NIELS BOHR (1885-1962)

1913 – Denmark

‘Electrons in atoms are restricted to certain orbits but they can move from one orbit to another’

Bohr’s was the first quantum model for the internal structure of the atom.

Bohr worked with RUTHERFORD in Manchester and improved upon Rutherford’s model, which said that electrons were free to orbit the nucleus at random.

Classical physics insisted that electrons moving around the nucleus would eventually expire and collapse into the nucleus as they radiated energy. Bohr resolved the issue surrounding Rutherford’s atomic structure by applying the concept of quantum physics set out by MAX PLANCK in 1900.
He suggested that the electrons would have to exist in one of a number of specific orbits, each being defined by specific levels of energy. From the perspective of quantum theory, electrons only existed in these fixed orbits where they did not radiate energy. The electrons could move to higher-level orbits if energy was added, or fall to lower ones if they gave out energy. The innermost orbit contains up to two electrons. The next may contain up to eight electrons. If an inner orbit is not full, an electron from an outer orbit can jump into it. Energy is released as light (a photon) when this happens. The energy that is released is a fixed amount, a quantum.

Quanta of radiation would only ever be emitted as an atom made the transition between states and released energy. Electrons could not exist in between these definite steps. This quantised theory of the electrons’ orbits had the benefits of explaining why atoms always emitted or absorbed specific frequencies of electromagnetic radiation and of providing an understanding of why atoms are stable.

Bohr calculated the amount of radiation emitted during these transitions using Planck’s constant. It fitted physical observations and made sense of the spectral lines of a hydrogen atom, observed when the electromagnetic radiation (caused by the vibrations of electrons) of the element was passed through a prism. The prism breaks it up into spectral lines, which show the intensities and frequencies of the radiation – and therefore the energy emissions and absorptions of the electrons.

Each of the elements has an atomic number, starting with hydrogen, with an atomic number of one. The atomic number corresponds to the number of protons in the element’s atoms. 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 that atoms will form bonds with others, and the ease with which they will do it, is determined by the configuration of electrons.
As elements are ordered in the periodic table by atomic number, it can be seen that their position in the table can be used to predict how they will react.

In addition to showing that electrons are restricted to orbits, Bohr’s model also suggested that

  • the orbit closest to the nucleus is lowest in energy, with successively higher energies for more distant orbits.
  • when an electron jumps to a lower orbit it emits a photon.
  • when an electron absorbs energy, it jumps to a higher orbit.

Bohr called the jump to another orbit a quantum leap.

Although it contained elements of quantum theory, the Bohr model had its flaws. It ignored the wave character of the electron. Work by WERNER KARL HEISENBERG later tackled these weaknesses.

Bohr’s theory of complementarity states that electrons may be both a wave and a particle, but that we can only experience them as one or the other at any given time. He showed that contradictory characteristics of an electron could be proved in separate experiments and none of the results can be accepted singly – we need to hold all the possibilities in mind at once. This requires a slight adjustment to the original model of atomic structure, we can no longer say that an electron occupies a particular orbit, but can only give the probability that it is there.

In 1939 he developed a theory of nuclear fission with Jon Archibald Wheeler (b.1911) and realised that the 235uranium isotope would be more susceptible to fission than the more commonly used 238uranium.
The element bohrium is named after him.

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WOLFGANG PAULI (1900- 58)

1925 – Austria

‘No two electrons in an atom can have the same quantum number’

A quantum number describes certain properties of a particle such as its charge and spin.

An orbital or energy level cannot hold more than two electrons, one spinning clockwise, the other anti-clockwise.

Electrons are grouped in shells, which contain orbitals. The shells are numbered ( n = 1,2,3 etc. ) outwards from the nucleus. These numbers are the ‘principle quantum numbers’.
An increase in n indicates an increase in energy associated with the shell, and an increase in the distance of the shell from the nucleus. The number of electrons allowed in a shell is 2n2. Each shell contains sub-shells or energy sub-levels. A shell can only have n sub-shells. A shell is given a number and a letter ( s,p,d,f,g,etc. ). For example, the electron shell structure of lithium is 1s22s1 (two electrons in ‘s’ sub-shell of the first shell, and one electron in ‘s’ sub-shell of the second shell; the superscript indicates the number of electrons in the shell).

The Pauli principle provided a theoretical basis for the modern periodic table.

1930 – Austria

‘The radioactive beta decay of an atomic nucleus in which a neutron turns into a proton and emits an electron does not seem to follow the law of conservation of energy. To account for the missing energy, Pauli postulated that a particle of zero charge and zero mass is released in such reactions’

A few years later ENRICO FERMI named the new particle a neutrino.
There are three known types of neutrino – muon, tau and electron.

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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. So two hydrogen atoms form the hydrogen molecule, H2, by each sharing their single electron. The two atoms are bound together by the shared electrons. This was proposed by Lewis and Irving Langmuir in 1916.

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.Water10_animation

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. 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.

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.

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.

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.

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’.
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.

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.

          

   

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.

diagram of CO2 bond length

CO2 bond length

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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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.

Photo portrait of EDWARD FRANKLAND ©

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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