2. Chapter 13: States of Matter 13.1 The Nature of Gases kinetic molecular theory (KMT), gas pressure (pascal, atmosphere, mm Hg), kinetic energy 13.2 The Nature of Liquids vaporization, evaporation, vapor pressure, boiling point 13.3 The Nature of Solids do NOT need to know crystalline structures 13.4 Changes of State sublimation, phase diagrams

3. 13.1 The Nature of Gases The skunk releases its spray! Within seconds you smell that all-too-familiar foul odor. You will discover some general characteristics of gases that help explain how odors travel through the air, even on a windless day. 13.1

4. What is KMT? * Based on the research of Robert Boyle (1627 – 1691) A theory that envisions molecules in motion Best describes properties and behaviors of gases * Imagines particles of a gas like billiard balls, moving and crashing into each and the walls of a container in a three-dimensional space.

5. Kinetic Theory and a Model for Gases Kinetic Molecular Theory (as stated in Prentice Hall text): »The particles in a gas are considered to be small, hard spheres with an insignificant volume. »The motion of the particles in a gas is rapid, constant, and random. »All collisions between particles in a gas are perfectly elastic. 13.1 The word kinetic refers to motion. –The energy an object has because of its motion is called kinetic energy. –According to the kinetic theory, all matter consists of tiny particles that are in constant motion.

6. Kinetic Theory and a Model for Gases a) Particles in a gas are in rapid, constant motion. b) Gas particles travel in straight-line paths c) The gas fills the container 13.1

7. Principles of KMT (as stated in Hein-Arena text) 1. Gases consist of tiny (submicroscopic) particles.

8. 2. Gas particles are very far apart; the volume occupied by a gas consists mostly of empty space (typically about 99.9% empty). 3. Gas particles are not attracted to each other (like ions are.)

9. 4. Gas particles move randomly in all directions, traveling in straight lines. * Gas molecules travel at very high speeds, about 6000 km/hr (~2700 mi/hr). http://www.chm.davidson.edu/ChemistryApplets/KineticMolecularTheory/BasicConcepts.html 5. The higher the average speed of the particles, the higher the temperature of that substance.  Absolute temperature: T based on average particle speed; measured in Kelvin (K)  Absolute zero (0 K): particles are not moving; there is no T below 0 K. http://www.chm.davidson.edu/ChemistryApplets/KineticMolecularTheory/PT.html K = °C + 273

10. 6. Gas particles collide with each other and with the walls of the container without losing energy. Pressure = the collision of gas particles with a surface

11. Gas Pressure Gas pressure results from the force exerted by a gas per unit surface area of an object. –An empty space with no particles and no pressure is called a vacuum. Atmospheric pressure results from the collisions of atoms and molecules in air with objects. Gas pressure is the result of simultaneous collisions of billions of rapidly moving particles in a gas with an object. 13.1

12. Gas Pressure A barometer is a device that is used to measure atmospheric pressure. 13.1

13. Gas Pressure –The SI unit of pressure is the pascal (Pa). –One standard atmosphere (atm) is the pressure required to support 760 mm of mercury in a mercury barometer at 25°C. 13.1

14. Kinetic Energy and Temperature 13.1 The particles in any collection of atoms or molecules at a given temperature have a wide range of kinetic energies. Most of the particles have kinetic energies somewhere in the middle of this range.

15. Kinetic Energy and Temperature Absolute zero (0 K, or –273.15°C) is the temperature at which the motion of particles theoretically ceases. Particles would have no kinetic energy at absolute zero. Absolute zero has never been produced in the laboratory. Average Kinetic Energy and Kelvin Temperature The Kelvin temperature of a substance is directly proportional to the average kinetic energy of the particles of the substance. 13.1

16. 13.1 Section Quiz. 1. According to the kinetic theory, the particles in a gas a)are attracted to each other. b)are in constant random motion. c)have the same kinetic energy. d)have a significant volume.

17. 13.1 Section Quiz. 2. The pressure a gas exerts on another object is caused by a)the physical size of the gas particles. b)collisions between gas particles and the object. c)collisions between gas particles. d)the chemical composition of the gas.

18. 13.1 Section Quiz. 3. The average kinetic energy of the particles in a substance is directly proportional to the a)Fahrenheit temperature. b)Kelvin temperature. c)molar mass of the substance. d)Celsius temperature.

19. 13.2 The Nature of Liquids Hot lava oozes and flows, scorching everything in its path, and occasionally overrunning nearby houses. When the lava cools, it solidifies into rock. The properties of liquids are related to intermolecular interactions. You will learn about some of the properties of liquids. 13.2

20. A Model for Liquids Substances that can flow are referred to as fluids. Both liquids and gases are fluids. 13.2

21. A Model for Liquids –The interplay between the disruptive motions of particles in a liquid and the attractions among the particles determines the physical properties of liquids. 13.2 What factors determine the physical properties of a liquid?

22. 13.2 Evaporation In an open container, molecules that evaporate can escape from the container. –The conversion of a liquid to a gas or vapor is called vaporization. –When such a conversion occurs at the surface of a liquid that is not boiling, the process is called evaporation. In a closed container, the molecules cannot escape. They collect as a vapor above the liquid. Some molecules condense back into a liquid

23. Evaporation –During evaporation, only those molecules with a certain minimum kinetic energy can escape from the surface of the liquid. 13.2 What is the relationship between evaporation and kinetic energy?

24. Vapor Pressure Vapor pressure is a measure of the force exerted by a gas above a liquid. 13.2

25. Vapor Pressure –In a system at constant vapor pressure, a dynamic equilibrium exists between the vapor and the liquid. The system is in equilibrium because the rate of evaporation of liquid equals the rate of condensation of vapor. 13.2 When can a dynamic equilibrium exist between a liquid and its vapor?

