JOSEPH LOUIS GAY-LUSSAC (1778-1850)

1808 – France

‘Volumes of gases which combine or which are produced in chemical reactions are always in the ratio of small whole numbers’

One volume of nitrogen and three volumes of hydrogen produce two volumes of ammonia. These volumes are in the whole number ratio of 1:3:2

N2 + 3H2 ↔ 2NH3

Along with his compatriot Louis Thenard, Gay-Lussac proved LAVOISIER’s assumption, that all acids had to contain oxygen, to be wrong.

portrait of GAY-LUSSAC ©

GAY-LUSSAC

Gay-Lussac re-examined JACQUES CHARLES’ unpublished and little known work describing the effect that the volume of a gas at constant pressure is directly proportional to temperature and ensured that Charles received due credit for his discovery.

Alongside JOHN DALTON, Gay-Lussac concluded that once pressure was kept fixed, near zero degrees Celsius all gases increased in volume by 1/273 the original value for every degree Celsius rise in temperature. At 10degrees, the volume would become 283/273 of its original value and at – 10degrees it would be 263/273 of that same original value. He extended this relation by showing that when volume was kept fixed, gas would increase or decrease the pressure exerted on the outside of the gas container by the same 1/273 factor when temperature was shifted by a degree Celsius. This did not depend upon the gas being studied and hinted at a deep connection shared by all gases. If the volume of a gas at fixed pressure decreased by 1/273 for every 1degree drop, it would reach zero volume at -273degrees Celsius. The same was true for pressure at fixed volume. That had to be the end of the scale, the lowest possible temperature one could reach. Absolute zero.

In an 1807 gas-experiment, Gay-Lussac took a large container with a removable divider down the middle and filled half with gas and made the other half a vacuüm. When the divider was suddenly removed, the gas quickly filled the whole container. According to caloric theory, temperature was a measure of the concentration of caloric fluid and removal of the divider should have led to a drop in temperature because the fluid was spread out over a greater volume without any loss of caloric fluid. (The same amount of fluid in a larger container means lower concentration).
Evidence linking heat to mechanical energy accumulated. Expenditure of the latter seemed to lead to the former.

Gay-Lussac was an experimentalist and his law was based on extensive experiments. The explanation of why gases combine in this way came from AVOGADRO.

Wikipedia-logo © (link to wikipedia)

NEXT button - JOHN DALTONTIMELINE

NEXT buttonGAS LAWS

Advertisements

JACQUES-ALEXANDRE-CESARE CHARLES (1746-1823)

1787 – France

‘The volume of a given mass of gas at constant pressure is directly proportional to its absolute temperature’

In other words, if you double the temperature of a gas, you double its volume. In equation form:  V/T = constant, or  V1/T1 = V2/T2,  where  V1 is the volume of the gas at a temperature  T1 (in kelvin) and  V2 the new volume at a new temperature  T2.

This principle is now known as Charles’ Law (although sometimes named after GAY-LUSSAC because of his popularisation of it fifteen years later – Gay Lussac’s experimental proof was more accurate than Charles’).
It completed the two ‘gas laws’.

A fixed amount of any gas expands equally at the same increments in temperature, as long as it is at constant pressure.

Likewise for a decline in temperature, all gases reduce in volume at a common rate, to the point at about -273degrees C, where they would theoretically converge to zero volume. It is for this reason that the kelvin temperature scale later fixed its zero degree value at this point.

CHARLES’ Law and BOYLE‘s Law may be expressed as a single equation, pV/T = constant. If we also include AVOGADRO‘s law, the relationship becomes pV/nT = constant, where n is the number of molecules or number of moles.

The constant in this equation is called the gas constant and is shown by R
The equation – known as the ideal gas equation – is usually written as pV = nRT

Strictly, it applies to ideal gases only. An ideal gas obeys all the assumptions of the kinetic theory of gases. There are no ideal gases in nature, but under certain conditions all real gases approach ideal behaviour.

Wikipedia-logo © (link to wikipedia)

NEXT buttonTIMELINE

NEXT buttonGAS LAWS

Related sites
Poster describing the combined Gas Laws

Combined Gas Laws