Start by using the ideal gas equation to figure out the the number of moles $O_2$ used.
$ P*V= n*R*T$
Where
$P" "$ is the pressure expressed in atm.
$V" "$ is the volume expressed in L.
$n" " $ is the number of moles.
$R" "$ is the universal gas constant, it has a value of $0.0821\ L* atm. mol^-1*K^-1$
$T" " $is the Kelvin temperature.
Now rearrange the ideal gas equation and solve for n.
$n = (P*V)/(R*T)$
$n = (12.9 \ Lxx1.2 \ atm)/(0.0821\ L* atm. mol^-1*K^-1 xx 297 \ K)$
$n = (12.9 \ cancel(L)xx1.2 \ cancel( atm))/(0.0821\ cancel(L) *cancel(atm) *mol^-1*cancel(K^-1) xx 297 \ cancel(K))$
$n = 0.635 \ mol.$
In the second step, write a balanced chemical equation for the reaction and use the of the equation to figure out the moles of ethylene used.
$C_2H_4 + 3O_2 -> 2CO_2 + 2H_2O$
$0.635 \ mol. O_2 xx (1\ mol. \ C_2H_4)/(3 \ mol. \ O_2) $
$0.635 \ cancel (mol. O_2) xx (1\ mol. \ C_2H_4)/(3 \ cancel( mol. \ O_2))$
$0.212 \ mol. C_2H_4$
