2. Conceptual Chemistry

3. Objective 1  Describe, at the molecular level, the difference between a gas, liquid, and solid phase.

4. Solids  Definite shape  Definite volume  Particles are vibrating and packed close together.  The particles do not flow.

5. Crystalline Solids  Particles are arranged in an organized pattern.  Example: Diamond

6. Amorphous Solids  Particles are not organized in an orderly fashion.  Example: Glass

7. Liquids  Indefinite shape  Definite volume  Liquids will take the shape of a container, but they maintain the same volume.  Particles are touching and packed close together.  The higher energy allows the particles to move around each other.

8. Viscosity  A liquid’s resistance to flow.

9. Gases  Indefinite shape  Indefinite volume  Gases take the shape of a container. They also occupy the volume of the container no matter how big or small it is.  High energy motion

10. Plasma  High energy matter  A common example is the sun.  Super high energy gas particles that lost electrons.  Plasma is the most common form of matter in the Universe.

11. States of Matter PropertySolid (s)Liquid (l)Gas (g) Particle SpacingClose Great EnergyLowMediumHigh MotionLowMediumHigh ShapeDefiniteIndefinite VolumeDefinite Indefinite

12. Objective 2  Describe states of matter using the kinetic molecular theory.

13. Kinetic Molecular Theory  The behavior of matter in its different states can be explained using Kinetic Molecular Theory.  Kinetic Molecular Theory – A theory explaining the states of matter based on the concept that the particles in all forms of matter are in constant motion.  Kinetic Energy – energy an object has due to its motion

14. Kinetic Energy and Kelvin Temperature  As particles are heated they absorb energy, thus increasing their average kinetic energy and their temperature.  Molecules at a specific temperature have a wide range of kinetic energies.  Theoretically, molecular motion stops at absolute zero (0 Kelvin).  Kelvin temperature scale reflects the relationship between temperature and average kinetic energy.  Kelvin temperature scale is directly proportional to the average kinetic energy.

15. Objective 3  Describe changes in states of matter with respect to kinetic energy and temperature.

16. Energy and Phase Changes  During a phase change, all the energy goes to motion until phase change is done.  The temperature does not change until the phase change is done.

17. Melting  Solid  Liquid Example 1 Example 2

18. Freezing  Liquid  Solid Example 1

19. Evaporation/Boiling  Liquid  Gas Example 1

20. Condensation  Gas  Liquid Example

21. Sublimation  Solid  Gas Example Opposite of Sublimation? Deposition Example

22. Objective 4  Describe the different variables that define a gas.

23. Kinetic Theory of Gases The basic assumptions of the kinetic molecular theory: Gases are mostly empty space. The molecules in a gas are separate, very small, and very far apart.

24. Kinetic Theory of Gases The basic assumptions of the kinetic molecular theory: Gas molecules are in constant, chaotic motion. Collisions between gas molecules are elastic (there is no energy gain or loss).

25. Kinetic Theory of Gases The basic assumptions of the kinetic molecular theory are: The average kinetic energy of gas molecules is directly proportional to the absolute temperature. Gas pressure is caused by collisions of molecules with the walls of the container.

26. Behavior of Gases Gases have weight. Gases take up space. Gases exert pressure. Gases fill their containers. Gases doing all of these things!

27. Variables that Describe a Gas Volume: measured in L, mL, cm 3 (1 mL = 1 cm 3 ) Amount: measured in moles (mol), grams (g) Temperature: measured in Kelvin (K)  K = ºC + 273 Pressure: measured in mm Hg, torr, atm, etc.  P = F / A (force per unit area) Exploding Trash Can Clip: http://www.youtube.com/watch?v=JsnjvFy7 aw8

28. Moderate Force (about 100 lbs) Small Area (0.0625 in 2 ) Enormous Pressure (1600 psi) P = F / A

29. Bed of Nails Large Surface Area (lots of nails) Moderate Force Small Pressure P = F / A

30. Units of Pressure 1 atm = 760 mm Hg 1 atm = 760 torr 1 atm = 1.013 x 10 5 Pa 1 atm = 101.3 kPa

31. Boyle’s Law As P , V  and vice versa…. Inverse relationship P 1 V 1 = P 2 V 2 For a given number of molecules of gas at a constant temperature, the volume of the gas varies inversely with the pressure.

32. Boyle’s Law and Kinetic Molecular Theory How does kinetic molecular theory explain Boyle’s Law?  Gas molecules are in constant, random motion.  Gas pressure is the result of molecules colliding with the walls of the container.  As the volume of a container becomes smaller, the collisions over a particular area of container wall increase…the gas pressure increases!

34. Pressure-Volume Calculations  Example: Consider the syringe. Initially, the gas occupies a volume of 8 mL and exerts a pressure of 1 atm.  What would the pressure of the gas become if its volume were increased to 10 mL?

