2. Chapter 10 Liquids and Solids

3. Topics  Intermolecular forces –Dipole-dipole forces  Hydrogen bonding –London Forces  The liquid state –Surface tension –Capillary action –Viscosity  An introduction to structures and types of solids –X-ray analysis of solids –Types of crystalline solids  Structure and bonding in metals –Bonding metals for metals –Meta alloys  Molecular solids  Ionic solids  Vapor pressure and changes of state  Phase diagrams 10.5 Section is self study

4. Intra- vs. Inter-molecular forces  intramolecular forces –inside molecules (bonding) –hold atoms together into molecule  intermolecular forces – –These are what hold the molecules together in the condensed states. – Forces between molecules –They get weaker as phase changes from S – L – G  When a substance changes state, molecule stays together but intermolecular forces are weakened 10.1 Intermolecular Forces

5. Intermolecular Forces Gases Gases – fill container, random rapid motion, never coming to rest or clumping together Motion is mainly translational Liquids Liquids – fixed volume, flow and assume shape of container, only slightly compressible, stronger forces hold molecules together Motion is mainly translational Motion is mainly translational Solids – fixed volume, definite shape, generally less compressible than liquids, forces hold particles in a fixed shape Motion is mainly vibrational

6. Intermolecular Forces Intermolecular forces are attractive forces between molecules Intramolecular forces hold atoms together in a molecule. Intermolecular vs Intramolecular 41 kJ to vaporize 1 mole of water (inter) 930 kJ to break all O-H bonds in 1 mole of water (intra) Generally, intermolecular forces are much weaker than intramolecular forces. “Measure” of intermolecular force boiling point melting point  H vap  H fus  H sub

7. + - + -  When molecules with dipole moments line up to minimize repulsion and maximize attraction  Very weak compared to covalent and ionic bonds + - + - + - + - + - + - + - + - Attractions Repulsion Dipole – Dipole Foces

8. u Molecules that line up in the presence of a electric field are dipoles. u The opposite ends of the dipole can attract each other so the molecules stay close together. u 1% as strong as covalent bonds u Weaker the covalent bonds with greater distance. u Small role in gases. u Molecules with these forces possess higher melting points and boiling points than nonpolar molecules of comparable molar mass

9. The strengths of i ii intermolecular forces are generally weaker than either ionic or covalent bonds. 16 kJ/mol (to separate molecules) 431 kJ/mol (to break bond) ++ ++ -- -- Polar molecules have dipole-dipole attractions for one another.

10. Types of intermolecular forces (between neutral molecules that posses dipole moment): Dipole-dipole forces: (polar molecules) S O O.. : : : ++ -- -- : -- -- S O O : : ++ dipole-dipole attraction What effect does this attraction have on the boiling point?

11. PolarNonpolar BPMMMoleculeBPMMMolecule -19228CO-19628N2N2 -8834PH 3 -11232SiH 4 -6278AsH 3 -9077GeH 4 97162ICl69160Br 2 Effect of polarity on boiling points Effect of polarity is usually small enough to be obscured by differences in molar mass HCl -85 BP ( o C) HBr -60 HI -30 BP increase although polarity decreases

12. Hydrogen Bonds A hydrogen bond is an intermolecular force in which a hydrogen atom covalently bonded to a nonmetal atom in one molecule is simultaneously attracted to a nonmetal nonmetal atom of a neighboring molecule The strongest hydrogen bonds are formed if the nonmetal smallhighly electronegative atoms are small and highly electronegative – e.g., N, O, F  very strong type of dipole-dipole attraction –because bond is so polar –because atoms are so small

13. Hydrogen bond  Cl(HCl) and S(H 2 S) do not form hydrogen bonding although they have electronegativity similar to N, why? bonding although they have electronegativity similar to N, why? –They are of bigger size to approach the hydrogen atom  Hydrogen bond is 5-10% as strong as the covalent bond

14. Hydrogen bonding is a weak to moderate attractive force that exists between a hydrogen atom covalently bonded to a very small and highly electronegative atom and a lone pair of electrons on another small, electronegative atom ( F, O, or N ).

