Chemical Forums
Chemistry Forums for Students => Physical Chemistry Forum => Topic started by: Adamcp898 on May 15, 2009, 03:22:57 PM
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Here's what looks like a simple enough kinetics problem but for whatever reason I can't quite seem to get it
"Ether decomposition is as follows:
(CH3)20 :rarrow: CH4 + H2 + CO
The rate of this reaction was determined by monitoring pressure changes in a constant volume container maintained at 450°C.
If the rate constant was 3.2x10-4 s-1 and the initial pressure was 0.350atm,
determine the pressure in the reaction flask after 8mins.
Comment on the magnitude of the pressure change and if any
particular precautions should be taken in carrying out this experiment."
Could the pressure be constant since the temperature and volume remain constant and the reaction ratios are all 1:1?
Or since, from the units used to describe the rate, the reaction is first order could you just use the rate equation for a first order reaction i.e. ln([A]f/[A]0) = -kt, to calculate it, using the pressure as an indication of concentration therefore having ln(x/0.350)? Doing it this way gives me an answer similar to the 0.350atm initial pressure.
Or should I calculate the number of moles from PV = nRT initially and use this value as the initial concentration and have ln(x/number of moles) and work it out using the above rate equation for a first order reaction?
Thanks to anyone who can shed some light on the matter for me.
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Just a start, how many moles of gas are there initially, how many moles when the reaction is complete? Do you still think that the pressure will be constant?
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Just a start, how many moles of gas are there initially, how many moles when the reaction is complete? Do you still think that the pressure will be constant?
Well since all the products are gases and we go from having 1 mole of one thing to having one mole each of three things then no I suppose the pressure shoud not be constant?