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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JOHN DALTON

1808 – Manchester, 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 he believed 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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EDWARD FRANKLAND (1825- 99)

1852 – England

‘The capacity of a given element to combine with other elements to form compounds is determined by the number of chemical bonds that element can form with other elements’

This ‘combining power’ is now termed valency or valence.

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EDWARD FRANKLAND

Valency is the number of electrons an atom of an element must lose or gain, either completely or by sharing, in order to form a compound. This leaves the atom with the stable electronic configuration of a noble gas (that is a completely full outer shell).
For example, in H2O, hydrogen has a valency of +1 (H+) and Oxygen -2 (O-2). Two hydrogen atoms lose one electron each; one oxygen atom gains these two electrons.

Every atom has a fixed number of bonds that it can form, and to be stable all of these must be employed. If a hydrogen atom bonds to another hydrogen atom, then the bonds on each atom will be fully used in forming H2, a molecule of hydrogen. The same can occur between two atoms of oxygen.
Alternatively, the two bonds on oxygen could be occupied by the bonds on two hydrogen atoms, forming water, H2O.
Frankland understood that only molecules in which atoms had all of their bonds occupied were stable. Most elements have a fixed valency, although some have more than one. The numerical values of valences represent the charge on the ion.

Lone pair shapes

Lone pair shapes

The concept of valence was further developed by FRIEDRICH AUGUST KEKULE who decided that the valence of carbon must be four which allowed carbon to form into chains of atoms or link into closed, six-atom rings. In the simplest such molecule, three of each carbon’s bonds are used to keep the ring together and the remaining bond on each carbon binds to a hydrogen atom. The resulting molecule of benzene contains six atoms of carbon and six atoms of hydrogen.

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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’

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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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SIR JOHN JOSEPH THOMSON (1856-1940)

1897 – England

’Not only was matter composed of particles not visible even with the modern microscope, as scientists from DEMOCRITUS to DALTON had surmised, but those particles were themselves composed of even smaller components’

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JJ THOMSON

By the end of the nineteenth century scientists had cleared up much of the confusion surrounding atomic theory. The discovery of the sub-atomic particle was made in April 1897. They believed that they now largely understood the properties and sizes of the atoms of elements; without question, hydrogen was the smallest of all.

When JJ Thomson announced the discovery of a particle one thousandth the mass of the hydrogen atom the particles were named ‘electrons’ and have been a fundamental part of the understanding of atomic science ever since.

Thomson was investigating the properties of cathode rays, now known to be a simple stream of electrons, but at the time the cause of widespread debate. The rays were known to be visible, like normal light, but they were quite clearly not normal light. He devised a series of experiments, which would apply measurements to the cathode rays and clarify their nature. The rays were created by passing an electric charge through an airless or gasless discharge tube.

By improving the vacuüm in the tube, it was demonstrated that the rays could be deflected by electric and magnetic fields. Thomson drilled a hole in the anode of the tube to allow the mysterious rays from the cathode to pass through. In the space after the anode, he arranged that a magnetic force field from a magnet would tug the cathode rays in one direction, and an electric force field between two electrically charged metal plates would tug them in the opposite direction. The rays would eventually strike the glass wall of the tube to create a familiar greenish spot of light on the phosphor-coated tube.

Thomson concluded that the rays were made up of particles, not waves. He saw that the properties of the particles were negative in charge and didn’t seem to be specific to any one element; they were the same regardless of the gas used to transport the electric discharge, or the metal used at the cathode. From his findings he concluded that cathode rays were made up of a jet of ‘corpuscles’ and, more importantly, that these corpuscles were present in all elements. Thomson devised a method of measuring the mass of the particles and found them to be a fraction of the weight of the hydrogen atom.

The position of the spot indicated how much the beam of cathode rays had been deflected. The deflection could be made zero by adjusting the magnetic and electric forces so that they perfectly balanced. In such a situation, Thomson could read off the strength of the electric force. He knew in theory how the magnetic force on a charged particle depends on its speed. By equating the two forces, he was able to deduce the speed of the cathode rays. The deflection was also influenced by the electric charge carried by the cathode ray particles, and their mass. The larger the charge, the greater the force the particles felt and the greater their deflection, the smaller the mass, the easier it was for any force to push the particles about and again, the greater their deflection.

