1. Unit 9: Liquids & Solids

2. Three States of Matter State ShapeVolumeWhy? Particles far apart; forces GasNoneNone Particles far apart; forcessmall Particles closer forces greater LiquidNoneFixed Particles closer forces greater Particles touch forces great SolidFixedFixed Particles touch forces great

3. Why Gas Laws and Not Solid or Liquid Laws? Gases are mainly empty space; have weak attractions between molecules Solids/liquids have particles which are closer together and have more varied forces between particles.

4. Phase Transitions Definition: Physical changes that result in changes of state. ALL phase changes involve energy (enthalpy). Review: Enthalpy = ΔH = heat at const P

5. Phase Changes that Require Energy If you have to put energy into a reaction to make it happen, it is an endothermic reaction. Endothermic Phase Changes: Melting (solid  liquid) a.k.a “fusion” Vaporization (liquid  gas) Sublimation (solid  gas)

6. Phase Changes that Release Energy If energy is released or given off by a reaction, it is an exothermic reaction. Exothermic Phase Changes: Condensation (gas  liquid) Deposition (gas  solid) Freezing (liquid  solid)

8. NOTE: What is constant at every phase change above????

9. Vaporization Definition: conversion of liquid to gas. (endothermic) We commonly call this evaporation. Condensation is the reverse process (exothermic) LIQUIDVAPOREvaporation condensation

10. Enthalpy of Vaporization Heat that is absorbed to vaporize a given amount of liquid at a constant temperature and pressure. Units: Energy (J, kJ or cal) per g or mol LIQUID + heat ---> VAPOR Compd.∆H vap (kJ/mol) H 2 O40.7 (at 100 o C) SO 2 26.8 (-47 o C) Xe12.6 (-107 o C)

11. Heat of Condensation is same quantity but opposite in sign! ∆H vap = - ∆H cond VAPOR ---> LIQUID + heat VAPOR ---> LIQUID + heat ∆H cond for water = - 40.7 kJ/mol

12. Example 9. 1: How many kilojoules of heat are required to vaporize 598.5 g of ethanol? The heat of vaporization is 43.3 kJ/mol.

13. Vapor Pressure Definition: Pressure exerted by a vapor in a closed flask in equilibrium with its liquid Equilibrium: two opposing processes occur at same rate. A: Rate vap> Rate cond B: Rate vap = Rate cond

14. Example 9.2 Liquid ethanol has a vapor pressure of 43.9 mm Hg at 20 deg. C. What is the minimum volume of a flask needed to vaporize 1.00 g of liquid ethanol?

15. Graphing Vapor Pressure As Temperature increases, vapor pressure increases. As attractive forces between molecules increase, vapor pressure decreases. Attractive Forces: Water > Ethanol > Ether Liquids with higher vapor pressure at a given T are said to be more volatile.

16. Questions: 1. What does this graph tell you about the relative attraction between molecules for substances a – e? 2. Which substance is most volatile?

17. Boiling Point A liquid boils when vapor pressure = atmospheric pressure. Normal Boiling Point: Temp. where P = 760 mm Hg (1 atm) on vapor pressure curve. Dependency on pressure As pressure increases, boiling point increases; as pressure decreases, boiling point decreases.

18. Normal Boiling Point

19. Melting and Freezing Point Conversion of a solid to a liquid is called melting, or fusion. Conversion of a liquid to a solid is called freezing. The freezing point = melting point Energy needed to melt a given quantity of solid is called the enthalpy of fusion, or Δ H fus.

20. Example 11.3 How much energy is required to melt 100.0 g of ice? The heat of fusion of water is 6.01 kJ/mol.

21. Example 11.4 How much energy in kJ is required to heat 100.0 g of liquid water from zero to 100 deg. C, and then vaporize all of it? Strategy: 1. Calculate energy needed to heat water (Q equation from Chapter 6) 2. Calculate how much energy is needed to vaporize the water. 3. Add the two amounts together.

22. Example 11.4 How much energy in kJ is required to heat 100.0 g of liquid water from zero to 100 deg. C, and then vaporize all of it?

23. Example 11.5 How much energy is required to heat 75.0 g of ice from 0.0 deg. C to 185.0 deg. C? The heat capacity of steam is 1.84 J/g deg. C. Strategy: 1. Calculate energy needed to melt ice. 2. Calculate how much energy is needed to heat the water to boiling. 3. Calculate how much energy is needed to vaporize the water. 4. Calculate how much energy is needed to heat the steam. 5. Add the four amounts together.

24. Example 11.5 How much energy is required to heat 75.0 g of ice from 0.0 deg. C to 185.0 deg. C? The heat capacity of steam is 1.84 J/g deg. C.

25. Cooling Curve: H 2 O (g)  H 2 O (s)

26. Intramolecular Forces The attractive forces that hold particles together in ionic, covalent and metallic bonds are called intramolecular forces “Intra-” prefix = within The forces inside a molecule holding the individual atoms together Ex.) Covalent bonds Ionic bonds Metallic bonds

27. Intermolecular Forces (IMF’s) : Bonds between Molecules 1. Much weaker than Chemical Bonding within molecules 2. Chemical Bonds (ionic and covalent) determine chemical properties 3. Intermolecular forces determine physical properties e.g. density, mp, bp, solubility, vapor pressure, etc.

