가스 혼합물: Dalton의 부분 압력 및 몰분율 법칙

Partial pressure is the pressure exerted by an individual gas component in a mixture, as if that gas occupied the container alone. Each gas in a mixture exerts its own pressure independently of other gases present, assuming no chemical reactions occur. The partial pressure depends on the number of molecular collisions from that specific gas component with container surfaces.

Dalton's law states that the total pressure of a gas mixture equals the sum of the partial pressures of all component gases. Since each gas occupies the same volume and temperature, their individual pressures add together without interference. This principle applies to ideal gas mixtures where gases do not chemically interact.

Mole fraction is the ratio of a component's moles to the total moles in a mixture. The partial pressure of any gas equals its mole fraction multiplied by the total pressure of the mixture. This relationship allows you to calculate individual gas pressures when you know the composition and total pressure.

Use the equation: partial pressure of component i equals its mole fraction times the total pressure. For example, if argon comprises 40% of a mixture (mole fraction 0.4) at 4 atm total pressure, argon's partial pressure is 0.4 × 4 atm = 1.6 atm. The remaining pressure belongs to other components.

Gas molecules are far apart and rarely collide with each other. Each gas component contributes pressure through its own molecular collisions with container walls, independent of other gases present. This independence holds true for ideal gases that do not chemically react, making pressure additive in mixtures.

The ideal gas law can be applied to each component separately or to the entire mixture. For a mixture, total pressure equals the total moles of all gases times RT divided by volume. Since all components occupy the same volume and temperature, the total pressure depends only on the combined number of moles present.

At constant temperature and volume, total pressure is determined solely by the total number of moles of gas present. Adding more gas molecules increases collisions with container walls, raising pressure proportionally. The identity of the gases does not matter, only their combined quantity.