1. 2006, Prentice Hall Unit 10 Liquids, Solids and Intermolecular Forces

2. 2 Interactions Between Molecules many of the phenomena we observe are related to interactions between molecules that do not involve a chemical reaction your taste and smell organs work because molecules in the thing you are sensing interact with the receptor molecule sites in your tongue and nose in this chapter we will examine the physical interactions between molecules and the factors that effect and influence them

3. 3 The Physical States of Matter matter can be classified as solid, liquid or gas based on what properties it exhibits Fixed = keeps shape when placed in a container, Indefinite = takes the shape of the container

4. 4 Structure Determines Properties the atoms or molecules have different structures in solids, liquid and gases, leading to different properties

5. 5 The Structure of Solids, Liquid and Gases

6. 6 Structure of Gases in the gas state, the particles have complete freedom from each other the particles are constantly flying around, bumping into each other and the container in the gas state, there is a lot of empty space between the particles on average

7. 7 Structure of Gases because there is a lot of empty space, the particles can be squeezed closer together – therefore gases are compressible because the particles are not held in close contact and are moving freely, gases expand to fill and take the shape of their container, and will flow

8. 8 Properties of Gases low densities compared to solids and liquids fluid the material exhibits a smooth, continuous flow as it moves take the shape of their container expand to fill their container can be compressed into a smaller volume

9. 9 Structure of Liquids the particles in a liquid are closely packed, but they have some ability to move around the close packing results in liquids being incompressible but the ability of the particles to move allows liquids to take the shape of their container and to flow – however they don’t have enough freedom to escape and expand to fill the container

10. 10 Properties of Liquids high densities compared to gases fluid the material exhibits a smooth, continuous flow as it moves take the shape of their container keep their volume, do not expand to fill their container can not be compressed into a smaller volume

11. 11 Structure of Solids the particles in a solid are packed close together and are fixed in position though they are vibrating the close packing of the particles results in solids being incompressible the inability of the particles to move around results in solids retaining their shape and volume when placed in a new container; and prevents the particles from flowing

12. 12 Structure of Solids some solids have their particles arranged in an orderly geometric pattern – we call these crystalline solids salt and diamonds other solids have particles that do not show a regular geometric pattern over a long range – we call these amorphous solids plastic and glass

13. 13 Properties of Solids high densities compared to gases nonfluid they move as entire “block” rather than a smooth, continuous flow keep their own shape, do not take the shape of their container keep their own volume, do not expand to fill their container can not be compressed into a smaller volume

14. 14 Why is Sugar a Solid But Water is a Liquid? the state a material exists in depends on the attraction between molecules and their ability to overcome the attraction the attractive forces between ions or molecules depends on their structure the attractions are electrostatic depend on shape, polarity, etc. the ability of the molecules to overcome the attraction depends on the amount of kinetic energy they possess

15. 15 Interactions Between Molecules many of the phenomena we observe are related to interactions between molecules that do not involve a chemical reaction your taste and smell organs work because molecules in the thing you are sensing interact with the receptor molecule sites in your tongue and nose in this chapter we will examine the physical interactions between molecules and the factors that effect and influence them

16. Intermolecular Attractive Forces

17. 17 Why are molecules attracted to each other? intermolecular attractions are due to attractive forces between opposite charges + ion to - ion + end of polar molecule to - end of polar molecule H-bonding especially strong larger charge = stronger attraction even nonpolar molecules will have a temporary induced dipoles (dispersion forces)

18. 18 Dispersion Forces also known as London Forces or Induced Dipoles caused by electrons on one molecule distorting the electron cloud on another all molecules have dispersion forces + -- - - -- -- - - + + + + - - - - - - - - - - - - - - - + + + + - - - - - - - - - - - - - - - + -- - - -- -- - -

19. 19 Instantaneous Dipoles

20. 20 Strength of the Dispersion Force depends on how easily the electrons can move, or be polarized the more electrons and the farther they are from the nuclei, the larger the dipole that can be induced strength of the dispersion force gets larger with larger molecules

21. 21 Attractive Forces and Properties stronger attractive forces between molecules = higher boiling point in pure substance stronger attractive forces between molecules = higher melting point in pure substance though also depends on crystal packing

22. 22 Dispersion Force and Molar Mass

23. 23 Dipole-to-Dipole Attraction polar molecules have a permanent dipole a + end and a – end the + end of one molecule will be attracted to the – end of another

24. 24 Polarity and Dipole-to-Dipole Attraction

25. 25 Attractive Forces + - + - + - + + + + _ _ _ _ Dispersion Forces – all molecules Dipole-to-Dipole Forces – polar molecules

26. 26 Attractive Forces and Properties “Like dissolves Like” miscible = liquids that do not separate, no matter what the proportions polar molecules dissolve in polar solvents water, alcohol, CH 2 Cl 2 molecules with O N or F have higher solubility in H 2 O due to H-bonding with H 2 O nonpolar molecules dissolve in nonpolar solvents ligroin (hexane), toluene, CCl 4 polar molecules do not mix with non polar molecules

27. 27 Immiscible Liquids When liquid pentane, a nonpolar substance, is mixed with water, a polar substance, the two liquids separate because they are more attracted to their own kind of molecule than to the other.

