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Chapter 13

Chapter 13 . Condensed States of Matter . States of Matter:. Solid - composed of particles packed closely together with little space between them. Solids maintain a specific shape.

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Chapter 13

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  1. Chapter 13 Condensed States of Matter

  2. States of Matter: Solid - composed of particles packed closely together with little space between them. Solids maintain a specific shape. Liquid - any substance that flows. (A fluid) - particles are free to slide past one another and continual change their positions. Particles are in constant motion and contact Gases - are fluids composed of particles in constant random motion. Gases are not touching most of the time.

  3. Kinetic-Molecular Description of Liquids & Solids • Solids & liquids are condensed states • atoms, ions, molecules are close to one another • highly incompressible • Liquids & gases are fluids • easily flow • Intermolecular attractionsin liquids & solids are strong

  4. Miscible liquids diffuse into one another • they are soluble in each other • for example: water/alcohol gasoline/motor oil • Immiscible liquids do not diffuse into each other • they are insoluble in each other • for example: water/oil water/cyclohexane

  5. Periodic Table Reminders Vertical columns- called groups or families. Elements in a group have similar chemical and physical properties Horizontal rows- called periods, elements within a period have properties that change progressively across the table

  6. Intermolecular forces, I.F. depend on the shape of the molecules and polarity (dipole moments), Lewis structures and electronegativity Table 13-1 Characteristics of solids, liquids, and gases Ex. Sulfur hexafluoride Ex. water

  7. I.F. types: Ion – ion • Ion – Ion attractions are the strongest. • makes up ionic bonding • tend to be crystalline solids (hard, but brittle) • very high melting points Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl- Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+ Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl- • ion-ion attractions

  8. I.F. types: Dipole – Dipole attractions (fairly strong) happens with polar molecules b/c they have permanent dipole moments H—Cl H—Cl δ+ δ- δ+ δ-

  9. I.F. types: Hydrogen bonding – a special dipole attraction, stronger than normal dipole-dipole attractions - very strong attraction ~ 2 conditions: must have a N, O, F, bonded to a H and at least one lone pair of electrons on N, O, F δ+ δ- H . . l H—N—H δ- :N –H δ+ δ+ l δ+ l H H δ+ δ+

  10. I.F. types: Induced Dipole London dispersion forces (the weakest), also called van der Waals attractions occurs in nonpolar molecules temporary dipole caused by interaction with another molecule boil and melt very easily δ+ δ- F—F F—F δ- δ+

  11. Properties of liquids Properties of liquids at constant temperature ~ no definite shape ~ definite volume ~ have surface tension, diffuse, medium density, viscosity, evaporation, capillary action, and vapor pressure. Strength of each depends on I.F.

  12. 1. Surface tension • Surface Tension - measure of the unequal attractions that occur at the surface of a liquid • molecules at surface of a liquid are only attracted in a down direction. Denser on top • At surface, molecules are attracted downward, thus liquid is denser on top • water bugs

  13. 2. Viscosity – pourability – how easily it flows ~ stronger I.F. more viscous the liquid is ~ geometry of molecule affects viscosity (more complex shapes = more viscous) ~ very long chains – more viscous b/c longer chains get tangled C—C—C—C—C—C—C—C —C—C—C—C more viscous less viscous molasses, honey oil water

  14. 3. Capillary action • – tendency of a liquid to be attracted or repelled by a very narrow tube - Stronger I.F. more cohesion • ~ when a molecule has attraction for itself, it’s called cohesion • ~ when a molecule has attraction for other molecules, it is adhesion • capillary rise implies adhesive > cohesive (water) • capillary fall implies cohesive > adhesive (mercury)

  15. 4. Evaporation – when a liquid changes to the vapor phase at a temp. that is less than it’s boiling point. Why? If the molecule can gain enough speed, they break through the liquid and go into the atmosphere

  16. 5. Vapor pressure – the pressure of a gas that exists over its solid or liquid state. High I.F. = low Pvap Does not depend on how much liquid/solid you have. Pvap depends on the temp. and type of substance. You have vapor pressure as long as there is evaporation of a liquid. Higher Temperature = Higher vapor pressure Boiling occurs when the Pvap of liquid = Patm

  17. The Liquid State • Vapor Pressure (High I.F. = low Pvapor )

  18. Boiling Points & Distillation ~ All liquids have different boiling points: based on I.F. Higher I.F., higher the normal boiling point. ~ can separate liquids on the basis of their b.p. (distillation) CH3OH has a lower boiling point than C2H5OH, so it changes to a gas first.

  19. Phase Diagrams – show the phase of matter at a variety of P and T. (pg 507) ~ Substances can be almost any phase, given the right P and T. H2o is less dense in solid state (ice has lots of space in it), water has a negative slope between the solid & liquid on the phase diagram. Water is densest at 4oC.

  20. Heat of Phase and Temp. change q = mcΔT q = heat m = mass c = specific heat ΔT = change in temp. specific heat is the amount of energy needed to raise 1 g of a substance 1oC If c is big, it’s hard to heat up or cool down. If less than one, easy to heat up or cool down. During a phase change, the temp. stays the same. Still heating/cooling, but no temp. change due to the breaking down or forming of I.F.

