an autodidact meets a dilettante…

I think it would be amusing – for me anyway – if my blog posts were connected by threads, one leading to another. For example, the last post on aerosinusitis resulted from comments on my previous one about my recent air travel, and this one results from comments about pressure differentials in the last one, and so on and on.

Well, anyway, Boyle’s law.

Boyle’s law describes how the pressure of a gas increases as its volume decreases.

Take a set amount of gas – that’s to say, a certain mass of gas – and decrease its volume. Then the pressure it exerts increases in proportion. That’s to say, the relationship between volume and pressure is inversely proportional, given a constant temperature and mass. This relationship can be expressed in the formula PV = k, where k is a constant. In ‘word’ terms, the product of pressure and volume is constant, controlling for the other factors. If the pressure is calculated as 6, and the volume 2, then if the volume is doubled to 4, then the pressure will be halved to 3, with, on both occasions, the constant being 12. Another way of expressing this relationship is P₁V₁ = P₂V₂.

The English chemist/physicist, or ‘natural scientist’, Robert Boyle, first published the relational law in 1662, though he wasn’t the first to notice a relationship between pressure and volume. Nor did he fully understand the reason for the relationship, because gases were not then seen as molecular, with the molecules in kinetic relationship to each other. However, Boyle’s thinking was moving in the right direction, as he theorised that air – the gas on which he experimented – was ‘a fluid of particles at rest in between invisible springs’. Edme Mariotte of France independently formulated the law a little over a decade after Boyle.

The best physical explanation for the law emerged more than two centuries later, with work on the kinetic theory of gases by James Clerk Maxwell and Ludwig Boltzmann. This theory explains pressure within a container as  a result of atoms or molecules colliding with the container at various rates and velocities. It provides a molecular, microscopic accounting of such macroscopic measurements as pressure, volume and temperature. Einstein’s work on Brownian motion, the motion of dust or pollen particles as seen under a microscope, helped confirm the theory, on a level kind of in between the molecular and the macroscopic. Interestingly, the idea that macroscopic conditions might be the result of microscopic bodies in collision was put forward by Lucretius nearly 2000 years ago.

Boyle’s law treats of an ideal gas, something not known or considered at the time because gases under standard conditions of temperature and pressure behave essentially like ideal gases. Other ideal gas laws include Charles’ law, which is a law of volumes, Gay-Lussac’s law, which treats pressure, and Avogadro’s law, which covers the proportional relationship between volume and the number of moles present (molar volume). As always, improvements in technology led to the observation of a wider range of conditions requiring new hypotheses, the confirmation of which led to new knowledge – in this case, the kinetic theory.