States of Matter Chapter 12

12.1 - Gases Kinetic molecular theory = describes how gases behave, has 3 major assumptions: Particle size = particles are small compared to volume of empty space, no forces between particles Particle motion = particles are in constant motion, move in a straight line until they collide with other particles or container walls, elastic collisions (no loss of kinetic energy)

12.1… 3. Particle energy = determined by mass and velocity of a particle (KE=1/2 mv2) In a sample of a gas, all particles have same mass but different velocity (so KE of particles varies) Temperature = measure of average KE of particles in a sample (so at a given temp., all particles have same average KE)

12.1 – Gas Behavior Low density = lots of space between gas particles Compression and expansion = empty space in a gas container can easily be pushed into a smaller volume (ex. Squishing memory foam) Diffusion and effusion = random motion of gas particles causes them to mix until they are evenly distributed

12.1 - Diffusion and Effusion Diffusion = movement of one material through another from area of high concentration to area of low concentration (ex. Spraying perfume) Effusion = gas particles escape through a tiny opening (ex. Popping a balloon) At the same temp., heavier particles effuse/diffuse slower than lighter ones (Graham’s law of effusion: rate is inversely proportional to molar mass.

12.1 – Gas Pressure Pressure = force exerted by gas particles colliding with the walls of the container Barometer = measures atmospheric pressure, height of mercury determined by gravity exerting downward force on mercury and upward force of air pressing down on surface of mercury (increase in air pressure = mercury rises, decrease in air pressure = mercury falls), air pressure depends on temp. and humidity

12.1 – Units of Pressure Pascal (Pa) = SI unit of pressure Average air pressure at sea level = 760 mm Hg 760 torr 101.3 kilopascals (kPa) 14.7 psi One atmosphere (1 atm)

Pressure Demos! How to get an egg inside a milk bottle (vacuum inside the bottle) Extreme Shaving Cream Monster Marshmallows

12.1 – Dalton’s Law of Partial Pressures Dalton’s Law of Partial Pressures = total pressure of a mixture of gases is equal to the sum of the pressures of all gases in the mixture Portion of the total pressure contributed by one gas is called its partial pressure Partial pressure of a gas depends on the number of moles of gas, size of the container, and temp.

12.2 – Forces of Attraction Intramolecular forces = attractive forces that hold particles together (bonds) Intermolecular forces = attractive forces between particles

12.2 – Intermolecular Forces Dispersion forces = weak force between two non-polar molecules (those with evenly distributed electrons) Electrons always moving, electron clouds of non-polar molecules repel each other Electron density around the nucleus shifts so one side of the molecule is more negative than the other Leads to weak attraction between these temporary dipoles

12.2 – Dispersion Forces

12.2 – Intermolecular Forces… Dipole-Dipole forces = forces between oppositely charged regions of polar molecules (permanent dipoles) Ex. HCl Partially positive H in one molecule of HCl lines up with partially negative Cl in another molecule

13.2 – Intermolecular Forces… Hydrogen bonds = a type of dipole-dipole attraction between molecules that contain hydrogen and a highly electronegative atom with at least one lone pair Explains why water is a liquid at room temperature and not a gas  life on Earth! Hydrogen atom has large partial positive charge and is attracted to oxygen’s large partial negative charge on another molecule

13.2 – Hydrogen Bonds

12.3 – Liquids and Solids

12.3 – Liquids and Solids… Liquids are much denser than gases due to intermolecular forces that hold particles together Liquids cannot be compressed as much as gases because particles are already closely packed together Molecules in a liquid

12.3 – Liquids and Solids… Fluidity = ability to flow, both gases and liquids are considered fluids because they can diffuse through one another Liquids are less fluid than gases due to intermolecular forces that slow down their movement Ex. Water pipe leak vs. natural gas leak

12.3 – Liquids and Solids… Viscosity = resistance of a liquid to flow Stronger intermolecular forces  higher viscosity Bigger molecules  higher viscosity Lower temperature  higher viscosity Ex. Motor oil = need less viscous oil in winter to make sure it keeps flowing at low temperatures, need more viscous oil in summer to make sure it stays thick enough to lubricate the engine at very high temperatures

