UNIT 6 GASES
OUTLINE States of Matter, Forces of Attraction Phase Changes Gases The Ideal Gas Law Gas Stoichiometry
STATES OF MATTER Remember that all matter exists in three physical states: Solid Liquid Gas
INTERMOLECULAR FORCES Why are some substances gases at room temperature (eg. CO 2 ) while some substances are solid at room temperature (eg. Iron)? Intermolecular forces of attraction is the general term to describe how particles are held together in a substance. 1. Dispersion 2. Dipole-Dipole 3. Hydrogen bonding
INTERMOLECULAR FORCES Strong intermolecular forces = higher melting point, boiling point. Intermolecular means between molecules Intramolecular means inside molecules (ie. chemical bonding)! Do not confuse these two things
1. DISPERSION FORCE Dispersion forces are also called London dispersion forces. Dispersion forces are weak forces that result from temporary changes in the density of electron clouds. Dispersion forces exist between ALL particles in ALL substances. Bigger electron cloud (ie. substance with more electrons) = stronger dispersion force. Eg. Gr.17
1. DISPERSION FORCE Dispersion forces are also called London dispersion forces. Dispersion forces are weak forces that result from temporary changes in the density of electron clouds.
2. DIPOLE-DIPOLE FORCES Dipole-dipole forces are the attractions between oppositely charged parts of polar molecules dipoles. Dipole-dipole forces ONLY exist between particles with a permanent dipole. Polar molecules only. Eg. HCl Dipole-dipole forces are slightly stronger than dispersion forces, but if that molecule is very large, dispersion forces are more significant.
2. DIPOLE-DIPOLE FORCES
3. HYDROGEN BOND NOT a chemical bond!! A hydrogen bond is a type of dipole-dipole attraction. A H that is chemically bonded to an O, N or F inside a molecule: positive. An O, N or F that is chemically bonded to a H inside a molecule: negative. On two different molecules (which contain H-O/N/F bond) one H and another O/N/F are electrostatically attracted. This is the strongest type of intermolecular force.
INTERMOLECULAR FORCES For each of the following compounds, determine the main intermolecular force. 1. Nitrogen 8. SiH 2 O 2. Carbon tetrachloride 3. H 2 S 4. Sulfur monoxide 5. N 2 H 2 6. BH 3 7. CH 4 O
INTERMOLECULAR FORCES: NOTES! Strong intermolecular forces = Stronger forces of attraction between particles = harder to separate = higher melting point, higher boiling points! Dispersion forces are weakest when comparing substances of similar size. Dispersion forces are stronger than other IMF is the substance is much larger than the other substances!
INTERMOLECULAR FORCES: NOTES! Dispersion forces are weakest when comparing substances of similar size. Dispersion forces are stronger than other IMF is the substance is much larger than the other substances! * See Question 5 on IMF Worksheet
PHASE CHANGES Phase Diagrams (12.4) Most substances exist in three states, depending on the temperature and pressure. Phase diagram is a graph of pressure versus temperature that shows the phases a substance exists for different T and P.
PHASE CHANGES Phase changes require energy: Melting: Ice has water molecules that are close together, held by H-bonding. Heat is transferred to the water molecules, and the molecules absorb enough energy to break these IMFs so that the molecules move further apart, into the liquid phase.
PHASE CHANGES Vaporization: liquid changing to gas Vapor Pressure: Pressure exerted by a vapor over the surface of a liquid. Boiling point: The temperature where the vapor pressure of a liquid equal the external atmospheric pressure.
PHASE CHANGES Phase changes that require energy: Melting Vaporization Sublimation: Solid directly to gas.
PHASE CHANGES Phase changes that release energy: Freezing Condensation Deposition
STOICHIOMETRY: SOLUTIONS CH.14
What is the critical temperature of compound X? If you were to have a bottle containing compound X in your closet, what phase would it most likely be in? At what temperature and pressure will all three phases coexist? If I have a bottle of compound X at a pressure of 45 atm and temperature of 100 C, what will happen if I raise the temperature to 400 C? Why can t compound X be boiled at a temperature of 200 C? If I wanted to, could I drink compound X?
