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Adv chem chapt 10

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  • 1. Liquids and Solids Advanced Chemistry Chapter 10
  • 2. Sections
    • Intermolecular Forces
    • Liquid state
    • Solid Structures
    • Metal Structures
    • Carbon and Silicon Networks
  • 3. Sections
    • Molecular Solids
    • Ionic Solids
    • Vapor pressure and State Change
    • Phase Diagrams
  • 4. States of Matter
    • When considering the three states of matter, properties of gases are strikingly different than solids and liquids. Liquids and solids share many similar characteristics
    • compressibility
    • density
    • intermolecular forces
  • 5. States of Matter
    • H 2 O(s)  H 2 O(l) ΔH ° fus = 6.02 kj/mol
    • H 2 O(l)  H 2 O(g) ΔH° vap = 40.7 kj/mol
    • Water densities:
      • 25°C and 1atm .99707g/cm 3
      • 25°C and 1065 atm 1.046g/cm 3
      • 400°C and 1atm 3.26x10-4 g/cm 3
      • 400°C and 242 atm .157g/cm 3
  • 6. Intermolecular Forces 10.1
  • 7. Intermolecular Forces
    • Electrons shared within the molecule are called intramolecular bonding.
    • In the condensed states of matter the attraction between molecules are called intermolecular forces.
  • 8. Intermolecular Forces
      • It is important to realize that when a molecule changes state, the molecule stays intact. The changes in state are due to the change in forces surrounding the molecule not from changes within the molecule.
      • 40.7kj needed to vaporize water
      • 934kj to break the O-H bond
  • 9. Dipole–Dipole Forces
    • Dipole-dipole forces occur when polar molecule (molecules with dipole moments) electrostatically attract each other by lining up the positive and negative ends of the dipoles.
    • Dipole-dipole forces are about 1% as strong as a covalent or ionic bond and rapidly become weaker when distances between the dipoles increases. The distances in a gas make these attractions relatively unimportant
  • 10. Dipole-Dipole Forces
    • In a condensed state, molecules line up dipoles to minimize repulsions and maximize attractions.
  • 11. Dipole-Dipole Forces
    • Some dipole-dipole forces are unusually strong. These usually form between H and another very electronegative atom.
    • These are stronger due to the high polarity of the bond and the closeness of the dipoles between the atoms.
      • These strong attractions have a strong impact on melting points and boiling points.
  • 12. Boiling Points of Covalent Hydrides
  • 13. Hydrogen bonds
    • Hydrogen bonds are the strongest in the smallest and lightest of the covalent molecules. This is primarily due to two factors:
    • large difference in electronegativities
    • small size of the atoms allows for close dipole interactions.
  • 14. Hydrogen bonds
  • 15. Hydrogen Bonds and Organics
    • Methanol (CH 3 OH) and ethanol (CH 3 CH 2 OH) have much higher boiling points than would be expected from their molar masses because of the O-H bonds that produce hydrogen bonding.
  • 16. London Dispersion Forces
    • Even without dipoles, molecules exert forces on each other.
    • The forces that exist among noble gas atoms and nonpolar molecules are called London dispersion forces.
  • 17. London Dispersion Forces
    • Usually it is assumed that electron dispersion is uniform throughout the molecule, but this is not always the case.
    • Since the movements of the electrons around the nucleus are somewhat random, a momentary nonsymmetrical electron distribution can develop that creates a temporary dipolar arrangement of charge.
  • 18. London Dispersion Forces
    • This temporary change in polarity can, in turn, temporary change the distribution of the neighboring molecule.
    • This phenomenon leads to an inter-atomic attraction that is relatively weak and short-lived, but can be significant in larger atoms at lower temperatures.
      • larger atoms have more electrons and increases the probability of a temporary dipole.
  • 19. London Dispersion Forces
  • 20. London Dispersion Forces
    • Polarizability is the ease at which an electron cloud can be distorted into a temporary dipole.
    • large atoms have a larger polarizability than smaller atoms
    • This also applies to molecules like H 2 , CH 4 , CCl 4 and CO 2 ; smaller molecules, but nonpolar.
