Question

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.

Answer 1

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"`

Check back if you're not with me.