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Chapter 12 Intermolecular Forces

Chapter 12 Intermolecular Forces. Earth: 15 ºC. Uranus: -214 ºC. What happens to molecules at the melting point?. Intra molecular forces versus Inter molecular forces (aka. van der Waals forces ). Chemical Change

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Chapter 12 Intermolecular Forces

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  1. Chapter 12Intermolecular Forces Earth: 15 ºC Uranus: -214 ºC

  2. What happens to molecules at the melting point? Intramolecular forces versus Intermolecular forces (aka. van der Waals forces) Chemical Change Breaks covalent, ionic and metallic bonds Physical Change Electrostatic forces between particles

  3. Intramolecular forces Intermolecular forces

  4. Section 12.1: Physical States and Phase Changes Increasing Energy Solid Liquid Gas Kinetic-Molecular View of Three States of Matter P.E. and K.E. together determine the physical state of any given substance. P.E. – draws molecules together Coulomb’s Law (Chap2) – particles with opposite charge attract each other. The energy of attraction between two particles is proportional to the product of the charges and inversely proportional to the distance between them. K.E. – separates or disperses molecules K.E. ~ f(absolute temperature) (in Kelvins) ºC = K - 273

  5. Section 12.1: Characteristics of Physical States Increasing Energy Solid Liquid Gas • Definite shape. • Definite volume. • Particles fixed & close. • Particle interaction v. strong. • Particle movement v. slow. • Takes shape of container. • Definite volume. • Particles (molecules) random & close. • Particle interaction strong. • Particle movement moderate. • Takes shape of container. • Fills container volume. • Particles (molecules) random and far apart. • Essentially no interaction. • Particle movement very fast. Ex: ice, iron, table salt Ex: water, oil, vinegar Examples: water vapor, helium gas

  6. Section 12.1: Phase Changes Condensation: Gas to liquid Vaporization: Liquid to gas Heat of vaporization Freezing: Liquid to solid Melting (Fusion): Solid to liquid Heat of fusion Sublimation: Solid to gas Deposition: Gas to solid Heat of sublimation

  7. Section 12.1: Energy and Phase Changes Enthalpy changes (∆H) accompany phase changes Exothermic Phase Changes (-∆H) Condensation: Gas to liquid Freezing: Liquid to solid Deposition: Gas to solid Enothermic Phase Changes (+∆H) Vaporization: Liquid to gas Melting (Fusion): Solid to liquid Sublimation: Solid to gas

  8. Section 12.1: Energy and Phase Changes Enthalpy change is different for different substances For a pure substance: ∆H is measured in change per mole of the substance and is specific to the pressure and temperature conditions Pressure is usually 1 atm, Temperature is that at which the phase change occurs Example: Phase changes of water H2O (l)  H2O (g) ∆H = ∆H º vap = 40.7 kJ/mol (at 100 ºC) H2O (s)  H2O (l) ∆H = ∆H º fus = 6.02 kJ/mol (at 0 ºC) H2O (g)  H2O (l) ∆H = ∆H º vap = -40.7 kJ/mol (at 100 ºC) H2O (l)  H2O (s) ∆H = ∆H º fus = -6.02 kJ/mol (at 0 ºC) Why is ∆H º vap (40.7 kJ/mol) greater than ∆H º fus (6.02 kJ/mol)? ∆H º subl = ∆H º fus +∆H º fus

  9. Section 12.2: Quantifying Phase Changes Temperature Heat Removed The Heating-Cooling Curve – shows how the temperature of a substance changes as heat is added or removed from a substance at a constant rate (at a constant P too)

  10. Interlude: Pressure Matters too (but we assume 1 atm in this class for phase change calculations)

  11. Interlude: Pressure Matters too (but we assume 1 atm in this class for phase change calculations) http://www.windows.ucar.edu/tour/link=/earth/Water/temp.html&edu=mid Methane hydrates http://pathways.fsu.edu/faculty/geeo/ Pressures ~ 1000 atm CH4 freezing point: -182.5 ºC

  12. Section 12.2: Quantifying Phase Changes But back to temperature……….. 5 heat-releasing stages Temperature Heat Removed Temp change: q = nC∆T where q is heat, n is # of moles, C is molar heat capacity Temp constant: q = n∆H where q is heat, n is # of moles, ∆H is heat released/absorbed

  13. Section 12.2: Quantifying Phase Changes In class problem: 12.20 Suggested problem: 12.2. 12.3, 12.12, 12.19, 12.27

  14. Section 12.2: Equilibrium and Phase Changes Closed Container Closed Container Open Container Systems can be closed to some things, but not others Heat source In a closed system, phases changes of many substances reach equilibrium. Open system – volume of liquid does not change – net direction of molecule movement is out of the liquid Closed system – volume of liquid does not change – net direction of molecule movement is out of the liquid Your system is defined by you: Is the ocean a closed system? Is Earth a closed system?