12.3 – Liquids and Solids Surface tension = energy required to increase surface area of a liquid by a given amount Stronger attractions between particles  higher surface tension Water has a very high surface tension due to hydrogen bonding, allows small insects to walk on water, why we need water AND soap to clean dishes/clothes Surface tension of water Magic milk

12.3 – Liquids and Solids… Capillary action = occurs when adhesion to walls of container is greater than cohesion in the liquid Adhesion = force of attraction between molecules that are different Cohesion = force of attraction between identical molecules Ex. Meniscus in graduated cylinder, absorbance of paper towels and diapers

12.3 – Liquids and Solids… Solids = not considered to be fluids, very ordered, more dense than liquids because particles are more closely packed together Exception = water, solid ice is less dense than liquid water due to three dimensional structure of water molecules (ice cubes and glaciers float, keeps aquatic life alive in the winter) Why ice floats...

12.3 – Liquids and Solids… Crystalline solids = atoms, ions, or molecules arranged in an orderly, geometric, 3-D structure Individual pieces are called crystals Unit cell = smallest part of a crystal retaining the crystal shape Molecules in solids

12.3 – Liquids and Solids 4 types of crystalline solids: Molecular solids = soft, poor conductors of heat and electricity, low to moderate melting points Allotropes = 2 different solid forms of the same element (ex. Carbon as graphite or diamond), very hard, poor conductors, high melting points Ionic solids = hard, brittle, poor conductors, high melting points Metallic solids = soft to hard, excellent conductors, low to high melting points

12.3 – Liquids and Solids… Amorphous solids = do not contain crystals, no repeating pattern, form when a molten material cools too quickly Ex. Glass, rubber, plastics

12.4 – Phase Changes Intro to phase changes

12.4 – Phase Changes Why do substances melt? As molecules are heated, they gain enough KE to overcome attractive forces that hold them together as a solid Melting point = temperature at which forces holding crystal lattice together are broken and solid becomes liquid REQUIRES energy

12.4 – Phase Changes Why do substances freeze? As heat is removed, molecules lose KE and slow down Freezing point = temperature at which a liquid is converted into a crystalline solid RELEASES energy

12.4 – Phase Changes Why do liquids become gases? When energy is added to a liquid, the temperature increases and molecules gain KE and escape the liquid Vaporization = process by which a liquid changes to a gas or vapor Evaporation = vaporization that occurs on the surface of a liquid that is not boiling REQUIRES energy

12.4 – Phase Changes… Vapor pressure = pressure exerted by a vapor over a liquid Boiling point = temperature at which vapor pressure of a liquid is equal to atmospheric pressure, molecules throughout the liquid have enough energy to vaporize and bubbles of vapor rise to the surface Water boiling in a vacuum Demo!

12.4 – Phase Changes… Why do gases become liquids? When a vapor molecule loses energy, it slows down and is more likely to collide with another molecule and bond, releasing energy and becoming a liquid Condensation = Process by which a gas or vapor becomes a liquid RELEASES energy Evaporation and Condensation

12.4 – Phase Changes… What happens when we sweat? Water molecules in sweat gain energy from heat coming off the body Some molecules gain enough energy to evaporate, taking energy with them Molecules left behind have less energy and therefore, a lower temperature

12.4 – Phase Changes… Sublimation = solid changes directly to a gas without passing through the liquid phase Ex. Dry ice REQUIRES energy Deposition = gas or vapor changes directly to a solid without passing through the liquid phase

12.4 – Phase Changes Phase diagram = Pressure vs. Temperature, shows relationship between pressure/temperature/phase of a substance Triple point = point on the diagram that shows the temp. and pressure at which all three phases can exist in equilibrium with eachother Critical point = point on the diagram that shows the temp. and pressure at which water can no longer exist as a liquid

12.4 - Phase Diagram