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States of Matter

States of Matter. I will use the kinetic-molecular theory to explain the physical properties of gases, liquids and solids I will compare types of intermolecular forces I will explain how kinetic energy and intermolecular forces combine to determine the state of a substance

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States of Matter

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  1. States of Matter • I will use the kinetic-molecular theory to explain the physical properties of gases, liquids and solids • I will compare types of intermolecular forces • I will explain how kinetic energy and intermolecular forces combine to determine the state of a substance • I will describe the role of energy in phase changes

  2. 13.1 Gases • I will use the Kinetic-molecular theory to explain the behavior of gases • I will describe how mass affects the rates of diffusion and effusion • I will explain how gas pressure is measured and calculate the partial pressure of a gas • Vocabulary kinetic-molecular theory elastic collision • temperature Graham’s Law of effusion diffusion • Pressure barometer pascal atmosphere • Dalton’s law of partial pressures

  3. Gases • Substances that are gases at room temperature • usually display similar physical preperties despite • their different compositions. • Why is there so little variation in behavior among gases?

  4. Kinetic-Molecular Theory

  5. Kinetic-Molecular Theory • Describes the behavior of gases in terms of • particles in motion • Makes several assumptions about gas particles: • Size • Motion • Energy

  6. Gas Assumptions • Particle Size • Separated by empty space • Volume of particle = small • Volume of empty space = large • NO significant attractive or repulsive forces • Particle Motion • Constant random motion • Move in a STRAIGHT line • Collide with walls or other particles • Collisions are elastic (NO kinetic energy is lost, just transferred)

  7. Gas Assumptions Cont... • Particle Energy • Affected by mass and velocity • KE = ½ mv2 • Single gas • Particles have same mass • Particles have different velocity • Particles have different KE • Temperature = measure of the average KE of the particles in a sample of matter • At a given temperature, ALL gases have the SAME average KE

  8. Gas Behavior

  9. Gas Behavior • Low Density • Compressible • Expandable • Diffuse • Effuse

  10. Gas Behavior • Low Density • Density = mass/volume • Large space between gas particles • Fewer gas molecules than solid or liquid molecules in the same volume

  11. Gas Behavior Compression & Expansion • Large amount of empty space between gas particles • Allows particles to be squished into a smaller volume • Stop squishing; random motion of particles fills the available space, expands to original volume

  12. Gas Behavior Diffusion Particles diffuse from areas of high concentration to areas of low concentration Rate of diffusion Depends on mass of particles Lighter = faster Heavier = slower • The movement of one material through another material • Gas particles have no significant forces of attraction • Particles can slide past each other • Mix until evenly distributed

  13. Gas Behavior

  14. Gas Behavior Effusion • Related to diffusion • When a gas escapes through a tiny opening • Ex. Puncture a balloon or tire • Inverse relationship between effusion rates and molar mass

  15. Gas Behavior Graham’s Law of Effusion Also applies to rate of diffusion Set up a proportion to compare the diffusion rates of 2 gases • The rate of effusion for a gas is inversely proportional to the square root of its molar mass

  16. Finding a Ratio of Diffusion Rates • Let’s Try!

  17. Practice Problems • RN2/RNe = 0.849 • RCO/RCO2 = 1.25 • 2.5mol/min Calculate the ratio of effusion rates for nitrogen (N2) and Neon (Ne). Calculate the ratio of diffusion rates for carbon monoxide (CO) and carbon dioxide (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 3.6 mol/min? Use to find ratios

  18. Gas Pressure

  19. Pressure Gas Pressure • Force per unit area • Gas particles exert pressure when they collide with: • The walls of their container • Each other • Exert pressure in ALL directions (because particles move in all directions) Individual particles can only exert little pressure Many particles colliding can exert substantial pressure pressure of gases is what keeps our tires inflated, makes our basketballs bounce, makes hairspray come out of the can, etc.

  20. Gas Pressure • Pressure increases when temperature increases because the molecules are moving with greater speed and colliding against the sides of their containers more often. • Therefore, the pressure inside that container is greater, because there are more collisions.

  21. Gas Pressure Atmospheric Pressure Varies by elevation Mountains = less Sea Level = more • Air pressure at Earth’s surface • Equal to: • Pressure exerted by 1 kg mass on a square centimeter

  22. Gas Pressure Measuring Air Pressure Air pressure varies because of: Changes in air temperature Changes in humidity Increase in air pressure = Hg rises Decrease in air pressure = Hg falls • Barometer • Measures atmospheric pressure • Mercury in it is always about 760mm • Exact amount determined by 2 forces • Gravity- downward force • Air pressure- upward force (air presses on surface of Hg)

  23. Gas Pressure Measuring Enclosed Gas Pressure • Manometer • Measures pressure of enclosed gas • The difference in height of the mercury in the 2 arms is used to calculate the pressure of the gas in the flask