  15. Concept of Open Systems and Steady-State Water in This system is not in steady-state if the volume changes with time. Water in Water in Water out Water out Water out Time When matter/energy is leaving and entering an open system, it can reach Steady-state Water in Definition of steady-state: FluxIN = FluxOUT Flux: Mass or Volume / time Ex: 100 L H2O / hr 20 g CaCO3 / day Water out

  16. Section 12.2: Liquid-Gas Equilibrium and Vapor Pressure Equilibrium vapor pressure – the pressure exerted by a vapor when it has reached equilibrium in a system that isclosed with respect to the vapor molecules When enough time passes, the system will reach equilibrium with respect to the vapor entering and exiting the liquid. Vapor is stuck in the container and will accumulate, putting pressure (P=Force/Area) on container walls. Universal Concept: When a system at equilibrium is disturbed, it counteracts the disturbance and eventually re-establishes a state of equilibrium (For chemical reactions, called Le Châtelier’s principle  Chap17, CHEM 163)

  17. Section 12.2: Liquid-Gas Equilibrium and Vapor Pressure Higher T = Higher V.P. Higher T  increases the fraction of molecules moving fast enough to escape the liquid  decreases the fraction of molecules moving slow enough to be captured

  18. Section 12.2: Liquid-Gas Equilibrium and Vapor Pressure Nonlinear relationship between T and P (in graph) expressed as linear relationship: Clausius-Clapeyron equation - mathematical relationship between T and P Know: P1, T1, ∆Hvap

  19. Section 12.2: Quantifying T – P Relationships In class problem: 12.22, 12.24 Suggested problem: 12.21, 12.23

  20. Section 12.2: Vapor Pressure and Boiling Point Why is my soup not as hot at Camp Muir?!??! Altitude = 10,000 feet Atmospheric pressure lower = 590 mm Hg Boiling Point = 90 °C Altitude = Sea level 760 mm Hg (=1 atm) Boiling Point = 100 °C People living in Denver, CO use pressure cookers to cook food at higher temperature. Boiling point – temperature at which the vapor pressure equals the external pressure Atmospheric Pressure Pressure exerted on the Earth by all the gas particles in the Earth’s atmosphere. Pressure = Force / Area From physics: F = ma F  force m  mass of particle a  acceleration (= g, acceleration due to gravity)

  21. Section 12.3: Types of Intermolecular Forces Bonding (Intramolecular) forces: Relatively strong involve large charges that are closer together Nonbonding (Intermolecular) forces: Relatively weak involve smaller charges that are farther together

  22. Section 12.3: Types of Intermolecular Forces Why are bonding (intramolecular) forces stronger than van der Waals (intermolecular forces)? Periodic Table trends are similar to those for bond length.

  23. Section 12.3: Types of Intermolecular Forces O Na Na Na Na Cl Cl Cl Cl H H +1 -1 + + + + + + + + + + + + + + Na + + Cl + + Cl Cl Cl Cl Na Na Na Na (1) Ion-dipole – an ion interacts with a partial charge Example: NaCl (table salt) dissolves in water Dissolution (NaCl “dissociates”)

  24. Section 12.3: Types of Intermolecular Forces (2) Dipole-dipole – polar molecules interact The greater the dipole moment of a molecule, the great the dipole-dipole forces between molecules of that type  more energy needed to separate them

  25. Section 12.3: Types of Intermolecular Forces O H H + + + + + + + + + + (2) Dipole-dipole – polar molecules interact Hydrogen bond – a special type of dipole-dipole force that arises between atoms that have a H atom bonded to a small, highly electronegative atom with lone electron pairs N, O, and F all fit this profile. (3) Charge-induced dipoles – a molecule with a full or partial charge induces a temporary dipole on a nonpolar molecule (4) London (dispersion) forces – caused by momentary oscillations of e- charge in atoms and, therefore, are present in all particles (atoms, ions, and molecules)

  26. Section 12.3: Trends in Polarizability Polarizability – the ease with which the e- cloud of a particle can be distorted Smaller atoms (ions) are less polarizable than larger ones  e-’s closer to the nucleus and, therefore, held more tightly Polarizability • Increases down a Group • Decreases from L  R • Cations less polarizable than their original atoms Anions are more polarizable than original atoms

