Chapter 10

1. Chapter 10 States of Matter

2. Section 1: The Kinetic-Molecular Theory of Matter • Based on the idea that particles of matter are always in motion. • Can be used to explain the properties of solids, liquids, and gases in terms of the energy of particles and the forces that act between them.

3. The Kinetic-Molecular Theory of Gases • The theory provides an understanding of the behavior of ideal gas molecules and the physical properties. • Ideal Gas- a hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory.

4. The Kinetic-Molecular Theory of Gases • Gases consist of large numbers of tine particles that are far apart relative to their size • Gases are much farther apart than molecules of liquids or solids • Most of the volume occupied by a gas is empty space • Gases have lower density than liquids and solids • They are easily compressed

5. The Kinetic-Molecular Theory of Gases 2. Collisions between gas particles and between particles and container walls are elastic collisions (no loss of total kinetic energy as long as the temperature is constant). • Kinetic energy is transferred between two particles during collisions.

6. The Kinetic-Molecular Theory of Gases • Gas particles are in continuous, rapid, random motion. They therefore possess kinetic energy, which is energy of motion. • There are no forces of attraction between gas particles, except near the temperature at which the gas condenses and becomes a liquid. • The temperature of a gas depend on the average kinetic energy of the particles of the gas. • Energy increases as temperature increases • Energy decreases as temperature decreases

7. Continued • Remember the Kinetic-Molecular Theory only applies to IDEAL GASES • They are not found in nature; however, there are many gases that behave nearly ideally

8. Expansion • Gases do not have a definite shape or volume. They fill the container. • If a gas is transferred from a 2-L container to a 3-L container, the gases will expand to fill the entire 3-L container.

9. Fluidity • Gas particles glide easily past one another. • This ability to flow causes gases to behave as liquids do; therefore, liquids and gases can be referred to fluids.

10. Low Density • The particles are so much farther apart in the gas state causing low density.

11. Compressibility • During compression, the gas particles are pushed closer together. The volume can be greatly decreased.

12. Diffusion and Effusion • Gases spread out and mix with one another, even without being stirred (the air: not just oxygen) • Diffusion-spontaneous mixing of two substances caused by their random motion (being able to smell perfume in the air after someone sprayed it) • Effusion- process by which gas particles pass through a tiny opening (smelling perfume when inside a balloon)

13. Real Gases vs Ideal Behavior • A real gas is a gas that does not behave completely according to the assumptions of the kinetic-molecular theory. • Noble gases and nonpolar gases react more the ideal gases

14. Section 2: Liquids • Yes, I know it sounds crazy, but liquids are the least common state of matter in the UNIVERSE. (On Earth they are extremely common) • This is due to the fact that liquids only exist at a very small range of temperatures

15. Properties of Liquids and the Kinetic-Molecular Theory • Definite volume and takes the shape of the container. • Liquid particles are in constant motion. • They are closer together causing them not to be able to move as fast as the gas particles. (more attractive forces) • Fluids(liquids & gases)-a substance that can flow and therefore take the shape of the container

16. Properties of Liquids and the Kinetic-Molecular Theory • Most liquids naturally flow downhill because of gravity (waterfalls) • Some liquids can flow in other directions. • Liquid helium near absolute zero has the unusual property of being able to flow uphill.

17. Relatively High Density • Most substances are hundreds of times denser in the liquid state than gas state • The higher the density, the closer the arrangement of particles (higher the attraction). • Most substances are only about 10% less dense as a liquid than a solid. • Water is one of the few substances that becomes less dense as a solid.

18. Density Column • The most dense on the bottom. • The least dense on the top

19. Relative Incompressibility • Liquids cannot be compressed very much at all due to the closeness of particles.

20. Ability to Diffuse • Liquids can diffuse (food coloring in water) • Diffusion happens much slower in liquids • Higher temps increase rate of diffusion • Lower temps decreases the rate of diffusion • ***Again Kinetic Energy-energy of motion

21. Surface Tension • A force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size. • Higher attraction, higher surface tension • Water is higher than most liquids **Surface tension-Penny

22. Surface Tension • Capillary action- the attraction of the surface of a liquid to the surface of a solid • Paper Chromatography (Y’all did this in the 8th grade. You used markers and you made dots on a piece of chromatography paper. Then put the edge of paper in the water. It separated the ink into the colors it was made of.) • Plant/Tree roots getting water • The meniscus to form on a graduation cylinder

23. Evaporation and Boiling • Vaporization- the process by which a liquid or solid changes to a gas • Boiling – change of liquid to bubbles of vapor which then go to a gas. • Evaporation- the process by which particles escape from the surface of a nonboiling liquid and enter the gas state. • All water that falls to the Earth (rain, sleet, snow, etc..) if from the evaporation of fresh water from the ocean • Evaporation of perspiration keeps you cool.

