Ideal gas molecules ignore each other. Put two gases in one vessel and each behaves exactly as if it owned the place, so the pressures simply add — that is Dalton's law of partial pressures. The practical form is $P_i = x_i\,P_{\text{total}}$, read aloud P-sub-i equals x-sub-i times P-total. The subscript $i$ means “whichever component you are asking about”: $P_i$ is that component's partial pressure in kPa, $P_{\text{total}}$ is what the gauge on the vessel reads, and $x_i$ is that component's mole fraction.

The mole fraction is built first: $x_1 = \dfrac{n_1}{n_1 + n_2}$, where $n_1$ is the amount of the component you want in moles and $n_2$ is the amount of everything else. Subscripts 1 and 2 here name the two COMPONENTS, not a before and an after — this is the one place in the chapter where they do. Moles over moles, so $x$ carries no unit at all, sits between 0 and 1, and all the mole fractions in a mixture add to exactly 1. That nakedness is what lets it multiply a pressure and hand the kilopascals straight back.

Read it backwards and it is a measurement: a probe reporting a partial pressure is reporting a composition. That is how a blood-gas analyser and a stack monitor both earn their keep — and why dry air's 21 % oxygen shows up as about 21 kPa of the 101 kPa pressing on you right now.