Chapter 5 Review: Gases. Section 1: Pressure. Pressure: the pressing of the particles of a gas against its surroundings Atmospheric pressure is the result of the masses of the gases in Earth’s atmosphere being pulled to Earth due to gravity
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1. A sample of hydrogen (H2) gas has a volume of 8.56 L at a temperature of 0oC and a pressure of 1.5 atm. Calculate the moles of H2 molecules present in this gas sample.
2. A sample of diborane gas (B2H6) has a pressure of 345 torr at a temperature of -15oC and a volume of 3.48 L. If conditions are changed so that the temperature is 36oC and the pressure is 468 torr, what will be the volume of the sample?
((1.000mol) x (.08026 L x atm/K x mol) x (273.2 K)/1.000 atm) = 22.42
This volume is the molar volume of an ideal gas. Of course, no gas actually has a molar volume of 22.42, but there are close deviations.
Molar mass of gas
n= (grams gas/molar mass) = m/molar mass
substitute into ideal gas equation-
(m/molar mass)RT/V = mRT/V(molar mass)
since mass/volume = density
P= (dRT)/molar mass
Consider the three flasks in the diagram below. Assuming the connecting tubes have negligible volume, what is the partial pressure of each gas and the total pressure when all the stopcocks are opened?
First flask: He --- 200. torr, 1.00L
Second flask: Ne --- 0.400 atm, 1.00L
Third flask: Ar --- 24.0 kPa, 2.00 L
Thomas Graham’s Law
No gas follows the ideal gas law exactly, however many gases at low pressure and high temperature come close.
To correct the assumptions made by the Kinetic Molecular Theory, Johannes van der Waals modified the ideal gas law for real gases :
P=((nRT)/(V-nb)) - a(n/V)2
(P + a(n/V)2)x (V-nb) = nRT
This takes into account molecular size(b) and molecular interaction forces(a).
b will be greater for large particles and a is higher for polar molecules.