Independent evidence from electrolysis (passing electricity through liquids) that electric charge came in discreet chunks, which he assumed to be carried by individual cathode ray particles, enabled Thomson to calculate their mass.
He arrived at a figure that was a thousand billion billion billionth of a kilogram – a 1000th of the mass of a hydrogen atom.

Knowing the deflection of the dot and the velocity of the particles (the slower the particles, the longer they were exposed to the electric force and the greater the deflection of the glowing dot), Thomson expected to be able to deduce their charge and mass. What he actually deduced was a combination of their charge and mass.

Atoms were made of smaller things, but the fundamental building-blocks were not hydrogen atoms, as had been maintained by PROUT.

Thomson’s particles were christened ‘electrons’ and were the first subatomic entities. Thomson visualized a multitude of tiny electrons embedded ‘like raisins in a plum pudding’ in a diffuse ball of positive charge.

‘The atom is a sphere of positively charged protons in which negatively charged electrons are embedded in just sufficient quantity to neutralise the positive charge’

This was the accepted picture of the atom at the start of the twentieth century until RUTHERFORD found a way to probe inside the atom in 1911.

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SOREN PETER SORENSEN (1868-1939)

1909 – Denmark

‘A scale of acidity and alkalinity. It runs from 0 (most acid) to 14 (most alkaline).
A neutral solution has a pH of 7’

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SORENSEN

A solution is acidic when the pH is less than 7 and basic (alkaline) when the pH is greater than 7.
The scale is logarithmic.

The pH measures the concentration of hydrogen ions, H+ , in water.

pH is defined as the negative logarithm of the hydrogen ion concentration.

The pH scale can only be used for solutions of acids and bases in water.

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Acid-Base Scale diagram

Acid-Base Scale

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LINUS PAULING (1901- 94)

1931 USA

‘A framework for understanding the electronic and geometric structure of molecules and crystals’

An important aspect of this framework is the concept of hybridisation: in order to create stronger bonds, atoms change the shape of their orbitals (the space around a nucleus in which an electron is most likely to be found) into petal shapes, which allow more effective overlapping of orbitals.

A chemical bond is a strong force of attraction linking atoms in a molecule or crystal. BOHR had already shown that electrons inhabit fixed orbits around the nucleus of the atom. Atoms strive to have a full outer shell (allowed orbit), which gives a stable structure. They may share, give away or receive extra electrons to achieve stability. The way atoms will form bonds with others, and the ease with which they will do it, is determined by the configuration of electrons.

Earlier in the century, Gilbert Lewis (1875-1946) had offered many of the basic explanations for the structural bonding between elements, including the sharing of a pair of electrons between atoms and the tendency of elements to combine with others to fill their electron shells according to rigidly defined orbits (with two electrons in the closest orbit to the nucleus, eight in the second orbit, eight in the third and so on).

Pauling was the first to enunciate an understanding of a physical interpretation of the bonds between molecules from a chemical perspective, and of the nature of crystals.

In a covalent bond, one or more electrons are shared between two atoms. The two atoms are bound together by the shared electrons. This was proposed by Lewis and Irving Langmuir in 1916. Two hydrogen atoms form the hydrogen molecule, H2, by each sharing their single electron.

In an ionic bond, one atom gives away one or more electrons to another atom. So in common salt, sodium chloride, sodium gives away its spare electron to chlorine. As the electron is not shared, the sodium and chlorine atoms are not bound together in a molecule. However, by losing an electron, sodium acquires a positive charge and chlorine, by gaining an electron, acquires a negative charge. The resulting sodium and chlorine ions are held in a crystalline structure.
Until Pauling’s explanation it was thought that they were held in place only by electrical charges, the negative and positive ions being drawn to each other.

