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Chapter 11

Chapter 11. Molecular Composition of Gases. 11.1 Volume-Mass Relationship. Early 1800’s Gay-Lussac notices something interesting… 2L H 2 + 1L O 2 → 2L H 2 O 2mL H 2 + 1mL O 2 → 2mL H 2 O 600L H 2 + 300L O 2 → 600L H 2 O

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Chapter 11

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  1. Chapter 11 Molecular Composition of Gases

  2. 11.1 Volume-Mass Relationship • Early 1800’s Gay-Lussac notices something interesting… 2L H2 + 1L O2 → 2L H2O 2mL H2 + 1mL O2 → 2mL H2O 600L H2 + 300L O2 → 600L H2O There is a 2:1:2 relationship! Just like in the balanced equn 2 H2 + O2 → 2 H2O • Gay-Lussac’s Law of Combining Volumes of Gases: At constant temperature and pressure, the volumes of gaseous reactants and products can be expressed as ratios of small whole numbers.

  3. Avogadro’s Law • 1811 : Equal volumes of gases at the same temperature and pressure contain an equal # of particles/moles.

  4. REMEMBER?!? • Molar Volume of a gas is 22.4 L / mol • No matter what the gas is as long as it’s at STP • EX: At STP, what is the volume of 14.36 mols of nitrogen gas?

  5. 11.2 The Ideal Gas Law • Ideal Gas Law: The mathematical relationship among pressure, volume, temperature and the # of moles of a gas. • Describes the behavior of an IDEAL gas. A gas described by the Kinetic-Molecular theory. PV= nRT • P- pressure V- volume • N- mole R- universal gas constant • T- temperature

  6. Universal gas constant (R) • It’s value depends on the Pressure units. P. 342

  7. Practice Problem • How many moles of gas at 100 °C does it take to fill a 2.5L flask at 1.50 atm? • Convert 100 °C to K … 100 + 273 = 373K PV= nRT • (1.50 atm) (2.5L) =n (0.0821 atm-L/mol-K) (373K) • = 0.122455777 moles

  8. More practice • What pressure, in atm, is exerted by 0.352 mol of H2 gas in a 4.08L container at 35°C? • A sample that contains 4.35 mol of a gas at 250 K has a pressure of 0.875 atm. What is the volume?

  9. Finding Molar Mass or Density from the Ideal Gas Law • When you are given GRAMS of a gas… you must convert to moles to use the Ideal Gas Law. • Use M for molar mass, m for mass & the equn: • PV = mRT/M or M = mRT/PV • Density (D) is mass/volume. Can use the Ideal Gas Law to solve for density. • D = MP/RT remember… M = molar mass

  10. Practice Problems • What is the density of 10.5g of argon gas at a pressure of 551 torr & a temperature of 25 °C? • What is the molar mass of a gas if 0.427g of the gas occupies a volume of 125 mL at 20.0 °C & 0.980 atm?

  11. 11.3 Stoichiometry of Gasesa review • By using what we’ve learned from Gay-Lussac and Avogadro… we can calculate Volume- Volume and Volume- Mass problems. • When converting Vol- Vol… use the coefficients • When converting Mass- Vol… use 22.4L/ mol , mole to mole, mole to mass.

  12. Practice Problems • What volume of O2 is needed to react completely with 0.699L of CO to form CO2 @ STP? • How many grams of water can be produced from the complete reaction of 3.77L of O2, at STP, with H2?

  13. 11.4 Effusion and Diffusion • REMEMBER… Diffusion: Gases can spread out and mix. The mixture will be spontaneous & uniform. Effusion: When a gas escapes a container though small holes. • Now the RATE at which a gas diffuses/effuses can be calculated…

  14. Graham’s Law of Effusion • Scottish Chemist, Thomas Graham studied the rates of diffusion and effusion. • Graham’s Law: the rates of effusion at the same temperature and pressure are inversely proportional to the square root of their molar masses.

  15. More massive particles move at a slower velocity than a less massive particle. Diffusion Effusion

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