The three main components of dry air and the percentage of...

The three main components of dry air and the percentage of each are nitrogen $(78.08 \%),$ oxygen $(20.95 \%)$ and argon $(0.93 \%)$ (a) Calculate the partial pressure of each gas in a sample of dry air at $760 \mathrm{mm}$ Hg. (b) Calculate the total pressure exerted by these three gases combined.

The three main components of dry air and the percentage of each are nitrogen $(78.08 \%),$ oxygen $(20.95 \%)$ and argon $(0.93 \%)$
(a) Calculate the partial pressure of each gas in a sample of dry air at $760 \mathrm{mm}$ Hg.
(b) Calculate the total pressure exerted by these three gases combined.

Key Concepts

Partial Pressure

Partial pressure refers to the pressure that each gas in a mixture would exert if it alone occupied the entire volume of the mixture at the same temperature. It can be calculated by multiplying the mole fraction (or percentage in the case of ideal gases) by the total pressure of the mixture.

Gas Mixtures

Understanding gas mixtures is essential in many fields of science and engineering. The concept involves analyzing how different gases interact in a mixture, where their individual behaviors conform to the ideal gas law and Dalton's Law, allowing for calculations of properties like total and partial pressures even when multiple gases are present.

Dalton's Law of Partial Pressures

This law states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of individual gases. Each gas contributes a pressure proportional to its mole fraction in the mixture, regardless of the presence of other gases. It is fundamental in calculating the pressure exerted by each component within a gaseous mixture.

Mole Fraction

The mole fraction is a way of expressing the concentration of a component in a mixture, defined as the ratio of the number of moles of that gas to the total number of moles of all gases present. In problems involving partial pressures, percentages by volume for ideal gases can be directly taken as mole fractions, facilitating the use of Dalton's Law.

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