  • 21. The Liquid State 10.2
  • 22. Liquid Characteristics
    • lack of rigidity
    • low compressibility
    • high density
    • rounded droplets
    • capillary action
    • viscosity
  • 23. Rounded Droplets
      • Occur due to the intermolecular forces of the liquid. The liquid molecules are subject to attraction from the side and from below, so liquid tends to form a shape with the minimum surface area – sphere.
      • The resistance of a liquid to increase surface area is from the energy that it takes to overcome intermolecular forces. This resistance is called surface tension .
  • 24. Rounded Droplets
      • Molecules that are polar and have stronger intermolecular forces have stronger surface tensions.
  • 25. Surface Tension
  • 26. Capillary Action
    • Capillary action is the spontaneous rising of a liquid in a narrow tube. This action is due to two forces
      • cohesive forces- the intermolecular forces among the molecules.
      • adhesive forces – the attractive forces between the liquid and the container.
  • 27. Adhesive forces
    • Adhesive forces happen when bonds within the container have polar bonds
        • For example: glass has O atoms that carry a partial negative charge that attracts the partial positive charge of the hydrogen in water. This balance between the strong cohesive forces and the strong adhesive forces produce a meniscus.
  • 28. Adhesive forces
      • A nonpolar substance, such as mercury, has a convex meniscus because the cohesive forces are stronger than the adhesive forces.
  • 29. Meniscus: Water vs. Mercury
  • 30. Viscosity
    • Viscosity is a fluids resistance to flow.
    • liquids with strong cohesive forces tend to be highly viscous.
      • Example: glycerol is highly viscous because of its ability to create hydrogen bonds.
  • 31. Viscosity
    • Molecular complexity also can affect viscosity because they can become entangled in each other.
      • Example: Gasoline has carbon chains from 3-8C long and is nonviscous. Grease is 20-25C long and is very viscous.
  • 32. Introduction to Structures and Types of Solids 10.3
  • 33. Types of Solids
    • Crystalline solids
    • Amorphous solids
  • 34. Crystalline Solids
    • Crystalline solids have a regular arrangement of components at a microscopic level and produce beautiful, characteristic shapes of crystals:
  • 35. Crystalline Solids
    • The positions of components are usually represented by a lattice.
    • lattice is a three dimensional system of units repeating in a pattern. The smallest repeating unit of the lattice is called the unit cell .
  • 36. Three types of Crystalline Solids
  • 37. Amorphous Solids
    • Amorphous solids have considerable disorder in their structures.
    • Example: Common glass looks like a solution frozen in place. It has a rigid shape but a great deal of disorder within its structure.
  • 38. X-ray Analysis of Solids
    • The structures of crystalline solids are commonly determine by X-ray diffraction.
    • This type of diffraction occurs when beams of light are scattered as they go through spaces between substances. Light scatters when the size of the spaces are similar to the wavelength of light.
  • 39. X-ray Analysis of Solids
  • 40. X-ray Analysis of Solids
    • A single wavelength is directed at the crystal and a diffraction pattern is obtained. The diffraction pattern is a series of light and dark areas on a photographic plate from constructive and destructive interference from waves of light.
    • The diffraction pattern can then be used to determine the interatomic spacings.
  • 41. X-ray Analysis of Solids
    • A diffractometer is a computer-controlled instrument used for carrying out the X-ray analysis of crystals
      • It rotates the crystal with respect to the X-ray beam and collects the data produced by the scattering. The techniques have been refined to the point that very complex structures can be determined, such as large biological enzymes.
  • 42. X-ray Analysis of Solids
    • The Bragg equation combines trigonometry and physics to determine the atomic spaces between crystals:
    • nλ = 2d sin θ
    • d is the distance between atoms and θ is the angle of incidence and reflection of the light. n is an integer, most commonly 1. (n is usually given)
  • 43. X-ray Analysis of Solids
  • 44. Example Problem
    • X-rays of wavelength 1.54 Â were used to analyze an aluminum crystal. A reflection was produced at θ = 19.3°. Assuming n=1, calculate the distance d between the planes of atoms producing this reflection
    • 2.33 Á
  • 45. Types of Solids
    • Ionic solids
      • ionic solids are made of ions
    • Molecular solids
      • Molecular solids have small units of covalently bonded molecules.
