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5.1 Pressure PowerPoint PPT Presentation

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5.1 Pressure. Pressure a. Devices used to measure pressure 1. Barometer: measures atmospheric pressure *Atmospheric pressure results from the mass of air being pulled toward the center of the earth by gravity.

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5.1 Pressure

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5.1 Pressure

  • Pressure

    a. Devices used to measure pressure

    1. Barometer: measures atmospheric pressure

    *Atmospheric pressure results from the mass of air being pulled toward the center of the earth by gravity.

    2. Manometer: measures the pressure of a gas in a container

5.1 Pressure

b. Units of Pressure

1. mmHg = Torr (invented the first barometer)

2. 1 atm = 760 mmHg = 760 torr = 1.013 x 105 Pa

3. Pressure = force/ area

5.2 Gas Laws

  • Gas Laws

    *pressure, temp, vol, amount

    A. Boyle’s Law: volume and pressure are inversely related at constant temp

    PV = k

    P1V1 = P2V2

5.2 Gas Laws

B. Charles’ Law: Kelvin temperature and volume vary directly with each other at constant temperature.

*if you graph a line of V vs. T all lines extrapolate to zero volume at 0K.

*V1/ T1 = V2/ T2

5.2 Gas Laws

C. Avogadro’s Law: for a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas present

* V1/ n1 = V2/ n2

5.3 The Ideal Gas Law

  • Ideal Gas Law

    A. PV = nRT, R = 0.08206 Latm/ Kmol

    B. Gas laws describe ideal gases, not real gases

    C. Real gas behavior approaches ideal at high temps and low pressures

5.4 Gas Stoichiometric

  • Gas Stoichiometry

    A. The moalr volume of a gas at STP is 22.4 L ideally

    B. STP 0oC, 1atm (273 K, 1.013 x 105Pa)

5.4 Gas Stoichiometric

  • CH4 (g) 2.80L @25oC 1.65 atm

  • O2(g) 35.0 L @31oC 1.25 atm

    CH4(g) + 2O2(g)  CO2(g) + 2H2O (g)

    CO2(g) __L @ 125oC 2.50 atm

5.4 Gas Stoichiometric

C. n = grams of gas/ molar mass

D = m/v

P = dRT/ molar mass

5.5 Dalton’s Law of Partial Pressure

  • Dalton: for a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone.

    A. Ptotal = P1 + P2 + P3 + …

    B. Calculate the partial pressure using the ideal gas law

    Ptotal = (n1 + n2…) RT/V

    C. Total pressure depends on the total number of moles of gas, not the identity of the particles

5.5 Dalton’s Law of Partial Pressure

D. Mole fraction is the ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture (x = chi)

x = n2/ ntotal = P2/ Ptotal

n = PV/RT if V, R, T are constant then n = P

*the partial pressure of a particular component of a gaseous mixture is the mole fraction of that component times the total pressure

5.6 The Kinetic Molecular Theory of Gases

  • Kinetic Theory

    -based on observations used to explain and predict the behavior of ideal gases

    A. Volume of particles can be assumed to be negligible

    B. Particles are in constant motion, pressure is created by the collision of particles with the walls of the container

    C. The particles are assumed to exert no forces on each other

    D. The average KE of a collection of gas particles is assumed to be directly proportional to the Kelvin temp. of the gas.

5.7 Effusion and Diffusion

  • Effusion: passage of a gas through tiny holes in the container

    A. Graham’s Law: the rate of effusion of a gas is inversely proportional to the square root of the mass of the particles of gas

    B. Diffusion: mixing of gases due to random motion of particles

5.8 Real Gases

  • Real Gases

    A. No gas exactly follows ideal gas behavior (close under high temp, and low pressure)

    B. Van der Waal’s equation: describes the behavior of a real gas by correcting for the actual volume of gas particles and forces between gas particles

5.8 Real Gases

[Pobs + a(n/v)2] x (V – nb) = nRT

Pobs = observed pressure

a(n/v)2 = pressure correction

V = volume of container

V – nb = volume correction

*a and b are constants obtained from observation of real gases

5.9 Chemistry in the Atmosphere

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