1. Unit 2 Matter, Energy, and Changes

2. Energy & Temperature

3. Energy Energy is the capacity to do work or produce heat.

4. Energy Forms:  Radiant ex. sunlight

5. Energy  Kinetic energy carried by objects in motion (includes mechanical & thermal)

6. Energy  Potential due to position/stored energy (includes electrical & chemical)

7. Energy Measuring: calories (cal) – amount of heat needed to raise temperature of 1 g H 2 O by 1 C   1 cal = 1 g × 1 C  Food energy is in Calories, 1 Cal = 1 kcal (or 1,000 cal)

8. Energy Measuring: SI unit is Joule (J), 1 J is about the energy to lift a medium-sized apple 1 meter from ground  1 cal = 4.184 J

9. Energy Law of Conservation of Energy: in any process, energy is neither created nor destroyed

10. Heat is a form of energy Some things heat easily – It takes little energy to change their temperature – How easily something heats is determined by the specific heat capacity of the substance (definition = amount of heat to change the temperature of 1 g of a substance by 1  C) Specific heat = heat (cal) mass (g) × temp. change (  C)

11. Heat Specific Heat (continued)… Water has a high specific heat (1 cal/g×  C) The amount of heat required to heat an object is the same as the amount of heat given off when that object cools

12. Heat Specific Heat Questions It takes 24.3 calories to heat 15.4 g of a metal from 22 ºC to 33ºC. What is the specific heat of the metal? 0.14 cal/(g×  C) Iron has a specific heat of 0.11 cal/gºC. How much heat will it take to change the temperature of 48.3 g of iron by 32.4ºC? 172 cal

13. Temperature What is temperature? comparison of how hot or cold an object is with some standard measure of the average kinetic energy of the particles in a sample of matter. The higher the average KE of the molecules the faster the molecules are moving. = Higher temperature

14. Temperature Temperature is different from heat Temperature is which way heat flows (always goes from hot to cold) Heat is a form of energy A drop of boiling water hurts, but a kilogram of boiling water can kill a person – same temperature, but very different values of heat

15. Temperature Units for Temperature Fahrenheit Celsius Kelvin (SI units)

16. Temperature Units for Temperature Kelvin (SI units)  Why don’t we use a degree mark with Kelvin temperatures?  The scale is an absolute scale, not based on water or some other substance  Lowest temperature reading in Kelvins is called Absolute Zero (NO negative values)  At this temperature KE is zero - ALL movement stops

17. Temperature Units for Temperature A quick look at the three scales.  0  C= 273 K = 32  F  100  C= 373 K = 212  F

18. Temperature Converting Fahrenheit/Celsius  C = 5/9 × (  F – 32)

19. Temperature Convert 98.6  F to  C Convert 25  C to  F

20. Temperature Converting Kelvin/Celsius  C = K – 273

21. Temperature Convert 399 K to  C Convert 25  C to K.

22. Matter & Changes

23. Matter Is anything that has mass and volume

24. States of Matter: Gas  no definite shape or volume  particles randomly scattered (high entropy)  particles moving quickly in constant motion  highly compressible  low density  rapid diffusion  high expansion on heating

25. States of Matter: Liquid  no set shape  definite volume  particles somewhat organized  particles free to move  slightly compressible  high density  slow diffusion  low expansion on heating

26. Properties of Liquids Viscosity – friction or resistance to motion, increases as temperature decreases Surface tension – molecules at the surface experience imbalanced attractive forces

27. States of Matter Both gasses and liquids are fluids. – This is because of weak intermolecular forces. – The molecules can slide easily over each other.

28. States of Matter: Solid  has particular shape  definite volume  particles very organized and close together (low entropy)  strong intermolecular forces  low compressibility  high density  slow diffusion

29. States of Matter: Solid  particles move only very slightly  can only vibrate and/or revolve in place  particles locked in place – they don’t flow  low expansion on heating  when heated, particles vibrate more rapidly until they shake themselves free of each other  energy goes into breaking bonds, not increasing motion  move differently, not faster

30. Basic Types of Solids 1.Crystalline solids 2.Amorphous solids

31. Crystalline Solids  are made of atoms arranged in highly ordered, patterns called unit cells  are made of regular repeating three dimensional arrangement of atoms in a solid  are the most common kind of solid  break at certain angles – Examples: table salt, table sugar, emeralds

32. Amorphous Solids  appear solid, but are more of a super-cooled liquid, have high viscosity, gradually soften as temperature increases  lack an orderly internal structure  are more of a super-cooled liquid than a solid  are rigid, but lack structure  do not melt, but get gradually softer  shatter at random angles – Examples: glass, plastics

33. Properties of Solids Some solids are good conductors due to particles being in contact with each other and passing the energy from one particle to another

34. Physical State of Matter State of matter at room temperature depends on strength of intermolecular (IM) forces For example, a substance with strong IM forces will be a solid while a substance with very weak IM forces will be a gas

35. Matter Intermolecular Forces  Are the forces between neighboring molecules

36. Kinetic Theory and States of Matter uKinetic theory says that all molecules are in constant motion. uPerfume molecules moving across the room are evidence of this.