    • Atomic solids
      • Atomic solids are made of elements such as carbon (graphite, diamond and the fullerenes), boron, silicon, and all metals.
  • 46. Types of Solids
  • 47. Atomic Solids
    • Atomic solids are broken down into subgroups depending on the bond that exists in the solid:
    • Metallic solids
      • Has delocalized nondirectional covalent bonding.
    • Network solids
      • atoms bond with strong directional covalent bonding that lead to giant molecules and networks
  • 48. Atomic Solids
    • Group 8A solids
      • noble gases are attracted to each other with London dispersion forces.
  • 49. Classification of Solids
  • 50. Structure and Bonding in Metals 10.4
  • 51. Metal Characteristics
    • Most of the properties that we see in metals is due to the nondirectional covalent bonding found in metal crystals.
    • High thermal conductivity
    • Electrical conductivity
    • Malleability
    • Ductility
  • 52. Metallic Crystals
    • Metallic crystals can be pictured as containing spherical atoms packed together that can be bonded to each other equally in all directions.
    • This arrangement is called closest packing.
  • 53. Closest Packing
    • The spheres pack in layers. Each sphere is surrounded by six others. These layers do not lie directly over those in the first layer, instead they fill the indentations of the layer below. The third layer is in the same position as the first. This is called aba arrangement.
  • 54. Closest Packing
    • The aba arrangement has the hexagonal unit cell and the resulting structure is called the hexagonal closest packed (hcp) structure.
    • The abc arrangement has a face-centered cubic unit cell and the resulting structure is called the cubic closest packed (ccp) structure . This has a repeating vertical placement every fourth layer.
  • 55. Closest Packing
  • 56. Closest Packing: Hexagonal
  • 57. Closest Packed: Cubic
  • 58. Closest Packing
    • Knowing the net number of atoms in a particular unit cell is important for many applications involving solids.
  • 59. Closest Packing
    • Example: A face centered cube (unit cell) is defined by the centers of the spheres on the cube’s corners. Therefore 8 cubes share a given corner sphere, so 1/8 of this sphere lies inside the unit cell. (8 corners x 1/8 sphere = 1sphere). The sphere at the center of each face is shared by two cubes. (6 faces x ½ sphere = 3 spheres). The total number of spheres for a face centered cube is 4.
  • 60. Closest Packing
    • Face – Centered Cubic Unit Cell
  • 61. Cubic Substances
    • Metals that form cubic closest packed solids are:
    • aluminum
    • iron
    • copper
    • cobalt
    • nickel
  • 62. Hexagonal Substances
    • Metals that form hexagonal closest packed solids are:
    • magnesium
    • zinc
  • 63. Other Metal Solids
    • Calcium and certain other metals can crystallize in either cubic or hexagonal solids.
    • Some metals, including many alkali metals, have structures that are characterized by a body-centered cubic (bcc) unit cell. In this structure, each sphere has 8 neighbors.
  • 64. Example Problem
    • Silver crystallizes in a cubic closest packed structure. The radius of a silver atom is 144pm. Calculate the density of solid silver?
    • 10.6 g/cm 3
  • 65. Bonding Models for Metals
    • In order to determine bonding for metals, one must account for the typical properties: durable, high melting point, malleable, ductile, and efficient in uniform conduction of heat and electricity in all directions.
    • These characteristics indicate that the bonds are strong and nondirectional. In other words, it is not easy to separate metal atoms but easy to move them.
  • 66. Electron Sea Model
    • Metal cations ‘swim’ in a sea of valence electrons that are mobile and shared.
    • This accounts for conduction and malleability and ductility.
  • 67. Electron Sea Model
  • 68. Band Model (MO Model)
    • In this model, the electrons are assumed to travel around the metal crystal in molecular orbitals formed from the valence atomic orbitals of the metal atoms.
