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Honors Chemistry

Honors Chemistry. Chapter 5 Gases. 5.1 Gases. Temperature vs. Intermolecular attraction Atomic Gases Noble Gases, H 2 , N 2 , O 2 , F 2 , Cl 2 Molecular Gases Usually light molecules with weak attraction forces Eg: HCl, CO 2 , NH 3 , H 2 S, NO 2 Ionic Compounds

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Honors Chemistry

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  1. Honors Chemistry Chapter 5 Gases

  2. 5.1 Gases • Temperature vs. Intermolecular attraction • Atomic Gases • Noble Gases, H2, N2, O2, F2, Cl2 • Molecular Gases • Usually light molecules with weak attraction forces • Eg: HCl, CO2, NH3, H2S, NO2 • Ionic Compounds • Strong forces; not normally gases

  3. 5.2 Pressure • Force per unit area • P = F/A • N/m2 unit defined as Pascal (Pa) • Standard air pressure = 101.325 kPa • Also called 1 atmosphere (atm) • Measured by unequal mercury levels • Manometers and barometers • Common unit called mmHg (or Torr) • Standard air pressure = 760 mmHg

  4. 5.2 Dimensional Analysis • Convert 75.0 kPa to mmHg • 75.0 kPa 760 mmHg----------- x --------------- = 563 mmHg 1 101.325 kPa • Try this one • Convert 1.25 atm to kPa

  5. 5.3 Boyle’s Law • Pressure is inversely proportional to volume • Hold temperature and amount of gas constant • V a 1/P • V = k x (1/P) • PV = k • Best used with changing conditions • P1V1 = P2V2

  6. 5.3 Boyle’s Law Problems • A 175 mL sample of methane is stored at 125 kPa. What pressure is needed to compress the gas to a volume of 50.0 mL? • P1V1 = P2V2 • (125 kPa) (175 mL) = P2 (50.0 mL) • P2 = 438 kPa • Try this one • A sample of argon occupies 476 mL at 650 Torr. Find the volume at 975 Torr.

  7. 5.3 Charles’ Law • Also credited to Gay-Lussac • Volume is directly proportional to temperature • Hold pressure and amount of gas constant • V a T • V = kT • Linear relationship • Must use Kelvins! • V1 V2--- = ---T1 T2

  8. 5.3 Charles’ Law Problems • A 5.00 L helium balloon is heated from 20oC to 75oC. Find its new volume. • V1/T1 = V2/T2 • 5.00 L V2--------- = --------293 K 348 K • V2 = 5.94 L • Try this one • A 670 mL sample of chlorine is stored at 50oC. At what temperature will its volume be 450 mL?

  9. 5.3 More Gas Laws • Another form of Charles’ Law • Pressure is directly proportional to temperature • P = kT • P1/T1 = P2/T2 • Avogadro’s Law • Volume is directly proportional to the amount of gas present • V a n • Volume relationships in chemical reactions

  10. 5.3 Avogadro’s Law Problems • How many liters of hydrogen are needed to completely react with 1 liter of oxygen? • 2 H2 + O2 2 H2O • 2 mol hydrogen react with 1 mol oxygen • V a n, so…. • 2 L hydrogen react with 1 L oxygen • Try this one • How many liters of ammonia are formed when 1 L of hydrogen reacts with excess nitrogen?

  11. 5.4 The Ideal Gas Equation • Ideal Gas • No intermolecular attraction forces • Particles have no volume • Combine Boyle’s, Charles, and Avogadro’s Laws • PV = nRT • STP = 1 atm, 273 K • Molar volume of a gas = 22.414 L at STP • R = 0.0821 atm L / mol K

  12. 5.4 Ideal Gas Equation Problems • A sample of fluorine occupies 3.65 L at 45oC and 2.50 atm. How many moles of fluorine are present? • PV = nRT • (2.50 atm)(3.65 L) = n (0.0821)(318 K) • n = 0.350 mol • Try this one • A 0.500 mol sample of propane occupies 2.15 L. If the temperature is 28oC, find the pressure.

