DMITRI MENDELEEV (1834-1907)

1869 – Russia

‘The properties of elements are periodic functions of their atomic weights’

Arrange the atoms in order of their atomic weight (relative atomic mass) and elements are also arranged in order of their properties. This arrangement of the elements is called the periodic table.

In the modern periodic table elements are no longer arranged by their atomic weight but by a more fundamental quantity; ‘atomic number’.

photo portrait of DIMITRY IVANOVICH MENDELEYEV© - link to Britannica online article

DIMITRY IVANOVICH MENDELEYEV

The atomic number of an element is the number of protons in the nucleus of one of its atoms; the number of neutrons, which contributes to atomic weight, is ignored. The modern periodic law is that ‘The properties of elements are periodic functions of their atomic numbers’.

In 1860 Dmitri Ivanovich Mendeleev attended a chemistry conference in Karlsruhe where the Italian Stanislao Cannizzaro’s speech announcing his rediscovery of the distinction between atoms and molecules (originally announced in 1811 by AVOGADRO) made a profound impression.

The German chemist Johann Wolfgang Döbereiner (1780-1849) had recognised mathematical patterns in elements that had similar properties. He found that adding the atomic weights of calcium (40) and barium (137) and dividing the total in two left a value close to the weight of strontium (88). Finding this same pattern repeated for lithium, sodium and potassium, and for chlorine, bromine and iodine confirmed the relationship, which he termed the Law of Triads.

In 1862, French scientist Alexandre Beguyer de Chancourtois developed a way of representing the elements by wrapping a helical list around a cylinder.

A repeating pattern in natural phenomena is a strong indication that there exists a simple, compact description.
The periodic table suggests that the distinct atoms of the elements may be described in terms of significantly fewer building blocks than the number of the individual elements. Atoms, then, were made of significantly fewer subatomic building blocks.

In 1869 the 35-year-old Mendeleev published a table of the 61 elements then known. His list of elements – ‘On the Relation of the Properties to the Atomic Weights of Elements’ – occupied a grid where the atomic weight increased as you went down a column (periods) and the elements in any particular row (groups or families) shared similar properties and valencies (metals and gases, for instance).

Mendeleev had to juggle the order of a few elements, assuming their weights to have been incorrectly measured, and predicted that some undiscovered elements would fill the gaps in the table, based on the properties of the elements surrounding the gaps.
The modern periodic table has been turned sideways.

By 1886, with the discoveries of gallium, scandium and germanium with the properties he had foretold, his prediction was fulfilled. By 1925, chemists had successfully identified all the 92 elements they believed to exist in nature. The first artificial element, neptunium, was synthesised in 1940. Many more elements have been made since then.

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‘The Genesis of a Law of Nature’ – Mendeleev

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JOHN DALTON

1808 – England

‘All matter is made up of atoms, which cannot be created, destroyed or divided. Atoms of one element are identical but different from those of other elements. All chemical change is the result of combination or separation of atoms’

Dalton struggled to accept the theory of GAY-LUSSAC because he believed, as a base case, that gases would seek to combine in a one atom to one atom ratio (hence believing the formula of water to be HO not H2O). Anything else would contradict Dalton’s theory on the indivisibility of the atom, which he was not prepared to accept.

The reason for the confusion was that at the time the idea of the molecule was not understood.
Dalton believed that in nature all elementary gases consisted of indivisible atoms, which is true for example of the inert gases. The other gases, however, exist in their simplest form in combinations of atoms called molecules. In the case of hydrogen and oxygen, for example, their molecules are made up of two atoms, described as H2 and O2 respectively.

Gay-Lussac examined various substances in which two elements form more than one type of compound and concluded that if two elements A and B combine to form more than one compound, the different masses of A that combine with a fixed mass of B are in a simple whole number ratio. This is the law of multiple proportions.

AVOGADRO’s comprehension of molecules helped to reconcile Gay-Lussac’s ratios with Dalton’s theories on the atom.