    • When metals atoms interact, the large number of resulting molecular orbitals become more closely spaced and finally form a virtual continuum of levels, called bands.
  • 69. Band Model (MO Model)
  • 70. Band Model
    • The electrons in partially filled MO’s are mobile. These conduction electrons are free to travel throughout the metal crystal. The MO occupied by these conducting electrons are called conduction bands.
  • 71. Band Model
  • 72. Metal Alloys
    • An alloy is best defined as a substance that contains a mixture of elements and has metallic properties. There are two types of alloys:
    • Substitutional alloy – some of the host metal atoms are replaced by other metal atoms of similar size.
    • Interstitial alloy – is formed when some of the interstices (holes) in the closest packed lattice are occupied by smaller atoms.
  • 73. Metal Alloys
  • 74. Substitutional Alloys
    • Example: brass: 1/3 of copper metal atoms are replaced by zinc atoms
    • Sterling silver- 93% silver and 7% copper.
    • Pewter- 85% tin, 7% copper, 6% bismuth and 2% antimony.
    • Plumbers solder – 95% tin and 5% antimony
  • 75. Interstitial Alloy
    • Example: Steel contains carbon atoms in the holes of an iron crystal. The presence of the interstitial atoms changes the properties of the host metal. Iron is relatively soft, ductile and malleable, but when carbon (which forms directional bonds), is introduced into the crystal, it makes the iron bonds stronger and less ductile.
  • 76. Interstitial Alloy
    • The amount of carbon directly affects the properties of steel:
    • Mild steels- contains less than .2% carbon: nails chains and cables.
    • Medium steels- contain .2-.6% carbon: rails and structural steel
    • High-carbon steel – .6-1.5% carbon: springs, tools and cutlery.
  • 77. Mixed Alloys
    • Some steels contain elements in addition to iron and carbon. These are called alloy steels and are viewed as being mixed interstitial and substitutional alloys.
    • Bicycle frames are usually constructed from a wide variety of alloy steels.
  • 78. Carbon and Silicon Network Atomic Solids 10.5
  • 79. Network Solids
    • Many atomic solids contain strong directional covalent bonds to form a solid that might be viewed as a “giant molecule.” These materials are typically brittle and do not efficiently conduct heat and electricity. Two examples of these network solids are carbon and silicon.
  • 80. Carbon
    • Two most common forms of carbon are diamond and graphite. They are typical network solids.
    • Diamond is the hardest naturally occurring substance.
    • Graphite is slippery, black and a conductor.
  • 81. Diamond
    • Each carbon is surrounded by a tetrahedral arrangement of other carbon atoms to form a large molecule. Diamond is an insulator not a conductor. Each carbon is sp 3 hybridized with localized bonding and therefore does not conduct.
    • Diamonds are often used for industrial cutting implements.
    • The application of 150,000 atm at 2800°C can break graphite bonds and rearrangement into a diamond structure.
  • 82. Graphite
    • The structure of graphite is based on layers of carbon atoms arranged in fused 6 C rings. The unhybridized p orbitals allow for delocalized electrons and therefore conductivity.
    • Graphite is used as a industrial lubricant. Because graphite has strong bonds within the layers and weak bonding between the layer, the layers slide past one another readily.
  • 83. Carbon
  • 84. Carbon: Graphite layers
  • 85. Silicon
    • Silicon is an important constituent of the compounds that make up the earth’s crust. Silicon is to geology what carbon is to biology and is fundamental to most rocks, sands and soils found in the earth’s crust.
    • Carbon compounds typically have long strings of C-C bonds
    • Silicon compounds typically involve chains of Si-O bonds.
  • 86. Silica
    • The fundamental silicon-oxygen compound is silica, which has the empirical formula SiO 2 . The structure that is formed is based on a network of SiO 4 tetrahedra with shared oxygen atoms rather than smaller SiO 2 units.
