Chapter 13 States of Matter 13.1 Gases The Kinetic-Molecular Theory Explaining the Behavior of Gases Gas Pressure

The Kinetic Molecular Theory Basic Assumptions Particle Size Gas particles have no volume (pin point particles) The space between particles is extremely large compared to the volume of the particles. Due to this distance, there is no significant attractive or repulsive force acting on the particles.

The Kinetic Molecular Theory Basic Assumption Particle Motion Gas particles are in constant random motion. Collisions between particles are elastic (Energy can be transferred from one particle to another during a collision, but no energy is lost when particles collide)

Basic Assumptions Particle Energy The mass and velocity of a particle determine the kinetic energy of a particle Temperature is a measure of the average kinetic energy of particles in a sample.

The Kinetic Molecular Theory Mass/Velocity Relationship Questions Condition #1: Two particles (one heavy and one light) traveling at the same velocity. Which exhibits the greatest kinetic energy? Condition #2: Two particles of the same size traveling at different velocities (fast and slow).

Explaining the Behavior of Gases Properties Low Density (pinpoint mass/volume of empty space) Random Motion Behaviors Compression Gases can be compressed due to the large space that exists between particles Gases expand to fill their containers due to constant random motion

Explaining the Behavior of Gases Properties No attractive or repulsive forces acting on particles Particles exhibit constant random motion Behaviors Particles can flow easily past each other in a process called diffusion. The rate of diffusion is dependent on the mass of the particles. Question: Based on this equation which particles diffuse faster, heavy or light particles?

Explaining the Behavior of Gases Property Particles exhibit constant random motion Behavior Effusion (similar to diffusion, where particles escape through a tiny opening) Graham’s Law of Effusion Rate of effusion

Questions/Problems What assumption of the Kinetic-Molecular Theory explains why a gas can expand to fill a container? How does the mass of a particle affect its rate of effusion? Calculate the ratio of diffusion rates for CO and CO2. What is the rate of effusion for a gas that has a molar mass twice that of a gas that effuses at a rate of 3.6mol/min?

The force that a gas exerts per unit area. Gas Pressure The force that a gas exerts per unit area. Measuring Air Pressure Using a Barometer Invented by Evangelista Torricelli A rise in air pressure will cause the height of mercury to rise. Two forces affect the height of the mercury column Gravity and Atmospheric Pressure A decrease in air pressure will cause the height of mercury to fall

Pressure and Altitude Question Which condition would cause the level of mercury in a barometer to fall below 760 mm? Being below sea level or being on the top of a mountain.

Units of Pressure The SI unit for pressure is the pascal (Pa) At sea level and 0oC conditions (STP-standard temperature and pressure) a barometer will read 760mm Hg. Equivalent pressure units 760 mm Hg = 101.3 kPa = 1atm = 760 torr = 14.7 psi

Factor-Label is in the Air Equivalent pressure units 760 mm Hg = 101.3 kPa = 1atm = 760 torr = 14.7 psi Convert 362 torr to kPa Convert 35.4 psi to torr Convert 48.9 kPa to psi

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 a a mixture” The partial pressure of a gas is dependent on the number of moles of gas, the size of the container and the temperature of the mixture.

Partial Pressure Problems 1. What is the partial pressure of hydrogen gas in a mixture of hydrogen and helium if the total pressure is 600 mm Hg and the partial pressure of helium is 439 mm Hg? 2. Find the total pressure in kPa for a mixture that contains three gases with partial pressures of 122 kPa, 35 psi, and 722 torr.

Chapter 13 States of Matter 13.2 Attractive Forces Dispersion Forces Dipole-dipole Forces Hydrogen Bonds

Intramolecular Forces Forces that occur between atoms, ions or molecules within a molecule. Bonding Attractive Parties Ionic Cations and anions Molecular Positive nuclei and shared electrons Metallic Metal cations and mobile electrons

Intermolecular Forces Forces that occur between molecules to hold them together Dispersion Dipole-Dipole Hydrogen Bonding

Dispersion Forces Weak forces that occur between non-polar molecules that result from a temporary shift in the density of electrons in electron clouds. (Butane) The electrons of two non-polar repulse one another which causes the temporary shift.

Dispersion Forces With increasing atomic number, the number of electrons in a molecule increases which results in a greater dispersion force. This explains why Cl2 is a gas, Br2 is a liquid and I2 is a solid. Greater Force = smaller distance between molecules

Dipole-dipole Forces Attractive force that occurs between molecules that have a permanent dipole. Stronger than dispersion forces as long as the two molecules have about the same mass.

Hydrogen Bonding For a hydrogen bond to form, hydrogen must be bonded to oxygen, fluorine or nitrogen.

Attractive Force Review Why are dipole-dipole forces typically stronger than dispersion forces? Which molecules listed below can form hydrogen bonds? Which ones could only experience dispersion forces (how would you know)? H2, NH3, HCl, HF 3. Predict the relative boiling points of the noble gases.

13.3 Liquids and Solids

Liquid Behavior Density- much denser than gases, not compressible Fluidity- diffuse slower than gases, still “flow” Viscosity- measure of resistance to flow Effect of temperature- higher temp, lower viscosity

Liquid Behavior Surface Tension Capillary Action Causes drops and meniscus Capillary Action Water can climb narrow tubes

Solid Behavior Density- more dense than gases and liquids and incompressible Crystalline Solids Unit Cells – smallest particle of a crystal that has the same shape as the crystal Crystal Structure

Types of Solids Molecular Solids-dispersion, dipole, or H-bonds (ex: sugar) Covalent Network Solids- covalent bonds with self (ex: diamond, graphite) Ionic Solids- ionic attraction (ex: salt) Metallic Solids- mobile electrons (ex: copper) Amorphous Solids- irregular pattern (ex: glass)

Phase Changes

Heating Curve

Endothermic Phase Changes Melting- solid absorbs energy until particles have enough speed to break free of IM forces holding them in place Vaporization-liquid absorbs energy until particles have enough speed to break free of IM forces holding them close together

Liquid to Gas Evaporation- occurs at surface Boiling- occurs throughout when vapor pressure equals atmospheric pressure

Sublimation-solid to gas

Exothermic Phase Changes Condensation Deposition Freezing

Phase Diagrams Triple Point