Atomic structure nucleus

Atomic Structure The Nucleus Atomic Structure Orbitals 4 Atomic Structure Electron Configurations 6 Development of Chemical Bonding Theory 7 The Nature of Chemical Bonds Valence Bond Theory sp Hybrid Orbitals and the Structure of Methane 12 sp Hybrid Orbitals and the Structure of Ethane 13 sp2 Hybrid Orbitals and the Structure of Ethylene 14 sp Hybrid Orbitals and the Structure of Acetylene 17 Hybridization of Nitrogen, Oxygen, Phosphorus, and Sulfur 18 The Nature of Chemical Bonds Molecular Orbital Theory 20 Drawing Chemical Structures 21 Summary 24... 

A representation of atomic structure. The various spheres are not drawn to scale. The lump of iron on the left would contain almost a million million million million (10 ) atoms, one of which is represented by the sphere in the top center of the page. In turn, each atom is composed of a number of electrons, protons, and neutrons. For example, an atom of the element iron contains 26 electrons, 26 protons, and 30 neutrons. The physical size of the atom is determined mainly by the number of electrons, but almost all of its mass is determined by the number of protons and neutrons in its dense core or nucleus (lower part of figure). The electrons are spread out around the nucleus, and their number determines atomic size but the protons and neutrons compose a very dense, small core, and their number determines atomic mass. 

Str-ucture determines properties and the properties of atoms depend on atomic structure. All of an element s protons are in its nucleus, but the element s electrons are distributed among orbitals of var ying energy and distance from the nucleus. More than anything else, we look at its electron configuration when we wish to understand how an element behaves. The next section illustrates this with a brief review of ionic bonding. 

This chapter builds an understanding of atomic structure in four steps. First, we review the experiments that led to our current nuclear model of the atom and see how spectroscopy reveals information about the arrangement of electrons around the nucleus. Then we describe the experiments that led to the replacement of classical mechanics by quantum mechanics, introduce some of its central features, and illustrate them by considering a very simple system. