According to Dalton's Law of Partial Pressures, what is true about the total pressure of a gas mixture?

The total pressure of a gas mixture, according to Dalton's Law of Partial Pressures, is indeed equal to the sum of the pressures of all the individual gases present in that mixture. This law states that each gas in a mixture exerts pressure independently of the others, and the total pressure is simply the sum of these partial pressures.

This principle is fundamental in understanding how gases behave in different environments, particularly within the respiratory system where various gases are present, such as oxygen and carbon dioxide. For example, if you have a mixture of oxygen and nitrogen, the total pressure will be the sum of the pressure exerted by the oxygen and the pressure exerted by the nitrogen, regardless of their individual amounts.

Other choices do not accurately represent this law. The pressure of the heaviest gas alone does not determine the total pressure, nor does it reflect the combined effects of all gases. The average pressure of the gases is also not relevant to the total pressure, as each gas contributes its unique pressure. Lastly, total pressure is influenced by temperature changes, as a change in temperature can affect the kinetic energy of gas molecules and consequently their individual pressures in the mixture.