1. States of Matter Chapter 9

2. 9.1 Gases Kinetic – movement Kinetic theory – explains the behavior particles in terms of their motions

3. 9.1 Kinetic Theory of Gases Kinetic Theory of Gases 1. Negligible volume 2. Move in rapid, constant straight-line motion 3. Collide elastically No net loss in kinetic energy 4. Far apart with no attractive or repulsive force

4. 9.3 Kinetic Energy and Temperature Kinetic Energy – energy of motion KE = ½ (m)(v) 2 m = mass v = velocity Temperature – measure of average kinetic energy. Molecules within a substance don’t all move at the same speed

5. Which has the higher average kinetic energy? Red

6. 9.3 Kinetic Energy and Kelvin Temperature The Kelvin Temperature scale reflects the relationship between temp. and average kinetic energy Directly Proportional Ex: The particles in helium gas at 200 K have twice the average kinetic energy of the particles in helium has at 100 K

7. 9.3 Kinetic Energy and Temperature The higher the temperature, the faster moving the particles At 0 K all molecular motion would stop Has never been produced in lab

8. Explaining the Behaviors of Gases Low Density D= m/v Gases are far apart Low mass High volume Compression –If gases are far apart, they can be pushed together

9. Expansion If gases are moving in constant straight-line motion, they will expand to fill any container. 13.1 Explaining the Behaviors of Gases

10. 9.2 Gas Pressure Pressure – simultaneous collisions of billions of molecules with an object Pressure = The force is created by the collisions More collisions  higher pressure Harder collisions  higher pressure When no gas particles are present, they do not collide and cause pressure. This resulting empty space is a vacuum. force area

11. 9.2 Gas Pressure Earth’s gravity holds air molecules in earth’s atmosphere Atmospheric pressure - air molecule collisions We live at the bottom of an ocean of air If you go up a mountain, atmospheric pressure decreases because the depth of air above you is less.

12. 9.2 Gas Pressure Units of pressure Pascal (Pa)-SI unit Atmosphere (atm)-Typical sea level pressure Millimeters mercury (mm Hg)-barometric reading 760 mm Hg = 1 atm = 101.3 kPa = 14.7 psi

13. Measuring Pressure – Barometer

14. 9.4 Forces of Attraction What separates gases from solids and liquids? (What do solids and liquids have that gases don’t) Attractive Forces

15. 9.4 Forces of Attraction There are two types of attractive forces 1. Intramolecular forces – forces within a compound that hold it together. 2. Intermolecular forces – forces between compounds that hold these compounds together.

16. Molecules are close together, but not so tightly packed they can’t move around. Unlike gases, liquids are held together by attractive forces (intermolecular forces – forces between molecules) This gives liquids a fixed volume Molecules are close together, but not so tightly packed they can’t move around. Unlike gases, liquids are held together by attractive forces (intermolecular forces – forces between molecules) This gives liquids a fixed volume 9.4 Nature of Liquids

17. Molecules are not moving fast enough to break attractive force and become a gas Not very compressible – too close together 9.4 Nature of Liquids

18. Density Intermolecular force holds liquids together causing a small volume – this means the density is higher than as a gas Fluidity Molecules aren't so close together they can’t move, but intermolecular forces make them less fluid than gases Viscosity – resistance of a liquid to flow Decreases with increase in temperature because the motion allows the molecules to overcome intermolecular forces

19. 9.4 Nature of Liquids Surface Tension Energy needed to increase the surface area of a liquid Caused by an imbalance of intermolecular force on the surface of a liquid

20. 9.4 Nature of Liquids Capillary Action Cohesion – force of attraction between identical molecules Adhesion – force of attraction between different molecules The adhesive force between water and glass is stronger than the cohesive force between water molecules Meniscus Capillary action

21. 9.5 Phase Changes Requiring Energy 2. Vaporization – process of a liquid changing to a gas or vapor Evaporation - Liquid to gas below boiling point Happens only on surface Only particles moving fast enough to overcome intermolecular forces escape to become a gas Cooling process - as high speed particles escape this lowers the average kinetic energy and therefore the temperature Heating liquid makes more high speed particles and rate of evaporation increases

22. What if this were the energy needed for vaporization

23. 9.5 Evaporation The conversion of a liquid to a gas or vapor below its boiling point is evaporation or vaporization In evaporation molecules at the surface of the liquid break away into the gas or vapor state The high energy molecules escape the surface.

24. 9.5 Evaporation Evaporation of a liquid in a closed container is different. No particles can escape H 2 O(g) molecules (water vapor) H 2 O(l) molecules

25. 9.5 Phase Change Equilibrium What happens to a liquid in a closed container? At first high evaporation and low condensation

26. 9.5 Phase Change Equilibrium Closed Container Continued Over time evaporation rate will stay constant (stays at constant temp in room) and condensation rate increases as the number of gaseous particles increases Eventually the rates will become equal and dynamic equilibrium is reached (particles are still traded but in equal amounts so no net change)

27. 1. Condensation – gas or vapor becomes a liquid slow moving gas particles that attract each other and become a liquid high speed particles that enter the liquid become trapped What is condensation?