  • 87. Silica
    • When silica is heated above its melting point (1600°c) and cooled rapidly, an amorphous solid called glass results. Glass has a lot of disorder as opposed to the crystalline nature of quartz. Glass, also homogeneous, more closely resembles a very viscous solution than it does a crystalline solid.
  • 88. Glass
    • The properties of glass can vary greatly depending on the additives.
    • Common glass results when substances like Na 2 CO 3 are added to the silica melt.
    • B 2 O 3 produce borosilicate glass which does not expand and contract during large temperature changes. (Pyrex)
    • K 2 O produces especially hard glass that can be ground into shapes for lenses and contacts.
  • 89. Glass
  • 90. Silicates
    • Compounds closely related to silica and found in most rocks, soils and clays are the silicates. Like silica, the silicates are based on interconnected SiO 4 tetrahedra, but instead of a O/Si ratio of 2:1, the ratio is typically higher. This higher ratio tends to make silicon-oxygen anions.
  • 91. Silicates
  • 92. Silicates
    • Ceramics are typically made from clays (which contain silicates) and hardened by firing at high temperatures. They tend to be strong, brittle and heat and chemical resistant.
    • Ceramic is heterogeneous and contain two phases: minute crystals of silicates that are suspended in a glassy cement.
  • 93. Clays
    • Clay comes from the weathering of feldspar, an Aluminosilicate (Na 2 O/K 2 O  Al 2 O 3  6SiO 2 ). This weathering produces kaolinite, that consists of tiny thin platelets of Al 2 Si 2 O 5 (OH) 4. When dry these platelets cling together and lock into place; when wet they can slide over one another. During firing, these platelets bind and form a glass.
  • 94. Ceramics
    • Ceramics constitute one of the most important classes of ‘high-tech” materials. Their stability at high temperatures and resistance to corrosion, make them an obvious choice for constructing jet and car engines.
    • Organoceramics are taking form by the addition of organic polymers to ceramics. This reduces some of the brittle nature of ceramics and allows them to be used for things such as flexible superconducting wire, microelectronic devices, prosthetic devices and artificial bones.
  • 95. Semiconductors
    • Elemental silicon has the same structure as diamond. The structure is different in that the energy gap between filled and empty MO’s is not as large and electrons can delocalize and make silicon a semi-conductor. At higher temperatures, more electrons get excited in the conduction bands and the conductivity of silicon increases.
  • 96. N-type Semiconductor
    • When small fraction of silicon atoms are replaced by arsenic atoms (one more valence electron), extra electrons become available for conduction and produce an n-type semi-conductor . These can conduct an electric current.
  • 97. P-type Semiconductor
    • When small fraction of silicon atoms are replaced by boron atoms (one less valence electron), an electron ‘vacancy’ is made. As electrons move, the fill the ‘hole’ and make a new one. This movement of electrons can therefore carry a current. This type of conductor (less electrons) is called a p-type semiconductor .
  • 98. Energy Level Diagrams for N-type and P-type Semiconductors.
  • 99. P-N Junction
    • Most important applications of semiconductors involve connection of a p-type and an n-type to form a p-n junction.
    • The red dots represent excess electrons in the n-type semiconductor and the white circles represent holes (electron vacancies.
  • 100. P-N Junction
    • At the junction a small number of electrons migrate from the n-type region into the p-type region. The effect of these migrations is to place a negative charge on the p-type region and a positive charge on the n-type region.
  • 101. P-N Junction
    • This charge buildup, called the contact potential or junction potential, prevents further migration of electrons. This transfer of electrons is therefore a ‘one-way’ transfer and under an external battery source will allow flow of electrons from the n to the p type regions.
  • 102. P-N Junction
    • When current is opposed it is said to be under reverse bias . When current flows easily, the junction is said to be under forward bias .
    • A p-n junction is a good rectifier , a device that produces a pulsating direct current from an alternating current.
  • 103. P-N Junction
    • When placed in a circuit where the current is constantly reversing, a p-n junction only transmits current under forward bias. Radios, computers and other electronic devices all use this rectifiers. This p-n junction revolutionized electronics.
