My textbook says that that for a mixture of #15.2*g# helium gas, and #30.6*g# dioxygen gas, the partial pressure is dominated by the helium. Can you find #P_"He"#, and #P_"dioxygen"# and #P_"Total"#?
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Well, I think the text book is right mate. Because #"helium gas"# is present in a large molar quantity, its partial pressure should dominate. Let's see what we get.
#"Dalton's Law of Partial Pressures"# states UNEQUIVOCALLY that in a gaseous mixture, the partial pressure exerted by a gaseous component is the SAME as it would exert if it ALONE occupied the container.
And thus #P_"He"=(n_"He"RT)/(V)=((15.2*g)/(4.0*g*mol^-1)xx0.0821*(L*atm)/(K*mol)xx295*K)/(5.00*L)#
#P_"He"=18.4*atm#
#P_(O_2)=(n_(O_2)RT)/(V)=((30.6*g)/(32.0*g*mol^-1)xx0.0821*(L*atm)/(K*mol)xx295*K)/(5.00*L)#
#P_"dioxygen"=4.63*atm#.
#"Dalton's Law of Partial Pressures"# further states that the total pressure, #P_"Total"# is the sum of the individual partial pressures.
So here #P_"Total"=P_"He"+P_(O_2)=(18.4+4.63)*atm=23.0*atm#, as required. Have another look at your calcuations. It is all too easy to make an error, and everybody has done this.
