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Chapter 13:. States of Matter. 13. 1 Gases. Kinetic Molecular Theory—attempts to explain the properties of gases. Assumes: Particles are small and separated by empty space. No significant attractive or repulsive forces. Particles are in constant, random motion.
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Chapter 13: States of Matter
13. 1 Gases • Kinetic Molecular Theory—attempts to explain the properties of gases. • Assumes: • Particles are small and separated by empty space. • No significant attractive or repulsive forces. • Particles are in constant, random motion. • Collisions are elastic (no kinetic energy lost) • Kinetic energy determined by KE = ½ mv2.
Temperature • Although all particles of the same gas have the same mass, they don’t have the same velocity or kinetic energy. • Temperature is a measure of the average kinetic energy of the particles. • At a given temperature, all gases have the same kinetic energy.
Behavior of Gases • KM theory explains the behavior of gases. • Low density: Due to the large amount of space between the particles in the gas. • Compression and expansion: Gases expand to fill their containers. Becomes more dense in a smaller container; less dense in a larger container. • Diffusion: As gas particles flow past one another, lighter particles travel more quickly (higher velocity). • Effusion: The escape of a gas through a tiny hole.
Graham’s Law of Effusion • States that the rate of effusion for a gas is inversely proportional to the square root of its mass. • To compare the rates of diffusion of two gases use the formula: RateA = sqrt(molar massB) RateB (molar massA)
Gas Pressure • Occurs because of the collision of gas particles with the walls of their container. • Pressure is defined as force per unit area. • The particles in Earth’s atmosphere exert pressure on Earth known as atmospheric pressure, or air pressure. • At higher elevations, air pressure is lower than at sea level.
Pressure • Air pressure is measured using a barometer or a manometer. • The SI unit for pressure is the pascal (Pa). 1 Pa = 1 N / 1 meter2 • There are other units to measure pressure, which include atmosphere (atm), torr, bar, pounds per square inch (psi) and kilopascals (kPa).
Dalton’s Law of Partial Pressures • Dalton’s law of partial pressures states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture. Ptotal = P1 + P2 + P3 + …Pn • The partial pressure of a gas depends on the number of moles of the gas, the size of the container, and the temperature of the mixture.
Collection of Gas Over Water • Dalton’s law can be used to determine the amount of gas produced by a reaction. • When the gas is collected in an inverted container over water, the gas will displace the water. • If the gas doesn’t react with water, the total pressure will be a combination of the gas and the water vapor, which has a fixed value.
Forces of Attraction • The forces that hold particles together in ionic, covalent and metallic bonds are called intramolecular forces. • Intermolecular forces are weaker than intramolecular forces. • There are three types of intermolecular forces: dispersion forces, dipole-dipole forces, and hydrogen bonds.
Dispersion Forces • These forces are also called London forces. • They are weak forces that result from temporary shifts in the density of electrons in electron clouds. • When two nonpolar molecules are close to one another the electron cloud in one molecule will repel the the electron cloud of the other. This causes the electron density to shift, creating a temporary dipole.
Dispersion Forces (cont’d) • Weak forces of attraction form between oppositely charged regions. • These forces are only important when no stronger force is acting. • They are noticeable between identical nonpolar molecules.
Dipole-Dipole Forces • Attractions between oppositely charged regions of polar molecules are called dipole-dipole forces. • Polar molecules have permanent dipoles and orient so that opposite charges are attracted. • These forces are stronger than dispersion forces.
Hydrogen Bonds • A hydrogen bond is a special type of dipole-dipole attraction that occurs between molecules containing hydrogen and either fluorine, oxygen or nitrogen.
Liquids and Solids • The KM theory also explains the behavior of liquids and solids; however they exhibit much stronger forces of attraction.
Liquids • They take the shape of their container but have a fixed volume. Their forces of attraction limit their range of motion. • Density: Higher in liquids because particles are held together. • Compression: requires enormous pressure, produces small volume change. • Fluidity: less fluid than gases.
Liquids (cont’d) • Viscosity: stronger the intermolecular forces, the higher the viscosity; increases with mass; decrease with temperature increase. • Surface tension: measure of the inward pull by particles to the interior of a liquid. Stronger attractions means higher surface tension. • Capillary action: movement up a narrow glass tube. Occurs when adhesion (intermolecular) is greater than cohesion (intramolecular).
Solids • Strong attractive forces that limit movement to vibrations. • Definite shape and volume. • Particles are more closely packed than in liquids, therefore a higher density. • Normally, liquids are less dense than solids, though water is an exception.
Crystalline Solids • A solid whose atoms, ions, or molecules are arranged in an orderly, geometric, three dimensional structure (lattice) are called crystalline solids. • Not all solids are crystalline. In amorphous solids there is no regular, repeating pattern.
Phase Changes • Most substances can exist in three states—solid, liquid, gas—depending on their temperature and pressure. • These states can coexist; however, if energy is added or taken away, one phase can change into another.
Phase Changes That Require Energy • The three phases that require energy are melting, vaporization and sublimation. • The temperature at which the forces holding the crystal lattice together are broken and a solid becomes a liquid is known as the melting point. • The process by which a liquid becomes a gas or vapor is known as vaporization. • Vaporization on the surface is called evaporation. • The temperature at which liquids become vapors is called the boiling point.
Sublimation • Sublimation is the process in which a solid changes directly to a gas without first becoming a liquid. • Examples of materials that sublime include solid air fresheners, dry ice, moth balls, etc.
Phase Changes That Release Energy • There are three phase changes that release energy: condensation, freezing and deposition. • Condensation is the process of a gas or vapor becoming a liquid. • The freezing point of a material is the temperature at which a liquid becomes a solid. • Going from a gas to a solid, without becoming a liquid is known as deposition.
Phase Diagrams • Temperature and pressure control the phase of a substance. • A phase diagram is a graph of pressure versus temperature that shoed in which phase a substance exists at different temperatures and pressures. • There are typically three regions and three curves. The curves represent where two phases coexist.
Triple Point • The triple point is the point on a phase diagram that represents the temperature and pressure at which the three phases of a substance can coexist. All 6 phases occur here. • The critical point shows the critical temperature and pressure, above which a substance cannot exist as a liquid.
