THE BIG IDEA: KINETIC THEORY. 1. Use the kinetic-molecular theory to account for the physical properties of states of matter. (13.

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1 HONORS CHEMISTRY - CHAPTER 13 STATES OF MATTER OBJECTIVES AND NOTES - V15 NAME: DATE: PAGE: THE BIG IDEA: KINETIC THEORY Essential Questions 1. What factors determine the physical state of a substance? 2. What are the characteristics that distinguish gases, liquids, and solids? 3. How do substances change from one state to another? Chapter Objectives 1. Use the kinetic-molecular theory to account for the physical properties of states of matter. (13.1) 2. Relate temperature to the average kinetic energy of the particles in a substance. (13.1) 3. Describe the relationship among kinetic energy, interparticle forces, and molecular mass as they relate to physical states of matter. (13.1) 4. Use the kinetic theory to explain gas pressure. (13.1) 5. Convert between units of pressure of kpa, atm, inches Hg, mm Hg, and psi. (13.1) 6. State the values of standard temperature and pressure (STP). (13.1) 7. Explain the significance of absolute zero, giving its value in degrees Celsius, Kelvin, and Fahrenheit. (13.1) 8. Describe the nature of a liquid in terms of the attractive forces between its particles. (13.2) 9. Differentiate between evaporation and boiling of a liquid. (13.2) 10. Describe condensed states. (13.2) 11. Describe a system at equilibrium. (13.2) 12. Describe vapor pressure, and how it relates to changes in temperature and pressure. (13.2) 13. Differentiate between a vapor and a gas. (13.2) 14. Describe boiling point and normal boiling point. (13.2) 1

2 15. Describe volatile and nonvolatile liquids. (13.2) 16. Read and be able to interpret data found on vapor pressure vs. temperature graphs. (13.2) 17. Describe melting point and normal melting point and how they relate to freezing and normal freezing point. (13.2) 18. Read and be able to interpret data found on phase change graphs (heating curves and cooling curves). (13.3) 19. Describe crystalline substances. (7.2, 7.3 & 13.3) 20. Describe and give examples of allotropes. (13.3) 21. Relate the bonding structure in ionic, metallic, network, and molecular substances to the properties they exhibit. (7.2, 7.3 & 13.3) 22. Describe sublimation and deposition. (13.4) 23. Name and describe six possible phase changes that matter can undergo. (13.1 through 13.4) 24. Demonstrate and be able to describe all aspects of laboratory safety rules and procedures. (Applicable every chapter) 13.1 The Nature of Gases A. Kinetic Theory and a Model for Gases 1. kinetic energy (KE): Energy of motion. Examples: boulder falling; air molecules moving faster as they are heated by the burning coal; vibrational, rotational, and translational motions are all forms of kinetic energy. a. KE = 1 2 mv2 b. In gas laws capital "V" is an abbreviation for volume, however in the kinetic energy equation note that the lowercase "v" represents velocity. 2 - HC - Chapter 13 - Objectives and Notes - V15

3 2. Postulates of the kinetic molecular theory: a. Gas particles (atoms or molecules) act like hard spheres that have mass. b. The distance between gas particles is relatively large. 1. The volume of the actual gas particles is so small as compared to the volume of their container that the gas particles can be considered to having a relative volume of zero. c. Gas particles are in constant, rapid, random motion. d. Collisions between gas particles are completely elastic. 1. elastic collision: A collision in which the particles involved do not lose any kinetic energy. Gas particles can transfer their kinetic energy to other gas particles, but as a whole, they do not slow down. If gas particles collided with each other in inelastic collisions, the particles would lose their kinetic energy and they would eventually stop moving; gravity would then pull all to the ground. e. Gas pressure is caused by gas particles colliding with the walls of its container. f. The distance between gas particles is so large, and their intermolecular forces are so small that gas particles are considered to exert no attractive forces between themselves. 2. States of matter a. Solids 1. Definite volume. 2. Definite shape. 3. Ordered arrangement of particles. 4. Particles very close together. 5. High density. 6. Nearly incompressible. 7. Very slightly affected by conditions of pressure and temperature. 3 - HC - Chapter 13 - Objectives and Notes - V15

