Usually drink, usually dance, usually bubble
EI has relations with a Street Style blog over at Stitsh.com, and recently a little vid caught our attention over there. Click here and scroll down to 28.06.08.
It reminded us of a little law that the examiners sometimes like to question, that is Henry’s Law:
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.
Okay, so what does that mean.
Most of the time we refer to Henry’s law by the formula p=kc (that’s one way of looking at it - Ed).
Another way is to say:
where:
is approximately 2.718, the base of the natural logarithm 
is the partial pressure of the solute (the gas being dissolved) above the liquid in which is being dissolved.
is the concentration of the solute in the solution
is the Henry’s Law constant, which has units such as L·atm/mol, atm/(mol fraction) or Pa·m3/mol (this is so that the dimensions all work out correctly - the funny thing about constants is that they usually can be expressed in many different units, depending on what units the rest of the equation is being calculated in….more on that another time).
(In other words, most of the time, we take the natural logarithms of both sides).
The pressure above a solution dictates how many collisions occur between the gas and the liquid. So if you increase the pressure above the solution, the partial pressure of the gas increases, the number of collisions increases, and more gas is dissolved. What will then happen is that an equillibrium will be achieved, where the number of molecules of gas crashing into the surface of the liquid will be the same as the number of molecules leaving the surface of the liquid.
The more observant amongst you will have realised that temperature hasn’t been mentioned yet except in the definition….
So what effect does temperature have?
Well, think of a can of “fizzy pop” (you’re showing your age there - Ed). When it comes out of the fridge, it’s not that fizzy, is it? However, the longer you leave it standing around, the closer it’s temperature comes to room temperature, and then when you go back to the can, first it will seem quite gassy, and then eventually it will go flat. This is because the gas in the drink is coming out of solution. The gas solubility relationship with temperature is very similar to the reason that vapor pressure increases with temperature. (This is Gay-Lussac’s Law: The pressure of a given number of moles (given amount) of gas, is directly proportional to its temperature in Kelvin (absolute temperature scale), when the volume is kept constant. Better known as P/T=k).
Increased temperature causes an increase in kinetic energy, which in a gas causes either expansion or an increase in pressure, or in this instance, more movement of the molecules, which break free of the surface of the solution! (The surface could be the gas side of a small bubble of gas trapped within the solution, which is one reason we get bubbles!)
If you want to see another demonstration of Henry’s law in action, look at a pan of water. As you warm the pan, small bubbles start to form, well before the pan reaches 100°C (373K). Those bubbles are air coming out of solution.
So why do the examiners like this concept: the Bends.
Decompression Sickness occurs when gas (specifically nitrogen) is breathed at higher than atmospheric pressure, and the diver then returns to atmospheric pressure without allowing the gas to come out of solution slowly, resulting in gas bubble formation, and hence, “the bends” (gas in the joints) and “the staggers” (gas bubbles in the brain causing confusion and ataxia) and “the chokes” (probably PE).
It is also a concept that comes into play when talking about Ostwald and Bunsen coefficients….(more on that another time).
(Equations courtesy of Wikipedia)
