 4 ions of Ca 2+ F 8 body ions x 1 = 8 ions of F Copyright McGraw-Hill 2009 46
Covalent crystals Held together by covalent bonds diamond graphite Copyright McGraw-Hill 2009 47
Molecular crystals Lattice points occupied by molecules Held together by intermolecular forces water Copyright McGraw-Hill 2009 48
Metallic crystals Lattice points occupied by atoms Generally bcc, fcc, hexagonal closest packed Very dense Bonding arises from delocalized electrons Copyright McGraw-Hill 2009 49
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12.5 Amorphous Solids Lack regular arrangement of atoms Glass is a familiar and important amorphous solid Transparent fusion of inorganic materials Chief component SiO 2 Behaves more as a liquid than a solid Copyright McGraw-Hill 2009 51
Comparison of crystalline quartz and amorphous quartz glass crystalline amorphous Copyright McGraw-Hill 2009 52
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12.6 Phase Changes Phase homogenous part of a system that is separated from the rest of the system by a well-defined boundary Phase change transition from one phase to another Caused by the removal or addition of energy Energy involved is usually in the form of heat Copyright McGraw-Hill 2009 54
The Six Possible Phase Changes Copyright McGraw-Hill 2009 55
Liquid-Vapor Phase Transition Boiling point the temperature at which the vapor pressure of liquid equals atmospheric pressure Molar heat of vaporization (DH vap ) the amount of heat required to vaporize one mole of a substance at its boiling point usually in kj/mol Dependent on the strength of intermolecular forces Condensation opposite of vaporization Copyright McGraw-Hill 2009 56
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Critical temperature (T c ) the temperature above which a gas cannot be liquified by application of pressure Critical pressure (P c ) the pressure that must be applied to liquefy a gas at T c. Supercritical fluid the fluid that exists above T c and P c. Copyright McGraw-Hill 2009 58
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Solid-Liquid Phase Transition Freezing transformation of liquid to solid Melting (fusion) opposite of freezing Melting point of solid (or freezing point of liquid) temperature at which the solid and liquid phases coexist in equilibrium Dynamic equilibrium in which the forward and reverse processes are occurring at the same rate Molar heat of fusion (DH fus ) energy to melt one mole of a solid usually in kj/mol Copyright McGraw-Hill 2009 60
Typical Heating Curve Copyright McGraw-Hill 2009 61
Solid-Vapor Phase Transition Sublimation process by which molecules go directly from the solid phase to the vapor phase Deposition reverse of sublimation Molar heat of sublimation (DH sub ) energy required to sublime one mole of solid usually in kj/mol iodine Copyright McGraw-Hill 2009 62
Heating Curve for Water Copyright McGraw-Hill 2009 63
Calculate the amount of energy (in kj) required to convert 125 g of ice at 10.0 o C to liquid water at the normal boiling point. Assume that the specific heat of ice is 2.050 J/g o C. Copyright McGraw-Hill 2009 64
Energy to warm ice from 10 o C to 0 o C DT 0.0 o C( 10.0 o C) 10.0 o C 2.050 J o q msdt 125 g 10.0 C o g C 3 kj 2.56310 J 2.563 kj 3 110 J Energy to melt ice at 0 o C 2.56310 3 kj q ndh mol 125 g 18.02 g vap 6.937 mol 6.937 6.01kJ mol mol 4.16910 kj Copyright McGraw-Hill 2009 65 1
Energy to warm water from 0.0 o C to 100.0 o C DT 100.0 o C0.0 o C 100.0 o C q msdt 125 g 4.184 g o C J 100.0 o C 5.23010 4 J 5.23010 4 1kJ J 110 3 J 5.23010 1 kj Total energy required 2.563 kj (4.16910 1 kj) (5.230 10 1 kj) 96.6kJ Copyright McGraw-Hill 2009 66
12.7 Phase Diagrams Phase diagram summarizes the conditions (temperature and pressure) at which a substance exists as a solid, liquid or gas Divided into three regions (solid, liquid, gas) Phase boundary line line separating any two regions Triple point the point at which all three phase boundary lines meet Copyright McGraw-Hill 2009 67
Phase Diagram of CO 2 Copyright McGraw-Hill 2009 68
Heating CO 2 Starting at 100 o c and 1 atm Copyright McGraw-Hill 2009 69
Phase Diagram of H 2 O Copyright McGraw-Hill 2009 70
What is a) the normal* melting point, b) the normal* boiling point and c) the physical state of the substance at 2.0 atm and 110 o C? *normal measured at 1.00 atm Copyright McGraw-Hill 2009 71
normal melting point normal boiling point ~135 o C ~200 o C Copyright McGraw-Hill 2009 72
solid physical state Copyright McGraw-Hill 2009 73
Key Points Intermolecular forces Dipole-dipole interactions Hydrogen bonding (London) dispersion forces Properties of liquids Surface tension Viscosity Vapor pressure Clausius-Clapeyron equation Copyright McGraw-Hill 2009 74
Crystal structure Unit cells Lattice point Packing spheres Coordination number Cubic unit cells Closest Packing Types of crystals Ionic Covalent Molecular Metallic Copyright McGraw-Hill 2009 75
Amorphous solids Phase changes Liquid-vapor transitions Boiling point Heat of vaporization Critical temperature and pressure Solid-liquid transitions Melting point Heat of fusion Phase diagrams Copyright McGraw-Hill 2009 76