16. Hydrogen bonding: It is very strong dipole- dipole interaction (bonds involving H-F, H-O, and H-N are most important cases).  + H-F  - ---  + H-F  - Hydrogen bonding

17. Hydrogen bonding between water molecules

18. Water ++ -- ++

20. Hydrogen Bonding Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen Hydrogen bonding between ammonia and water

21. Examples of hydrogen bond The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. IT IS NOT A BOND. A H … B A H … A or A & B are N, O, or F

22. Hydrogen Bonding Effects Solid water is less dense than liquid water due to hydrogen bonding Hydrogen bonding is also the reason for the unusually high boiling point of water

23. The larger the molecule the larger the Van der Waals attraction due to more electrons in the molecule. The stronger the attraction, the higher the boiling point. Boiling Points for Some Non Polar Molecules

24. 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 Molar mass

25. Hydrogen Bonding in other molecules Many organic acids can form dimers due to hydrogen bonding Certain organic molecules can also form an intramolecular hydrogen bond

26. Ethanol shows hydrogen bonding Do these compounds show hydrogen bonding?

28. Hydrogen bonding and solubility  Some compounds containing O, N & F show high solubilities in certain hydrogen containing solvents. high solubilities in certain hydrogen containing solvents.  NH 3 & CH 3 OH dissolves in H 2 O through the formation of H-bonds

29. 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.

30. London Dispersion Forces The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. Fritz London 1900-1954 London forces increase with the size of the molecules.

31. London Dispersion Forces  They exist in every molecular compound  They are significant only for nonpolar molecules and noble gas atoms  They are weak, short-lived  Caused by formation of temporary dipole moments

32. Instantaneous polarization causes instantaneous dipole

33. “Electrons are shifted to overload one side of an atom or molecule”. ++ ++ -- -- attraction

34. London Dispersion Forces -Relatively weak forces that exist among noble gas atoms and nonpolar molecules. (Ar, C 8 H 18 ) -Caused by instantaneous dipole, in which electron distribution becomes asymmetrical. -The ease with which electron “cloud” of an atom can be distorted is called polarizability.

35. Polarizability: the ease with which an atom or molecule can be distorted to have an instantaneous dipole. “squashiness” In general big molecules are more easily polarized than little ones.

36. Intermolecular Forces Polarizability Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted. Polarizability increases with: greater number of electrons more diffuse electron cloud Dispersion forces usually increase with molar mass.

37. 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.  Much, much weaker than other forces.  Also called.  Also called Van der Waal’s forces.

38. Relative Magnitudes of Forces The types of bonding forces vary in their strength as measured by average bond energy. Covalent bonds (400 kcal/mol) Hydrogen bonding (12-16 kcal/mol ) Dipole-dipole interactions (2-0.5 kcal/mol) London forces (less than 1 kcal/mol) Strongest Weakest

39. Which one(s) of the above are most polarizable? Hint: look at the relative sizes.

40. London Forces in Hydrocarbons

41. Practice  which has highest boiling pt? –HF, HCl, or HBr?  Identify the most important intermolecular forces : –BaSO 4 –H2S–H2S–H2S–H2S –Xe –C2H6–C2H6–C2H6–C2H6 –P4–P4–P4–P4 –H2O–H2O–H2O–H2O –CsI ionic dipole-dipole H-bonding London Dispersion

42. Which has stronger intermolecuar forces?  CO 2 or OCS –CO 2 : nonpolar so only LD –OCS: polar so dipole-dipole  PF 3 or PF 5 –PF 3 : polar so dipole-dipole –PF 5 : nonpolar so only LD  SF 2 or SF 6 –SF 2 : polar so dipole-dipole –SF 6 : nonpolar so only LD  SO 3 or SO 2 –SO 3 : nonpolar so LD only –SO 2 : polar so dipol-dipole

43. S O O What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH 4 CH 4 is nonpolar: dispersion forces. SO 2 SO 2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO 2 molecules. 11.2

44. 10.2 The Liquid state Properties of Liquids  Low compressibility  Lack of rigidity  High density compared to gases  Beading (beads up as droplets)  Surface tension  Capillary action  Viscosity  Stronger intermolecular forces cause each of these to increase.