Gay-Lussac’s ratio for water could be explained by two molecules of hydrogen (four ‘atoms’) combining with one molecule of oxygen (two ‘atoms’) to result in two molecules of water (2H2O).

2H2 + O2 ↔ 2H2O

When Dalton had considered water, he could not understand how one atom of hydrogen could divide itself (thereby undermining his indivisibility of the atom theory) to form two particles of water. The answer proposed by Avogadro was that oxygen existed in molecules of two and therefore the atom did not divide itself at all.

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JOHN DALTON (1766-1844)

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.

Graph demonstrating the varying solubility of gases

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.

A_New_System_of_Chemical_Philosophy - DALTON's original outline

A_New_System_of_Chemical_Philosophy

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WILLIAM PROUT (1785-1850)

1815 – England

‘Atoms are not the smallest thing’

After ANTOINE LAVOISIER had compiled his list of the then known elements, another 32 were added in the years following his death. Fifty kinds of fundamental building blocks for matter seemed excessive. In 1815 Prout, using AVOGADRO’s method of comparing the relative densities and weights of gases, proposed that all atoms appeared to have weights that were exact multiples of the weight of the lightest atom, hydrogen, and that the different atomic weights of elements are whole-number multiples of the atomic weight of hydrogen (Prout’s hypothesis).

portrait of Scottish physician WILLIAM PROUT ©

WILLIAM PROUT

He took this as proof that all atoms were actually made from hydrogen atoms and the idea was adopted as atomic theory and used for later investigations of atomic weights and the classification of the elements.

If all atoms are made from atoms of hydrogen, then it could be possible to transform an atom of one element into an atom of another.
If atoms had been assembled from other things, then they could not themselves be the smallest things in creation.

Apart from the method of weighing atoms being controversial, there are exceptions to the rule. Chlorine is 35.5 times as heavy as hydrogen.

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AMEDEO AVOGADRO (1776-1856)

1811 – Italy

‘Equal volumes of all gases at the same temperature and pressure contain the same number of molecules’

In 1811, when Avogadro proposed his HYPOTHESIS, very little was known about atoms and molecules. Avogadro claimed that the same volume of any gas under identical conditions would always contain the same number of fundamental particles, or molecules. A litre of hydrogen would contain exactly the same number of molecules as a litre of oxygen or a litre of carbon dioxide.

Drawing of AVOGADRO ©

In 1814 ANDRE AMPERE was credited with discovering that if a gas consisted of a single element, its atoms could clump in pairs. The molecules of oxygen consisted of pairs of oxygen atoms, and the molecules of chlorine, pairs of chlorine atoms.
Diatomic gases possess a total of six degrees of simple freedom per molecule that are related to atomic motion.

This provides a way of comparing the weights of different molecules. It was only necessary to weigh equal volumes of different gases and compare them. This would be exactly the same as comparing the weights of the individual molecules of each gas.

Avogadro realised that GAY-LUSSAC‘s law provided a way of proving that an atom and a molecule are not the same. He suggested that the particles (molecules) of which nitrogen gas is composed consist of two atoms, thus the molecule of nitrogen is N2. When one volume (one molecule) of nitrogen combines with three volumes (three molecules) of hydrogen, two volumes (two molecules) of ammonia, NH3, are produced.

N2 + 3H2 ↔ 2NH3

However, the idea of a molecule consisting of two or more atoms bound together was not understood at that time.

Avogadro’s law was forgotten until 1860 when the Italian chemist STANISLAO CANNIZZARO (1826-1910) explained the necessity of distinguishing between atoms and molecules.

Avogadro’s constant
From Avogadro’s law it can be deduced that the same number of molecules of all gases at the same temperature and pressure should have the same volume. This number has been determined experimentally: it’s value is 6.022 1367(36) × 1023AVOGADRO’S NUMBERAvogadro's_number_in_e_notation

That at the same temperature and pressure, equal volumes of all gases have the same number of molecules allows a simple calculation for the combining ratios of all gases – by measuring their percentages by volume in any compound. This in turn facilitates simple calculation of the relative atomic masses of the elements of which it is composed.

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