Each atom consists of a very small nucleus composed of protons and neutrons, which is encircled by moving electrons. Both electrons and protons are electrically charged, the charge magnitude being 1.60 x10-19 C, which is negative in sign for electrons and positive for protons; neutrons are electrically neutral. Masses for these subatomic particles are infinitesimally small; protons and neutrons have approximately the same mass, 1.67 x10-27 kg, which is significantly larger than that of an electron, 9.11 x10-31 kg.
Each chemical element is characterized by the number of protons in the nucleus, or the atomic number (Z). For an electrically neutral or complete atom, the atomic number also equals the number of electrons. This atomic number ranges in integralunits from 1 for hydrogen to 92 for uranium, the highest of the naturally occurring elements.
The atomic mass (A) of a specific atom may be expressed as the sum of the masses of protons and neutrons within the nucleus. Although the number of protons is the same for all atoms of a given element, the number of neutrons (N) may be variable. Thus atoms of some elements have two or more different atomic masses,the most common isotope of carbon, carbon 12 (12C) (A = 12.00000). Within this scheme, the masses of protons and neutrons are slightly greater than unity, and A = Z + N
The atomic weight of an element or the molecular weight of a compound may be specified on the basis of amu per atom (molecule) or mass per mole of material. In one mole of a substance there are 6.023 x 1023 (Avogadro�s number) atoms or molecules. These two atomic weight schemes are related through the following equation:
1 amu/atom (or molecule) = 1 g/mol
For example, the atomic weight of iron is 55.85 amu/atom, or 55.85 g/mol. Sometimes use of amu per atom or molecule is convenient; on other occasions g (or kg)/mol is preferred; the latter is used in this object.
ELECTRONS IN ATOMS
ATOMIC MODELS
Schematic representation of the Bohr atom.
During the latter part of the nineteenth century it was realized that many phenomena involving electrons in solids could not be explained in terms of classical mechanics. What followed was the establishment of a set of principles and laws that govern systems of atomic and subatomic entities that came to be known as quantum mechanics. An understanding of the behavior of electrons in atoms and crystalline solids necessarily involves the discussion of quantum-mechanical concepts. However, a detailed exploration of these principles is beyond the scope of this description, and only a very superficial and simplified treatment is given.
One early outgrowth of quantum mechanics was the simplified Bohr atomic model, in which electrons are assumed to revolve around the atomic nucleus in discrete orbitals, and the position of any particular electron is more or less well defined in terms of its orbital. This model of the atom is represented in Figure above. Another important quantum-mechanical principle stipulates that the energies of electrons are quantized; that is, electrons are permitted to have only specific values of energy. An electron may change energy, but in doing so it must make a quantum jump either to an allowed higher energy (with absorption of energy) or to a lower energy (with emission of energy). Often, it is convenient to think of these allowed electron energies as being associated with energy levels or states.
To be continued.....
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