45. Surface tension  The resistance to an increase in its surface area  Polar molecules and liquid metals show high surface show high surface tension tension

46. Surface tension  Molecules in the middle are attracted in all directions. u Molecules at the the top are only pulled inside. u Minimizes surface area.

47. Surface Tension  One water molecule can hydrogen bond to another because of this electrostatic attraction.  Also, hydrogen bonding occurs with other molecules surrounding them on all sides. H H O ++ ++ -- H H O ++ -- ++

48. Surface Tension  A water molecule in the middle of a solution is pulled in all directions.

49. Surface Tension  This is Not true for molecules at the surface.  Molecules at the surface are only pulled down and to each side.  This holds the molecules at the surface together tightly.  This causes surface tension.

50. Surface tension  All liquids have surface tension –water is just higher than most others  How can we decrease surface tension? –Use a surfactant - surface active agent –Also called a wetting agent, like detergent or soap –Interferes with hydrogen bonding

51. Beading  Water drops are rounded, because all molecules on the edge are pulled to the middle- not outward to the air!

52. Adhesive forces Adhesive forces are intermolecular forces between unlike molecules Cohesive forces Cohesive forces are intermolecular forces between like molecules

53. Beading  If a polar substance is placed on a non-polar surface. –There are cohesive, –But no adhesive forces.  And Visa Versa

54. Meniscus is the interface between a liquid and the air above it Adhesion attracted to glass Cohesion attracted to each other

55. Capillary Action   Capillary action results from intermolecular interactions  Liquids spontaneously rise in a narrow tube.  Glass is polar.  It attracts water molecules (adhesive forces)

57.  Glass has polar molecules.  Glass can also hydrogen bond.  This attracts the water molecules.  Some of the pull is up a cylinder.

58.  Water curves up along the side of glass.  This makes the meniscus, as in a graduated cylinder  Plastics are non- wetting; no attraction

59. Meniscus In Glass In Plastic

60. Viscosity Viscosity is a measure of a liquid’s resistance to flow –strong inter molecular forces  highly viscous –large, complex molecules  highly viscous –Cyclohexane has a lower viscosity than hexane. –Because it is a circle- more compact.

61. Model  Can’t see molecules so picture them as-  In motion but attracted to each other  With regions arranged like solids but –with higher disorder. –with fewer holes than a gas. –Highly dynamic, regions changing between types.

62. 10.3 An introduction to structures and types of solids Types of Solids  Crystalline Solids: highly regular three dimensional arrangement of their components [table salt (NaCl)]  Amorphous solids: considerable disorder in their structures (glass: components are frozen in place before solidifying and achieving an ordered arrangement)  The positions of components in a crystalline solid are usually represented by a lattice

63. Crystalline solids Amorphous solids

64. An amorphous solid does not possess a well-defined arrangement and long-range molecular order. A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing Crystalline quartz (SiO 2 ) Non-crystalline quartz glass

65. Representation of Components in a Crystalline Solid  Lattice: A 3-dimensional system that describes the locations of components (atoms, ions, or molecules) that make up the unit cells of a substance.  Unit Cell: The smallest repeating unit in the lattice.  Three  There are Three common types of unit cells: –simple cubic –body-centered cubic –face-centered cubic

66. Crystal Structures - Cubic Simple Face-Centered Body-Centered

67. Unit Cells The simple cubic cell is the simplest unit cell and has structural particles centered only at its corners The body-centered cubic (bcc) structure has an additional structural particle at the center of the cube The face-centered cubic (fcc) structure has an additional structural particle at the center of each face

68. Cubic Unit cells in 3 dimensions At lattice points: Atoms Molecules Ions lattice points The simple cubic cell is the simplest unit cell and has structural particles centered only at its corners

