Examples of Dalton's Law of Partial Pressures Example

Explore practical examples to understand Dalton's Law of Partial Pressures.
By Jamie

Introduction to Dalton’s Law of Partial Pressures

Dalton’s Law of Partial Pressures states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of each individual gas. This principle is essential in various scientific fields, including chemistry, physics, and engineering, allowing for accurate calculations and predictions in gas behavior.

1. Breathing and Partial Pressures in the Lungs

In the context of human respiration, Dalton’s Law helps explain the behavior of different gases in the lungs. When we inhale, our lungs fill with air, which consists of various gases, including oxygen (O₂) and nitrogen (N₂).

In a healthy adult, the atmospheric pressure at sea level is approximately 760 mmHg. The composition of air is about 21% oxygen and 78% nitrogen. To calculate the partial pressures:

  • Partial pressure of O₂: 0.21 * 760 mmHg = 159.6 mmHg
  • Partial pressure of N₂: 0.78 * 760 mmHg = 593 mmHg

Thus, the total pressure in the lungs can be derived from these partial pressures:
Total Pressure = 159.6 mmHg + 593 mmHg = 752.6 mmHg

This example illustrates how Dalton’s Law applies to biological systems, helping us understand gas exchange during respiration.

Notes

  • Variations in altitude can affect these calculations, as atmospheric pressure decreases with elevation.

2. Gas Mixtures in Chemical Reactions

In a laboratory setting, chemists often deal with mixtures of gases when conducting experiments. For instance, consider a reaction involving hydrogen (H₂) and oxygen (O₂) gases in a closed container. If both gases are at a pressure of 200 mmHg, Dalton’s Law allows us to determine the total pressure in the system after they are mixed.

  • Partial pressure of H₂: 200 mmHg
  • Partial pressure of O₂: 200 mmHg

Applying Dalton’s Law:
Total Pressure = Partial pressure of H₂ + Partial pressure of O₂ = 200 mmHg + 200 mmHg = 400 mmHg

This example is crucial for predicting the behavior of gas reactions, particularly in stoichiometric calculations.

Notes

  • Reactions involving gases often change partial pressures, especially if gases react to form new products.

3. Understanding the Behavior of Gases in Scuba Diving

Scuba divers must understand the principles of gas laws, including Dalton’s Law, to ensure safety underwater. When a diver descends, the pressure increases, affecting the partial pressures of the gases in their breathing mixture (usually air). At a depth of 10 meters, the pressure is approximately 2 atmospheres (or about 1520 mmHg).

Assuming the diver breathes air, the partial pressures at this depth are:

  • Partial pressure of O₂: 0.21 * 1520 mmHg = 319.2 mmHg
  • Partial pressure of N₂: 0.78 * 1520 mmHg = 1187.2 mmHg

Total Pressure = 319.2 mmHg + 1187.2 mmHg = 1506.4 mmHg (close to 1520 mmHg due to minor variations)

Understanding these pressures is vital for avoiding nitrogen narcosis and managing decompression.

Notes

  • Divers often use special gas mixtures, such as nitrox, to alter the proportions of oxygen and nitrogen, influencing the partial pressures involved.