ch 12 acad.notebook January 12, 2016 Ch 12 States of Matter (solids, liquids, gases, plasma, Bose Einstein condensate)

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1 Ch 12 States of Matter (solids, liquids, gases, plasma, Bose Einstein condensate) BIG IDEA The kinetic molecular theory explains the different properties of solids, liquids and gases. I CAN: 1) use the kinetic molecular theory to explain the physical properties of gases, liquids and solids. 2) compare types of intermolecular forces. 3) describe the role of energy in phase changes. BEC link Kinetic Molecular Theory (kinetos "to move") Particles are in constant motion in all kinds of matter (s, l, g, p) Apply the KMT to Gases (gases Greek word for "chaos") 1. Particle Size very small particles separated by empty space volume of particles is very small compared to the volume of the empty space no significant attractive or repulsive forces between the particles 2. Particle Motion particles are in constant, random motion collisions are elastic no loss of kinetic energy collisions with walls of container create gas pressure 3. Particle Energy kinetic energy is the energy a particle has due to its motion KE = 1/2 mv 2 (m = mass, v = velocity) In a gas, the particles have the same mass but velocities are different. The average speed of an oxygen molecule is 1027 mi/hr at 20 o C. Temperature is a measure of the average kinetic energy of the particles in a substance. As temperature increases, KE increases. As temp. decreases, KE decreases. At Absolute Zero, all motion of particles stops. 0 Kelvin or 273 o C 1

2 Behavior of Gases 1. Low Density: Cl 2 (g) g/ml Au(s) 19.3 g/ml a great deal of space exists between gas particles 2. Compression and Expansion: 3. Diffusion the movement of particles from high concentration to low concentration. Lighter particles diffuse more rapidly than heavier particles. Demo: perfume 4. Effusion a gas escapes from a container through a tiny opening (puncture in a bike tire) Graham's Law of Effusion and Diffusion There is an inverse relationship between the rate of effusion (or diffusion) and the square root of the molar mass of the gas. Lighter particles diffuse more rapidly than heavier particles. How does density change as a gas is compressed or expanded? Gas Pressure Pressure: Force per unit Area: P = F/A Practice: Compare rates of diffusion for N 2 and Ne. Compare rates of diffusion for CO and CO 2. Examples: bed of nails, high heeled shoes, walking on a frozen pond Gas pressure: results from collisions of particles on an object or with walls of container. Atmospheric pressure: results from collisions of air molecules with objects. Air pressure is lower at higher altitudes that at sea level. Barometers measure atmospheric pressure (invented by Torricelli, 1600s) 2

3 Units of Pressure Practice: 1. kilopascal (kpa) 2. atmosphere (atm) 3. Torr 4. millimeters of Hg (mm Hg) 5. pounds per square inch (psi) At sea level: 1 atm = kpa = 760 torr = 760 mm Hg = 14.7 psi Dalton's Law of Partial Pressures the total pressure of a mixture of gases is equal to the sum of their partial pressures P TOTAL = P A + P B + P C +... Practice: 3

4 Dispersion Force Dipole Dipole Force Hydrogen Bonding 4

5 Practice: Liquids have a fixed volume take the shape of their container KMT particles in constant motion Properties: Liquids cont. 4. Viscosity the resistance of a liquid to flow; affected by intermolecular forces and size of particles and temperature 1. Density: much denser than gases due to stronger intermolecular forces 2. Compressibility: much than gases 3. Fluidity: fluid than gases compare a water leak to a gas leak 5

6 Liquids cont. 5. Surface Tension a measure of the inward pull by particles in the interior of a liquid H 2 O high surface tension due to Surfactants lower the surface tension of H 2 O by breaking the H bonds Liquids cont. 6. Cohesive vs Adhesive Forces Cohesion attractive forces between identical molecules Adhesion attractive forces between different molecules Crystalline Solids Solids particles are in constant motion because they are vibrating have strong intermolecular forces particles are arranged in an orderly, geometric 3D structure Unit Cell the smallest 3D repeating unit in a crystal lattice 1. Density solids are usually more dense than liquids EXCEPT for. Simple Body centered Face centered 6

7 Type Unit particles Characteristics Examples Molecular Ionic Metallic molecules covalent bonds Ions Amorphous not a crystalline none solid Types of Crystalline Solids Soft: poor conductivity Ice, dry ice (CO 2), sugar Hard; brittle; NaCl, KBr, poor conductors CaCO 3 Soft to hard; Positive ions malleable and surrounded by ductile; good moving electrons conductors Fe, Cu, Ag glass, rubber, plastic 7

8 Vaporization a liquid changes to a gas or vapor 1) Sublimation Examples of Phase Changes Dry Ice (solid CO 2 ) sublimes at room temp Solid Air Fresheners Ice Cubes left in a freezer 2) Condensation Glasses steam up Water droplets form on outside of a glass containing a cold drink 3) Deposition Evaporation takes place only at the surface of a liquid Vapor Pressure the pressure exerted by a vapor over a liquid Boiling Point the temperature at which the vapor pressure equals the external or atmospheric pressure. At higher elevations, atmospheric pressure is lower so the boiling point is lower. Water vapor changes to frost on a window Phase Diagram a graph of temperature (x axis) vs pressure (yaxis) Shows whether a substance is a solid, liquid, or gas under different conditions of T and P Triple Point represents where all 3 phases of a substance can coexist (all 6 phase changes occur here) Critical Point above this point, only a vapor exists 8

9 Phase Diagram for Water Phase Diagram for CO 2 9