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1914 – Manchester, England

‘Moseley’s law – the principle outlining the link between the X-ray frequency of an element and its atomic number’

ca. 1910s --- Physicist Henry Gwyn Jeffreys MOSELEY --- Library Image by © Bettmann/CORBIS


Working with ERNEST RUTHERFORD’s team in Manchester trying to better understand radiation, particularly of radium, Moseley became interested in X-rays and learning new techniques to measure their frequencies.
A technique had been devised using crystals to diffract the emitted radiation, which had a wavelength specific to the element being experimented upon.

In 1913, Moseley recorded the frequencies of the X-ray spectra of over thirty metallic elements and deduced that the frequencies of the radiation emitted were related to the squares of certain incremental whole numbers. These integers were indicative of the atomic number of the element, and its position in the periodic table. This number was the same as the positive charge of the nucleus of the atom (and by implication also the number of electrons with corresponding negative charge).

By uniting the charge in the nucleus with an atomic number, a vital link had been found between the physical atomic make up of an element and its chemical properties, as indicated by where it sits in the periodic table.
This meant that the properties of an element could now be considered in terms of atomic number rather than atomic weight, as had previously been the case – certain inconsistencies in the MENDELEEV version of the periodic table could be ironed out. In addition, the atomic numbers and weights of several missing elements could be predicted and other properties deduced from their expected position in the table.

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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


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ROBERT BOYLE (1627- 91)

1662 – England

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

If you double the pressure of a gas, you halve its volume. In equation form: pV = constant; or p1V1 = p2V2, where the subscripts 1 & 2 refer to the values of pressure and volume at any two readings during the experiment.

Born at Lismore Castle, Ireland, Boyle was a son of the first Earl of Cork. After four years at Eton College, Boyle took up studies in Geneva in 1638. In 1654 he moved to Oxford where in 1656, with the philosopher John Locke and the architect Christopher Wren, he formed the experimental Philosophy Club and met ROBERT HOOKE, who became his assistant and with whom he began making the discoveries for which he became famous.

Robert Boyle. New Experiments Physico-Mechanical. Oxford: Thomas Robinson, 1662

New Experiments Physico-Mechanical 1662

In 1659, with Hooke, Boyle made an efficient vacuum pump, which he used to experiment on respiration and combustion, and showed that air is necessary for life as well as for burning. They placed a burning candle in a jar and then pumped the air out. The candle died. Glowing coal ceased to give off light, but would start glowing again if air was let in while the coal was still hot. In addition they placed a bell in the jar and again removed the air. Now they could not hear it ringing and so they found that sound cannot travel through a vacuum.

He proved Galileo’s proposal that all matter falls at equal speed in a vacuum.

He established a direct relationship between air pressure and volumes of gas. By using mercury to trap some air in the short end of a ‘J’ shaped test tube, Boyle was able to observe the effect of increased pressure on its volume by adding more mercury. He found that by doubling the mass of mercury (in effect doubling the pressure), the volume of the air in the end halved; if he tripled it, the volume of air reduced to a third. His law concluded that as long as the mass and temperature of the gas is constant, then the pressure and volume are inversely proportional.

The Skeptical Chymist

Boyle appealed for chemistry to free itself from its subservience to either medicine or alchemy and is responsible for the establishment of chemistry as a distinct scientific subject. His insistence on experimental analysis as the arbiter of elemental status promoted an area of thought which influenced the later breakthroughs of ANTOINE LAVOISIER (1743-93) and JOSEPH PRIESTLY (1733-1804) in the development of theories related to the chemical elements.

Boyle extended the existing natural philosophy to include chemistry – until this time chemistry had no recognised theories.

The idea that events are component parts of regular and predictable processes precludes the action of magic.
Boyle sought to refute ARISTOTLE and to confirm his atomistic (or ‘corpuscular’) theories by experimentation.

In 1661 he published his most famous work, ‘The Skeptical Chymist’, in which he rejected Aristotle’s four elements – earth, water, fire and air – and proposed that an element is a material substance consisting at root of ‘primitive and simple, or perfectly unmingled bodies’, that it can be identified only by experiment and can combine with other elements to form an infinite number of compounds.

The book takes the form of a dialogue between four characters. Boyle represents himself in the form of Carneades, a person who does not fit into any of the existing camps, as he disagrees with alchemists and sees chemists as lazy hobbyists. Another character, Themistius, argues for Aristotle’s four elements; while Philoponus takes the place of the alchemist, Eleutherius stands in as an interested bystander.

In the conclusion he attacks chemists.

