Chemical Bond
The historical development of the idea of the chemical bond and its electron nature is briefly outlined. The ionic, covalent and metallic bonds are presented as ideal types in a continuum of intermediate cases. Thus, bonds of polar-covalent nature are int
- PDF / 2,621,891 Bytes
- 107 Pages / 439 x 666 pts Page_size
- 115 Downloads / 234 Views
Chemical Bond
2.1
Historical Development of the Concept
The concept of the chemical bond is central to modern chemistry. Its classical form, which gradually and painstakingly developed in the course of the 19th century, described molecules as a combinations of linked atoms. The idea proved extremely useful for interpreting, systematizing and predicting chemical facts, although for a long time it developed without any understanding of the underlying physics. This ‘black box’ situation began to change towards the close of the century. G. J. Stoney in 1881 calculated the elementary charge of electricity and in 1891 named it ‘electron’. In 1894, W. Weber suggested that the atom consists of positive and negative electric charges. In 1897, W. Wiechert, J. J. Thomson, and J. S. Townsend measured the charge of the electron. In 1902–1904, William Thomson (Lord Kelvin) and J. J. Thomson developed the ‘plum cake’ atomic model, with electrons distributed within the homogenous sphere of positive electricity. In 1904, H. Nagaoka suggested that the positive charge is located in the center of the atom, the electrons orbiting around it. Finally, in 1911 E. Rutherford proved this planetary model experimentally. In 1904, R. Abegg proposed that the valence of an atom corresponds to the number of electrons it lost or gained, the sum of which must be equal to 8 and the highest positive valence to the Group (column) number in the Periodic Table. In 1908, J. Stark postulated that chemical properties of an atom are defined by its outer (‘valence’) electrons, and W. Ramsay in his essay Electron as the element already mentioned the electronic nature of the bond between atoms. Finally, in 1913, N. Bohr proposed the model where the majority of the electrons in a molecule are located around the nuclei as in isolated atoms, and only their outer electrons rotate around the axes connecting atoms, forming the chemical bond. In 1916 W. Kossel explained the formation of ions by the transfer of electrons from one atom to another to complete the outer electronic shells of both to the stable 8-electron configurations; he also introduced the important idea that there is a gradual transition from purely polar compounds (e.g., HCl) to typically non-polar ones (e.g., H2 ) [1]. In the same year Lewis described the formation of the covalent bond by two identical atoms sharing their electrons to acquire stable octets [2]. Langmuir developed the theory of Lewis, postulating that electrons in the atom are distributed in layers, with the ‘cells’ for
S. S. Batsanov, A. S. Batsanov, Introduction to Structural Chemistry, DOI 10.1007/978-94-007-4771-5_2, © Springer Science+Business Media Dordrecht 2012
51
52
2 Chemical Bond
2 electrons in the first layer, 8 in the second, 18 in the third and 32 in the fourth [3, 4, 5]. For a long time, the octet rule was regarded as the norm of chemical bonding, and deviations from it as exceptions. However, later these exceptions became more and more numerous, until their explanation required the introduction of new ideas w
Data Loading...