  21. Phase changes s  ℓ melting (or fusion) ℓ  s freezing ℓ  g boiling g  ℓ condensation (liquefaction if forced to occur by pressure) s  g sublimation g  s deposition

  22. melt or evaporating (boiling) + qfreeze or condensing -q

  23. q = mHf enthalpy of fusion q = mHv enthalpy of vaporization q = mHs enthalpy of sublimation (A-12) q is + if substance is melting or evaporating q is – if substance is freezing or condensing Ex. How much heat is required to raise the temp. of 50.0 g of ice at –12.0oC to 120.0oC

  24. Trends in boiling points of Liquids GasMWBP(oC) The boiling point increases in response to molecular size

  25. The boiling point increases in response to molecular size

  26. In the Liquid State HF has the highest B.P. b/c of Hydrogen bonding. The rest increases in response to molecular size.

  27. In the Liquid State Water has the highest B.P. because of Hydrogen bonding. The rest increases b/c of increase in molecular size.

  28. Various boiling points Arrange the following substances in order of increasing boiling points. C2H6, NH3, Ar, NaCl, AsH3 Ar < C2H6 < AsH3 < NH3 < NaCl nonpolar, nonpolar, polar, very polar, ionic (H-bonding)

  29. Amorphous & Crystalline Solids • Amorphous solids do nothave a well ordered structure. Particles are irregularly arranged so IF vary in strength within a sample Ex. paraffin, glasses • Crystalline solids have well defined structures that consist of extended array of repeating units. Have defined IF. give X-ray difractionpatterns ~ see Bragg equation in book Ex. Ice, salt

  30. Structure of Crystals • unit cell - smallest repeating unit of a crystal Ex. bricks are repeating units for buildings • 7 basic crystal systems • We do not need to learn these 7 now – just an FYI for your future…

  31. Types of Solids 4 Types of Solids: Ionic, Metallic, Molecular, Covalent Ionic: positive and negative ions arranged in a specific structure. Electrostatic attractions are strong. Metallic: metals where each valence electron is thought to belong to the entire structure. So metals are seen as a positive nuclei with a sea of electrons. The mobility of electrons helps explain the electrical conductivity of metals.

  32. Molecular solids: are solids made up of molecules that are next to each other in unit cells. The attractive forces between individual molecules are relatively weak. They are volatile and insulators. Simple covalent compounds usually form molecular solids Covalent solids: Network solids or giant molecules – individual atoms are covalently bonded to other atoms and those atoms are bonded to other atoms, etc. This makes covalent solids very hard with very high melting points. Most are nonconductors.

  33. Examples of Bonding in Solids • Ionic Solids • ions occupy the unit cell • examples: • CsCl, NaCl, ZnS

  34. Examples of Bonding in Solids • Metallic Solids • positively charged nuclei surrounded by a sea of electrons • positive ions occupy lattice positions • examples: • Na, Li, Au, Ag, ……..

  35. Examples of Bonding in Solids • Molecular Solids • molecules occupy unit cells • low melting points,volatile & insulators • examples: • water, sugar, carbon dioxide, benzene

  36. Examples of Bonding in Solids • Covalent Solids • atoms that are covalently bonded to one another • examples: • SiO2 (sand), diamond, graphite, SiC

  37. Bonding in Solids - Variations in Melting PointsIonic Solids Melt at fairly high temps b/c the attraction between ions are much stronger than in molecular solids but weaker than in covalent solids. Attractive forces increase as charges on ions increase & their radii decrease. CompoundMelting Point (oC) LiF 842 LiCl 614 LiBr 547 LiI 450 CaF2 1360 CaCl2 772 CaBr2 730 CaI2 740

  38. Bonding in Solids - Variations in Melting PointsMetallic Solids Melting points vary widely b/c there are large variations in the strengths of metallic bonding. Most metals have fairly high melting points but Mercury is a liquid at room temp. MetalMelting Point (oC) Na 98 Pb 328 Al 660 Cu 1083 Fe 1535 W 3410

  39. Bonding in Solids - Variations in Melting PointsMolecular Solids Have low melting points (most < 300o C ) because the attractive forces between the molecules are rather weak. CompoundMelting Point (oC) ice 0 ammonia -77.7 benzene, C6H6 5.5 napthalene, C10H8 80.6 benzoic acid, C6H5CO2H 122.4

  40. Bonding in Solids - Variations in Melting PointsCovalent Solids Melt at high temps (most > 1500o C ) because the attractive forces between the individual particles are very strong. SubstanceMelting Point (oC) sand, SiO2 1713 carborundum, SiC ~2700 diamond >3550 graphite 3652-3697

  41. Brief summary Intermolecular attractions from strongest to weakest • Ion-Ion: ionic compounds (metal/nonmetal) • Hydrogen bonding: H attached to a N, O, or F and lone pair of e- on the central atom • Dipole – Dipole: polar compounds • London Dispersion Forces (induced dipole): all compounds exhibit this, but it is most important with non-polar compounds.

  42. Effects of intermolecular attractions • The compound that has the highest boiling pt, melting pt, and heat of vaporization corresponds to the compound with the strongest I.F. • The highest vapor pressure corresponds to the lowest intermolecular attractions. • If two compounds are nonpolar, the one with the greatest molecular mass has the greater London Forces. • If two compounds are ionic, the one with the greatest charge ions has the greater I.F. If same charge, the smallest ions have the greatest I.F.

  43. For many years, the world’s record for flying gliders was 60,000 ft. It was set by a Texan who flew into an updraft in front of an approaching storm. The pilot had to fly out of the updraft and land, not because he was our of air (there was still plenty of air in his compressed air bottle) but because he was not wearing a pressurized suit. What would have happened to the pilot’s blood if he had continued to fly higher?

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