24. Formation of Solids • When liquids cool enough for the attractive forces to become great enough to have order. This reduces the kinetic energy. • Freezing (solidification) - physical change of a liquid to a solid by the removal of heat. • Liquid water to ice • Liquid Wax (burning a candle) to paraffin

25. Section 3: Solids • Definite shape and volume • Particles are packed more closely than in liquids (stronger forces) • Constant motion (just can move very much)

26. Crystalline solids & Amorphous solids • Crystalline- consist of crystals (diamonds, quartz, salt, ice) • Amorphous- particles are randomly arranged (plastics, glass)

27. Definite Melting Point • Melting Point- the temperature at which a solids changes to a liquids by the addition of heat. • The melting point (solid to liquid) and the freezing point (liquid to solid) are the same temperature. • This point can be used to identify substances.

28. High Density & Incompressibility • In general substances are most dense in the solid state (exception: water) • Higher density is due to closeness of particles • Considered to incompressible

29. Low Rate of Diffusion • Diffusion can occur. However, it is extremely slow

30. Section 4: Changes of State • Phase (or state)- any part of a system that has uniform composition and properties. • Look at Table 2 on page 342

31. Equilibrium • Dynamic condition in which two opposing changes occur at equal rates in a closed system. • Condensation is happing at the same rate as evaporation therefore no mass change

32. Volatile Liquids • Liquids that evaporate readily due to weak forces of attraction. • Ether • Rubbing alcohol Nonvolatile liquids do not evaporate readily -water

33. Boiling • Conversion of a liquid to a vapor within the liquid as well as at its surface. • Boiling point- the temperature at which this conversion happens • The lower the pressure, the lower the boiling point • This is why on the back of a cake mix it says if you are at a certain elevation, then you need to bake it at a lower temp. • Pressure cookers increase pressure; therefore they boil and cook quicker

34. Energy and Boiling • Energy must be added continuously in order to keep a liquid boiling • The temperature at the boiling point remains constant despite the continuous addition of energy. The temp will change when the conversion has finished.

35. Molar Enthalpy of Vaporization • The amount of energy as heat that is needed to vaporize one mole of liquid at the liquid’s boiling point at constant temperature. • The stronger the attraction of molecules the more energy will be needed to break those bonds. • Water has an extremely high molar enthalpy. This is why it is such a great coolant.

36. Freezing and Melting • Freezing involves a loss of energy • Melting involves absorbing (gaining energy) • Molar enthalpy of fusion- the amount of energy as heat required to melt on mole of slid at the solid’s melting point

37. Sublimation and Deposition • Sublimation- solid going to gas phase, skipping the liquid phase • Dry ice • Iodine • Ice can but slowly (snow disappearing even when it has gotten above freezing) • Deposition- gas to solid phase, skipping the liquid phase • Frost

38. Phase Diagrams: Look on page 347 in your book • Triple point- indicates the temp and pressure in which a solid, liquid, and gas can all be present • Critical point- critical temp and pressure • Critical temp- above this temp the substance can’t exist as a liquid regardless of pressure • Critical pressure- the lowest pressure at which the substance can exist as a liquid at the critical temp.

39. Section 5: Water • On Earth, water is the most abundant liquid • Makes up about 75 % of Earth’s surface • Water is about 70-90% of all living things mass

40. Physical Properties of Water • At room temperature: it is a liquid, transparent, odorless, tasteless, almost colorless • Freezes and Melts at 0oC • Water expands as it freezes • Solid water is less dense than liquid water • Boils at 100oC • Molar enthalpy of fusion is 6.009 kJ/mol • Molar enthalpy of vaporization is 40.79 kJ/mol

41. Example • What quantity of energy is released when 506 grams of liquid water freezes? • What mass of steam is required to release 4.97 x 105 kJ of energy on condensation?