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Liquids

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Liquids

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  1. Liquids Chapter 10

  2. Review: Gases • Indefinite shape • Indefinite volume • Take the shape and volume of container • Particles are far apart • Particles move fast • Low Density • Easy to expand and compress

  3. Review: Solids • Definite shape • Definite volume • Particles close together, fixed • Particles move very slowly • High density • Hard to expand/compress

  4. Liquids: in between • Closer to properties of solids • Slow diffusion • High attraction between particles • Medium amount of energy

  5. Forces of Attraction • Intramolecular forces: Hold atoms together within a molecule • covalent and ionic bonds • Intermolecular forces: Hold molecules to each other • 3 types

  6. Dipole-Dipole Attraction • Dipole: molecule with a separation of charge (polar covalent) • Due to differences in electronegativity • ~1% as strong as a covalent bond

  7. Hydrogen Bond • Very strong dipole-dipole attraction (10% of a covalent bond’s strength) • Occurs when H is bonded to O, N, F in a verypolar bond O H 3.5 2.1 • Gives water its unusual properties

  8. H-Bonding Affects Boiling Points • Strong attraction requires much energy to overcome, so water is a liquid at normal temperatures

  9. London Dispersion Forces • Occur in all substances-polar and nonpolar • Due to formation of instantaneous dipoles as electrons moving around nucleus concentrate on 1 side of molecule or atom

  10. This induces a dipole in neighboring atoms or molecules • These are the weakest intermolecular forces • These are the only forces of attraction in nonpolar substances

  11. Importance of Water • Covers 70% earth’s surface • Necessary for reactions in living cells • Moderates earth’s temperature • Coolant for engines & nuclear power plants • Transportation • Growth medium for many organisms

  12. Properties of Water • Colorless • Tasteless • At 1 atm, water freezes at 0°C and vaporizes completely at 100°C •  Liquid phase occurs from 0-100°C

  13. Special Properties of Water • Surface Tension • Liquids tend to form a “skin” making the surface less penetrable by solids

  14. Unequal attraction at surface, mainly down Equal attraction in all directions

  15. Surface Tension • Detergents can interfere with the attractions and will cause the paperclip to fall

  16. Adhesion • Attraction of the surface of a liquid to the surface of a solid • Depends on the material • Water is attracted to glass • Mercury is not No adhesion

  17. Cohesion • Molecular attractions within a material • (water molecules to water molecules, for example) • Here, cohesion causes water to form beads; ad- hesion causes it to stick to the web

  18. Capillary Action • A liquid rises in a narrow tube when it breaks the surface tension • Movement of water through paper

  19. Ice Floats! • Molecules in a liquid have more movement than a solid and more energy (particles move apart • Generally, solids are more dense than their ir liquids • Liquid water

  20. Ice • When water becomes fixed points in a solid, hydrogen bonds hold molecules in place • Gives ice an open hexagonal structure • Greater volume means lower density than a liquid

  21. Higher K.E. causes distance between molecules to be more Max. density at 4 C Lower K.E. causes distance between molecules to be less

  22. Phase Changes of Water • At 1 atm, water freezes at 0°C and vaporizes completely at 100°C •  Liquid phase occurs from 0-100°C • Changes from one phase to another will either require energy or release energy • Solid Liquid • Liquid Gas • Solid Gas Melting/Freezing Vaporization/Condensation Sublimation/Deposition

  23. From Solid to Liquid • As energy is added, K.E. increases • Solid warms up • At 0°C, solid begins to melt and temperature remains at 0 until all solid is turned to liquid • When all is liquid, temperature begins to rise

  24. From Liquid to Gas • As heat is added, K.E. increases (increase in temperature) • At 100C, bubbles form in liquid • Temperature remains the same until all liquid is converted to a gas. • Once all is a gas, added energy causes temperature to increase

  25. Ice melting/water vaporizing • When a substance is in phase, increasing the energy increases the temperature • When a substance is changing phase, increasing the energy does not increase the temperature but is used to break forces of attraction between molecules

  26. Phase diagram • Temperature vs. Energy

  27. Heating a Solid

  28. Melting a Solid

  29. Heating a Liquid

  30. Vaporizing a Liquid

  31. Heating a Gas

  32. Calculating Energies • Energy is measured in calories or Joules • 1.00 cal = 4.184 J • The amount of energy needed to change states depends on: • Type of matter • Quantity of matter

  33. Type of matter • Molar heat of fusion (Hfusion)= energy needed to melt 1 mole of a substance • Molar heat of vaporization (Hvap)= energy needed to vaporize 1 mole of a substance

  34. For Water • (Hfusion) = 80.0 cal/g = .334 kJ/g • (Hvap) = 540. cal/g = 2.26 kJ/g

  35. Finding Energy in a Phase Change • Change from a solid liquid • q = mHfusion • Change from a liquid gas • q = mHvap q = energy (cal or J) m = mass (g) Hfusion=heat of fusion (cal/g) q = energy (cal or J) m = mass (g) Hvap=heat of vaporization

  36. Example How much heat in calories is needed to melt 15.0 g of water? q = mHfusion 15.0 g water x 80.0 cal = 1.20 x 103 cal 1 g water

  37. Energy and Being In Phase • When all of a substance is in one phase, (e.g. all liquid), the amount of energy required to cause a temperature change depends on: • type of substance • amount of substance • range of temperature change

  38. Type of Matter • Specific Heat (c): Energy required to change the temperature of 1 gram of a substance by 1 Celsius degree • cg = specific heat of a gas • cl= specific heat of a liquid • cs = specific heat of a solid

  39. For Water • cg = .480 cal/gC or 2.01 J/gC • cl = 1.00 cal/gC or 4.18 J/gC • cs = .500 cal/gC or 2.09 J/gC

  40. Finding Energy When In Phase • q = mcT • T = Tfinal – Tinitial q = energy (cal or J) m = mass (g) c = specific heat (cal/gC) T=change in temperature (C)

  41. Example • How much energy is required to heat 50.0 g of water from 20.0 C to 85.0 C? • q = mclT • T = 85.0 – 20.0 = 65.0 C • q = (50.0g)(1.00 cal/gC)(65.0 C) • q = 3,250 cal

  42. Phase changes

  43. In Phase

  44. Phase diagram • Temperature vs. Energy

  45. Phase Change Problems • Draw graph. • Mark start and stop points. • Every corner means a new equation is needed. • Flat sections will use q = mH ( no T means no slope). • Find each energy (q1, q2, q3..). • Add all energies to get the total energy.

  46. Phase Changes • Vaporization (evaporation): molecules of a liquid escape the liquid’s surface • Requires energy to overcome intermolecular forces Maxwell Boltzman distribution Molecules with enough energy to evaporate

  47. To evaporate, a particle must: • Be at the surface • Have sufficient energy • Be moving in the right direction I’m free!! Moving in the wrong direction Not enough energy Not at surface

  48. Evaporation produces Vapor Pressure • A closed container with a vacuum has a liquid added to it. • Molecules begin to evaporate • Some particles are recaptured by the liquid • Eventually rate of particles leaving = rate of being recaptured

  49. Equilibrium Vapor Pressure • When rate of evaporation = rate of condensation the pressure becomes constant (Equilibrium vapor pressure)

  50. Vapor pressure and temperature • As temperature increases, more particles evaporate