Ch. 14/15 – Solids, Liquids and Solutions

1. Ch. 14/15 – Solids, Liquids and Solutions

2. Intermolecular Forces Forces of attraction between different molecules rather than bonding forces within the same molecule. • Dipole-dipole attraction Between particles with charged sides http://dwb4.unl.edu/ChemAnime/DIPOLED/DIPOLED.ht ml Hydrogen bonds - Strongest Between H’s & F, O, N on other molecules • Dispersion forces - Weakest • Caused by electrons shifting towards one end of a molecule.

3. “Water: The Magnificent Dipole” One side of water is negatively charged because the oxygen atom keeps the shared electrons longer than the hydrogen atoms. As a result the oxygen side is negatively charged and the hydrogen side of water is positively charged. O

4. It is its polarity and hydrogen bonding that give water its many unusual Properties! i.e. high boiling point, expansion upon freezing, and surface tension

5. Hydrogen bonds - Strongest Between H’s & F, O, N on other molecules Ice Liquid

6. Water is always trying to pull itself into a tight ball as long as there is nothing nearby that has a charge on it. Therefore, this surface is not repelling water; it’s simply not attracting it and keeping water from doing what it does naturally. Water pulls on itself so much that it forms a “skin.” It’s called surface tension.

7. O O O O O Wax does not repel water We’ve heard that wax or oils repel water. But that isn’t true. Water is so attracted to other water molecules that anything between them is squeezed out of the way. Oildroplet

8. Forces and Phases • Substances with very little intermolecular attraction exist as gases • Substances with strong intermolecular attraction exist as liquids • Substances with very strong intermolecular (or ionic) attraction exist as solids

9. “V” is for Vocabulary! Vapor – gaseous state of a substance that is not normally a gas at room temperature. Volatile – evaporates rapidly (due to weak intermolecular forces) Ex. ‘thin’ liquids Viscous – evaporates slowly (due to strong intermolecular forces) Ex. ‘thick’ liquids

10. Phase Differences

11. Three Phases of Matter

12. Liquids • A decrease in the average kinetic energy of gas particles causes the temperature to decrease. • As it cools, the particles tend to move more slowly, if they slow down enough, attractive forces - called van der Waal’s forces –pull them very close together so they can only slip & slide past each other. Condensation – Change of a gas to a liquid It is now in liquid form!

13. The Nature of Liquids • The conversion of a liquid to a gas or vapor below its boiling point is called Vaporization • (occurs at the surface of a liquid) • In an open container, this process is called Evaporation Particles near the surface with enough kinetic energy that happen to bounce in the right direction escape!

14. Microscopic view of a liquid near its surface.

15. The Nature of Liquids • Eventually the particles will lose energy and return to the liquid state, or condense. • What are the odds that they will return to the original liquid? • What if we cover the container? • So, the particles begin to evaporate, then some begin to condense. Eventually, the number of particles evaporating will equal the number condensing & the space above the liquid will be saturated with vapor

16. A dynamic equilibrium now exists where Rate of evaporation = rate of condensation • Note that there will still be particles that evaporate and condense • But, there will be no NET change • It will not look like there are changes taking place!

17. Evaporation is a cooling processCooling occurs becauseparticles with the highest energy escape first • Particles left behind have lower average kinetic energies; thus the temperature decreases • Similar to removing the fastest runner from a race- the remaining runners have a lower average speed • Evaporation helps to keep our skin cooler on a hot day!

18. The Nature of Liquids • A liquid will evaporate faster when heated because the added heat increases the average kinetic energy needed to overcome the attractive forces so more particles have enough energy to ‘escape’!

19. If you were to add a drop of water below the tube to the left what would happen? It would rise to the top & evaporate. What would it do to the surface of the mercury? (push it down!) Vapor Pressure – pressure exerted by vapor!

20. ( Since different liquids evaporate at different rates, they have different vapor pressures (at the same temps)

21. ( Which liquid is the most volatile? Yes! Diethyl Ether! It depressed the mercury the most (highest vapor pressure!)

