HENRI LOUIS LE CHATELIER (1850-1936)

1888 – France

‘When a system in equilibrium is subjected to a change in conditions, it adjusts itself so as to try to oppose that change’

photo portrait of HENRI LOUIS LE CHATELIER ©

The principle is a consequence of the law of conservation of energy.

Le Chatelier’s principal is valuable in understanding how to control the industrial production of chemicals such as ammonia.
Nitrogen and hydrogen react to form ammonia. When the pressure of this system is increased, more ammonia is produced, but when the pressure is lowered, ammonia is decomposed into hydrogen and nitrogen. Thus by controlling pressure and temperature, ammonia can be produced with the minimum of waste.

Le Chatelier was a chemist and is remembered for inventing thermocouples for measuring high temperatures (1877) and oxyacetylene welding (1895).

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JOSEF VON FRAUNHOFER (1787-1826)

1823 – Germany

‘The spectroscope’

A significant improvement on the apparatus used by Newton. Sunlight, instead of passing through a pinhole before striking a prism, is passed through a long thin slit in a metal plate. This creates a long ribbon-like spectrum, which may be scanned from end to end with a microscope.

image of the visible portion of the electromagnetic spectrum showing a series of dark fraunhofer lines

Cutting across the ribbon of rainbow colours are thin black lines. The lines are present even when a diffraction grating is used instead of a prism, proving that the lines are not produced by the material of a prism, but are inherent in sunlight.

An equivalent way of describing colours is as light waves of different sizes.
The wavelength of light is fantastically small, on average about a thousandth of a millimeter, with the wavelength of red light being about twice as long as that of blue light.

Fraunhofer’s black lines correspond to missing wavelengths of light.

By 1823 Fraunhofer had measured the positions of 574 spectral lines, labeling the most prominent ones with the letters of the alphabet. The lines labeled with the letters ‘H’ and ‘K’ correspond to light at a wavelength of 0.3968 thousandths of a millimeter and 0.3933 thousandths of a millimeter, respectively. The lines are present in the spectrum of light from stars, usually in different combinations.

Fraunhofer died early at the age of 39 and it was left to the German GUSTAV KIRCHHOFF to make the breakthrough that explained their significance.

Astronomers today know the wavelengths of more than 25,000 ‘Fraunhofer lines’.

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ROBERT BUNSEN (1811- 99) GUSTAV KIRCHHOFF (1824- 87)

1860 – Germany

‘Each chemical element, when heated to incandescence, produces its own characteristic lines in the spectrum of light’

For example, sodium produces two bright yellow lines.
Bunsen developed the Bunsen burner in 1855.
In the flame test the Bunsen burner’s non-luminous flame does not interfere with the coloured flame given off by the sample.

Kirchhoff was a professor of physics at Heidelberg. Bunsen and Kirchhoff together developed the first spectroscope, a device used to produce and observe a spectrum. They used their spectroscope to discover two new elements, caesium (1860) and rubidium (1861).

In 1860 Kirchhoff made the discovery that when heated to incandescence, each element produces its own characteristic lines in the spectrum.

This means that each element emits light of a certain wavelength – sodium’s spectrum has two yellow lines (wavelengths about 588 and 589 nanometres). The Sun’s spectrum contains a number of dark lines, some of which correspond to these wavelengths.

The Swedish scientist ANDERS ANGSTROM had, four years earlier, found that a gas always absorbs light at the same wavelength that it emits light. If the gas is hotter than the light source, then more light is emitted by the gas than absorbed, creating a bright line in the spectrum of the light source. If the gas is cooler than the light source the opposite happens; more light is absorbed by the gas than is emitted, creating a dark line.
The dark solar D lines told Kirchhoff that sodium is present in the relatively cool outer atmosphere of the Sun. This could be tested in the laboratory by burning a piece of chalk in a hot oxygen-hydrogen torch. The intensely bright limelight that is produced may be passed through a cooler sodium flame and the light emerging examined through a spectroscope. Crossing the spectrum of the artificial light occur black lines at the same wavelength that a sodium flame emits light. This solved the mystery of the FRAUNHOFER LINES.

Scientists now had a means to determine the presence of elements in stars. By comparing the dark lines in the spectra of light from the stars with the bright lines produced by substances in the laboratory, Kirchhoff had been able to identify the elements that made up a celestial body millions of miles away in space.


      

In England the astronomer William Huggins recorded the spectra of hundreds of stars and showed the unmistakable fingerprints of familiar elements that are found on the Earth’s surface. The stars are made of exactly the same kind of atoms as the Earth.

In 1868 Norman Lockyer described a spectral line in the yellow region very close to the wavelength of the two ‘D’ spectral lines of sodium. After repeated attempts to discover a substance that produced the same line on Earth, it appeared that the line did not correspond to any hitherto known element. Lockyer gave the element the name ‘helium’, the gas later to be found associated with radioactive decay in ores containing uranium.