28. 28 Hydrogen Bonding Molecules that have HF, OH or NH groups have particularly strong intermolecular attractions unusually high melting and boiling points unusually high solubility in water this kind of attraction is called a Hydrogen Bond

29. 29 Properties and H-Bonding Name Form- ula Molar Mass (g/mol) Structure Boiling Point, °C Melting Point, °C Solubil- ity in Water EthaneC2H6C2H6 30.0-88-172immisc EthanolCH 4 O32.064.7-97.8 misc- ble

30. 30 Intermolecular H-Bonding

31. 31 Hydrogen Bonding When a very electronegative atom is bonded to hydrogen, it strongly pulls the bonding electrons toward it. Since hydrogen has no other electrons, when it loses the electrons, the nucleus becomes deshielded exposing the proton The exposed proton acts as a very strong center of positive charge, attracting all the electron clouds from neighboring molecules

32. 32 H-Bonds vs. Chemical Bonds hydrogen bonds are not chemical bonds hydrogen bonds are attractive forces between molecules (intermolecular force) chemical bonds are attractive forces that make molecules (intramolecular force)

33. 33 Attractive Forces & Properties

34. 34 Types of Intermolecular Forces Type of Force Relative Strength Present inExample Dispersion Force weak, but increases with molar mass all atoms and molecules H2H2 Dipole – Dipole Force moderate only polar molecules HCl Hydrogen Bond strong molecules having H bonded to F, O or N HF

35. 35

36. Phase Changes

37. 37 Escaping from the Surface the process of molecules of a liquid breaking free from the surface is called evaporation also known as vaporization evaporation is a physical change in which a substance is converted from its liquid form to its gaseous form the gaseous form is called a vapor

38. 38 Evaporation over time, liquids evaporate – the molecules of the liquid mix with and dissolve in the air the evaporation happens at the surface molecules on the surface experience a smaller net attractive force than molecules in the interior but all the surface molecules do not escape at once, only the ones with sufficient kinetic energy to overcome the attractions will escape

39. 39 Factors Effecting the Rate of Evaporation increasing the surface area increases the rate of evaporation increasing the temperature increases the rate of evaporation weaker attractive forces between the molecules = faster rate of evaporation liquids that evaporate quickly are called volatile liquids, while those that do not are called nonvolatile

40. 40 Escaping the Surface the average kinetic energy is directly proportional to the kelvin temperature but not all molecules in the sample have the same kinetic energy those molecules on the surface that have enough kinetic energy will escape raising the temperature increases the number of molecules with sufficient energy to escape

41. 41 Escaping the Surface since the higher energy molecules from the liquid are leaving, the total kinetic energy of the liquid decreases, and the liquid cools the remaining molecules redistribute their energies, generating more high energy molecules the result is the liquid continues to evaporate

42. 42 Reconnecting with the Surface when a liquid evaporates in a closed container, the vapor molecules are trapped the vapor molecules may eventually bump into and stick to the surface of the container or get recaptured by the liquid – this process is called condensation a physical change in which a gaseous form is converted to a liquid form

43. 43 Dynamic Equilibrium evaporation and condensation are opposite processes eventually, the rate of evaporation and condensation in the container will be the same opposite processes that occur at the same rate in the same system are said to be in dynamic equilibrium

44. 44 Evaporation and Condensation When water is just added to the flask and it is capped, all the water molecules are in the liquid. Shortly, the water starts to evaporate. Initially the speed of evaporation is much faster than speed of condensation Eventually the condensation and evaporation reach the same speed. The air in the flask is now saturated with water vapor.

45. 45 Vapor Pressure once equilibrium is reached, from that time forward, the amount of vapor in the container will remain the same as long as you don’t change the conditions the partial pressure exerted by the vapor is called the vapor pressure the vapor pressure of a liquid depends on the temperature and strength of intermolecular attractions

46. 46 Boiling in an open container, as you heat a liquid the average kinetic energy of the molecules increases, giving more molecules enough energy to escape the surface so the rate of evaporation increases eventually the temperature is high enough for molecules in the interior of the liquid to escape – a phenomenon we call boiling

47. 47 Boiling Point the temperature at which the vapor pressure of the liquid is the same as the atmospheric pressure is called the boiling point the normal boiling point is the temperature required for the vapor pressure of the liquid to be equal to 1 atm the boiling point depends on what the atmospheric pressure is the temperature of boiling water on the top of a mountain will be cooler than boiling water at sea level

48. 48 Temperature and Boiling as you heat a liquid, its temperature increases until it reaches the boiling point once the liquid starts to boil, the temperature remains the same until it all turns to a gas all the energy from the heat source is being used to overcome the attractive forces in the liquid