35. Equation for Boyle’s Law  P 1 V 1 =P 2 V 2 where: P 1 = initial pressure V 1 = initial volume P 2 = final pressure V 2 = final volume

36. P1V1=P2V2P1V1=P2V2  Using the same syringe example, just “plug in” the values: P 1 V 1 = P 2 V 2 (1 atm) (8 mL)=(P 2 ) (10 mL)

37. P1V1=P2V2P1V1=P2V2 (1 atm) (8 mL)=P 2 (10 mL) P 2 = 0.8 atm

38. Example: A sample of gas occupies 12 L under a pressure of 1.2 atm. What would its volume be if the pressure were increased to 3.6 atm? (assume temp is constant) P 1 V 1 = P 2 V 2 (1.2 atm)(12 L) = (3.6 atm)V 2 V 2 = 4.0 L

39. Example: A sample of gas occupies 28 L under a pressure of 200 kPa. If the volume is decreased to 17 L, what be the new pressure? (assume temp is constant) P 1 V 1 = P 2 V 2 (200 kPa)(28 L) = (P 2 )(17 L) P 2 = 329 kPa

40. Temperature – Volume Relationships  What happens to matter when it is heated?  It EXPANDS.  What happens to matter when it is cooled?  It CONTRACTS.  Gas samples expand and shrink to a much greater extent than either solids or liquids.

41. Charles’ Law The volume of a given number of molecules is directly proportional to the Kelvin temperature. As T , V  and vice versa…. Direct relationship Video Clip 1Video Clip 1, Clip 2

42. Temperature – Volume Relationship  Doubling the Kelvin temperature of a gas doubles its volume.  Reducing the Kelvin temperature by one half causes the gas volume to decrease by one half…  WHY KELVIN? The Kelvin scale never reaches “zero” or has negative values.

43. Converting Kelvin  To convert from Celsius to Kelvin: add 273. Example: What is 110 ºC in Kelvin? 110 ºC+273=383 K

44. Converting Kelvin  To convert from Kelvin to Celsius: subtract 273. Example: 555 K in Celsius? 555 K-273=282 ºC

45. Example: A sample of nitrogen gas occupies 117 mL at 100.°C. At what temperature would it occupy 234 mL if the pressure does not change? V 1 = 117 mL;T 1 = 100 + 273 = 373 K V 2 = 234 mL;T 2 = ??? V 1 / T 1 = V 2 / T 2 (117 mL) / (373 K) = (234 mL) / T 2 T 2 = 746 K

46. Example: A sample of oxygen gas occupies 65 mL at 28.8°C. If the temperature is raised to 72.2°C, what will the new volume of the gas? V 1 = 65 mL;T 1 = 28.8 + 273 = 301.8 K V 2 = ??? mL;T 2 = 72.2 + 273 = 345.2 K V 1 / T 1 = V 2 / T 2 (65 mL) / (301.8 K) = (V 2 ) / 345.2 K V 2 = (65 mL) (345.2 K) / (301.8 K) V 2 = 74.3 mL

47. Temperature – Pressure Relationships  Picture a closed, rigid container of gas (such as a scuba tank) – the volume is CONSTANT.  So, what would happen to the kinetic energy of the gas molecules in the container if you were to heat it up?  How would this affect pressure? Egg in a Bottle:! Video Clip States of Matter Interactive

48. Temperature – Pressure Relationships  Raising the Kelvin temperature of the gas will cause an INCREASE in the gas pressure.  WHY?  With increasing temperature, the K.E. of the gas particles increases – they move faster!  They collide more often and with more energy with the walls of the container…

50. Objective 5  Understand the law of conservation of energy in terms of gaining and losing energy.

51. Law of Conservation of Energy  Energy is neither created nor destroyed.

52. Energy  Potential Energy(PE) = energy of position; stored energy Kinetic Energy(KE) = energy related to motion

53. Example: a toy “wind-up” car  When you wind up the car and tighten the spring, you supply the car with potential energy.  As the car moves, the spring unwinds, and provides kinetic energy to the moving parts.  In this example, POTENTIAL ENERGY has been converted to KINETIC ENERGY.

55. Example: a snowboarder poised at the top of a hill  At the top of the hill, the snowboarder has potential energy (energy of position)  as the snowboarder races down the hill, the stored PE is converted to KE

58. Which type of energy is shown in each picture?

59. How does this apply to Chemical Energy? Energy stored within chemical bonds is stored chemical energy (a form of PE).

60. Chemical Energy When a fuel is burned, these chemical bonds break; reactant atoms reorganize to form new bonds; the products formed have more stable atom arrangements, and; some of the energy stored in the reactants is released in the form of heat and light.

62.  Example: the burning of methane gas (CH 4 ) – the gas we use in our laboratory burners!

63. Example: Methane, CH 4 CH 4 + 2O 2  CO 2 + 2H 2 O + ENERGY

64. Chemical reactions or changes may be: ENDOTHERMICorEXOTHERMIC

65. Endothermic Change – an energy-requiring process; absorbs energy  Example: breaking chemical typically requires an input of energy

66.  EXOTHERMIC CHANGE: an energy-releasing process; gives off energy  Example: the formation of stable, chemical bonds

67. Endothermic or Exothermic?  We compare the input of energy to break the bonds (start the reaction) to the output of energy when new bonds are formed

68. **IF…  more energy is given off than was added, the process is EXOTHERMIC;  more energy is added than is given off, the process is ENDOTHERMIC.

69. Graph of Exothermic Reaction : REACTANTS PRODUCTS

70. Graph of Endothermic Reaction: REACTANTS PRODUCTS