1801 – England
‘The total pressure of a mixture of gases is the sum of the partial pressures exerted by each of the gases in the mixture’
Partial pressures of gases:
Dalton stated that the pressure of a mixture of gases is equal to the sum of the pressures of the gases in the mixture. On heating gases they expand and he realised that each gas acts independently of the other.
Each gas in a mixture of gases exerts a pressure, which is equal to the pressure it would exert if it were present alone in the container; this pressure is called partial pressure.
Dalton’s law of partial pressures contributed to the development of the kinetic theory of gases.
His meteorological observations confirmed the cause of rain to be a fall in temperature, not pressure and he discovered the ‘dew point’ and that the behaviour of water vapour is consistent with that of other gases.
He showed that a gas could dissolve in water or diffuse through solid objects.
The varying solubility of gases
Further to this, his experiments on determining the solubility of gases in water, which, unexpectedly for Dalton, showed that each gas differed in its solubility, led him to speculate that perhaps the gases were composed of different ‘atoms’, or indivisible particles, which each had different masses.
On further examination of his thesis, he realised that not only would it explain the different solubility of gases in water, but would also account for the ‘conservation of mass’ observed during chemical reactions – as well as the combinations into which elements apparently entered when forming compounds – because the atoms were simply ‘rearranging’ themselves and not being created or destroyed.
In his experiments, he observed that pure oxygen will not absorb as much water vapour as pure nitrogen – his conclusion was that oxygen atoms were bigger and heavier than nitrogen atoms.
‘ Why does not water admit its bulk of every kind of gas alike? …. I am nearly persuaded that the circumstance depends on the weight and number of the ultimate particles of the several gases ’
In a paper read to the Manchester Society on 21 October 1803, Dalton went further,
‘ An inquiry into the relative weight of the ultimate particles of bodies is a subject as far as I know, entirely new; I have lately been prosecuting this enquiry with remarkable success ’
Dalton described how he had arrived at different weights for the basic units of each elemental gas – in other words the weight of their atoms, or atomic weight.
Dalton had noticed that when elements combine to make a compound, they always did so in fixed proportions and went on to argue that the atoms of each element combined to make compounds in very simple ratios, and so the weight of each atom could be worked out by the weight of each element involved in a compound – the idea of the Law of Multiple Proportions.
When oxygen and hydrogen combined to make water, 8 grammes of oxygen was used for every 1 gramme of hydrogen. If oxygen consisted of large numbers of identical oxygen atoms and hydrogen large numbers of hydrogen atoms, all identical, and the formation of water from oxygen and hydrogen involved the two kinds of atoms colliding and sticking to make large numbers of particles of water (molecules) – then as water has an identity as distinctive as either hydrogen or oxygen, it followed that water molecules are all identical, made of a fixed number of oxygen atoms and a fixed number of hydrogen atoms.
Dalton realised that hydrogen was the lightest gas, and so he assigned it an atomic weight of 1. Because of the weight of oxygen that combined with hydrogen in water, he first assigned oxygen an atomic weight of 8.
There was a basic flaw in Dalton’s method, because he did not realise that atoms of the same element can combine. He assumed that a compound of atoms, a molecule, had only one atom of each element. It was not until Italian scientist AMADEO AVOGADRO’s idea of using molecular proportions was introduced that he would be able to calculate atomic weights correctly.
In his book of 1808, ‘A New System of Chemical Philosophy’ he summarised his beliefs based on key principles: atoms of the same element are identical; distinct elements have distinct atoms; atoms are neither created nor destroyed; everything is made up of atoms; a chemical change is simply the reshuffling of atoms; and compounds are made up of atoms from the relevant elements. He published a table of known atoms and their weights, (although some of these were slightly wrong), based on hydrogen having a mass of one.
Nevertheless, the basic idea of Dalton’s atomic theory – that each element has its own unique sized atoms – has proved to be resoundingly correct.
If oxygen atoms all had a certain weight which is unique to oxygen and hydrogen atoms all had a certain weight that was unique to hydrogen, then a fixed number of oxygen atoms and a fixed number of hydrogen atoms combined to form a fixed weight of water molecules. Each water molecule must therefore contain the same weight of oxygen atoms relative to hydrogen atoms.
Here then is the reason for the ‘law of fixed proportions’. It is irrelevant how much water is involved – the same factors always hold – the oxygen atoms in a single water molecule weigh 8 times as much as the hydrogen atoms.
Dalton wrongly assumed that elements would combine in one-to-one ratios as a base principle, only converting into ‘multiple proportions’ (for example from carbon monoxide, CO, to carbon dioxide, CO2) under certain conditions. Each water molecule (H2O) actually contains two atoms of hydrogen and one atom of oxygen. An oxygen atom is actually 16 times as heavy as a hydrogen atom. This does not affect Dalton’s reasoning.
The law of fixed proportions holds because a compound consists of a large number of identical molecules, each made of a fixed number of atoms of each component element.
Although the debate over the validity of Dalton’s thesis continued for decades, the foundation for the study of modern atomic theory had been laid and with ongoing refinement was gradually accepted.
<< top of page