How do you calculate the partial pressures of this problem?

A mixture containing 2.53 g each of CH4(g), C2H4(g) and C4H10(g) is contained in a 1.50 L flask at a temperature of 25°C.
(a) Calculate the partial pressure of each of the gases in the mixture.
(b) Calculate the total pressure of the mixture.
Answers in atmospheres.

1 Answer
Aug 13, 2016

Answer:

In a gaseous mixture, the pressure exerted by a component is the same as if it ALONE occupied the container............

Explanation:

.........and the total pressure is the sum of the individual partial pressures.

The above was a restatement of Dalton's Law of Partial Pressures, which was a very early experimental finding. From this, to find the pressure exerted by a component, all we need to do is find the total number of moles of each gas, and proceed from there.

In your problem there are #(2.53*g)/(16.04*g*mol^-1)" methane; "(2.53*g)/(28.05*g*mol^-1)" ethylene"; (2.53*g)/(58.12*g*mol^-1)" butane".#

Now we may use Dalton's Law to get the individual partial pressures:

#P_"methane"=(n_"methane"RT)/V# #=# #(2.53*gxx0.0821*L*atm*K^-1*mol^-1xx298*K)/(16.04*g*mol^-1xx1.50*L)# #=# #??atm#

And #P_"Total"=P_"methane"+P_"ethylene"+P_"butane"#

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