69. Body-Centered Cubic Unit cells in 3 dimensions The body-centered cubic (bcc) structure has an additional structural particle at the center of the cube

70. Face-Centered Cubic The face-centered cubic (fcc) structure has an additional structural particle at the center of each face

71. Sample is powdered X-rays of single wavelength is used Distance between planes of atoms in the crystal are calculated from the angles at which the rays are diffracted using Bragg equation X-Ray analysis of solids

72. Spots from diffracted X-rays Spot from incident beam X-Ray analysis of solids X-Ray diffraction

74. Extra distance traveled by lower ray = BC + CD = n = 2d sin  Reflection of X-rays from two layers of atoms 1 st layer of atoms 2 nd layer of atoms

75. n = 2d sin 

76. Bragg Equation u n 2 = d sin  u d = distance between atoms u n = an integer u = wavelength of the x-rays

77. X rays of wavelength 0.154 nm are diffracted from a crystal at an angle of 14.17 0. Assuming that n = 1, what is the distance (in pm) between layers in the crystal? n = 2d sin  n = 1  = 14.17 0 = 0.154 nm = 154 pm d = n 2sin  = 1 x 154 pm 2 x sin14.17 = ____________

78. When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 408.7 pm. Calculate the density of silver. d = m V V = a 3 = (408.7 pm) 3 = 6.78X10 7 pm 3 ___ atoms/unit cell in a face-centered cubic cell m = 4 Ag atoms 107.9 g mole Ag x 1 mole Ag 6.022 x 10 23 atoms x = 1.79X10 -22 g d = m V 7.17 x 10 -22 g 6.83 x 10 -23 cm 3 = = 1.79X10 -22 g / 6.78X10 7 pm 3 = 2.6X10 -30 g/pm 3

79. Types of crystalline Solids u Ionic solids (ionic compounds) Ions (held by electrostatic attraction) at point in lattice They conduct electric current when dissolved in water u Molecular solids (molecular compounds) Molecules (held by: dispersion and/or dipole-dipole forces) at each point in lattice. Ice is a molecular solid H 2 O u atomic solids (metals, nonmetals, noble gases) Elements (C, B, Si) that are composed of atoms at lattice points. Three types: Metallic– metallic bond Network – strong covalent bonding Group 8A –London Dispersion Forces

80. atomic- networkionicmolecular

82. Types of Crystals Ionic Crystals – Ion-Ion interactions are the strongest Lattice points occupied by cations and anions Held together by electrostatic attraction Hard, brittle, high melting point Poor conductor of heat and electricity CsClZnSCaF 2

83. Types of Crystals Molecular Crystals Lattice points occupied by molecules Held together by intermolecular forces (dipole-dipole, and/or London dispersion forces Soft, low melting point Poor conductor of heat and electricity

84. Types of Crystals Atomic solids/Network – Stronger than IM forces but generally weaker than ion-ion Lattice points occupied by atoms Held together by covalent bonds Hard, high melting point Poor conductor of heat and electricity diamond graphite carbon atoms

85. Types of Crystals Metallic Crystals – Typically weaker than covalent, but can be in the low end of covalent Lattice points occupied by metal atoms Held together by metallic bonds Soft to hard, low to high melting point Good conductors of heat and electricity 11.6 Cross Section of a Metallic Crystal nucleus & inner shell e - mobile “sea” of e -

86. Types of Crystals 11.6

87. 10.4 Structure and Bonding of Metals u Physical properties of metals u Ionization energy  E is small (outer electrons move relatively free); this results in High electrical conductivity High thermal conductivity They are –Ductile: can be drawn oust into wires –Malleable: can be hammered into thin sheets u Electrons act like a glue holding atomic nuclei u Crystals of nonmetals break into small pieces if it is hammered (brittle) u They have luster (reflect light) u They form alloys