“I think I may presume that what I have hitherto Discursed will induce you to think, that Chymists have been much more happy finding Experiments than the Causes of them; or in assigning the Principles by which they may be best explain’d”
He pushes the point further: “me thinks the Chymists, in the searches after truth, are not unlike the Navigators of Solomon’s Tarshish Fleet, who brought home Gold and Silver and Ivory, but Apeas and Peacocks too; For so the Writings of several (for I say not, all) of your Hermetick Philosophers present us, together with divers Substantial and noble Experiments, Theories, which either like Peacock’s feathers made a great show, but are neither solid nor useful, or else like Apes, if they have some appearance of being rational, are blemished with some absurdity or other, that when they are Attentively consider’d, makes them appear Ridiculous”

The critical message from the book was that matter consisted of atoms and clusters of atoms. These atoms moved about, and every phenomenon was the result of the collisions of the particles.

He was a founder member of The Royal Society in 1663. Unlike the Accademia del Cimento the Royal Society thrived.

Like FRANCIS BACON he experimented relentlessly, accepting nothing to be true unless he had firm empirical grounds from which to draw his conclusions. He created flame tests in the detection of metals and tests for identifying acidity and alkalinity.

It was his insistence on publishing chemical theories supported by accurate experimental evidence – including details of apparatus and methods used, as well as failed experiments – which had the most impact upon modern chemistry.

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EDWIN McMILLAN (1907- 90) GLENN SEABORG (1912- 99)

1940 – USA

‘Elements heavier than uranium in the periodic table (transuranium elements) are made artificially. Uranium (U, atomic number 92) is the heaviest element known to exist naturally in detectable amounts on the Earth’

Photograph of SEABORG & McMILLAN together ©


In 1933 ENRICO FERMI showed that the nucleus of most elements would absorb a neutron.
In 1940 McMillan, a nuclear physicist, produced and identified the first artificial element, neptunium (Np, 93). In 1943 Seaborg, a chemist, succeeded in creating plutonium (Pu, 94).

So far more than 20 synthetic elements have been created. All are unstable, decaying with half-lives ranging from a year to a few milliseconds.
At least thirteen transuranium elements have been named after scientists:-
curium (Cm, 96: Marie and Pierre Curie [1944]), einsteinium (Es, 99: Albert Einstein [1952]), fermium (Fm, 100: Enrico Fermi [1952]), mendelevium (Md, 101: Dmitri Mendeleev [1955]), nobelium (No, 102: Swedish chemist Alfred Nobel (1833-96), known for his bequest for the foundation of the Nobel Prizes [1956]), lawrencium (Lr, 103: Ernest O. Lawrence, a physicist best known for development of the cyclotron [1961]), rutherfordium (Rf, 104: Ernest Rutherford [1968]), seaborgium (Sg, 106: Glenn Seaborg [1974]), bohrium (Bh, 107: Niels Bohr [1981]), meitnerium (Mt, 109: Lise Meitner [1982]); roentgenium (Rg, 111: Wilhelm Conrad Röntgen was first created in 1994 by the GSI Helmholtz Centre for Heavy Ion Research near Darmstadt in Germany [1994]), copernicium (Cn, 112: named after astronomer Nicolaus Copernicus [1996]), flerovium (Fl, 114 named after Soviet physicist Georgy Flyorov [2012]).

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HANS BETHE (1906-2005)

1938 – USA

‘Energy in stars is produced by hydrogen fusion reactions’

In nuclear fusion the nuclei of light atoms combine at very high temperatures and release enormous amounts of energy that is radiated from the surface of the star as heat and light.

photo of HANS BETHE &copy:


An ordinary star is one of the simplest entities in nature; it is a sphere of gas that is by mass 73 percent hydrogen, 25 percent helium and 2 percent other elements. The temperature in the centre is high enough to fuse four nuclei of hydrogen together to form one helium nucleus.

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OTTO HAHN (Germany 1879-1968) LEIS MEITNER (Austria 1878-1968) FRITZ STRASSMANN (Germany 1902- 80)

1938 – Germany

‘Nuclear Fission. The breaking up of the nucleus of a heavy atom into two or more lighter atoms. Energy is released during the process’

A reinterpretation of the results of the mid 1930s neutron-bombarding experiments of ENRICO FERMI with uranium offered an alternative explanation to Fermi’s own idea that the uranium had transmuted into new heavier elements. The three German scientists offered the explanation that the uranium nucleus had in fact been broken down into a number of smaller nuclei

with the release of potentially huge amounts of energy under the rules of Einstein’s formula E = mc2.


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