  24. Gas Pressure Units of Pressure SI Unit = Pascal (Pa) Atmosphere = atm (used to report air pressure)

  25. Gas Laws

  26. Gas Laws • Dalton’s Law of Partial Pressures • Gay-Lussac’s Law • Boyle’s Law • Charles’s Law • Ideal Gas Law • Combined Gas Law

  27. Dalton’s Law of Partial Pressures • P total = P1 + P2 + P3 +…Pn • The total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture. • Partial pressure = the portion of the total pressure contributed by a single gas Depends on • # of moles of gas • Size of container • Temperature of the mixture Does NOT depend on identity of gas (at a given temperature and pressure, the partial pressure of 1 mole of ANY gas is the same)

  28. Dalton’s Law of Partial Pressures • Practice Problem • A mixture of oxygen (O2), carbon dioxide (CO2), and nitrogen (N2) has a total pressure of 0.97 atm. What is the partial pressure of O2, if the partial pressure of CO2 is 0.70 atm and the partial pressure of N2 is 0.12 atm? PO2 = 0.97 atm – 0.70 atm -0.12 atm PO2 = 0.15 atm

  29. Boyle’s Law • P1V1 = P2V2 • The volume of a given amount of gas held at a constant temperature varies inversely with pressure • As pressure increases, volume decreases and vice versa

  30. Boyle’s Law • Practice Problem • A sample of helium gas in a balloon is compressed from 4.0 L to 2.5 L at a constant temperature. If the pressure of the gas in the 4.0 L volume is 210 kPa, what will the pressure be at 2.5 L? 210 kPa (4.0 L) = P2 (2.5 L) P2 = 340 kPa

  31. Charles’s Law • V1V2 • T1= T2 • The volume of a given mass of a gas is directly proportional to its Kelvin temperature at constant pressure. • When Kelvin temperature increases, volume increases and vice versa • Tk = 273 + Tc

  32. Charles’s Law • Practice Problem • A gas sample at 40 degrees C occupies a volume of 2.32 L. If the temperature is raised to 75.0 degrees C, what will the volume be, assuming the pressure remains constant? 40.0 degrees C + 273 = 313K 75 degrees C + 273 = 348K 2.32 L = V2 313 K 348 K V2 = 2.58 L

  33. Gay-Lussac’s Law • P1P2 • T1= T2 • The pressure of a given mass of gas varies directly with the Kelvin temperature when the volume remains constant. • As Kelvin temperature increases, pressure increases and vice versa

  34. Gay-Lussac’s Law • Practice Problem • The pressure of a gas in a tank is 3.20 atm at 22.0 degrees C. If the temperature rises to 60.0 degrees C, what will be the gas pressure in the tank? 22.0 degrees C + 273 = 295 K 60.0 degrees C + 273 = 333 K 3.20 atm= P2 295 K 333 K P2 = 3.61 atm

  35. The Combined Gas Law • P1V1 = P2V2 • T1 T2 • The relationship among pressure, volume, and temperature of a fixed amount of gas. • Pressure increases, volume decreases • Pressure increases, temperature increases • Volume increases, temperature increases

  36. The Combined Gas Law • Practice Problem • A gas at 110 kPa and 30.0 degrees C fills a flexible container with an initial volume of 2.00 L. If the temperature is raised to 80.0 degrees C and the pressure increased to 440 kPa, what is the new volume? 30.0 degrees C + 273 = 303 K 80.0 degrees C + 273 = 353 K 110(2.00) = 440(V2) 303 353 V2 = 0.58 L

  37. Ideal Gas • Ideal Gas • One whose particles • Take up NO space • Have NO intermolecular attractive forces • Follows ideal gas laws under ALL conditions of temperature and pressure

  38. Real Gas • NO gas in the real world is truly ideal • particles have some volume • some attractive forces • Real gases deviate most from ideal behavior at low temperatures and high pressures • Examples: • Liquid nitrogen is used to store biological tissues at low temps • Increased pressure allows a larger mass of propane to fit into a smaller volume for easier transport

  39. Ideal Gas Law • PV = nRT • Describes the physical behavior of an ideal gas in terms of the pressure, volumes, temperature, and number of moles of gas present. • R = ideal gas constant

  40. Ideal Gas Law • Calculate the number of moles of gas contained in a 3.0 L vessel at 3.00 x 102 K with a pressure of 1.50 atm. V = 3.0 L T = 3.00 x 102 K P = 1.50 atm R = 0.0821 L atm/ mol K n = ? 1.50 (3.0) = n (0.0821)(3.00 x 102 K) n = 0.18 mol

  41. Density of a Gas • D = MP • RT • Practice Problem • What is the density of a gas at STP that has a molar mass of 44.0 g/mol? STP = 273 K, 1 atm D = 44.0(1) 0.0821(273) D = 1.96 g/L

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