88. Metallic Crystals u Can be viewed as containing atoms (spheres) packed together in the closest arrangement possible u The spheres are packed in layers u Closest packing- when each sphere has 12 neighbors 6 on the same plane 3 in the plane above 3 in the plane below

89. Packing in Crystals “Open” packing has larger voids in between particles compared to close-packed crystals

90. Closest packing in Metals

91. Closest packed structures It has (aba) arrangements that occur when the spheres of the third layer occupy positions so that each sphere in the third layer lies directly over a sphere in the first layer Hexagonal closest-packed (hcp) structure

92. Closest packed structures It has (abc) arrangement that occurs when the spheres of the third layer occupy positions that NO sphere lies over one in the first layer Cubic closest-packed (ccp) structure An atom in every fourth layer lies over an atom in the First layer

93. Net number of spheres (atoms) in a unit cell, length of the edge of the cell, density of the closest packed solid Face centered cubic cell u Atoms occupy corners and centers of the faces u Atoms at the corners do not touch each other u Atoms contact is made at the face diagonal u 74% of the space is occupied u Ca, Sr, transition metals u An atom at the center of the face of cube is shared by another cube that touches that face. u Only atom is assigned to a given cell u An atom at the center of the cube is a part of 8-different cubes touching that point. Only cornenr atom belongs to the cell

94. Density of closest packed solid u # of spheres (atoms) per unit cell = u Density of closes packed solid

95. Body centered cubic cell u Atoms contact along a body diagonal u Atoms occupy the corners and one at the center u In a unit cell, 8 atoms occupy the corners plus one in the center u 68% of the space is occupied u Available in Group I elements + Ba

96. Number of atoms assigned to each type of cell u Simple cube u Body centered cube u Face centered cube

97. 1 atom/unit cell (8 x 1/8 = 1) 2 atoms/unit cell (8 x 1/8 + 1 = 2) 4 atoms/unit cell (8 x 1/8 + 6 x 1/2 = 4)

98. Relationship between the atomic the radius and the edge length in different unit cells

99. When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 408.7 pm. Calculate the density of silver. d = m V V = a 3 = (408.7 pm) 3 = __________________ ___ atoms/unit cell in a face-centered cubic cell m = 4 Ag atoms 107.9 g mole Ag x 1 mole Ag 6.022 x 10 23 atoms x = ___________ d = m V 7.17 x 10 -22 g 6.83 x 10 -23 cm 3 = = ___________________ 12.4 Compare Ex 12.3, p.384

100. Bonding models for Metals u Shape of pure metals con be changed but most metals have high melting points u Thus, bonding in most metals is strong and non-directional (although difficult to separate metal atoms, it is easy to move them provided they stay in contact of each other

101. u The highest energy level for most metal atoms does not contain many electrons u These vacant overlapping orbitals allow outer electrons to move freely around the entire metal u Metallic crystal is an array of positive ions (cations) in a sea of roaming valence electrons Electron Sea Model

102. Metallic Bonding: “sea of e - ’s”

103. Bonding models for Metals u These roaming electrons form a sea of electrons around the metal atoms u Malleability and ductility bonding is the same in every direction one layer of atoms can slide past another without friction u Conductivity of heat and electricity from the freedom of electrons (mobile electrons) to move around the atoms

104. Metallic bonding: Molecular orbital model for metals (Band model) u Electrons are assumed to travel around the metal crystal in molecular orbitals formed from the valence atomic orbitals of metal atoms u When many metal atoms interact in a crystal a large number of resulting molecular orbitals become more closely spaced and form a continuum of levels called bands

105. Metallic bonding: Molecular orbital model for metals (Band model) 1s 2s 2p 3s 3p Filled Molecular Orbitals Empty Molecular Orbitals Magnesium Atoms

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

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

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

109. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms This makes metals conductors Malleable because the bonds are flexible.