37. Kinetic Theory and States of Matter Kinetic Theory helps to explain why a gas behaves as it does. It also helps us understand the changes in physical states of matter

38. The Kinetic Theory of Gases Makes three descriptions of gas particles 1. A gas is composed of particles – molecules or atoms – particles are considered to be hard spheres that are far enough apart that we can ignore their volume – Between the particles is empty space

39. 2. The particles are in constant random motion. – They move in straight lines until they bounce off each other or the walls. – The molecules don’t travel very far without hitting each other so they move in random directions. – The average speed of an oxygen molecule is 1656 km/hr at 20°C

40. 3.All collisions are perfectly elastic – no energy is “lost” during collisions.

41. The Kinetic Molecular Theory Explains properties of liquids 1. Particles are in constant motion, but are closer than those in a gas – Attractive forces (aka IM forces) are more effective than those between gas particles This makes the particles more organized than in a gas, but they are still free to move past each other

42. The Kinetic Molecular Theory Explains properties of liquids 2. High density – Due to relatively close arrangement of particles 3. Low compressibility – Due to particles being close together 4. Ability to diffuse – Happens because particles are free to move – Is slower in liquids because particles are close and attractive forces slow movement

43. The Kinetic Molecular Theory Explains properties of liquids 5. Surface tension – Due to attractive forces between particles Because particles at the surface have no liquid particles above them, a spherical shape occurs

44. The Kinetic Molecular Theory Explains properties of solids 1. Particles are in constant motion, but are closer than those in a gas or a liquid – Attractive forces (aka IM forces) are strongest between particles of a solid This makes the particles very organized and able to move only by vibrating or rotating

45. The Kinetic Molecular Theory Explains properties of solids 2. Definite shape and volume – regardless of container – Because of high IM forces between particles 3. Definite melting point (crystalline solids) – Melting happens when particles have enough kinetic energy to overcome IM forces that hold them in place

46. The Kinetic Molecular Theory Explains properties of solids 4. High density (and low compressibility) – Happens because particles are more closely packed than in gas or liquid This is what limits compressibility 5. Low rate of diffusion – Millions of times slower than in liquids Particles mostly fixed, so moving past one another is extremely slow when it occurs

47. Changes of State Energy and change of state… To change states, particles must overcome the attractive forces holding them together (the number of particles does not change)

48. Changes of State Energy and change of state… There are 6 changes of state 1.Vaporization 2.Condensation 3.Solidification 4.Liquefication 5.Sublimation 6.Deposition

49. Changes of State Vaporization Liquid changes to a gas (at temperature lower than the boiling point) Requires energy input

50. Changes of State Vaporization Sometimes called evaporation – Evaporation technically happens only when liquid vaporizes from an uncontained sample (example – leaving the lid off of a perfume bottle)

51. Changes of State Vaporization Happens when a rapidly moving particle near surface of liquid gains enough energy to escape attractive forces of other particles

52. Changes of State Vaporization Related vocabulary: Volatile liquid – one that readily evaporates

53. Changes of State More about evaporation… Molecules at the surface break away and become gas. Only those with enough KE escape Evaporation is a cooling process. It requires energy.

54. Changes of State Boiling point Temperature at which vapor pressure becomes equal to the atmospheric pressure

55. Changes of State More about boiling… Making bubbles of substance in gaseous state Forces liquid level to rise Liquid must push against air pressure on the liquid.

56. Changes of State Still more about boiling… A liquid boils when the vapor pressure = the external air pressure Temperature where this happens is the boiling point Normal boiling point is the temperature a substance boils at 1 atm pressure. – The normal boiling point of water is 100 o C The temperature of a liquid can never rise above it’s boiling point Energy goes into breaking attractive forces, not moving faster.

57. Changes of State Changing the boiling point… Lower the pressure (like going up into the mountains) – Lower external pressure requires lower vapor pressure – Easier to make bubbles – Lower vapor pressure means lower boiling point – Food cooks slower

58. Changes of State Changing the boiling point… Raise the external pressure (like using a pressure cooker) – Raises the vapor pressure needed – Harder to make bubbles – Raises the boiling point – Food cooks faster

59. Changes of State Different boiling points… Different substances boil at different temperatures because they have different intermolecular forces – Weaker IM forces gives lower boiling point Different vapor pressures – Low vapor pressure gives high boiling point

60. Changes of State Heat of vaporization Amount of heat necessary to vaporize a given amount of liquid

61. Changes of State Condensation Gas changes to a liquid Molecules are attracted to one another Releases energy.

62. Changes of State Solidification (aka Freezing) Particles get closer together and are more organized than in the liquid state Releases energy

63. Changes of State Liquefication (aka Melting) Particles become less organized and farther apart, requires energy input Melting point - temperature at which solid and liquid form of substance exist in equilibrium, also called freezing point

64. Changes of State Liquefication (aka Melting) Heat of fusion - amount of heat needed to convert a given amount of solid into a liquid

65. Changes of State Sublimation Solid changes directly to a gas, requires energy input

66. Changes of State Deposition Gas changes directly to a solid, releases energy

67. Liquid Solid Gas Require energy to break IM forces Release energy Sublimation Deposition Melting Vaporization Freezing Condensation

68. Changes of State Temperature and Phase Change The temperature doesn’t change during any phase change. – If you have a mixture of ice and water, the temperature is 0º C The temperature will not go higher (or lower) until the change of state is complete

69. Changes of State Heating curves Describe changes of state of matter, a graph of sample temperature as a function of time

70. Changes of State

71. Phase diagrams Relates states of matter to temperature and pressure

72. Changes of State