  27. Why dry ice (solid CO2) sublimates

  28. Biodiesel Lab: Bomb Calorimeter Combustion reaction – heat flows from the system to the surroundings = exothermic Heat is lost to: (1) water in the calorimeter (2) the calorimeter itself

  29. Section 12.4: Zooming in on Liquids Increasing Energy Solid Liquid Gas Liquids are least understood at the molecular level. Orderly & random at different times Randomness of particles  any region is pretty much identical to any other Orderliness of particles  Different regions identical Macroscopic properties of liquids are well understood: • Surface tension • Capillarity • Viscosoty

  30. Section 12.4: Surface Tension Intermolecular forces exert different effects on a molecule at the surface of a liquid than at the interior: A liquid tends to minimize the # of molecules at the surface. Interior molecules – attracted by water molecules on all sides Surface molecules –attracted to water molecules below and on sides  Experience a net downward attraction How insects walk on water. (water strider) • At the surface of a liquid, water molecules behave as a thin, elastic membrane or “skin”  surface tension – energy required to increase the surface area (J/m2 of surface area increased)

  31. Section 12.4: Surface Tension Surfactants (surface- active agents)—destroy surface tension by congregating at the surface and disrupting the hydrogen bonds between water molecules Example: Needle on water. Example: Respiratory distress syndrome (RDS) in infants. Occurs when infant does not produce a surfactant that breaks H-bonds and does not allow O2 and CO2 exchange between alveoli and capillaries in the lungs. H-bonds break, needle sinks. H-bonds hold needle on water surface. The stronger the forces are between the particles in a liquid, the greater the surface tension.

  32. Section 12.4: Capillarity Capillary action – the rising of a liquid through a narrow space against the pull of gravity  due to competition between intermolecular forces in a liquid (cohesive forces) And those between the liquid and the tube walls (adhesive forces) TLC and plant pigment lab Meniscus on a test tube Glass = SiO2 Mercury (Metallic bonds stronger than any interaction with SiO2) Water (H-bonds with SiO2)

  33. Section 12.4: Viscosity triglyceride + methanol 3 methyl ester + glycerol A liquid’s resistance to flow  resistance decreases as Temp increases Molecular shape plays a role – Biodiesel lab Larger molecules – make more contact = higher viscosity Smaller molecules – make less contact = lower viscosity

  34. Section 12.5: Uniqueness of Water • • Na Na Na Na Cl Cl Cl Cl O + + H H Cl Cl Cl Cl Na Na Na Na The water molecules is bent and highly polar  due to this structure and charge distribution, water can engage in four H bonds with its neighbors. (1) Water is the “universal solvent” (solvent = the compound that does the dissolving) Dissolves a range of solutes ( = the compounds that are dissolved) Polar Covalent substances Nonpolar Covalent substances Ionic substances CH3CH2OH C6H12O6 N2 gas

  35. Section 12.5: Uniqueness of Water Water has a high heat of vaporization – heat from Sun results in vaporization of ocean water  heat stored in water vapor carried poleward  heat released when water vapor condenses back to liquid water – called latent heat transport oceanmotion.org/html/background/climate.htm (2) Water has a high specific heat capacity (the measure of the heat absorbed by a substance for a given rise in temperature – Section 6.3) In other words, water can absorb a lot of heat with relatively small changes in temp. Earth: Daily temperature changes = 40 ºC (in deserts – most extreme) Waterless Moon: 250 ºCdaily fluctuations

  36. Section 12.5: Uniqueness of Water Solids (minerals) Air Water (3) Surface properties are crucial to living things Trees get water due to capillary action in soils and in xylem (veins of trees) Plant debris floating on water surface provides shelter and nutrients

  37. Section 12.5: Uniqueness of Water (4) Density of solid and liquid water Large spaces in the ice due to the hexagonal crystal structure result in solid water being more dense than liquid  lake surfaces freeze in winter (organisms live below)

  38. Section 12.5: Uniqueness of Water Summarized

  39. Heating-Cooling Curve Practice How much heat would need to be added to heat 50.0 g of water ice at -50.0 ºC to water vapor at 135 ºC? Given: Cice = 37.6 J/mol ºC Cliquid = 75.4J/mol ºC Cgas = 33.1 j/mol ºC ∆Hfusion = 6.02 kJ/mol ∆Hvaporization = 40.7 kJ/mol Answer: 1.59 x 105 J

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