Chapter 13

 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

 Kinetic Theory and a Model for Gases ◦ 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. 13.1

 Kinetic Theory and a Model for Gases ◦ What are the assumptions of the kinetic theory as it applies to gases? 13.1

 Assumptions: 1.Gases contain particles in constant, random, straight-line motion. 2.When gas particles collide with one another and the walls of the container energy is transferred but there is no net loss of energy: all collisions are elastic. 3.Gas particles are very far apart, therefore the volume of a sample of gas is the volume of the container: gas particles have negligible volume. 4.Gas particles have no attraction for each other. 13.1

 Particles in a gas are in rapid, constant motion. 13.1

 Gas particles travel in straight-line paths. 13.1

 Gases fill their container. 13.1

 Gas pressure is measured as the force exerted by a gas per unit surface area of the container.  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. 13.1

 Gas Pressure ◦ How does kinetic theory explain gas pressure?  Gas pressure is the result of simultaneous collisions of billions of rapidly moving gas particles with the walls of the container. 13.1

 A barometer is a device that is used to measure atmospheric pressure. 13.1

◦ The SI unit of pressure is the pascal (Pa).  Other units are the atmosphere (atm), psi (pounds per square inch), and millimeters of mercury (mmHg)  One standard atmosphere (atm) is the pressure required to support 760 mm of mercury in a mercury barometer at 25°C. 13.1

 Kinetic Energy and Temperature ◦ What is the relationship between the temperature in Kelvin and the average kinetic energy of particles?  The Kelvin temperature of a substance is directly proportional to the average kinetic energy of the particles of the substance.  As temperature increases, average kinetic energy increases. The size of the sample is irrelevant 13.1

◦ Absolute zero (0 K, or –273.15°C) is the temperature at which the motion of particles theoretically ceases. 13.1

 Gas samples are often referred to at STP.

 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

 A Model for Liquids ◦ What factors determine the physical properties of a liquid?  Kinetic energy and intermolecular forces of attraction. 13.2

◦ The intermolecular forces of attraction are strong enough to hold the particles of a liquid close together, but weak enough to allow them to flow past one another.

 Evaporation ◦ What is the relationship between evaporation and kinetic energy?  During evaporation, those molecules with “above average” kinetic energy can overcome IMFs and escape from the surface of the liquid.  Increasing the average kinetic energy will increase the rate of evaporation. 13.2

 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. 13.2

 In an open container, molecules that evaporate can escape from the container.

 In a closed container, the molecules cannot escape. They collect as a vapor above the liquid. Some molecules condense back into a liquid. 13.2

 Vapor Pressure ◦ When can a dynamic equilibrium exist between a liquid and its vapor?  When the liquid is in a closed container (system). 13.2

◦ In a closed 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

◦ 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 required kinetic energy necessary to escape the surface of the liquid. 13.2

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

 Manometer 13.2

 Boiling Point ◦ Under what conditions does boiling occur?  When a liquid is heated to a temperature at which particles throughout the liquid have enough kinetic energy to overcome all IMFs and vaporize, the liquid begins to boil. 13.2

 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). 13.2

◦ Boiling Point and Pressure Changes  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

 Altitude and Boiling Point 13.2

◦ “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 kPa. 13.2

Ref. table H

 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

 A Model for Solids ◦ How are the structure and properties of solids related?  Particles of a solid have such strong IMFs that they arrange themselves in a regular geometric pattern – a crystal.  The general properties of solids reflect the orderly arrangement of their particles and the fixed locations of their particles. 13.3

◦ When the kinetic energy of the particles increases enough they will overcome some of the IMFs and begin to slide past one another. The solid melts.  This destroys the crystal lattice structure.  The melting point (mp) is the temperature at which a solid changes into a liquid.  In a closed system at the melting point a dynamic equilibrium exists between a solid and liquid. 13.3

 Can elements exist in more than one form? ◦ Allotropes are two or more different molecular forms of the same element in the same physical state.  Allotropes have different properties because their structures are different.  Only a few elements have allotropes: carbon, oxygen and phosphorus 13.3

 Carbon Allotropes 13.3

 Do all solids have a crystalline structure?  An amorphous solid lacks an ordered internal structure.  Rubber, plastic, asphalt, and glass are amorphous solids.  A glass is a transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing. 13.3

 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. 13.4

 Sublimation ◦ When can sublimation occur?  Sublimation occurs in solids with vapor pressures that exceed atmospheric pressure at or near room temperature. 13.4

 The change of a substance from a solid to a vapor without passing through the liquid state is called sublimation.  These solids have weak IMFs. 13.4

 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 (deposition). 13.4