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States of Matter Chemistry 1-2 Enriched Mr States of Matter Chemistry 1-2 Enriched Mr. Chumbley Modern Chemistry: Chapter 10 THE KINETIC-MOLECULAR THEORY OF MATTER Section 1 p. 310 – 314 Kinetic-Molecular Theory • It is generally not possible to study individual particles, so scientists study the behavior of large groups of particles • The kinetic molecular theory was developed in the late 19th century to account for the different ways matter behaves as a solid, liquid, or gas • The kinetic-molecular theory explains that the behavior of physical systems depends on the combined actions of the molecules constituting the system KMT and Gases • Kinetic-molecular theory can be used to describe the properties and behavior of all states of matter, but is well demonstrated by gases • An ideal gas is a hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory Kinetic-Molecular Theory of Gases • Gases consist of large numbers of tiny particles that are far apart relative to their size • Collisions between gas particles and between particles and container walls are elastic collisions • An elastic collision is one in which there is no net loss of kinetic energy • Gas particles are in continuous, rapid, random motion and possess kinetic energy • Kinetic energy is the energy of motion • There are no forces of attraction between gas particles • The temperature of a gas depends on the average kinetic energy of the particles of the gas Properties of Gases • Expansion • Gases do not have a definite shape or volume, but rather completely fill their container • Fluidity • Gas particles slide easily past one another • Fluids are substances whose particles can easily flow around one another • Low Density • The density of a gas at atmospheric pressure is about 1/1000 the density of the same element in its liquid or solid state Properties of Gases • Compressibility • Since gases particles exist very far apart, they can be compressed together to be closer together • Diffusion and Effusion • Diffusion is the spontaneous mixing of particles of two substances caused by random motion • Effusion is the process by which gas particles pass through a tiny opening Real Gases • While the properties of ideal gases are specifically described, they do not really exist • A real gas, is a gas that does not behave completely according to the assumptions of the kinetic-molecular theory LIQUIDS Section 2 p. 315 – 318 Properties of Liquids • Fluidity • Liquid particles also slide easily past one another • Fluids are substances whose particles can easily flow around one another and take the shape of its container • High Density • The difference in density between solids and liquids is very small, usually within about 10% • Relative Incompressibility • Exerting pressures on liquids results in a very small change in volume Properties of Liquids • Diffusion • Like gases, liquids can diffuse and equally mix • Diffusion in liquids is much slower than in gases • Surface Tension • Surface tension is a force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing the surface are a to the smallest possible size • Capillary action is the attraction of the surface of a liquid to the surface of a solid Liquid to Other States • Vaporization is the process by which a liquid or solid changes to a gas • Evaporation is the process by which particles escape from the surface of a nonboiling liquid and enter the gas state • Boiling is another way to turn a liquid into a gas, but will be discussed in more detailed later • Freezing, or solidification, is the physical change of a liquid to a solid by the removal of energy as heat SOLIDS Section 3 p. 319 – 323 Properties of Solids • Most of the properties of solids are a result of the fact that the particles in a solid hold relatively fixed positions • There are two types of solids • Crystalline solids are those that consist of crystals • A crystal is a substance in which the particles are arranged in an orderly, geometric, repeating pattern • Amorphous solids are those in which the particles are arranged randomly Properties of Solids • Solids can be described as having a definite shape and volume • Solids also have a definite melting point • Melting is the physical change of a solid into a liquid by the addition of heat • The melting point is the temperature at which a solid becomes a liquid • Solids have a high density and incompressibility CHANGES OF STATE Section 4 p. 324 – 330 Changes of State Gas freezing Solid Liquid melting Changes of State and Equilibrium • A phase is any part of a system that has uniform composition and properties • Equilibrium is a dynamic condition in which two opposing changes occur at equal rates in a closed system • Changes of phase at equilibrium occur