Henry's Law: Solubility, Temperature, and Pressure

Hello! Welcome back to your course on the theory of technical scuba diving.

In our previous lessons, we explored two fundamental physical laws. First, we used Dalton's Law to understand how the partial pressure of each gas in your breathing mix increases with depth. Then, we examined how this increasing pressure makes the gas you breathe denser, increasing the work of breathing and creating a risk of CO₂ retention.

Today, we'll investigate the third and perhaps most critical piece of the physics puzzle: what happens to these gases once they are in your lungs. This lesson focuses on how gases move from your lungs into your body's tissues, the very process that governs decompression.

Lesson 3: The Dissolving Gas

Learning Outcome: By the end of this lesson, you will be able to analyze the relationship between gas solubility, temperature, and pressure using Henry's Law.

Understanding this law is not just academic; it is the theoretical foundation for preventing decompression sickness ("the bends") and for planning every technical dive you will ever do.

1. The Core Principle: Pressure and Solubility

At its heart, Henry's Law is surprisingly simple. It describes how much gas can dissolve into a liquid. The most common real-world example is a carbonated beverage.

To get a clear, intuitive grasp of this concept, let's start with a short video.

As the video explains, Henry's Law states:

The amount of a given gas that will dissolve in a liquid is directly proportional to the partial pressure of that gas in contact with the liquid.

This means if you double the partial pressure of a gas over a liquid, you can dissolve twice as much of that gas into the liquid. This is a direct relationship.

• High Pressure -> High Solubility (Gas dissolves into the liquid)
• Low Pressure -> Low Solubility (Gas comes out of the liquid)

For a diver, the "liquid" is your blood and body tissues, and the "gas" is the inert gas you're breathing (primarily nitrogen when breathing air). The pressure is determined by your depth.

2. On-gassing and Off-gassing: Henry's Law in a Diver's Body

Let's formalize the process described in the video. To do this, we need to introduce a few key terms that are central to decompression theory.

Based on the reading, let's clarify these crucial concepts:

• Gas Tension: This is the pressure of a gas that is currently dissolved in a tissue. Think of it as the internal pressure of that gas within your body.
• Partial Pressure: This is the pressure of the gas in your lungs, as determined by your breathing mix and depth (Dalton's Law).
• Pressure Gradient: This is the difference between the partial pressure in your lungs and the gas tension in your tissues. This gradient is the driving force that causes gas to move. Gas always moves from a higher pressure to a lower pressure.
• Saturation: This is the equilibrium state where the gas tension in the tissues equals the partial pressure in the lungs. The pressure gradient is zero, and there is no net movement of gas.

Let's trace these concepts through a dive:

1. On the Surface: Your tissues are saturated with nitrogen. The nitrogen tension in your body is equal to the partial pressure of nitrogen in the air you're breathing (~0.79 ATA). The pressure gradient is zero.

2. Descent (On-gassing): As you descend, the ambient pressure increases, and so does the partial pressure of nitrogen in your lungs. Now, the partial pressure in your lungs is higher than the nitrogen tension in your tissues. This creates a pressure gradient, forcing nitrogen from your lungs into your blood and then into your tissues.

3. Ascent (Off-gassing): As you ascend, the ambient pressure drops. The partial pressure of nitrogen in your lungs becomes lower than the nitrogen tension in your tissues. The pressure gradient is now reversed. Nitrogen moves out of your tissues, into your blood, and is expelled through your lungs.

The danger, as the video explained, is when the ascent is too fast. The pressure drops so quickly that the dissolved nitrogen comes out of solution faster than the blood can transport it to the lungs. It forms bubbles directly in your tissues and bloodstream, causing Decompression Sickness (DCS).

3. Factors Influencing Solubility

Pressure is the primary driver, but it's not the only factor. The learning outcome also requires us to consider temperature and the properties of the gas itself.

Temperature

For gases dissolving in a liquid, the relationship is generally inverse:

• Higher Temperature -> Lower Solubility
• Lower Temperature -> Higher Solubility

This is the opposite of what we experience when dissolving a solid like sugar in water. A colder liquid can hold more dissolved gas.

While a diver's core body temperature is fairly stable, this principle is a fundamental property of the system. It also has implications in extreme cold-water diving or if a diver becomes hypothermic, but for now, it's important to know the basic relationship.

Gas Type and Liquid Type (Solubility Coefficient)

Henry's Law includes a "solubility coefficient," which is a way of saying that different gases dissolve differently in different liquids.

The UCSB reading Chapter 2 - Physics of Diving made a critical point: "nitrogen is five times more soluble... in fat than in water."

Since the human body is composed of many different types of tissues (blood, muscle, bone, fat), each with a different makeup and blood supply, they will all absorb and release inert gas at different rates.

• Tissues with a rich blood supply (like the brain and lungs) will saturate and de-saturate quickly.
• Tissues with a poor blood supply (like fat and dense cartilage) will saturate and de-saturate very slowly.

This is the entire basis for decompression models (like the Bühlmann model we will study later), which model the human body as a series of theoretical "tissue compartments" that all on-gas and off-gas at different half-times.

To reinforce the diving application and the role of different gases, let's watch one more short clip.

Conclusion

In this lesson, you've learned about Henry's Law and its profound implications for divers. We've moved beyond the gas in your cylinder and lungs to understand how it interacts with your body itself.

Key Takeaways:

• Henry's Law: The amount of gas that dissolves in a liquid (your tissues) is directly proportional to the partial pressure of that gas (in your lungs).
• Pressure Gradients Drive Gas Exchange: The difference between the partial pressure in your lungs and the gas tension in your tissues dictates whether you are on-gassing or off-gassing.
• Decompression Sickness: DCS is caused by a rapid decrease in pressure, which forces dissolved inert gases to come out of solution as bubbles within the body.
• Controlling Ascent is Key: A slow, controlled ascent with appropriate decompression stops keeps the pressure gradient small, allowing dissolved gas to be safely eliminated via the lungs.
• Other Factors: Gas solubility is also influenced by temperature (inversely) and the specific properties of the gas and tissues involved (e.g., nitrogen is highly soluble in fat).

You now have the three pillars of diving physics in place: Dalton's Law (partial pressure), the relationship between pressure and density, and Henry's Law (gas solubility). These principles work together to define the physiological challenges of diving.

In our next lesson, we will apply these concepts to manage the physiological effects of breathing different gas mixtures at depth by learning to calculate Equivalent Air Depth (EAD) and Equivalent Narcotic Depth (END).