A chemist has 2. 0 mol of methanol (CH3OH). The molar mass of methanol is 32. 0 g/mol. What is the mass, in grams, of the sample
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The question is incomplete, here is the complete question:
Ammonia has been studied as an alternative "clean" fuel for internal combustion engines, since its reaction with oxygen produces only nitrogen and water vapor, and in the liquid form it is easily transported. An industrial chemist studying this reaction fills a 5.0 L flask with 2.2 atm of ammonia gas and 2.4 atm of oxygen gas at 44.0°C. He then raises the temperature, and when the mixture has come to equilibrium measures the partial pressure of nitrogen gas to be 0.99 atm.
Calculate the pressure equilibrium constant for the combustion of ammonia at the final temperature of the mixture. Round your answer to 2 significant digits.
<u>Answer:</u> The pressure equilibrium constant for the reaction is 32908.46
<u>Explanation:</u>
We are given
Initial partial pressure of ammonia = 2.2 atm
Initial partial pressure of oxygen gas = 2.4 atm
Equilibrium partial pressure of nitrogen gas = 0.99 atm
The chemical equation for the reaction of ammonia and oxygen gas follows:
<u>Initial:</u> 2.2 2.4
<u>At eqllm:</u> 2.2-4x 2.4-3x 2x 6x
Evaluating the value of 'x':
So, equilibrium partial pressure of ammonia = (2.2 - 4x) = [2.2 - 4(0.495)] = 0.22 atm
Equilibrium partial pressure of oxygen gas = (2.4 - 3x) = [2.4 - 3(0.495)] = 0.915 atm
Equilibrium partial pressure of water vapor = 6x = (6 × 0.495) = 1.98 atm
The expression of for above equation follows:
Putting values in above equation, we get:
Hence, the pressure equilibrium constant for the reaction is 32908.46