Unit 6 – States of Matter

1. Conceptual Chemistry Unit 6 – States of Matter

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

3. Solids • Definite shape • Definite volume • Particles are vibrating and packed closetogether. • The particles do not flow.

4. Crystalline Solids • Particles are arranged in an organized pattern. Example: Diamond

5. Amorphous Solids • Particles are not organized in an orderly fashion. Example: Glass

6. 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.

7. Viscosity • A liquid’s resistance to flow.

8. 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

9. Plasma • High energy matter • A common example is the sun. • Super high energy gas particles that lostelectrons. • Plasma is the mostcommon form of matter in the Universe.

10. States of Matter

11. Objective 2 • Describe states of matter using the kinetic molecular theory.

12. Kinetic Molecular Theory • can explain the behavior of matter in its different states. Kinetic Molecular Theory: Explains 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.

13. Kinetic Energy and Kelvin Temperature Temperature: the average kinetic energy of the particles in a material • As particles are heated, they absorb energy, thus increasing their average kineticenergy and their temperature. • Motion stops at absolutezero (0 Kelvin). • Kelvin temperature scale reflects the relationship between temperature and average kinetic energy. It is directlyproportional.

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

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

16. Melting • Solid  Liquid Example 1 Example 2

17. Freezing • Liquid  Solid Example 1

18. Evaporation/Boiling • Liquid  Gas Example 1

19. Condensation • Gas  Liquid Example

20. Sublimation • Solid  Gas Example Opposite of Sublimation? Deposition Example

21. Objective 4 • Describe the different variables that define a gas.

22. Kinetic Theory of Gases • Gases are mostly emptyspace. • The molecules in a gas are separate, very small, and very far apart.

23. Kinetic Theory of Gases • Gas molecules are in constant, chaotic motion. • Collisions between gas molecules are elastic (there is no energy gain or loss).

24. Kinetic Theory of Gases • 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.

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

26. Variables that Describe a Gas • Volume: measured in L, mL, cm3 (1 mL = 1 cm3) • 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)

27. Moderate Force (about 100 lbs) Small Area (0.0625 in2) P = F /A Enormous Pressure (1600 psi)

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

29. Units of Pressure • 1 atm = 760 mm Hg • 1 atm = 760 torr • 1 atm = 1.013 x 105 Pa • 1 atm = 101.3 kPa

30. Boyle’s Law • For a given number of molecules of gas at a constant temperature, the volume of the gas varies inversely with the pressure. • As P, V and vice versa…. • Inverse relationship • P1V1 = P2V2

31. Boyle’s Law and Kinetic Molecular Theory How does kinetic molecular theory explain Boyle’s Law? • Gas molecules are in constant, randommotion. • 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!

33. 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?

34. Equation for Boyle’s Law • P1V1 = P2V2 where: P1 = initial pressure V1 = initial volume P2 = final pressure V2 = final volume

35. P1V1 = P2V2 • Using the same syringe example, just “plug in” the values: P1V1 = P2V2 (1 atm) (8 mL) = (P2) (10 mL)

36. P1V1 = P2V2 (1 atm) (8 mL) = P2 (10 mL) P2 = 0.8 atm

37. 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) • P1V1 = P2V2 • (1.2 atm)(12 L) = (3.6 atm)V2 • V2 = 4.0 L

38. 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) • P1V1 = P2V2 • (200 kPa)(28 L) = (P2)(17 L) • P2 = 329 kPa

39. 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.

40. 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 1, Clip 2

41. 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.

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

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

44. 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? • V1 = 117 mL; T1 = 100 + 273 = 373 K • V2 = 234 mL; T2 = ??? • V1 / T1= V2 / T2 • T2 = 746 K

45. 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? • V1 = 65 mL; T1 = 28.8 + 273 = 301.8 K • V2 = ??? mL; T2 = 72.2 + 273 = 345.2 K • V1 / T1= V2 / T2 • V2 = 74.3 mL

46. Temperature – Pressure Relationships • Picture a closed, rigid container of gas (such as a scuba tank) – the volume is CONSTANT. • 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? States of Matter Interactive Egg in a Bottle:! Video Clip

47. 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.