I need help to understand something about the Haber process ...
Why is it that text books say that a change in pressure does not change the position of equilibrium and does not affect the value of Kp ?
Yet quite plainly if you increase the pressure in the Haber process you will increase the yield of Ammonia as the system compensates for increased pressure by shifting to the side with least moles of gas thereby reducing vol/pressure effects ... ??
But how can Kp therefore stay the same ? Here is how it is worked out ...and following that I have gathered what data I can. But it is many years since I studied Physical Chemistry at University and apart from anything else I have hunted the internet for an answer and cannot find one. So barring buying a text book about gas laws, I am hoping there is someone out there who can help ! Thanks !
Total Pressure measured in kPa or atm or Nm-2 ....
(conversion chart for changing between units http://www.convert-me.com/en/convert/pressure/nmsq.html ...)
The total pressure of a mixture of gases equals the sum of the individual gas pressures. Each partial pressure is the same percent of the total pressure as the percent each gas is of the total volume.
This table shows percentage of NH3 at various temps and pressure. At each temp you will see that the percentage of NH3 increases with increased pressure :
Atmospheres of Pressure
10 25 50 100 200 300 400 500 1000
100ºC 91.7 94.5 96.7 98.4 99.4
200ºC 50.7 63.6 74 81 89 94.6 98.3
300ºC 14.7 27.4 39.5 52.5 66.7 79.7 92.6
400ºC 3.9 8.7 15.3 25.2 38.8 48 55.4 61 79.8
500ºC 1.2 2.9 5.6 10.6 18.3 26 32 38 57.5
600ºC 13 17 21
700ºC 0.2 1.1 2.2 12.9
at 500ºC Kp = 1.45x10ˉ5 or 0.48 NH3
at 300ºC Kp = 4.34 x 10-3
at 400ºC Kp = 1.64 x10-4
at 450ºC Kp = 4.51 x 10 -5
at 472ºC Kp = 2.79 x 10-5 (7.38 atm H2 2.46 atm N2 0.166 atm NH3)
at 500ºC Kp = 1.45 x 10 -5
at 550ºC Kp = 5.38 x 10 -6
at 600ºC Kp = 2.25 x 10-6
Partial pressure of a gas in a mixture of gases is defined as the pressure the gas would exert if it alone occupied the entire volume.
pV=nRT where p is partial pressure
R=0.0820575 atm dm3K-1mol-1 or 8.3142 JK-1mol-1 (the more common SI unit)
The effect of changing the pressure on a gas phase reaction depends on the stoichiometry of the reaction
In this case there are 4 on the left (reactants) and 2 moles on the right (products).
If we compress a system by a factor of 10 that is at 500ºC where Kp =1.4 x 10-5 we get the following:
Before Compression After Compression
PNH3=0.12 atm PNH3 = 8.4 atm
PN2 = 2.4 atm PN2 = 21 atm
PH2 = 7.3 atm PH2 = 62 atm
Before the system was compressed the partial pressure of NH3 was only about 1% of the total pressure. After the system is compressed the partial pressure of NH3 is almost 10% of the total. The system was under stress from the pressure and it has now shifted in the direction which minimises the effect of this stress. It shifts towards the products because this reduces the number of particles of gas, thereby decreasing the total pressure on the system.
But if we slot this data into the Kp equation how can the Kp value possibly stay the same ? Surely the number must increase using these new partial pressures ?
Boiling point of :
N2 = -196ºCso ammonia will liquefy first if cool mixture down.
H2 = -253ºC
NH3 = -33ºC
Gas boiling points are higher at high pressure....
To get Kp from Kc
Kp = Kc(RT) to the power delta n where n is change is no of moles =2 in this occassion
Ideal Gas Law....
PV=nRT is initial pressure of one gas
pV=nRT to get partial pressure of one
This has loads of further info ...
- A Case Study of the Haber Process
- James Mungall tuition
- Le Chatelier's Principle