110. Metal Alloys u An alloy is a mixture of elements and has metallic properties u substitutional alloy host metal atoms are replaced by other metal atoms happens when they have similar sizes u interstitial alloy metal atoms occupy spaces created between host metal atoms happens when metal atoms have large difference in size

111. Examples u Brass substitutional 1/3 of Cu atoms replaced by Zn u Steel interstitial Fe with C atoms in between makes harder and less malleable

112. Metal Alloys  Substitutional Alloy: some metal atoms replaced by others of similar size. brass = Cu/Znbrass = Cu/Zn

113. Metal Alloys (continued)  Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbonsteel = iron + carbon

114. 10.6 Molecular solids. u Molecules occupy the corners of the lattices. u Common examples: ice, dry CO 2, S 8, P 4, I 2 u Different molecules have different forces between them (H-bonds, or dipole-dipole or London forces, or a combination of all these forces) u These forces depend on the size of the molecule. u They also depend on the strength and nature of dipole moments.

115. Molecular solids with nonpolar molecules (without dipoles): H 2, CCl 4 u Most are gases at 25ºC. u The only forces are London Dispersion Forces. u These depend on size of atom. u Large molecules (such as I 2 ) can be solids even without dipoles.

116. Molecular solids with polar molecules (with dipoles): HCl, NH 3 u Dipole-dipole forces are generally stronger than L.D.F. u Hydrogen bonding is stronger than Dipole- dipole forces. u No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds. u Stronger forces lead to higher melting and freezing points.

117. Water is special u Each molecule has two polar O-H bonds. u Each molecule has two lone pair on its oxygen. u Each oxygen can interact with 4 hydrogen atoms. H O H  

118. Water is special H O H   H O H   H O H   u This gives water an especially high melting and boiling point.

119. Examples of molecular solids

120. 10.7 Ionic Solids u They comprise the extremes in dipole dipole forces- (ionic forces) u Atoms are actually held together by electrostatic attractions of opposite charges. u They possess huge melting and boiling points. u Atoms are locked in lattice so they are hard and brittle. u Every electron is accounted for so they are poor conductors-good insulators.

121. u In most of binary ionic compounds, larger ions are arranged in closest packing arrangement, hexagonal (hcp) or cubic (ccp) closest packing u smaller ions fit in the holes created by the larger ions

122. Closest Packing Holes The hole is formed by 3 spheres in the Same layer The hole formed when a sphere occupies a dimple formed by three spheres in an adjacent layer The holes formed between two sets of three spheres in adjoining layers of closest packed structure

123. Closest packing holes Tetrahedral holes : are located above a sphere in the bottom layer Octahedral holes : are located above a void in the bottom layer

124. Examples u Trigonal holes are so small that they are never occupied in binary ionic compounds u The type of the hole whether tetrahedral or octahedral depends mainly on Relative sizes of cations and anions u In ZnS S 2-, ions are arranged in ccp with the smaller Zn 2+ ions in the tetrahedral holes u In NaCl, ions are arranged in ccp with Na+ ions in the octahedral holes.

126. Two Examples

127. Ionic Crystal Structures Smaller cations can fill the voids between the larger anions Tetrahedral hole filling occurs when the radii ratio is: 0.225 < r c /r a < 0.414 Octahedral hole filling occurs when the radii ratio is: 0.414 < r c /r a < 0.732 The arrangement is cubic if r c /r a > 0.732

129. 10. 7 Vapor Pressure and changes of state u vapor- gas phase above a substance that exists as solid or liquid at 25°C and 1 atm. u Vaporization or Evaporation - change from liquid to gas at or below the boiling point. (Endothermic process) u Condensation is the change of a gas to a liquid (Exothermic process)

130. u Heat or enthalpy of vaporization, ∆H vap : Energy required to vaporize 1 mole of a liquid at 1 atm u water has a large ∆H vap (40.7 kJ/mol), (because of hydrogen bonding)

131. Molar Heats of Vaporization for Selected Liquids

132. Vapor pressure u Initially, liquid in a closed container decreases as molecules enter gaseous phase u When equilibrium is reached, no more net change occurs u Rate of condensation and rate of vaporization become equal  Molecules still are changing phase but no net change (Dynamic equilibrium) Gas liquid independent of volume Vapor pressure is independent of volume of container as long as some liquid is present (liquid-vapor equilibrium)

133. Evaporation and condensation H 2 O (l) H 2 O (g) Rate of condensation Rate of evaporation = Dynamic Equilibrium

134. Equilibrium vapor pressure or vapor pressure It is the pressure exerted by the vapor when it is in dynamic equilibrium with a liquid at a constant temperature.