26. Vapor Pressure Vapor Pressure and Temperature Change –An increase in the temperature of a contained liquid increases the vapor pressure. –The particles in the warmed liquid have increased kinetic energy. As a result, more of the particles will have the minimum kinetic energy necessary to escape the surface of the liquid. 13.2

27. Vapor Pressure –Vapor Pressure Measurements The vapor pressure of a liquid can be determined with a device called a manometer. 13.2

28. Boiling Point –When a liquid is heated to a temperature at which particles throughout the liquid have enough kinetic energy to vaporize, the liquid begins to boil. The temperature at which the vapor pressure of the liquid is just equal to the external pressure on the liquid is the boiling point (bp). Because a liquid boils when its vapor pressure is equal to the external pressure, liquids don’t always boil at the same temperature. –At a lower external pressure, the boiling point decreases. –At a higher external pressure, the boiling point increases. 13.2 Under what conditions does boiling occur?

29. Boiling Point Altitude and Boiling Point 13.2

30. Boiling Point Normal Boiling Point Because a liquid can have various boiling points depending on pressure, the normal boiling point is defined as the boiling point of a liquid at a pressure of 101.3 kPa. 13.2

31. 13.2 Section Quiz 1. In liquids, the attractive forces are a)very weak compared with the kinetic energies of the particles. b)strong enough to keep the particles confined to fixed locations in the liquid. c)strong enough to keep the particles from evaporating. d)strong enough to keep particles relatively close together.

32. 13.2 Section Quiz 2. Which one of the following is a process that absorbs energy? a)freezing b)condensation c)evaporation d)solidifying

33. 13.2 Section Quiz 3. In a sealed gas-liquid system at constant temperature eventually a)there will be no more evaporation. b)the rate of condensation decreases to zero. c)the rate of condensation exceeds the rate of evaporation. d)the rate of evaporation equals the rate of condensation.

34. 13.2 Section Quiz 4. Where must particles have enough kinetic energy to vaporize for boiling to occur? a)at the surface of the liquid b)at the bottom of the container c)along the sides of the container d)throughout the liquid

35. 13.2 Section Quiz 5. The boiling point of a liquid a)increases at higher altitudes. b)decreases at higher altitudes. c)is the same at all altitudes. d)decreases as the pressure increases.

36. 13.3 The Nature of Solids In 1985, scientists discovered a new form of carbon. They called this form of carbon buckminsterfullerene, or buckyball for short. You will learn how the arrangement of particles in solids determines some general properties of solids. 13.3

37. A Model for Solids The general properties of solids reflect the orderly arrangement of their particles and the fixed locations of their particles The melting point (mp) is the temperature at which a solid changes into a liquid. 13.3

38. Crystal Structure and Unit Cells In a crystal, the particles are arranged in an orderly, repeating, three-dimensional pattern called a crystal lattice. The shape of a crystal reflects the arrangement of the particles within the solid 13.3

39. Crystal Structure and Unit Cells Crystal Systems (for information only) A crystal has sides, or faces. Crystals are classified into seven crystal systems. 13.3

40. 13.3 Section Quiz 1. A solid will melt when a)all the particles have the same kinetic energy. b)bonds form between the particles. c)disruptive vibrations overcome attractive forces. d)attractions overcome disruptive vibrations.

41. 13.4 Changes of State Familiar weather events can remind you that water exists on Earth as a liquid, a solid, and a vapor. As water cycles through the atmosphere, the oceans, and Earth’s crust, it undergoes repeated changes of state. You will learn what conditions can control the state of a substance.

42. Sublimation The change of a substance from a solid to a vapor without passing through the liquid state is called sublimation. –Sublimation occurs in solids with vapor pressures that exceed atmospheric pressure at or near room temperature. When solid iodine is heated, the crystals sublime, going directly from the solid to the gaseous state. When the vapor cools, it goes directly from the gaseous to the solid state.

43. Phase Diagrams A phase diagram is a graph that gives the conditions of temperature and pressure at which a substance exists as solid, liquid, and gas (vapor). The triple point describes the only set of conditions at which all three phases can exist in equilibrium with one another

44. Six Phase Changes SOLID LIQUID GAS B.P. Temp (°C) M.P. C.P. F.P.

45. 13.4 Section Quiz. 1. Identify the change of state that occurs when solid CO 2 changes to CO 2 gas as it is heated. a)condensation b)freezing c)vaporization d)sublimation

46. 13.4 Section Quiz. 2. Sublimation occurs in solids if the vapor pressure at or near room temperature a)exceeds atmospheric pressure. b)equals atmospheric pressure. c)is less than atmospheric pressure. d)is less than half the atmospheric pressure.

47. 3. What is the significance of a line in a phase diagram? a)Only one phase is present. b)Two phases are in equilibrium. c)Three phases are in equilibrium. d)The distinction between two phases disappears. 13.4 Section Quiz.

48. 4. What is the significance of the triple point in a phase diagram? a)Temperature and pressure are equal. b)Two phases are in equilibrium. c)Three phases are in equilibrium. d)The distinction among three phases disappears. 13.4 Section Quiz.

49. Vocabulary Checklist 13.1– kinetic energy, kinetic molecular theory, gas pressure, atmospheric pressure, vacuum, barometer, pascal, standard temperature (1 atm=101.3 kPa= 760 mm Hg) 13.1 concept– average kinetic energy of particles is directly proportional to the Kelvin temperature of the substance 13.2– vaporization, evaporation, vapor pressure, boiling point, normal boiling point 13.3– melting point 13.4– sublimation, triple point 13.4 concept– phase diagram