Understanding Sea Divers' Gas Mixtures and High Pressure

Sea divers operate under significant pressure as they descend into the water. This high pressure affects the behavior of gases, particularly their solubility in liquids like blood. The air we normally breathe is about 21% oxygen and 78% nitrogen. While nitrogen is relatively inert at normal atmospheric pressure, its behavior changes drastically under the high pressure experienced during deep dives.

The Challenge of High Pressure Underwater

According to certain scientific principles, the solubility of gases in a liquid increases with the partial pressure of the gas above the liquid. Underwater, the pressure is much higher than at the surface. This increased pressure leads to a higher partial pressure of the gases in the diver's breathing mixture, which in turn causes more gas to dissolve in the diver's blood and tissues. When divers ascend, the pressure decreases. If the ascent is too rapid, the dissolved gases, particularly nitrogen, become less soluble and can come out of solution as bubbles. These bubbles can block blood vessels and cause severe pain, tissue damage, and even death. This condition is known as decompression sickness, or "the bends". Also, high partial pressure of nitrogen can cause nitrogen narcosis, affecting the diver's judgment and coordination.

How Gas Mixtures Address the Problem

To mitigate these risks, sea divers often use special gas mixtures instead of regular air for deep dives. The mixture mentioned in the question contains Helium (He), Nitrogen (N₂), and Oxygen (O₂) in specific percentages (11.7% He, 56.2% N₂, and 32.1% O₂). The key component added is Helium. Helium is used because it is much less soluble in blood and tissues than nitrogen. By replacing a significant portion of the nitrogen with helium, the amount of gas that dissolves in the diver's body at high pressure is reduced. This lowers the risk of bubble formation during ascent.

Applying the Relevant Gas Law

The principle that explains why the solubility of a gas in a liquid is proportional to its partial pressure above the liquid is known as Henry's Law. This law is fundamental to understanding the behavior of gases dissolving in blood under varying pressures, which is exactly the scenario faced by sea divers.

Henry's law can be expressed mathematically as:

\(C = k_H P\)

Where:

• \(C\) is the solubility of the gas in the liquid (e.g., concentration).
• \(k_H\) is Henry's law constant (specific to the gas, solvent, and temperature).
• \(P\) is the partial pressure of the gas above the liquid.

This equation shows that increasing the partial pressure (\(P\)) of a gas above a liquid increases its solubility (\(C\)) in the liquid.

Evaluating the Options

Let's look at why Henry's Law is the correct explanation and why the other options are not:

• Graham's law of diffusion: This law deals with the rate at which gases diffuse, which is related to their molar mass. It does not explain the solubility of gases in liquids under pressure.
• Raoult's law: This law relates the vapor pressure of a solution to the vapor pressure of the pure solvent and the mole fraction of the solute. It is primarily used for liquid-liquid solutions or ideal solutions and is not directly applicable to the solubility of gases in liquids under external pressure in this context.
• Dalton's law of partial pressure: This law states that the total pressure exerted by a mixture of gases is the sum of the partial pressures of the individual gases. While useful for calculating the partial pressures of He, N₂, and O₂ in the breathing mixture at a given depth, it does not explain how these gases dissolve in the diver's blood.
• Henry's law: This law directly relates the solubility of a gas in a liquid to its partial pressure above the liquid. The increased pressure underwater leads to higher partial pressures of the breathing gases, increasing their solubility in the diver's blood. The use of a mixture with less soluble helium is a direct application of understanding Henry's law to minimize dissolved gas volume at depth.

Therefore, the use of the specific gas mixture by sea divers to cope with high pressure under water is an application of Henry's law, which governs the solubility of gases in liquids under pressure.

Revision Table: Key Concepts

Additional Information: Diving and Gas Solubility

Diving physiology is a complex field that relies heavily on the principles of gas laws. Nitrogen narcosis occurs because nitrogen at high partial pressures affects the central nervous system. Helium is much less narcotic than nitrogen at the same partial pressure, making it safer for deep dives. The specific percentages of gases in a diver's mixture (like Trimix, which includes Helium, Nitrogen, and Oxygen) are carefully calculated based on the planned depth and duration of the dive to ensure sufficient oxygen for metabolism while minimizing the risk of narcosis and decompression sickness, all governed by the solubility principles described by Henry's Law.

Oxygen toxicity is another factor; while essential for life, oxygen at very high partial pressures (which occur when diving deep with standard air or rich oxygen mixtures) can also become toxic. This is why the oxygen percentage in the breathing mixture might be reduced for very deep dives, compensated by increasing the inert gas portion (Nitrogen and/or Helium).

Understanding the relationship between pressure, gas solubility, and partial pressures is crucial for safe diving practices. Henry's Law provides the fundamental basis for this understanding regarding gas solubility in the diver's body.