Helium had not previously been found on Earth because it is both inert and lighter than air, ironic because after hydrogen, helium is the second most common element in the universe.

In 1904 RUTHERFORD would declare that the presence of helium in the Sun was evidence that sunlight was a product of radioactive processes. The absence of any FRAUNHOFER lines in sunlight that corresponded to radium dealt a blow to this hypothesis. Was there another way of releasing atomic energy than radioactivity?

  

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WILLARD LIBBY (1908- 80)

1946 – USA

‘Radiocarbon can be used to estimate the age of any organic material. The radioactive isotope of carbon,14C (carbon-14) is present in all living things. When life stops 14C begins to decay. From the rate of decay the age (or time of death) of an organism can be calculated’

The two most common forms of carbon 12C and 13C, make up virtually all types of carbon and are stable – 12C is the simplest form and is made up of 6 protons and 6 neutrons; 13C is slightly heavier because it has one more neutron. 14C, known as radiocarbon has the unstable combination of 6 protons (defining it as carbon) and 8 neutrons.

In the late 1940s Libby led the team at the University of Chicago, USA, that developed radiocarbon dating using the radioactive isotope 14C.

Living things go on absorbing 14C until the time of their death. The half-life of 14C is 5730 years – once an organism dies, 14C begins to decay. As a result the ratio of 12C to 14C changes with time. By measuring this ratio, it can be determined when the organism died.

Libby suggested that minute amounts of radiocarbon come from the upper part of the atmosphere. He put forward the idea that when high-energy particles formed in deep space – cosmic rays – reach the atmosphere, they interact with nitrogen gas to form radiocarbon. He argued that the newly formed radiocarbon is rapidly converted to carbon-dioxide, CO2, and is taken up by plants during photosynthesis; with the result that the radiocarbon enters the food chain. Everything alive should therefore have the same radiocarbon concentration as the atmosphere.

Once an individual dies, some of the 14C atoms begin to disintegrate and give off an electron to reform nitrogen. Libby argued that if the original radiocarbon content is known. it should be possible to measure the remaining 14C in a sample of tissue to back-calculate its age, in a similar way to estimating how much time has passed by measuring the amount of sand left in the top of an egg timer.
By the end of the 1940s, Libby and his team had shown that the radiocarbon content of the air was the same around the world and that 14C could be used to date anything organic.

The crucial principle is the half-life of the unstable atom, the rate at which it will break down. The longer the half-life of a material, the further back in time a dating method can go. With radiocarbon, the dating range is 40,000 to 60,000 years.

When Libby originally measured the half-life of radiocarbon, he calculated it to be just over 5720 years. During the 1950s a new estimate of 5568 years was made by other researchers, who assumed that Libby had got his figures wrong and the 5568-year half-life was adopted by the scientific community.
It is now known that the half-life of radiocarbon is 5730 years, virtually identical to Libby’s original estimate. As a result of the large number of samples that had already been dated, the incorrect value of 5568-years is used in estimates – confusingly this is now termed the ‘Libby half-life’. As all labs use the same half-life value, all ages are directly comparable.

With radiocarbon dating the assumptions made are:

  1. that the atmosphere has had the same 14C content in the past as today
  2. that all things alive have the same radiocarbon content as one-another and as the atmosphere
  3. that no more radiocarbon is added to a sample after death

To obtain a final radiocarbon age, we have to use a point in time to compare against. 1950 is used as year zero and all ages are described relative to this as ‘before present’ (BP). Radiocarbon dating does not give a precise date and estimates are given within a range of uncertainty.

Libby received numerous awards for this work,including the 1960 Nobel Prize for Chemistry. Libby also worked on the Manhattan Project during World War II, helping to enrich the uranium used in the atomic bombs.

picture of the Nobel medal - link to nobelprize.org

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HENRY CAVENDISH (1731-1810)

1766 – England

‘Three Papers Containing Experiments On Factitious Airs’

(gases made from reactions between liquids and solids)

1798 – Density of the earth
Using a torsion balance and the application of NEWTON’s theory of gravity, Cavendish concluded that the earth’s density was 5.5 times that of water.

Born of the English aristocracy and inheritor of a huge sum of money half way through his life, Cavendish is remembered for his work in chemistry.
He demonstrated that hydrogen (inflammable air) and carbon dioxide (fixed air) were gases distinct from ‘atmospheric air’. His claim to the discovery that water was not a distinct element – a view held since the time of ARISTOTLE – but a compound made from two parts hydrogen to one part oxygen, became confused with similar observations made by ANTOINE LAVOISIER.

Full length drawing of Henry Cavendish  &copy:

CAVENDISH

1871 – England

Almost all his discoveries remained unpublished until the late nineteenth century when his notes were found and JAMES CLERK MAXWELL dedicated himself to publishing Cavendish’s work, a task he completed in 1879.
By then many potential breakthroughs, significant at the time, had been surpassed by history.

In 1871 the endowment of the Cavendish Laboratory was made to Cambridge University by Cavendish’s legatees.

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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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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.

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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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