4 b. Liquids 1. Definite volume. 2. No definite shape; take the shape of the container. 3. Disordered particles. 4. Particles close together. 5. Intermediate density. 6. Very slightly compressible. 7. Diffuse slowly through other liquids. 8. Slightly affected by conditions of pressure and temperature. c. Gases 1. No definite volume; completely fill container. 2. No definite shape; take the shape of the container. 3. Extremely disordered particles. 4. Particles far apart. 5. Extremely low density. 6. Very compressible. 7. Diffuse rapidly. 8. Are greatly affected by conditions of pressure, temperature, and volume. 4. On the molecular level there are always competing forces that exist between kinetic energy and interparticle attractions. The force of the kinetic energy wants to make the particles move apart and the interparticle forces want to pull them together. The state of a substance depends on the temperature (average kinetic energy) of the substance, the strength of its interparticle forces, and its molar mass. a. Substances that have high molar mass and/or have high interparticle forces tend to be solids. b. Substances that have medium molar mass and/or have medium interparticle forces tend to be liquids. c. Substances that have low molar mass and/or have low interparticle forces tend to be gases. 4 - HC - Chapter 13 - Objectives and Notes - V15

5 B. Gas Pressure 1. pressure: Force per unit area exerted by moving particles. The pressure exerted by a substance is based on the number of collisions that occur and the force of those collisions. a. Pressure = Force Area b. Increasing the temperature of a substance increases its kinetic energy. Since the mass of the substance does not change, increasing temperature increases the velocity of a substance; this increases the pressure of the substance. 1. The increase in speed means that each collision occurs with more force. 2. The increase in speed also increases the number of collisions that occur. c. The pressure exerted by a gas depends on two factors. 1. The concentration of the gas, i.e., the number of gas particles per unit volume. 2. The average kinetic energy of the gas particles. 2. air pressure: The pressure exerted by the air in the atmosphere. The pressure is caused by the gravity of the earth attracting the gas particles. a. Air pressure in Denver is normally lower than air pressure at sea level because Denver's altitude is much higher than sea level, therefore, there is much less air over Denver being pulled to the ground by gravity than would occur at sea level. 3. standard pressure: 101 kilopascals (101 kpa) = 29.9 inches Hg = 760. mm Hg = 1.00 atmosphere (atm) = 14.7 pounds per square inch (14.7 psi). a. Note: A conversion factor for converting between any two units of pressure can be created using the above factors. Examples: 101 kpa 29.9 inches 760. mm Hg or, etc. 101 kpa 4. Standard temperature and pressure (STP) for gases is 0 o C and 1 atmosphere pressure. 5 - HC - Chapter 13 - Objectives and Notes - V15

6 5. barometer: A device used for measuring atmospheric pressure by measuring the height of the mercury column supported by the atmosphere. Barometer P Gas Sample (Atmosphere) = Height Mercury C. Kinetic Energy and Temperature 1. standard temperature: 0 o C = 273 K. a. K = o C b. o C = o F c. A temperature increase from 5 K to 10 K is a doubling in temperature and a doubling in heat content. A temperature increase from 5 o C to 10 o C or an increase from 5 o F to 10 o F is not (got that, is not) a doubling of heat content or kinetic energy. (Some students will ignore me and get this wrong on the test and the final, but you can make sure you get it right!) 6 - HC - Chapter 13 - Objectives and Notes - V15

7 13.2 The Nature of Liquids B. Evaporation 1. vapor: A term that represents the gaseous state of a substance that is a liquid or solid at normal conditions. a. When water evaporates it is referred to as water vapor. 2. gas: A term used for a substance that is a gas under normal conditions. a. Oxygen at standard conditions is referred to as a gas. 3. equilibrium/dynamic equilibrium (D ): A state in which two or more processes occur at the same rate so that no net change occurs in the system 4. New Jersey: A state in which the greatest chemistry teacher works. (Modesty prevents me from stating who that person happens to be). 5. vaporization: The process in which a liquid becomes a gas or vapor. a. The process is increased by an increase in temperature and/or a decrease in atmospheric pressure. b. It can take place above or below the boiling temperature. c. It is the reverse of condensation. d. evaporation: The process by which molecules on the surface of a liquid leave the liquid and become a gas 1. This process takes place below the boiling temperature. 2. The process is increased by an increase in temperature and/or a decrease in pressure. 6. condensation: The process by which gaseous molecules become a liquid. a. The process is increased by a decrease in temperature and/or an increase in pressure. b. It is the reverse of evaporation. C. Vapor Pressure 1. vapor pressure/equilibrium vapor pressure: The partial pressure of a vapor at the surface of its parent liquid; it is dependent on intermolecular forces and temperature. 2. volatile: A liquid that is easily vaporized; it has a high vapor pressure and therefore a low boiling point. Examples: alcohol and gasoline. 3. nonvolatile: A liquid that is not easily vaporized; it has a low vapor pressure and therefore a high boiling point. Examples: molasses and mercury. 7 - HC - Chapter 13 - Objectives and Notes - V15