Pauling’s work provided a value for the energy involved in the small, weak hydrogen bond.
When a hydrogen atom forms a bond with an atom which strongly attracts its single electron, little negative charge is left on the opposite side of the hydrogen atom. As there are no other electrons orbiting the hydrogen nucleus, the other side of the atom has a noticeable positive charge – from the proton in the nucleus. This attracts nearby atoms with a negative charge. The attraction – the hydrogen bond – is about a tenth of the strength of a covalent bond.
In water, attraction between the hydrogen atoms in one water molecule and the oxygen atoms in other water molecules makes water molecules ‘sticky’. It gives ice a regular crystalline structure it would not have otherwise. It makes water liquid at room temperature, when other compounds with similarly small molecules are gases at room temperature.Water_animation

He devised the electronegativity scale, which ranks elements in order of their electronegativity – a measure of the attraction an atom has for the electrons involved in bonding ( 0.7 for caesium and francium to 4.0 for fluorine ). The electronegativity scale lets us say how covalent or ionic a bond is.

One aspect of the revolution he brought to chemistry was to insist on considering structures in terms of their three-dimensional space. Pauling showed that the shape of a protein is a long chain twisted into a helix or spiral, now known as an alpha-helix. The structure is held in shape by hydrogen bonds.
He also explained the beta-sheet, a pleated sheet arrangement given strength by a line of hydrogen bonds.

1922 – while investigating why atoms in metals arrange themselves into regular patterns, Pauling used X-ray diffraction at CalTech to determine the structure of molybdenum.

When X-rays are directed at a crystal, some are knocked off course by striking atoms, while others pass straight through as if there are no atoms in their path. The result is a diffraction pattern – a pattern of dark and light lines that reveal the positions of the atoms in the crystal.
Pauling used X-ray and electron diffraction, magnetic effects and measurements of the heat of chemical reactions to calculate the distances and angles between atoms forming bonds. In 1928 he published his findings as a set of rules for working out probable crystalline structures from the X-ray diffraction patterns.

Pauling’s application of quantum theory to structural chemistry helped to establish the subject. He took from quantum mechanics the idea of an electron having both wave-like and particle-like properties and applied it to hydrogen bonds. Instead of there being just an electrical attraction between water molecules, Pauling suggested that wave properties of the particles involved in hydrogen bonding and those involved in covalent bonding overlap. This gives the hydrogen bonds some properties of covalent bonds.

1939 – ‘The Nature of the Chemical Bond and the Structure of Molecules’
Pauling suggests that in order to create stronger bonds, atoms change the shapes of their waves into petal shapes; this was the ‘hydridisation of orbitals’.

Pauling developed six key rules to explain and predict chemical structure. Three of them are mathematical rules relating to the way electrons behave within bonds, and three relate to the orientation of the orbitals in which the electrons move and the relative position of the atomic nuclei.

      

      

Describing hybridisation, he showed that the labels ‘ionic’ and ‘covalent’ are little more than a convenience to group bonds that really lie on a continuous spectrum from wholly ionic to wholly co-valent.

1951 – published his findings one year after WILLIAM LAWRENCE BRAGG’s team at the Cavendish Laboratory.

CARBON BONDING
As carbon has four filled and four unfilled electron shells it can form bonds in many different ways, making possible the myriad organic compounds found in plants and animals. The concept of hybridisation proved useful in explaining the way carbon bonds often fall between recognised states, which opened the door to the realm of organic chemistry.

X-ray diffraction alone is not very useful for determining the structure of complex organic molecules, but it can show the general shape of the molecule. Pauling’s work showed that physical chemistry at the molecular level could be used to solve problems in biology and medicine.

A problem that needed resolving was the distance between particular atoms when they joined together. Carbon has four bonds, for instance, while oxygen can form two. It would seem that in a molecule of carbon dioxide, which is made of one carbon and two oxygen atoms, two of carbon’s bonds will be devoted to each oxygen.

Well-established calculations gave the distance between the carbon and oxygen atoms as 1.22 × 10-10m. Analysis gave the size of the bond as 1.16 Angstroms. The bond is stronger, and hence shorter. Pauling’s quantum .3-2. explanation was that the bonds within carbon dioxide are constantly resonating between two alternatives. In one position, carbon makes three bonds with one of the oxygen molecules and has only one bond with the other, and then the situation is reversed.

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