  • 104. Molecular Solids 10.6
  • 105. Network Solids
    • Sometimes network solids can be considered to be one giant molecule or have large discrete molecular units in a lattice-type position. These molecules have strong bonds within the molecules but relatively weak between the molecules.
    • Common examples: Ice, dry ice (solid carbon dioxide), Sulfur (S 8 ), Phosphorus (P 4 )
  • 106. Network Solids
  • 107. Network Solids
    • When molecules do have dipole moments, their intermolecular forces are significantly greater, especially when hydrogen bonding is possible.
    • Water not only has polar bonds, a dipole moment, has hydrogen bonds, but it also can have a total of four hydrogens associated with every oxygen atom.
  • 108. Ionic Solids 10.7
  • 109. Ionic Solids
    • Ionic solids are stable, high melting substances held together by the strong electrostatic forces that exist between oppositely charged ions.
  • 110. Ionic Solids
    • Most binary ionic solids can be explained by the closest packing of spheres. Typically the larger ions, usually anions, are packed in one of the closest packed arrangements (hcp and ccp).
    • The smaller cations fit into the holes among the closest packed anions. This packing maximizes the electrostatic attractions among oppositely charged ions and minimizes the repulsion of like charges.
  • 111. Ionic Solids
    • There are three types of holes in closest packed structures:
    • Trigonal holes are formed by three spheres in the same layer
  • 112. Ionic Solids
    • There are three types of holes in closest packed structures:
    • Tetrahedral holes are formed when a sphere sits in the dimple of three spheres in an adjacent layer.
  • 113. Ionic Solids
    • There are three types of holes in closest packed structures:
    • Octahedral holes are formed between two sets of three spheres in adjoining layers of the closest packed structures.
  • 114. Ionic Solids
    • The holes increase in size in the order:
    • trigonal < tetrahedral < octahedral
    • The trigonal holes are so small that they are never occupied in binary ionic compounds. Tetrahedral and octahedral holes are occupied if the relative sizes of the ions allow.
  • 115. Ionic Solids
    • Example: Zinc Sulfide (ZnS) creates a ccp structure. The Zn 2+ has a radius of 70pm and the S 2- ion has an ionic radius of 180pm. There are 4 spheres (atoms/anions) in a face-centered cubic unit cell and 8 tetrahedral holes. So only half of the holes in the ccp unit are filled with cations.
  • 116. Ionic Solids
    • Example: Sodium chloride can be described in terms of a ccp structure. Na + resides in octahedral holes. The locations of the octahedral holes in the face-centered cubic unit is marked by X. The number of spheres (anions) in the structure is the same number of octahedral holes. Since NaCl is a 1:1 binary compound. All octahedral holes are used.
  • 117. Example Problem
    • Determine the net number of Na + and Cl - ions in the sodium chloride unit cell.
    • 4 Na + and 4 Cl -
  • 118. Example Problem
    • Classify each of the following substances according to the type of solid it forms:
    • gold
    • carbon dioxide
    • lithium fluoride
    • krypton
    • metallic
    • molecular
    • ionic
    • 8a
  • 119. Types and Properties of Solids
  • 120. Vapor Pressure and Changes of State 10.8
  • 121. Vaporization
    • Vaporization, or evaporation, is the process of liquid molecules escaping the liquid’s surface and forming a gas.
    • Vaporization is endothermic because energy is required to overcome the relatively strong intermolecular forces in the liquid.
      • Water has strong intermolecular forces and this increases the energy required to vaporize. Also making it a great coolant.
  • 122. Vaporization
    • The energy required to vaporize 1 mole of liquid at 1 atm is called the heat of vaporization or the enthalpy of vaporization .
      • The symbol for this is ΔH vap.
  • 123. Vapor Pressure
    • Condensation is the process by which vapor molecules re-form a liquid.
    • The evaporation process occurs at a constant rate at a given temperature, and once an equilibrium has been reached, the rate of condensation will equal the rate of evaporation.
  • 124. Vapor Pressure
    • Molecules in a given system are constantly escaping from and entering the liquid at high rate. However, there is not net change because the two opposite processes just balance each other.
