PowerPoint to accompany Chapter 9b-2 Gases. Brown, LeMay, Bursten, Murphy, Langford, Sagatys:...

PowerPoint to accompany

Chapter 9b-2

Gases

Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia

Dalton’s Law of Partial Pressures

The total pressure of a mixture of gases equals the sum of the pressures that each would exert if it were present alone.

In other words,

Ptotal = P1 + P2 + P3 + …


Partial Pressures and Mole Fraction

Since P1 = n1RT/V

Then P1 = n1RT/V

P2 = n2RT/V

=n1

n2

= X (mole fraction)


Partial Pressures and Mole Fraction

So, when one collects a gas over water, there is water vapour mixed in with the gas.

To find only the pressure of the desired gas, one must subtract the vapour pressure of water from the total pressure.

Figure 9.13


Kinetic-Molecular Theory1. Gases consist of large numbers of

molecules that are in continuous, random motion.

2. The combined volume of all the molecules of the gas is negligible relative to the total volume in which the gas is contained.

3. Attractive and repulsive forces between gas molecules are negligible. (See note*)


N2 =28g/M

Maxwell-Boltzmann Distribution


Maxwell-Boltzmann Distribution


Kinetic-Molecular Theory

4. Energy can be transferred between molecules during collisions, but the average kinetic energy of the molecules does not change with time, as long as the temperature of the gas remains constant.

5. The average kinetic energy of the molecules is proportional to the absolute temperature.


Molecular Effusion and Diffusion

Effusion

The escape of gas molecules through a tiny hole into an evacuated space.

Diffusion

The spread of one substance throughout a space or throughout a second substance.

Figure 9.17

Gas Molecules spread uniformly

You cannae hide a fert!


Tae a Fert, Roy Williamson

Oh what a sleekit horrible beastieLurks in yer belly efter the feastieJust as ye set doon among yer kinThere sterts to stir an enormous wind.

The neeps and tatties and mushy peasStert workin like a gentle breezeBut soon the puddin wi the sauncie faceWill have ye blawin’ all ower the place.

Nae matter whit the hell ye daeA’bodys gonnae have tae payEven if ye try to stifle,It’s like a bullet oot a rifle.

Hawd yer bum tight tae the chairTae try and stop the leakin airShift yersel frae cheek tae cheekPrae tae God it disnae reek.

But aw yer efforts go assunderOot it comes like a clap o’ thunderRicochets aroon the roomMichty me, a sonic boom!

God almighty it fairly reeks;Hope I huvnae shit ma breeksTae the bog I better scurryAw whit the hell, its no ma worry.

A’body roon aboot me chokin,Wan or two are nearly bokinI’ll feel better for a whileCannae help but raise a smile.

Wis him! I shout with accusin glower,Alas too late, he’s just keeled owerYe dirty bugger they shout and stareI dinnae feel welcome any mair.

Where ere ye go, let yer wind gan freeSoonds like just the job fur meWhit a fuss at Rabbie's pertyOwer the sake o won wee ferty.


Graham’s Law of Effusion & Diffusion

The ratio of the rates of effusion of two gases is equal to the square root of the inverse ratio of their molecular masses or densities.

The effusion rate of a gas is inversely proportional to the square root of its molecular mass.

Mathematically, this can be represented as:

Rate1 / Rate2 = square root of (Mass2 / Mass 1)


Graham’s Law of Effusion & Diffusion

For the 2 common gases in air: the rate off effusion of Nitrogen and Oxygen.

N2, Nitrogen, has a molecular mass of 28.0 g. O2, Oxygen, has a molecular mass of 32.0 g. Therefore, to find the ratio, the equation would be:

RateN2/RateO2 = square root of 32.0 g / 28.0 g.

This works out to: RateN2/RateO2 = 1.069


Gas X2 effuses at 0.355 times that of O2

rx2 = 0.355 rO2

rx2 / rO2 = 0.355 = SQRT [(32.0 g/Mol)/Mx2]

(32.0 g/Mol)/ Mx2) = (0.355)2 = 0.126

Mx2 = (32.0 g/Mol)/0.126 = 254 g/Mol X g/Mol = 254/2 = 127 g/Mol Iodine is 126.9 so the gas was I2

Since WWII we have separated 2 Uranium isotopes 235 from 238 as UF6 gas

r235 / r238 = sqrt (352.04/349.03) = 1.0043

Applying Graham’s Law of Effusion


Real Gases Figure 9.21

In the real world, the behaviour of gases only conforms to the ideal-gas equation at: relatively high temperature and low pressure, low gas concentrations, low molecular weights and simple symmetric molecules or atoms.

For n=1,

for H2 (blue) and heavier gases (red, green). Note that heavy gases are attractive at low pressure & all are repulsive at high pressure. Why?


Deviations from Ideal Behaviour

The assumptions made in the kinetic-molecular model break down at high pressure and/or low temperature.

Figure 9.22 Figure 9.23


Corrections for Non-ideal Behaviour

The ideal-gas equation can be adjusted to take these deviations from ideal behaviour into account.

The corrected ideal-gas equation is known as the van der Waals equation.

nRT n2a P = --------- - -------- V – nb V2


The van der Waals Equation

) (V − nb) = nRTn2aV2(P +

Table 9.2


Examples of Gas Chemistry& Physics in the Real World

Atmospheric pollution: industry, engines, cows O3 generating & destroying reactions Volcanic aerosols Warfare: Cl2, Sarin, Mustard Gas, Phosgene Photosynthesis, O2 and the Carbon cycle Rocket & Jet Propulsion Gas Shocks, Combustion Engines Perfume and pheromones Clathrates and Gas Hydrates Explosives


End of Part 9b2:

Real Gas Applications

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