at equal rates • Equilibrium vapor pressure is the pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature • Volatile liquids are liquids that evaporate quickly due to relatively weak forces of attraction between their particles Boiling and Condensation • Equilibrium vapor pressure is used to explain and define boiling • Boiling is the conversion of a liquid to a vapor within the liquid as well as at its surface • The boiling point is the temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure • Condensation is the process by which a gas changes into a liquid Sublimation and Deposition • Sublimation is the change of state from solid directly to gas • Deposition is the change of state from gas directly to a solid Freezing and Melting • Freezing is the physical change of liquid to solid • The freezing point is the temperature at which the solid and liquid are in equilibrium at standard atmospheric pressure • Melting is the physical change of a solid into a liquid by the addition of heat • The melting point is the temperature at which a solid becomes a liquid Phase Diagrams • A phase diagram is a graph of pressure versus temperature that shows the conditions under which phases of a substance exist • The triple point of a substance indicates the temperature and pressure conditions at which the solid, liquid, and vapor of the substance can exist at equilibrium • The normal melting point is the temperature at which a substance turns from a solid to a liquid at 1 atm of pressure • The normal boiling point is the temperature at which a substance turns from a liquid to a gas at 1 atm of pressure • The critical point indicates the point at which a phase diagram ends and a substance’s liquid and gas phase are indistinguishable Phase Diagram Molar Enthalpy • When a substance changes from one state to another, there is a measurable amount of energy required to do so • This energy is typically in the form of heat, and it causes a change in the internal energy of the substance • Molar enthalpy of vaporization (ΔHv) is the amount of energy as heat that is needed to vaporize one mole of liquid at the liquid’s boiling point at constant pressure • Molar enthalpy of fusion (ΔHf) is the amount of energy as heat that is required to melt one mole of a solid at the solid’s melting point Molar Enthalpy • Molar enthalpy of vaporization indicates both the amount of energy absorbed for vaporization and the amount of energy released for condensation • Molar enthalpy of fusion indicates both the amount of energy absorbed for melting and the amount of energy released for freezing • Molar enthalpies of sublimation also exist, and indicate the amount of energy absorbed for sublimation and the amount of energy released for deposition Sample Problem 10A (p. 333) How much energy is absorbed when 47.0 g of ice melts at standard temperature and pressure (STP)? How much energy is absorbed when the same mass of liquid water boils? ΔHv = 40.65 kJ/mol ΔHf = 6.01 kJ/mol Specific Heat • Specific heat (c) is the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius, or one Kelvin 푞 푐 = 푚∆푇 • Specific heat is a property of any substance, just like molar enthalpy • However, finding the energy gained or lost through a temperature change is often more useful 푞 = 푐푚∆푇 Heating Curve Sample 16A (p. 503-504) • A 4.0 g sample of glass is heated from 1˚C to 41 ˚C, and was found to have absorbed 32 J of heat. • What is the specific heat of this type of glass? • How much energy will the same glass sample gain when it is heated from 41 ˚C to 71 ˚C? WATER Section 5 p. 331 – 333 Properties of Water and Hydrogen Bonds • Many of the properties of water are a result of the structure of its molecules and the hydrogen bonding between them • Substances with similar sizes and masses to water are generally gases at room temperature, because they lack hydrogen bonds • The particularly strong intermolecular forces of water also explain why water has high surface tension and capillary action • The hydrogen bonds also cause ice to form rigid, hexagonal structures as liquid water freezes • This is why ice is less dense then liquid water Properties of Water and Hydrogen Bonds Molar Enthalpy of Water • The molar enthalpy of fusion for water is high compared to other chemicals with similar molecular masses • This is related to the hydrogen bonding causing a more rigid structure • The molar enthalpy of vaporization of water (and boiling point) is also high compared to chemicals with similar molecular masses • This is because a significant amount of energy is required to break the hydrogen bonds of liquid water • These properties allow water to be a great regular of temperature because large quantities of water are large stores of heat.
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