29. 9.6 Boiling Point of a Liquid Boiling point is the temperature at which the vapor pressure of the liquid is just equal to the external pressure Bubbles form throughout the liquid

30. Water Molecules in Liquid and Steam

31. Vapor pressure - vaporized particles will exert an outward pressure increasing temperature will increase vapor pressure Vapor Pressure

32. Why do different substances have different vapor pressures? Stronger Intermolecular Forces!

33. 9.6 Boiling Point Normal boiling point- the boiling point of a liquid a standard pressure The boiling point of a liquid will change as the pressure changes Decreased Pressure= Decreased boiling point In Denver water boils at 95 o C The temperature of a boiling liquid can never rise above its boiling point. The more heat that is supplied, the faster the liquid will burn until it all boils away.

36. Boiling vs. Evaporation Boiling point: atmospheric pressure = vapor pressure Evaporation: molecules go from liquid to gas phase Evolutionary process - slow Revolutionary process - fast Lyophilization – freeze drying AIR PRESSURE 15psi VAPOR PRESSURE 15 psi liquid gas

37. 9.7 The Nature of Solids Tightly packed - held together by stronger intermolecular forces than liquids or gases Do not flow - molecules vibrate around fixed points Incompressible

38. 9.7 The Nature of Solids Density Generally the solid form of a substance is the most dense ~10% more dense than liquids Packing tightly together decreases volume H 2 O is the exception

39. 9.7 The Nature of Solids – Types of Solids Crystalline Most solids are crystalline Crystal - the atoms, ions or molecules are arranged in a very orderly, repeating 3D pattern known as the crystal lattice

40. 9.7 The Nature of Solids – Types of Solids Unit Cell Smallest group of particles that retains the shape of the crystal. Allotropes 2 or more different molecular forms of the same element in the same state. Carbon- Diamond and graphite

41. GraphiteDiamond Two allotropes of carbon http://www.creative-chemistry.org.uk/molecules/carbon.htm

44. 9.7 Types of Crystalline Solids TypeUnit Particle Characteristics of Solid Example Atomic atoms Soft Very low melting point Poor conductivity Noble gases Molecular molecules Fairly soft Low to medium-high mp. Poor conductivity I 2, H 2 O, CO 2 Covalent molecular Atoms connected by covalent bonds Very hard Very high mp. Poor conductivity Diamond (C) and Quartz (SiO 2 ) Ionic Ions Hard and brittle High mp. Poor conductivity NaCl metallic Metal cations surrounded by mobile valence electrons Soft to hard and malleable Low to high mp. Excellent conductivity Any metal

45. 9.7 The Nature of Solids – Types of Solids Amorphous No ordered internal structure No definite melting point, melt over a range or gradually “soften” Example Glass, rubber, and many plastics

46. Difference between crystalline and amorphous solids When a crystalline solid is shattered, the fragments have the same surface angles as the original solid When an amorphous solid is shattered the fragments have irregular angles and jagged edges.

47. 9.7 The Nature of Solids Melting Point- The temperature at which the solid turns into a liquid. Freezing Point- The temperature at which the liquid turns into a solid. Ionic solids have high melting points because they are held together by stronger forces. Not all solids melt Ex: Wood- Decomposes

48. 9.8 Sublimation Sublimation – change of a solid directly to a gas Solid CO 2 is called “dry ice” because it sublimates from a solid to a gas without becoming a liquid Ice sublimes and this is how “frost free” refrigerators work Ice cubes left in freezer for a time get smaller The ice sublimes inside the freezer. Deposition- change of a gas directly to a solid - Snowflakes

49. 9.9 Phase Diagram Phase Diagram – diagram that shows relationship between solid, liquid and gas states for a substance in a closed container at different temperatures and pressures

50. Phase Diagram of Water

52. 9.9 Phase Diagram Triple point - temp and pressure where all three states of matter co-exist in equilibrium Critical point – above this temperature, the substance can only exist as a gas

53. Plasma: The Fourth State of Matter Gaseous mixture of electrons and positive ions Generally occurs only at high temperatures where the kinetic energy is enough to strip electrons off of gaseous atoms

54. Plasma: The Fourth State of Matter Happens to a small degree (very unstable and only lasts for a brief moment in time) at lower temperatures in fluorescent lights, neon signs, lightning and the “northern lights.” Only few of the atoms present are ionized at any moment

55. 9.10 Plasma: The 4 th state of matter A tremendous amount of energy is required to create highly ionized plasmas. Cold Plasmas- Gas can be converted to plasma between temperature of 50,000 and 100,000 K Hot Plasmas- Gas is converted to plasma between 10,000,000 and 1,000,000,000. Stars

56. solid liquid gas Heat added Temperature ( o C) A B C D E Heating Curve for Water 0 100 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 487

57. solid liquid gas vaporization condensation melting freezing Heat added Temperature ( o C) A B C D E Heating Curve for Water 0 100 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 487

58. Heating Curve