135. Vapor Pressure’s measurement u Liquid can be injected under inverted tube u part of the liquid evaporates to the top of tube u P vap can be determined u by height of Hg P atm = P vap + P Hg u The pressure of the vapor phase at equilibrium: P vap u can be measured when using a simple barometer

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

137. 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.

138. 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

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

140. Vapor pressure and nature of liquids u Vapor pressure depends upon the nature of the liquid u Liquids with high vapor P (volatile liquids) liquids with a high vapor pressure evaporate quickly weak intermolecular forces u Liquids with low vapor P Strong IMFs, London dispersion forces (large molar masses) or dipole-dipole forces

141. Vapor pressure and temperature u Vapor pressure increases with T u More molecules have enough KE to overcome IMFs

142. VAPOR PRESSURE CURVES A liquid boils when its vapor pressure =‘s the external pressure.

143. Temperature Effect Kinetic energy # of molecules T1T1 Energy needed to overcome intermolecular forces in iquid

144. Kinetic energy # of molecules T1T1 T1T1 T2T2 u At higher temperature more molecules have enough energy - higher vapor pressure. Energy needed to overcome intermolecular forces in liquid

145. Molar heat of vaporization (  H vap ) is the energy required to vaporize 1 mole of a liquid. ln P = -  H vap RT + C Clausius-Clapeyron Equation P = (equilibrium) vapor pressure T = temperature (K) R = gas constant (8.314 J/K mol) Vapor pressure and Temperature

146. Mathematical relationship u ln is the natural logarithm ln = Log base e e = Euler’s number an irrational number like  u  H vap is the heat of vaporization in J/mol

147. u R = 8.3145 J/K mol. Mathematical relationship

148. Vapor Pressure for solids u Solids also have vapor pressure u Sublimination solid gas directly Example: dry ice: CO 2 u heat of fusion (∆H fus ) enthalpy of fusion enthalpy change at melting point

149. Heating Curve u as energy is added, it is used to increase the T u when it reaches melting point, the energy added is used to change molecules from (s) to (l)  plot of T vs. time where energy is added at constant rate

150. Changes of state u What happens when a solid is heated? u The graph of temperature versus heat applied is called a heating curve. u The temperature a solid turns to a liquid is the melting point. u The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion  H fus

151. Heating Curve for Water Ice Water and Ice Water Water and Steam Steam mp bp Temp Time (Heat added)

152. Heating Curve for Water Heat of fusion Heat of vaporization Slope is Heat Capacity Time (Heat added) Temp  H vap =2260 J/g  Hfus=334 J/g

154. Heating curve for 1 gram of water

155.  H fus =334 J/g Specific Heat of ice = 2.09 J/gK Specific Heat of water = 4.184 J/gK  H vap =2260 J/g Specific Ht. Steam = 1.84 J/gK

156. Calculate the enthalpy change upon converting 1 mole of water from ice at -12 o C to steam at 115 o C. solid -12 o C solid 0 o C liquid 0oC0oC 100 o C gas 100 o C gas 115 o C  H 1 +  H 2 +  H 3 +  H 4 +  H 5 =  H total Sp. Ht. +  H fusion + Sp. Ht. +  H Vaporization + Sp. Ht. =  H total Specific Heat of ice = 2.09 J/gK  H fus =334 J/g Specific Heat of water = 4.184 J/gK Specific Ht. Steam = 1.84 J/gK  H vap =2260 J/g