  • 125. Rates of Condensation and Evaporation.
  • 126. Vapor Pressure
    • The pressure of the vapor present at equilibrium is called the equilibrium vapor pressure, or more commonly, the vapor pressure of the liquid.
    • A simple barometer can measure the vapor pressure of a liquid.
  • 127. Vapor Pressure
    • Liquid is injected at the bottom of the tube of mercury and floats to the surface. A portion of the liquid evaporates at the top of the column, producing a vapor whose pressure pushes some mercury out of the tube.
  • 128. Vapor Pressure
    • When the system reaches equilibrium, the vapor pressure can be determined from the change in the height of the mercury column
  • 129. Vapor Pressure
    • P atmosphere = P vapor + P Hg column
    • P vapor = P atmosphere - P hg column
  • 130. Vapor Pressure
    • The vapor pressure of liquids vary widely. Liquids with high vapor pressures are said to be volatile. They evaporate rapidly in an open dish.
  • 131. Vapor Pressure
    • The vapor pressure of a liquid is principally determined by the size of the intermolecular forces in the liquid.
      • Liquids with strong molecular forces have relatively low vapor pressures because it takes so much energy for the molecules to escape.
      • In general, substances with large molar masses have relatively low vapor pressures due of large dispersion forces.
  • 132. Vapor Pressure
    • Vapor pressure increases significantly with temperature.
      • In order to break intermolecular forces, a sufficient amount of kinetic energy is needed.
      • As temperature of the liquid increases, so does kinetic energy of the liquid.
  • 133. Vapor Pressure
  • 134. Vapor Pressure of Water
  • 135. Vapor Pressure
    • The nature of the temperature, vapor pressure relationship is quadratic. P vap vs. 1/T (Kelvin) gives a direct relationship.
  • 136. Vapor Pressure
    • Vapor pressure equation:
    • R is the universal gas constant (8.3145 J/K), and C is a constant characteristic of a given liquid (y=intercept). In means natural logarithm.
  • 137. Example Problem
    • Using the graph, determine whether water or diethyl ether has the larger enthalpy of vaporization.
    • Ether has the smaller slope therefore smaller ΔH vap
  • 138. Example Problem
    • The vapor pressure of water at 25°C is 23.8 torr, and the heat of vaporization of water at 25°C is 43.9 kj/mol. Calculate the vapor pressure of water at 50°C.
    • 93.7 torr
  • 139. Sublimation
    • Like liquids, solids have vapor pressures. When a solid sublimes , it goes directly from the solid to the gaseous state without passing through the liquid state.
    • Example: Dry Ice.
  • 140. Changes of State
    • Typically when a solid is heated, it will form a liquid and then boil to form a vapor. This process can be represented by a heating curve .
    • Temperature vs. time when energy is added a constant rate.
  • 141. Heating Curve
  • 142. Heating Curve
    • The plateaus in the heating curve represent the positions of phase change.
    • At the melting point, the temperature remains constant until the solid has completely changed to liquid.
    • At the boiling point the temperature remains constant as the added energy is used to vaporize the liquid
  • 143. Heating Curve
    • The energy change that occurs at the melting point when a solid melts is called heat of fusion or enthalpy of fusion .
      • Note that changes of state are physical changes, although intermolecular forces have been broken, no chemical bonds have been broken.
  • 144. Enthalpy of Fusions and Melting Points
  • 145. Melting and Boiling
    • The melting and boiling points for a substance are determined by the vapor pressure of the solid and liquid states.
  • 146. Melting and Boiling
    • At 0°C the vapor pressure of ice is less than that of liquid. Vapor pressure of ice increases more rapidly than water.
  • 147. Melting and Boiling
    • A point is reached when the liquid and solids have identical vapor pressures. This is the melting point .
  • 148. Freezing Point
    • At a temperature at which the vapor pressure of the solid is greater than that of the liquid, the solid would sublime and the vapor would be added to the water.