49. 49 Energetics of Evaporation as it loses the high energy molecules through evaporation, the liquid cools then the liquid absorbs heat from its surroundings to raise its temperature back to the same as the surroundings processes in which heat flows into a system from the surroundings are said to be endothermic as heat flows out of the surroundings, it causes the surroundings to cool as alcohol evaporates off your skin, it causes your skin to cool

50. 50 Energetics of Condensation as it gains the high energy molecules through condensation, the liquid warms then the liquid releases heat to its surroundings to reduce its temperature back to the same as the surroundings processes in which heat flows out of a system into the surroundings are said to be exothermic as heat flows into the surroundings, it causes the surroundings to warm

51. 51 Heat of Vaporization the amount of heat needed to vaporize one mole of a liquid is called the heat of vaporization  H vap it requires 40.7 kJ of heat to vaporize one mole of water at 100°C endothermic   H vap depends on the initial temperature since condensation is the opposite process to evaporation, the same amount of energy is transferred but in the opposite direction  H cond = -  H vap

52. 52 Heats of Vaporization of Liquids at their Boiling Points and at 25°C Liquid Chemical Formula Normal Boiling Point, °C  H vap at Boiling Point, (kJ/mol)  H vap at 25°C, (kJ/mol) water H2OH2O10040.744.0 isopropyl alcohol C 3 H 7 OH82.339.945.4 acetone C3H6OC3H6O56.129.131.0 diethyl ether C 4 H 10 O34.526.527.1

53. 53 Temperature and Melting as you heat a solid, its temperature increases until it reaches the melting point once the solid starts to melt, the temperature remains the same until it all turns to a liquid all the energy from the heat source is being used to overcome the attractive forces in the solid that hold them in place

54. 54 Energetics of Melting and Freezing when a solid melts, it absorbs heat from its surroundings, it is endothermic as heat flows out of the surroundings, it causes the surroundings to cool as ice in your drink melts, it cause the liquid to cool when a liquid freezes, it releases heat into its surroundings, it is exothermic as heat flows into the surroundings, it causes the surroundings to warm

55. 55 Heat of Fusion the amount of heat needed to melt one mole of a solid is called the heat of fusion  H fus fusion is an old term for heating a substance until it melts, it is not the same as nuclear fusion since freezing is the opposite process to melting, the same amount of energy transferred is the same, but in the opposite direction  H crystal = -  H fus in general,  H vap >  H fus because vaporization requires breaking all attractive forces

56. 56 Heats of Fusion of Several Substances Liquid Chemical Formula Melting Point, °C  H fusion, (kJ/mol) water H2OH2O0.006.02 isopropyl alcohol C 3 H 7 OH-89.55.37 acetone C3H6OC3H6O-94.85.69 diethyl ether C 4 H 10 O-116.37.27

57. Heat of Fusion and Vaporization 57

58. 58 Sublimation sublimation is a physical change in which the solid form changes directly to the gaseous form without going through the liquid form like melting, sublimation is endothermic

59. Crystalline Solids

60. 60 Types of Crystalline Solids

61. 61 Molecular Crystalline Solids Molecular solids are solids whose composite units are molecules Solid held together by intermolecular attractive forces dispersion, dipole-dipole, or H-bonding generally low melting points and  H fusion

62. 62 Ionic Crystalline Solids Ionic solids are solids whose composite units are formula units Solid held together by electrostatic attractive forces between cations and anions cations and anions arranged in a geometric pattern called a crystal lattice to maximize attractions generally higher melting points and  H fusion than molecular solids because ionic bonds are stronger than intermolecular forces

63. 63 Atomic Crystalline Solids Atomic solids are solids whose composite units are individual atoms Solid held together by either covalent bonds, dispersion forces or metallic bonds melting points and  H fusion vary depending on the attractive forces between the atoms

64. 64 Types of Atomic Solids

65. 65 Types of Atomic Solids Covalent Covalent Atomic Solids have their atoms attached by covalent bonds effectively, the entire solid is one, giant molecule because covalent bonds are strong, these solids have very high melting points and  H fusion because covalent bonds are directional, these substances tend to be very hard

66. 66 Types of Atomic Solids Nonbonding Nonbonding Atomic Solids are held together by dispersion forces because dispersion forces are relatively weak, these solids have very low melting points and  H fusion

67. 67 Types of Atomic Solids Metallic Metallic solids are held together by metallic bonds metal atoms release some of their electrons to be shared by all the other atoms in the crystal the metallic bond is the attraction of the metal cations for the mobile electrons often described as islands of cations in a sea of electrons

68. 68 Metallic Bonding the model of metallic bonding can be used to explain the properties of metals the luster, malleability, ductility, electrical and thermal conductivity are all related to the mobility of the electrons in the solid the strength of the metallic bond varies, depending on the charge and size of the cations – so the melting points and  H fusion of metals vary as well