28. Three General Types of Intermolecular Forces (IMF’s) Dispersion (London, van der Waals) Dipole/dipole Hydrogen Bonding

29. London Dispersion Forces Definition: IMF between two non-polar molecules formed by temporary positive and negative attractions due to the shifting of electron cloud. Found in all substances, but become important when they are the only IMF present. Strength increases as molar mass increases.

30. Formation of a dipole in two nonpolar I 2 molecules: London Dispersion Forces

31. Higher molar mass ---> larger dispersion forces MoleculeBoiling Point ( o C) MoleculeBoiling Point ( o C) CH 4 (methane) - 161.5 CH 4 (methane) - 161.5 C 2 H 6 (ethane)- 88.6 C 2 H 6 (ethane)- 88.6 C 3 H 8 (propane) - 42.1 C 3 H 8 (propane) - 42.1 C 4 H 10 (butane) - 0.5 C 4 H 10 (butane) - 0.5 Higher boiling point means GREATER IMF’s! London Dispersion Forces

32. Example 11.6. Account for the fact that chlorine is a gas, bromine is a volatile liquid, and iodine is a volatile solid at room temperature.

33. Dipole-Dipole Forces Definition: Attractions between oppositely charged regions of polar molecules. Caused by attraction of one dipole for another. Present in all polar substances! Solubility and dipole-dipole forces: “like dissolves like”

34. Dipole-Dipole Forces

35. Hydrogen Bonding A special type of dipole-dipole force occurring only between molecules with a H atom bonded to either a F, O, or N atom. How to recognize: F, O, or N directly bonded to H Two reasons why hydrogen bonds are stronger than dipole-dipole forces: a. F, O, N very electronegative b. H is a small atom Hydrogen bonding is FON!

36. Hydrogen Bonding

37. Hydrogen Bonding in H 2 O H-bonding is especially strong in water because The O—H bond is very polar There are 2 lone pairs on the O atom Accounts for many of water’s unique properties.

38. H bonds ---> abnormally high specific heat capacity of water (4.184 J/gK) This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy. Hydrogen Bonding in H 2 O

39. Hydrogen Bonding H bonds lead to abnormallyhighboiling point of water. See Screen 13.7

40. Hydrogen Bonding in DNA  Hydrogen bonding plays a key role in maintaining the double helix structure of DNA

41. Example: Identify the types of intermolecular forces present in compounds of: Hydrogen Fluoride Pentane (C 5 H 12 ) Hydrochloric Acid Ethanol (Ethyl Alcohol)

42. Example  Rank these substances in terms of increasing boiling point. N 2 CCl 4 CH 3 Cl NH 3

43. Liquids  Viscosity – a measure of the resistance of a liquid to flow The particles in a liquid are close enough together that their attractive forces slow their movement as they flow past one another The stronger the attractive forces (intermolecular forces), the more viscous the liquid is.  As temperature increases, viscosity decreases.

44.  Surface tension – an inward force that tends to minimize the surface area of a liquid A measure of the inward pull by particles in the interior The stronger the intermolecular forces, the higher the surface tension Liquids In water, this is due mainly to hydrogen bonding!

45. Liquids  Surfactant – any substance that interferes with the hydrogen bonding between water molecules & reduces surface tension

46.  Surfactants used to clean up oil spills as well  Exxon Valdez oil spill in 1989 spilled over 700,000 barrels of oil into the water near Alaska

47. Network Covalent Solids Giant molecules connected by strong covalent bonds Properties: hard, high mp, nonconductors

48. Network Covalent Solids Atoms that can form multiple covalent bonds (look for C, Si, and other Group 14 elements) are able to form network covalent solids. All atoms in the entire structure are bonded together with covalent chemical bonds.

49. Physical Properties of Graphite vs. Diamond PropertyGraphiteDiamond Density (g/mL) 2.273.51 Hardness Very softVery hard Color Shiny blackColorless/transparent Electrical Conductivity HighNone  H comb (kJ/mol) -393.5-395.4

50. Metallic and Ionic Solids: See Chem Act. 24 and 25

51. Ionic Solids A compound where each cation is simultaneously attracted to an anion. Review: How can we identify an ionic compound?

52. Properties of Ionic Solids 1. Molecules, atoms or ions locked into a CRYSTAL LATTICE 2. Particles are CLOSE together 3. STRONG IM forces 4. Highly ordered, rigid, incompressible ZnS, zinc sulfide

53. Metallic Solids A solid consisting of entirely metals. Characteristics: Electrons are “delocalized” (they can move freely) Good conductors, malleable, ductile

54. Metallic Solids  Metallic solids – positive metal ions surrounded by a sea of mobile electrons Mobile electrons make metals malleable and ductile because electrons can shift while still keeping the metal ions bonded in their new places  Metallic solids are good conductors of heat and electricity Metallic Bonds

55. Amorphous Solids  Amorphous solid – a solid in which the particles are not arranged in a regular, repeating pattern “Amorphous” = “without shape” Often form when a molten material cools too quickly to allow enough time for crystals to form  Common examples: glass, rubber, many plastics

56. Types of Solids TYPEEXAMPLEFORCE Ionic NaCl, CaF 2, ZnSIon-ion MetallicNa, FeMetallic MolecularIce, I 2 Dipole Ind. dipole NetworkDiamondExtended Graphitecovalent