Henry’s Law

William Henry was an English physician and chemist who in 1803 proposed what is now called Henry’s law, which states that “At a constant temperature, the amount of a given gas dissolved in a given type and volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid. “

The Law essentially has two parts; one states that as pressure increases, solubility of gasses in liquids increases. The Second part states that as temperature increases, solubility of gasses in liquids decreases.

In simplistic terms:

• Under more pressure, a greater quantity of gas can be absorbed by a liquid
• Secondly, the cooler the liquid, more gas can be absorbed by it, as a liquid warms up, the gas escapes from it.

Mathematically the Law is expressed as:

P=KC     Where

P = the partial pressure of the gas

C = Concentration of the gas

K= Henry’s Law Constant

When trying to understand Henry’s Law, it helps to draw parallels to a bottle of Soda. Before the bottle is opened, its contents are under pressure, which causes the carbon dioxide in it to be soluble in the soda. As soon as you open the cap, you release the pressure causing the carbon dioxide gas to lose its solubility and escape in the form of bubbles or fizz.

In the same manner, as a diver descends, nitrogen inhaled has nowhere to escape and under pressure gets soluble in the bloodstream, muscles and tissues. This is no problem, until the diver begins his ascent.  As a diver ascends, the pressure is released and like the soda bottle, the nitrogen in the body tries to escape and may form bubbles if the diver ascends too quickly causing DCS or Decompression Sickness also known as the bends. This is the reason why divers ascend gradually, to allow the nitrogen to dissipate rather than form bubbles.

Henry’s Law also explains the reasons why divers are asked to not to take hot baths after a dive, or asked to abstain from doing strenuous activities or exercise. Based on the second portion of Henry’s Law, the increase in temperature caused by the exercise or hot bath may cause the nitrogen to become less soluble and increase the off-gassing possibly cause DCS or Decompression Sickness.

In the same manner, while diving in colder water, the on-gassing/absorption of nitrogen is greater, which should be taken into consideration, as it will allow the diver shorter dive times, and shallower dives.

Also Read: An Introduction to Scuba Gas Laws – Part 1 : Boyle’s Law
An Introduction to Scuba Gas Laws – Part 2: Charles’ Law

Photos Gaetan Lee, azteca90

© Aquaviews - SCUBA Blog, 2009. | An Introduction to Scuba Gas Laws – Part 3: Henry’s Law is an original article from Aquaviews.net | 2 comments

Charles’ Law

If you have read my first article in this series, An Introduction to Scuba Gas Laws – Part I: Boyle’s Law, you will already be familiar with Boyle’s Law which studies the relation between pressure and volume of a gas while temperature remains constant.  Charles’ law also studies volume of gas however taking into consideration the effects of temperature on the volume and Pressure of the gas.

Jacques Alexander Charles was a French scientist, mathematician, inventor and a balloonist who first studied the effects of temperature on the volumes of a gas and formulated Charles’ Law in 1787. The law states that “At constant pressure, the volume of a given mass is of an ideal gas increases or decreases by the same factor as its temperature increases or decreases.”

Mathematically Charles’ Law is expressed as:

P1 x V1     P2 x V2
______ = ________
T1               T2

Where’ T’ is temperature,’ P’ is pressure and ‘V’ volume.  ‘1’ indicates the initial temperature and volume while the’ 2’ indicates the final temperature and volume.

According to Charles’ Law, if we were to take a balloon filled with air and increase the temperature of the air inside, the volume of air would increase causing the balloon to expand.  This is caused by the heating of the molecules of air inside the balloon causing them to move rapidly. In the same manner if we cooled the balloon in a freezer, the volume of air decrease, making the balloon look partially deflated.

In the same manner a Scuba tank filled to 3000 PSI if heated, would increase the pressure in the tank as the volume of the air inside the tank is at a constant.

So how is this applicable to Scuba Diving? For starters, Charles’ law helps divers understand the hazards of leaving scuba tanks out in the hot sun, or why we should never leave tanks in the trunk of a hot car. The gas under pressure subjected to heat can cause the tank to explode. A scuba tank filled to capacity with compressed air at 3000 psi could just as easily go up to 3400-3500 psi if heated. Proper storage of air tanks on the dive boat too is crucial to ensure the tanks aren’t left to bake in the sun.

The law explains why the pressure in a scuba tank recently filled may appear to drop or increase slightly depending on whether the tanks were filled outdoors, or indoors. The law also explains why scuba tanks when being filled with compressed air, get hot.

I’ll admit knowledge of Charles’ Law isn’t as useful to us as perhaps Boyle’s Law, nevertheless it helps us understand everyday phenomenon and helps us better understand the remaining gas laws as well.

© Aquaviews - SCUBA Blog, 2009. | An Introduction to Scuba Gas Laws – Part 2: Charles’ Law is an original article from Aquaviews.net | One comment