  13. 5.4 Gas Density • Since n = m/M…. • PV = (m/M) RT • MPV = mRT • Divide by V to get density (m/V) • MP = rRT • Gas density expressed in g/L

  14. 5.4 Gas Density Problems • Find the density of nitrous oxide at STP. • First, find molecular mass of N2O • MP = rRT • (44.0 g/mol)(1.00 atm) = r (0.0821)(273 K) • r = 1.96 g/L • Try this one…. • A gas is found to have a density of 2.54 g/L at 15oC and 1.50 atm. Find its molecular mass.

  15. 5.5 Gas Stoichiometry • Mass-Mass problems (review) • Volume-Volume problems • Volume is proportional to moles, so…. • Mol relationship from reaction can be used directly • No conversions needed!

  16. 5.5 Volume-Volume Problem • 2 H2 + O2 2 H2O • If 3.25 L of oxygen react, how many liters of water vapor are formed? • 3.25 L O2 2 L H2O------------- x ------------ = 6.50 L H2O 1 1 L O2 • Volume-Volume is just Avogadro’s Law!

  17. 5.5 Mass-Volume Problems • Key step – get to moles! • Mass conversion – use molecular mass • Volume conversion – use gas equation • Need to know temperature and pressure conditions

  18. 5.5 Mass-Volume Problems • 25.0 g of sodium react with excess water at STP. How many liters of hydrogen are produced? • 2 Na + 2 H2O  2 NaOH + H2 • 25.0 g Na 1 mol Na 1 mol H2------------ x ----------- x ----------- = 0.543 mol 1 23.0 g Na 2 mol Na • Now use ideal gas equation to get volume

  19. 5.5 Mass-Volume Problems • PV = nRT • (1.00 atm) V = (0.543 mol)(0.0821)(273 K) • V = 12.2 L • Try this one • Potassium chlorate decomposes into potassium chloride and oxygen gas. How many grams of KClO3 are needed to produce 5.00 L of oxygen at 0.750 atm and 18oC? • Hint: This one is backwards!

  20. 5.6 Dalton’s Law • Partial pressure – the pressure of an individual gas in a mixture of gases • Total pressure of a mixture equals the sum of the partial pressures of each gas • Pt = P1 + P2 + P3 + ... • Partial pressure is proportional to the mol fraction (X1 = n1 / nt) • P1 = X1 Pt

  21. 5.6 Dalton’s Law • 2.00 mol He is mixed with 1.00 mol Ar. Find the partial pressure of each at 1.75 atm pressure. • XHe = 2.00 mol / 3.00 mol = 0.667 • XAr = 1.00 mol / 3.00 mol = 0.333 • PHe = (0.667) (1.75 atm) = 1.17 atm • PAr = (0.333) (1.75 atm) = 0.583 atm • Try this... • Find the partial pressure of oxygen in air if it makes up 21% of the Earth’s atmosphere by volume. (Note: The volume gives you the mole ratio because of Avogadro’s law.)

  22. 5.7 Kinetic Molecular Theory • Explains gas behavior in terms of molecular motion • Energy • Work done by a moving object • Measured in SI unit Joule (J) • Kinetic energy • Energy due to motion • K = ½ mv2 • KMT is a simplification of reality (ideal gas)

  23. 5.7 Kinetic Molecular Theory • Gas molecules are separated by great distances • They can be treated as “point masses” • Gas molecules are in constant random motion • Frequent elastic collisions (no energy lost) • No attractive or repulsive forces • Average K is proportional to Temperature

  24. 5.7 Distribution of Molecular Speeds • Maxwell-Boltzmann Distribution • Molecular speeds distributed around average • Peak velocity depends on temperature and on molec. mass • Root Mean Square Speed • _____vrms = √3RT/M • Rate of diffusion

  25. 5.8 Deviations from Ideal Behavior • We made approximations! • Point masses • No intermolecular forces • These approximations become bad at... • High pressure • Low temperature • Liquefaction • van der Waals Equation • (P + an2/V2) (V – nb) = nRT

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