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Chapter 12

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Judith C Poe

Chapter 12 Intermolecular Forces, Liquids, Solids, and Phase Changes 12.1:Overview of Physical States and Phase Changes  Phase – a physically distinct, homogenous part of system. Each physical state is a phase  Interactions b/w the potential energy and the kinetic energy of the particles give rise to the properties of each phase: o the potential energy in the form of intermolecular forces tends to draw the molecules together. According to Coulomb’s Law, the electrostatic potential energy depends on the charges of the particles and the distances b/w them o the kinetic energy is associated w. the random motion of the molecules tends to disperse them. It is related to the their average speed and is proportional to the absolute temperature  These interactions explain phase changes (changes in physical state from one phase to another) A Kinetic-Molecular View of the Three States 1. Intramolecular forces exist within each molecule, the chemical behaviour of all three states is the same b/c each consists of the same molecule 2. Intermolecular Forces exist b/w the molecules, they physical behaviour of these states is different b/c the strengths of the forces differ  Each phase is accompanied by a standard enthalpy change, given in kilojoules per mole  ΔHº fus Ep(energy of attraction) -as the Ekdecreases (due to decrease in temp) E pecome more important (attractions) The heat (q) is equal to the amount (moles) x the molar heat capacity (C)(gas) x ∆T [ q = n xwater ∆T] Stage 2: Gaseous Water Condenses -intermolecular attractions cause the slowest of the molecules to form microdroplets then liquid -during phase changes, the temp and E ake constant, at the same temp the molecules move farther b/w collisions but their average speed is the same -has a lower Epand changes from gas to liquid, the heat is the amount x the negative heat of vaporization [q = n x - ∆H vap -this stage contributes the greatest portion of the total heat released b/c of the large decrease pn E Type Exo/Endo Definiton Condensation Exothermic Gas  Liquid (temp drops, gas particles come - ΔHº vap together to form a liquid) Vaporization Endothermic Liquid Gas (increase in temperature = Heat of Vaporization: higher K.E) ΔHº vap Freezing Exothermic Liquid Solid (decrease in temp, decrease in - ΔHº fus K.E) Melting(Fusion) Endothermic Solid  Liquid (increase in temp and K.E) Heat of Fusion: ΔHº fus Sublimation Endothermic Solid  Gas Deposition Exothermic Gas  Solid Stage 3:Liquid Water Cools -the molecules in the liquid state continue to lose heat, and decreases in K.E (goes from 100ºC to 0ºC) - The heat (q) is equal to the amount (moles) x the molar heat capacity (C)(liquid) x ∆T [ q = n x water ∆T] Stage 4: Liquid Water Freezes -at 0ºC, the sample loses Epas increasing intermolecular attractions cause the molecules to align - q = n x -∆fus Stage 5: Solid Water Cools -the solid furthers solidifies as temperature decreases - q = n x C x ∆T (C = for solid water) water -According to Hess’ Law the total amount of heat released is the sum of the heats released for the individual stages -Within a phase, the heat flow is accompanied by a change in temperature, thus increase in E k -During a phase change, heat flow occurs at a constant temperature which is associated w. a change in E p The Equilibrium Nature of Phase Changes -in a closed container phase changes are reversible and reach equilibrium Liquid Gas Equilbria 1. Open System: nonequilbrium process, as the temperature increases the molecules have high enough kinetic energy to evaporate 2. Closed System: two processes take place: some molecules escape and then collide w. the surface (the slower ones condense) 3. Disturbing the System at Equilibrium  Decrease in Pressure = increase in volume, the rate of condensation temporarily falls below the rate of vaporization (forward process is faster-liquid to gas) b/c fewer molecules enter the liquid than leave it  Increase in Pressure: decreasing volume, rate of condensation exceeds rate of vaporization b/c more molecules enter the liquid than leave it (reverse process is faster-gas to liquid)  When a system in equilibrium is disturbed, it counteracts the disturbance & equilib is re-established How does a liquid boil?  in an open container, the weight of the atmosphere bears down on a liquid surface. As the temp rises, molecules move more quickly through the liquid. At some temp, the average E of tke molecules in the liquid is great enough for them to form bubbles of vapour in the interior and the liquid boils.  once boiling begins, the temp of the liquid remains constant until all of the liquid is gone Solid-Liquid Equilbria  Melting Point – the temperature at which the melting rate equals the freezing rate  pressure has a little effect b/c liquids & solids are incompressible Solid-Gas Equilibria  a substance sublimes rather than melts b/c the intermolecular attractions are not great enough to keep the molecules near each other when they leave the solid state 12.3: types of Intermolecular Forces A. How Close Can Molecules Approach Each Other?  Bond Length and Covalent Radius: o Bond length: the shorter distance, is b/w two bonded atoms (i.e two Cl atoms) in the same molecule o Covalent Radius: ½ the bond length distance  Van der Waals distance and Radius: o Distance: the longer distance is b/w two nonbonded atoms in adjacent molecules (i.e Cl). It is called the VDW distance. At this distance, intermolecular attractions balance electron-cloud repulsions, thus the VDW distance is a close as one Cl molecule can approach another 2 o Radius: is ½ the closest distance b/w the nuclei of identical nonbonded atoms. The VWD radius of an atom is always larger than its covalent radius. Like covalent radii, VDW radii decrease across a period and increase down a group  Ion-Dipole Forces: occur when an ion and nearby polar molecule (dipole) attract each other. The most important e
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