22. Vapor Pressures of Liquids Which is the most volatile liquid here? Diethyl ether It has the highest vapor pressure at any Temp. Which has the strongest forces of attraction? Water

23. A liquid boils when its vapor pressure equals the external pressure, so the boiling point changes if the external pressure changes. • Bubbles form throughout the liquid, rise to the surface, and escape into the air • Normal boiling point- is when the vapor pressure of a liquid equals standard pressure. (1 atm)

24. The boiling point (bp) is the temperature at which the vapor pressure of the liquid is equal to the external pressure on the liquid VP BP Direct relationship!

25. Normal bp of water = 100 oC However, in Denver = 95 oC, since Denver is 1600 m above sea level and average atmospheric pressure is about 85.3 kPa(Recipe adjustments?) • In pressure cookers, which reduce cooking time, water boils above 100 oC due to the increased pressure

27. Vapor Pressures of Liquids Normal bp when crossing here At any pt. on a curve line, liquid is boiling

28. SOLIDS • If you cool a liquid, the particles lose kinetic energy and slow down. • If they slow down enough, extra forces of attraction pull them in so close together that they can only vibrate in place. • Freezing – change of a liquid to a solid.

29. Types of Solids • Molecular solids • Metallic solids • Ionic solids • Covalent network solids (diamonds)

30. Crystals or Crystalline Solids • Particles of crystals are arranged in repeating geometric patterns NaCl

31. Representation of Components in a Crystalline Solid Crystal Lattice: The orderly, regular, 3-dimensional arrangement of particles (atoms, ions, etc.) in a crystal.

32. Unit Cell The smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice Ex. A cubic lattice system has 3 types of unit cells

33. Table salt crystals are shaped like cubes.

34. Diamond, a form of carbon, is also a crystalline solid. • the crystals are shaped something like pyramids.

35. Non-crystalline solids • Many solids do not form crystals- Amorphous • Their molecules do not arrange into repeating patterns • often because they are too large. • No definite melting point • Examples: • Glass - also called a super-cooled liquid • many plastics, soot, asphalt, butter

36. PHASE CHANGES • PHASE CHANGES – change is physical state (melting, freezing, boiling, condensing, sublimation, deposition) • BOTH PHASES present during a phase change • Temperature remains constant during a phase change. • Sublimation – change of a solid directly to a gas (dry ice, iodine, snow) • Deposition – change of a gas directly to a solid.

37. Temperature (C°) 60° 20° 0° -20° 0 Heat (kilojoules)

38. Heating and cooling curve for water heated at a constant rates. A-B = Solid ice, temperature is increasing. Particles gain kinetic energy, vibration of particles increases. Ice

39. B-C = Solid starts to change state from solid to liquid. Temperature remains constant as energy is used to break inter-molecular bonds. H2O (s)  H2O (l) 0ºC

40. C-D = temperature starts to rise once all the solid has melted. Particles gain kinetic energy. Liquid water

41. D-E = Liquid starts to vaporize, turning from liquid to gas. The temperature remains constant as energy is used to break inter-molecular forces. H2O () H2O (g) 100ºC

42. E-F = temperature starts to rise once all liquid is vaporized. Gas particles gain kinetic energy. steam

43. The heating/cooling curve for water

44. constant Temperature remains __________ during a phase change. Water phase changes KE not changing during phase changes Boiling condensation Melting freezing Kinetic energy increasing on slopes

45. Phase Diagram • Represents phases as a function of temperature and pressure.

46. Recall Classification of Matter Also called solutions

47. Aqueous Solutions – water solutions Solute The part of a solution that gets dissolved – the part in lesser quantity Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent The part of a solution that does the dissolving – the part in greater quantity Water in salt water Water in soda

48. Solutions occur in all 3 phases!

49. Two substances with similar intermolecular forces are likely to be soluble in each other. “like dissolves like” • non-polar molecules are soluble in non-polar solvents • Ex. Grease in gasoline • Ionic compounds & polar molecules are soluble in polar solvents • Ex. Ethanol in water; salt in water 12.2

50. When a salt dissolves in water, the positive ions are attracted to the slightly negative ends of the water molecules and the negative ions are attracted to the slightly positive ends of the water molecules. They ‘dissociate’ (separate) The ions are more strongly attracted to each other. but they become surrounded by water molecules and can’t get back together! Solvation – where solvent molecules surround solute particles.