8 D. Boiling Point See my website for a Vapor Pressure and Boiling Point Graph of various substances. 1. boiling point (condensation point): The temperature at which the vapor pressure of a liquid is equal to the atmospheric pressure. It is also the temperature at which the liquid and vapor exist at equilibrium. a. While evaporation occurs only at the surface of a liquid, boiling occurs throughout the entire liquid. 2. normal boiling point (normal condensation point): The temperature at which the vapor pressure of a liquid is equal to standard atmospheric pressure. 3. phase change graphs (phase change curves): A type of graph that shows the relationship between energy and temperature at a constant pressure as a sold changes to a liquid and then a gas/vapor (heating curve) or as a gas/vapor changes to a liquid and then a solid (cooling curve). See my website for examples of phase change graphs (Referred to as Heating and Cooling Graphs). Each example focuses in on specific relationships seen on phase change graphs The Nature of Solids A. A Model for Solids 1. melting point (freezing point): The temperature at which the solid and liquid exist at equilibrium under a given pressure. It is also the temperature at which the vapor pressures of the solid and liquid are equal. 2. normal melting point (normal freezing point): The temperature at which the solid and liquid exist at equilibrium at standard atmospheric pressure. It is also the temperature at which the vapor pressures of the solid and liquid are equal at standard atmospheric pressure. B. Crystal Structure and Unit Cells 1. allotropes: Two or more different molecular forms of the same element in the same physical state. Their properties are different because they have different (drum roll... ) bonding structures. a. Examples: Red and white phosphorus, diamond, graphite, and buckminsterfullerene, oxygen gas and ozone (O 3 ). 8 - HC - Chapter 13 - Objectives and Notes - V15

9 There Should Have Been a C. Types of Solids 1. Ionic a. Particles of unit cell: anions and cations. b. Strongest intraparticle/ interparticle forces: ionic bonds. c. Properties: In nature, found as solids at room temperature. 1. Hard but brittle, very poor electrical conductors (excellent electrical conductors in the molten form or dissolved in water). 2. Melting points: usually very high, d. Examples: NaCl, CaBr 2, K 2 SO Metallic a. Particles of unit cell: cations (surrounded by a "sea" of electrons). b. Strongest intraparticle/interparticle forces: metallic bonds. c. Properties: Found as crystalline solids at room temperature (Hg is weird exception, don t worry about it). 1. Soft to very hard, malleable, ductile, good heat and electrical conductors. d. Examples: Li, K, Fe, and Cu. 3. Macromolecules a. Particles of unit cell: atoms. b. Strongest intraparticle/interparticle forces: covalent. c. Properties: Found as solids at room temperature. d. Examples: C (diamond), SiO 2 (sand) 4. Molecular a. Particles of unit cell: molecules, atoms (monatomic molecules), usually composed of nonmetals. b. Strongest intraparticle forces: covalent. c. Strongest interparticle forces: Van der Waals. d. Properties: Found as gases, liquids, and soft solids at room temperature. e. Examples: He, O 2, I 2, S 8, P 4, H 2 O, and CO 2 f. Modified molecules: A combination of both molecular and network. In this type of solid the layers are macromolecules, i.e., the forces holding each individual layer together are covalent. The individual layers then attach to each other in a molecular fashion, that is, the forces holding one layer to another layer are Van der Waals forces. Graphite is the classic example of a modified molecule. 9 - HC - Chapter 13 - Objectives and Notes - V15

10 13.4 Changes of State A. Sublimation 1. sublimation: The process of a solid changing directly to a gas without first becoming a liquid; the process is increased by an increase in temperature and/or a decrease in pressure. a. The opposite of deposition. 2. deposition: The process of a gas changing directly to a solid without first becoming a liquid. a. The opposite of sublimation b. Substances that sublime/deposit are usually molecular solids that have weak Weak Forces. c. Examples of substances that sublime/deposit: dry ice, solid room air fresheners, mothballs, and I HC - Chapter 13 - Objectives and Notes - V15