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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JANS JACOB BERZELIUS (1779-1848)

1840 – Sweden

‘An element can exist in two or more forms with different properties’

The various forms are known as allotropes. Graphite, diamond and buckyballs are three crystalline allotropes of carbon.

Berzelius contributed more than just allotropes to chemistry. When DALTON revived the idea of the atom as the unit of matter, he used circular symbols to represent atoms. Berzelius discarded Dalton’s cumbersome system and in its place introduced a rational system of chemical shorthand.

He declared ‘I shall take as the chemical sign the initial letter of the Latin name of each element. If the first two letters be common to two elements I shall use both the initial letter and the first letter they have not in common’.

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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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FRIEDRICH KEKULE (1829- 96)

1865 – Belgium

portrait of AUGUST STRADONITZ KEKULE &copy:

‘Carbon is tetravalent and is capable of forming ring-type organic molecules’

In the 1860s scientists knew about the molecular formula of benzene, C6H6 , but they did not know how the six atoms are arranged in space. Kekule was the first chemist to suggest that carbon is tetravalent, that is, one carbon atom can combine with four atoms – he saw the possibility that the benzene molecule could be ring-shaped.

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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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ALEKSANDER OPARIN (1894-1980)

1936 – Russia

‘In the Earth’s early atmosphere simple inorganic compounds combined to form complex organic compounds, which formed the first living cell’

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ALEKSANDER OPARIN

Viewed correctly, life is compatible with the basic principles of physics and chemistry.

Life is possible, but entropy increases. The late nineteenth century way of looking at biological systems fits into this scheme – looking for patterns in large numbers – in assigning precise properties to classes and groups, not to individuals. At the beginning of the nineteenth century, an organism expended vital force in order to perform its work of synthesis and morphogenesis; at the end of the century, the belief was that it consumed energy.

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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.

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