1. Honors Unit 9: Liquids & Solids

2. Three States of Matter State of Matter ShapeVolumeWhy? Solid Definite shape Definite volume Particles close together; strong intermolecular forces Liquid Not definite (takes shape of container) Fixed Particles still close; reasonable amount of intermolecular forces Gas Not definite (takes shape of container) No set volume (takes volume of container) Particles are far apart; intermolecular forces are negligible

3. Why Gas Laws and not Solid/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. These forces between particles (intermolecular forces) play a major role in the behavior of liquids and solids, whereas they are negligible in gases!

4. Intermolecular Forces (IMF) Bonds between Molecules  “Inter-” prefix = between  Short range forces between different molecules in a sample  There are 3 main types of intermolecular forces London Dispersion forces Dipole-dipole forces Hydrogen bonding

5. Covalent vs. Intermolecular Forces Intermolecular forces = London Dispersion, Dipole- dipole, Hydrogen bond density, mp, bp, solubility, vapor pressure Determine physical properties like density, mp, bp, solubility, vapor pressure Intramolecular force = covalent bond, metallic bond, ionic bond Determine chemical properties like reactivity, flammability, etc. All three intermolecular forces are weak relative to the strength of a covalent bond!  Attractive energy in ice is 50 kJ/mol  Covalent bond in water is 928 kJ/mol

6. London Dispersion Forces IMF between two non-polar molecules formed by temporary positive and negative attractions due to the shifting of electron cloud. Weakest intermolecular force (they’re temporary!) Found in all substances, but become important when they are the only IMF present. Strength increases as molar mass (number of electrons) increases.

7. Formation of an induced dipole in two nonpolar helium molecules: London Dispersion Forces

8. **Higher molar mass = larger dispersion forces Higher boiling point means GREATER IMF’s! (More forces holding particles in liquid state – can’t get to gas state as easily) MoleculeBoiling Point (in °C) CH 4 (methane)- 161.5 C 2 H 6 (ethane)- 88.6 C 3 H 8 (propane)- 42.1 C 4 H 10 (butane)- 0.5

9. Example #1 Account for the fact that chlorine is a gas, bromine is a volatile liquid, and iodine is a volatile solid at room temperature.

10. Dipole-Dipole Forces Definition: Attractions between oppositely charged regions of polar molecules.  Caused by attraction of δ + for δ -  Present in all polar substances! Compounds w/ DDF have higher melting and boiling points  The molecules “stick together” and require much more energy for the phase change Solubility:  “like dissolves like”

11. Dipole-Dipole Forces

12. Hydrogen Bonding A special type of dipole-dipole force  Unusually strong IMF  Must have an H attached directly to an F or O or N in Lewis structure  The δ + H from one molecule is strongly attracted to the negative end of the dipole of another Remember: Hydrogen bonding is FON!

13. Hydrogen Bonding

14. Hydrogen bonding raises melting & boiling points because more energy is required to break the forces between molecules. Hydrogen Bonding

15. H-bonding is especially strong in water because  The O—H bond is very polar  There are 2 lone pairs of electrons on the O atom Accounts for many of water’s unique properties such as: Hydrogen Bonding  High specific heat capacity of water (4.184 J/g · K)  The reason lakes/oceans control climate  Increases volume upon freezing (floats)  High surface tension

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

17. Example #2 (a) What types of intermolecular forces are present in the following substances? (b) Rank these substances in order of increasing boiling point. N 2 HF SiCl 4 CH 3 Cl NH 3

18. Example #2 (a) What types of intermolecular forces are present in the following substances? (b) Rank these substances in order of increasing boiling point. N 2 HF SiCl 4 CH 3 Cl NH 3

19. Relative Strength of Intermolecular Forces (Weakest)(Strongest)

20. Phase Changes What is a phase change? A physical change in a substance’s state of matter ALL phase changes involve energy (enthalpy, ΔH). What are some examples of phase changes?

21. 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 a.k.a “fusion”(solid  liquid)  Vaporization/Evaporation(liquid  gas)  Sublimation (solid  gas)

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

24. Sample Heating Curve Questions 1. Where does evaporation occur? 1. _______________ 2. What is the melting point of this sample? 2. _______________ 3. Where is there only a liquid present? 3. _______________ 4. Where would the molecules have the most kinetic energy? 4. _______________ 5. At what time does boiling begin? 5. _______________ 6. When does melting end? 6. _______________ 7. Where would ­freezing happen? 7. _______________

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

26. Cooling Curve: H 2 O (g)  H 2 O (s) *** Note that cooling curves are just the opposite of a heating curve!

27. Phase Transitions Melting/Freezing (+/-)  Heat of fusion (ΔH fus ) – energy required to melt/freeze 1 mole of a substance Vaporization/Condensation(+/-)  Heat of vaporization (ΔH vap ) – energy required to vaporize/condense 1 mole of a substance  ∆H cond = - ∆H vap Sublimation/Deposition  Phase change between solid and gas

28. Heating Curve Equations q = m c ∆T (use for temperature changes) q = n ∆H (use for phase changes) Note that temperature DOES NOT change during a phase change!