157. Calculate the enthalpy change upon converting 1 mole of water from ice at -12 o C to steam at 115 o C. solid -12 o C solid 0 o C liquid 0 o C liquid 100 o C gas 100 o c gas 115 o c  H 1 +  H 2 +  H 3 +  H 4 +  H 5 =  H total Sp. Ht. +  H fusion + Sp. Ht. +  H Vaporization + Sp. Ht. =  H total Specific Heat of ice = 2.09 J/gK  H fus =334 J/g Specific Heat of water = 4.184 J/gK Specific Ht. Steam = 1.84 J/gK

158. Normal melting Point u Melting point is determined by the vapor pressure of the solid and the liquid. u Melting point is the temp at which the vapor pressure of the solid = vapor pressure of the liquid where the total pressure is 1 atm.

159. Solid Water Liquid Water Water Vapor Vapor Apparatus that allows solid and liquid water to interact only through the vapor state

160. Solid Water Liquid Water Water Vapor Vapor u A temp at which the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium.

161. Solid Water Liquid Water Water Vapor Vapor u While the molecules of water condense to a liquid to achieve equilibrium.

162. u This can only happen if the temperature is above the melting point since solid is turning to liquid. Solid Water Liquid Water Water Vapor Vapor

163. u A temperature at which the vapor pressure of the solid is less than that of the liquid, the liquid will release molecules to achieve equilibrium. Solid Water Liquid Water Water Vapor Vapor

164. Solid Water Liquid Water Water Vapor Vapor u While the molecules of water condense to a solid.

165. u The temperature must be below the melting point since the liquid is turning to a solid. Solid Water Liquid Water Water Vapor Vapor

166. u Temperature at which the vapor pressure of the solid and liquid are equal, the solid and liquid are vaporizing and condensing at the same rate. This is the Melting (freezing) point (Temp at which solid and vapor can coexist) Solid Water Liquid Water Water Vapor Vapor

167. Normal Boiling Point u Temp when vapor pressure inside the bubbles equals 1 atm u If P vap < 1 atm, no bubbles can form, there is too much pressure on surface

168. Normal Boiling Point u Boiling occurs when the vapor pressure of liquid becomes equal to the external pressure. u Normal boiling point is the temperature at which the vapor pressure of a liquid is exactly 1 atm pressure. u Super heating - Heating above the boiling point. u Supercooling - Cooling below the freezing point.

169. Exceptions u supercooling Material can stay liquid below melting point because doesn’t achieve level of organization needed to make solid u superheating when heated too quickly, liquid can be raised above boiling point causes “bumping” Changes of state do not always occur exactly at bp or bp

170. 10.9 Phase Diagrams u A plot Representing phases of a substance in a closed system (no material escapes into the surroundings and no air is present) as a function of temperature and pressure.

171. Phase changes Gas and liquid are indistinguishable. Critical temperature and critical pressure (all 3 phases exists here)

172. Phase Diagrams fusion curve triple point critical point vapor pressure curve sublimation curve

173. Critical Point  where if the T is increased, vapor can ’ t be liquefied no matter what P is applied  at the end of liquid/gas line  after this point, only one fluid phase exists that is neither gas nor liquid  called supercritical fluid

174.  Critical temperature Temperature above which the vapor can Temperature above which the vapor can not be liquefied. not be liquefied.  Critical pressure pressure required to liquefy gas AT the pressure required to liquefy gas AT the critical temperature. critical temperature.  Critical point critical temperature and pressure (for critical temperature and pressure (for water, T c = 374°C and 218 atm). water, T c = 374°C and 218 atm).

175. Phase diagram for Water Normal mp Normal bp Critical temp Water Expands upon freezing -ve slope of S/L boundary line means that mp of ice decreases as the external P increases

176. Phase Diagram for H 2 O

177.  Most substances have a positive have a positive slope of slope of solid/liquid line solid/liquid line  because solid is usually more dense than liquid  water has a negative slope

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

179. Solid Liquid Gas Triple Point Critical Point Temperature Pressure

180. Solid Liquid Gas u This is the phase diagram for water. u The density of liquid water is higer than solid water. Temperature Pressure

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