GASES Gases look and behave very differently from liquids and solids. To try and explain the properties of gases, scientists came up with the kineticmolecular theory of gases. This theory describes the behaviour of matter in terms of particles in motion.
KINETIC-MOLECULAR THEORY *IMPORTANT* 1. Gases are made up of small particles that are separated from one another by empty space. 2. Volume of the particles is small compared with the volume of the empty space between particles. 3. There are no forces of attraction between particles.
KINETIC-MOLECULAR THEORY 4. Gas particles are in constant, random motion. 5. Collisions between gas particles are elastic (do not lose energy)
KINETIC-MOLECULAR THEORY The theory can explain properties of gases. 1. Gases expand because: -constant, random motion -no attractive forces between particles
KINETIC-MOLECULAR THEORY 2. Gases can contract because: -gas particles are tiny and far apart 3. Gases have low density because: -gas particles are tiny and far apart -no attractive forces between particles
KINETIC-MOLECULAR THEORY The theory can explain why gases have low density (mass per volume) and can be compressed or expanded (random motion of particles fills the available space)
KINETIC-MOLECULAR THEORY Diffusion: movement of one substance through another Effusion: gas escaping through an opening. Gases can diffuse and effuse because: constant, random motion
KINETIC-MOLECULAR THEORY Graham s law of effusion: Heavier gases effuse more slowly than lighter gases.
KINETIC-MOLECULAR THEORY The temperature is a measure of the average kinetic energy of the gas particles in a sample. KE = ½ m v 2
KINETIC-MOLECULAR THEORY The temperature is a measure of the average kinetic energy of the gas particles in a sample. KE = ½ m v 2
KINETIC-MOLECULAR THEORY Gas Pressure: Gas particles exert pressure when they collide with the walls of their container. Units: atmosphere (atm), pascal (Pa)
KINETIC-MOLECULAR THEORY Dalton s Law of Partial Pressures: The total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture.
KINETIC-MOLECULAR THEORY
KINETIC-MOLECULAR THEORY Partial pressures can be used to find the amount of gas produced by a reaction. Eg. Gas collected becomes a mixture. Total pressure inside container is the sum of the partial pressures of water vapor and new gas.
KINETIC-MOLECULAR THEORY
KINETIC-MOLECULAR THEORY Knowing the pressure, volume of the container, and temperature of a gas allows you to calculate the number of moles of gas!
KINETIC-MOLECULAR THEORY Real gases: 1. Particles of real gases do have a physical volume 2. Particles of real gases can exert attractive forces on each other
KINETIC-MOLECULAR THEORY Real gases behave like an ideal gas : 1. Low pressure (particles far apart) 2. High temperature (lots of kinetic energy) 3. Weak attraction to each other
THE GAS LAWS Boyle s Law: At constant temperature, increasing pressure decreases volume. Charles s Law: At constant pressure, increasing temperature increases volume. Guy-Lussac s Law: At constant volume, increasing temperature increases pressure.
THE IDEAL GAS LAW Avogadro thought that because particles are so small, 1000 large krypton gas particles would occupy the same volume as 1000 small helium gas particles. Equal volumes of gases at the same temperature and pressure contain equal numbers of particles.
THE IDEAL GAS LAW So at 0 o C and 1.00 atm of pressure, 1 mole of gases (1 mol = 6.02 x 10 23 particles) occupy 22.4 L of volume. This is true for ANY gas (no matter what size the gas particle actually is) Eg. You have 3.50L of a gas. How many moles of this gas do you have?
THE IDEAL GAS LAW Combining all the laws! The ideal gas constant, R = 0.0821 L atm/mol K PV = nrt