Chapter 12 Liquids and Solids

1. Chapter 12LiquidsandSolids 2006, Prentice Hall

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

5. Properties of the States of MatterGases • 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

6. Properties of the States of MatterLiquids • 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

7. Properties of the States of MatterSolids • 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

8. The Structure of Solids, Liquid and Gases

9. 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

10. 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

11. 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

12. 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

13. 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

14. Why is Sugar a Solid ButWater 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. Properties of LiquidsViscosity • some liquids flow more easily than others • the resistance of a liquid to flow we call viscosity • larger the attractive forces between the molecules = larger the viscosity • also, molecules whose shape is not round will have a larger viscosity

16. Properties of LiquidsSurface Tension • liquids tend to minimize their surface – a phenomenon we call surface tension • this tendency causes liquids to have a surface that resists penetration

17. Surface Tension • molecules in the interior of a liquid experience attractions to surrounding molecules in all directions • but molecules on the surface experience an imbalance in attractions, effectively pulling them in • to minimize this imbalance and maximize attraction, liquids try to minimize the number of molecules on the exposed surface by minimizing their surface area • stronger attractive forces between the molecules = larger surface tension

18. Forces of Attraction within a Liquid • Cohesive Forces = forces that try to hold the liquid molecules to each other • surface tension • Adhesive Forces = forces that bind a substance to a surface • capillary action • meniscus

19. 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

20. 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

21. 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

22. 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

23. 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

24. 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

25. 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

26. 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. When water is just added to the flask and it is capped, all the water molecules are in the liquid. Evaporation and Condensation

27. 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

28. 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

29. 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

30. 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

31. 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

32. 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

33. Heat of Vaporization • the amount of heat needed to vaporize one mole of a liquid is called the heat of vaporization • DHvap • it requires 40.7 kJ of heat to vaporize one mole of water at 100°C • endothermic • DHvap 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 • DHcond = -DHvap

34. Heats of Vaporization of Liquidsat their Boiling Points and at 25°C

35. Apply the Solution Map: Information Given: 155 kJ Find: g H2O CF: 40.7 kJ = 1 mol; 18.02 g = 1 mol SM: kJ → mol → g Example:Calculate the amount of water in grams that can be vaporized at its boiling point with 155 kJ of heat. = 68.626 g H2O • Sig. Figs. & Round: = 68.6 g H2O

36. 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

37. 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

38. Heat of Fusion • the amount of heat needed to melt one mole of a solid is called the heat of fusion • DHfus • 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 • DHcrystal = -DHfus • in general, DHvap > DHfus because vaporization requires breaking all attractive forces

39. Heats of Fusion of Several Substances

40. 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

41. IntermolecularAttractive Forces

42. 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

43. - - - - - - - - - - - - - - + - - + - - - - - - - - + - + + - - + - + + - - - - - - - - - - - - - - + + - - - - - - - - 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

44. Instantaneous Dipoles

45. 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

46. 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

47. Dispersion Force and Molar Mass

49. Permanent Dipoles • because of the kinds of atoms that are bonded together and their relative positions in the molecule, some molecules have a permanent dipole • all polar molecules have a permanent dipole

50. 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