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Gases & Atmospheric Chemistry

Gases & Atmospheric Chemistry

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Gases & Atmospheric Chemistry

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  1. Gases & Atmospheric Chemistry

  2. Read 11.1 (p. 516 – 519) • Draw 3 particle pictures of the different states of matter. Include a container in your diagram. • Add symbols to illustrate the motion of the particles. • Volunteers? • Vgas=Vcontainer • Define: KE, T • Explain pop can demo. • Define: P

  3. Introduction… Demo 1: Water and Cue Card Why does the water stay in the glass? Demo 2: Pop Can What is happening in terms of STATES of water? Liquid  Gas  Liquid What is happening in terms of VOLUME of water? Small  EXTREMELY LARGE  Small 1 mol H2O(l)= 18 mL or 0.018L 1 mol of H2O(g)= 22.4 L

  4. We are surrounded by gases! • Air (or the Atmosphere) is made up of: • This air is colliding with us…this is called PRESSURE …Or specifically the pressure we feel is “Atmospheric Pressure” What would the pressure be like on top of Mount Everest? Pressure: force exerted on an object per unit of surface area Air pressure= 1 atmosphere= 1 atm

  5. States of Matter • How do intermolecular forces of attraction differ between states? • Generally, how is particle size related to states of matter? • CH4 vs. C5H12 • Larger molecules= higher boiling points= liquid @ room temp • Why???? The bigger the molecule the more opportunities for temporary dipoles to form…dispersion forces add up and this increase the IMF of attraction Are all states compressible? • Only gases…what is condensation?

  6. Types of Kinetic Energy Every moving particle has energy= kinetic energy Kinetic=motion Solid= vibrational motion Liquid= rotational and vibrational motion Gas= Translational, rotational, and vibrational motion

  7. Kinetic Molecular Theory of Gases • The volume of an individual gas molecule is negligible compared to the volume of the container.* • No attractive or repulsive forces between gas molecules.* • Gas molecules move randomly in all directions, in straight lines (translational, rotational, vibrational). • Perfect elastic collisions between gas molecules (i.e. no loss of kinetic energy). • An increase in temperature will increase the motion of molecules. This means there is an increase in the average kinetic energy. * Assumptions for an IDEAL GAS

  8. What can affect kinetic energy? TEMPERATURE! How?  Temperature= Kinetic energy How is pressure affected?  Temperature= Kinetic energy =  Pressure How is volume affected?  Temperature= Kinetic energy =  Pressure=  Volume of Gas* *space that the gas take up BUT usually Vgas= Volumecontainer * Temperature is a measure of the average kinetic energy

  9. Pressure • The force exerted on an object per unit of surface area P= F = N = Pa A m2 Pressure: kinetic motion and collisions with surroundings Atmospheric Pressure is what we feel around us. - Air molecules have a mass. -They are pulled by gravity (acceleration of objects on earth= 9.81 m/s2 or 32.2 ft/s2) to exert pressure on Earth.

  10. Units of Pressure Pressure can be measured by: • Atmospheres= 1 atm • KiloPascals (SI Unit)= 101.3 kPa • Millimeters of Mercury= 760 mmHg (1st mercury barometer) • Torricelli’s= 760 torr • Pounds per square inch (Imperial)= 14.7 psi 1 atm= 101.3kPa= 760 mmHg= 760 torr= 14.7 psi How many kPa are in 3.57 atm of pressure?

  11. Discovering Pressure Barometer= 1st instrument to measure air pressure Torricelli’s Barometer” Torricelli found that air pressure at 0° C 1 atm= 101.3kPa= 760 mmHg= 760 torr= 14.7 psi Standard Temperature and Pressure (STP) 0°C 101.3kPa Standard Ambient Temperature and Pressure (SATP) 25°C 100kPa 760mm ~ 30 inch

  12. K: No negative values! Kelvin Scale • Charles found that the x-intercept would always be -273°C • Kelvin (c.1800) inferred that at -273°C VOLUME WOULD BE ZERO (molecular motion would cease, NO kinetic energy) • p. 549 #1,2 0 K= -273°C= “Absolute zero” TK= °C + 273.15

  13. Charles’ Law Gay-Lussac (c.1800) referenced Charles’ work, and it became known as… Charles’ Law: the volume of a fixed mass of gas is proportional to its temperature (K) when the pressure is kept constant V1 = V2 T1 T2 T MUST be K! Practice Pg. 552 #2

  14. Boyle’s Law Mathematically: P1V1= P2V2 p. 559 #1,2 • Boyle’s Law (1662): the volume of a given amount of gas at constant temperature, varies INVERSELY with the applied pressure • If we change the pressure by a factor of x, then the volume will change inversely by that same factor

  15. Gay-Lussac’s Law Recall: 1. Temperature is a measure of average kinetic energy 2. Vgas= Volume of container Gay- Lussac’s Law: the pressure of a fixed amount of gas, at constant volume, is directly proportional to its Kelvin temperature. P1 = P2 T1 T2 p. 559 #3 MUST use K! Practice Pg.559Q:1-3

  16. Combined Gas Law Recall: STP 0°C/273K & 101.3kPa SATP 25°C/298K & 100kPa Combine: 1. Bolye’s Law PiVi= PfVf 2. Charles Law Vi = Vf TiTf 3. Gay-Lussac’s Law Pi = Pf TiTf Ex. A sample of gas has a volume of 150mL at 260K and 92.3kPa. What will the new volume be at 376k and 123 kPa? P1V1 = P2V2p. 560 #1-3 T1 T2 T MUST be K!

  17. Dalton’s Law of Partial Pressures Imagine mixing 3 different gases each having a different pressure… What would the final pressure be? Dalton’s Law of Partial Pressures: the total pressure of a mixture of gases is the sum of the pressures of each of the individual gases Ex. What is the pressure of O2 in the atmosphere? Ptotal= P1 + P2 + P3 + … +Pn Practice Pg. 594 #1-4 Pg. 596 #1-3