  • 149. Freezing Point
    • At a temperature at which the vapor pressure of the solid is less than that of the liquid, the liquid would evaporate and the vapor would be added to the ice.
  • 150. Freezing Point
    • At a temperature at which the vapor pressures of the solid and liquid are identical, the vapor is in equilibrium. This is the freezing point of water.
  • 151. Melting Point and Boiling Point
    • The normal melting point of a liquid is the temperature at which the solid and liquid states have the same vapor pressure under conditions where the total pressure is 1 atmosphere.
    • The normal boiling point of a liquid is the temperature at which the vapor pressure of the liquid is exactly 1 atmosphere.
  • 152. Supercooled and Superheated
    • Changes of state do not always occur exactly at the boiling point or melting point.
    • Water can be supercooled below 0°C at 1 atm and remain in the liquid state. At some point the correct ordering of molecules occurs and ice forms, releasing energy in the exothermic process and bringing the temperature back up to the melting point.
  • 153. Supercooled and Superheated
  • 154. Supercooled and Superheated
    • Changes of state do not always occur exactly at the boiling point or melting point.
    • A liquid can also be superheated, or raised to temperatures above its boiling point, especially if it is heated rapidly. Boiling requires high-energy molecules to gather in the same vicinity for bubble formation. This may not happen at the boiling point.
  • 155. Supercooled and Superheated
    • Once a bubble does form, when a liquid is superheated, its internal pressure is greater than the atmospheric pressure. This bubble can burst before rising to the surface, blowing the surrounding liquid out of the container. This is called bumping and is a common experimental problem.
  • 156. Supercooled and Superheated
    • Boiling chips are often added to prevent bumping. These are bits of porous ceramic material containing trapped air that escapes on heating, forming tiny bubbles that act as ‘starters’ for the vapor bubble formation. This allows for smooth onset of boiling.
  • 157. Phase Diagrams 10.9
  • 158. Phase Diagram
    • A phase diagram is a convenient way of representing the phases of a substance as a function of temperature and pressure. It shows which state exists at a given temperature and pressure.
    • Conditions for these phase diagrams are assumed to be a closed system and is plotted as temperature vs pressure.
  • 159. Phase Diagram
  • 160. Phase Diagram
    • The solid/liquid boundary has a negative slope.
      • Melting point of ice decreases as external pressure increases.
      • This is different for most substances other than water because the density of ice is less than that of liquid water at the melting point.
  • 161. Boiling Points of Water vs. Pressure
  • 162. Example 1
    • Pressure is 1 atm. Water moves through the changes of state according to the vapor pressure at the corresponding temperatures.
  • 163. Example 2
    • Pressure is 2 torr. Water will sublime at -10°C. This is when the vapor pressure of the ice is equal to the external pressure of 2 torr. Vapor pressure of liquid water is always greater than 2 torr and therefore will not form.
  • 164. Example 3
    • Pressure is 4.58 torr. When temperature reaches .01°C (273.16K), water reaches the triple point. Solid and liquid water have identical vapor pressures and all three states of water exist. This is the only condition in a closed system that allows this.
  • 165. Example 4
    • Pressure is 225 atm. Liquid water can be present at this temperature because of the high external pressure. As temperature increases, liquid gradually turns to vapor, but goes through a ‘fluid’ region.
  • 166. Example 4
    • The fluid region is neither true liquid or vapor. This unusual behavior occurs because the conditions are beyond the critical point for water.
  • 167. Critical Point
    • The critical temperature can be defined as the temperature above which the vapor cannot be liquefied no matter what pressure is applied.
    • The critical pressure is the pressure required to produce liquefaction at the critical temperature
    • Together, the critical temperature and critical pressure define at the critical point.
  • 168. Critical Point
    • The critical point for water is 374°C and 218 atm. Anything beyond this point, involves the intermediate “fluid” region.
  • 169. Phase Diagrams
    • The phase diagram for CO 2 shows the liquid state does not exist at a pressure of 1 atm. The solid/liquid line has a positive slope, since the density of solid CO 2 is greater than that of liquid CO 2
  • 170. THE END!!