29. Example #3 Note there is no temperature change during a phase change!! So, use the thermochemical equation: q = nΔH How many kilojoules of heat are required to completely vaporize 598.5 g of ethanol? The heat of vaporization is 43.3 kJ/mol.

30. Example #4: How much energy in kJ is required to heat 100.0 g of liquid water from zero to 100°C, and then vaporize all of it? ΔH vap, water = 40.79 kJ/mole

31. Example #5: How much energy is required to heat 75.0 grams of ice from 0.0°C to 185.0°C? ΔH vap, water = 40.79 kJ/mole ΔH fus, water = 6.01 kJ/mole C ice = 2.09 J/g°C C steam = 1.84 J/g°C C water = 4.18 J/g°C

32. Example #5: How much energy is required to heat 75.0 grams of ice from 0.0°C to 185.0°C? ΔH vap, water = 40.79 kJ/mole ΔH fus, water = 6.01 kJ/mole C ice = 2.09 J/g°C C steam = 1.84 J/g°C C water = 4.18 J/g°C

33. Vapor Pressure (Review) Vapor pressure – pressure due to force of gas particles above a liquid colliding with walls of container Higher temp. =higher vapor pressure

34. Liquid-Vapor Equilibrium  Equilibrium = two opposing processes occur at same rate. (A) Rate vap > Rate cond (B) Rate vap = Rate cond

35. Liquid-Vapor Equilibrium  Equilibrium = two opposing processes occur at same rate.  Dynamic Equilibrium with vapor pressure = when the rate at which the liquid vaporizes is equal to the rate at which the vapor condenses  The liquid level in the container does not change  Molecules are constantly moving between phases with no net change

36. Vapor Pressure Curves (see pg. 11 in ref. book) 1.What does this graph tell you about the relationship between vapor pressure and temperature?

37. Vapor Pressure Curves As temperature increases, vapor pressure increases.  As attractive forces between molecules increase, vapor pressure decreases.  Attractive (IMF) Forces: Water > Ethanol > Ether  Liquids with higher vapor pressures at a given T are said to be more volatile

38. 1.What does this graph tell you about the IMF between molecules in substances a – e? 2.Which substance is most volatile? 3.At a specific temperature, where would the boiling point of a substance be?

39. Boiling Point  A liquid boils when vapor pressure = atmospheric pressure.  At this point, the liquid molecules overcome atmospheric pressure and jump into the gas phase.  Normal Boiling Point = temperature where vapor pressure equals standard pressure (760 mm Hg, 1 atm) on vapor pressure curve  Dependency on pressure:  As pressure increases, boiling point increases; as pressure decreases, boiling point decreases.

40. 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 intermolecular (attractive) forces between molecules, the more viscous the liquid is.  As temperature increases, what happens to the viscosity?

41.  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!

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

43. Classes of Solids Crystalline Network Covalent Metallic Molecular Ionic

44. Crystalline Solids  Crystalline solid – a solid in which the atoms, ions, or molecules are arranged in an orderly, geometric, three-dimensional structure  Unit cell – the smallest arrangement of connected points that can be repeated to form the lattice  A.k.a. The representative part of the whole crystal

45. Crystalline Solids Pyrite (cubic) Rutile (tetragonal)

46. Crystalline Solids Borax (monoclinic) Copper sulfate (triclinic)

47. Network Covalent Solids Atoms that can form multiple covalent bonds are able to form network covalent solids. (such as C, Si, and other Group 14 elements) Giant molecule connected by strong covalent bonds.  Properties: hard, high mp, nonconductors

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

49. Metallic Solids Metallic solids – positive metal ions surrounded by a sea of mobile electrons Electrons are “delocalized” (they can move freely) 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 Video Clip – Metals & Metallic Bonds

50. Molecular Solids  Formed when atoms are bonded together with covalent bonds  Forms a molecule  Properties of Molecular Solids:  Formed from 2 nonmetals  Do not conduct electricity (can be used as insulators)  Low melting & boiling points

51. Ionic Solids  Formed from attraction between cation & anion  Properties of Ionic Solids:  Formed from metal (cation) & nonmetal (anion)  Strong electrolytes – will conduct electricity when dissolved in water (or when molten)  High melting & boiling points  Forms crystal lattice

52. Ionic Solids Forms crystal lattice made of formula units Ionic formula units show the ratio of cation to anion in the crystal lattice.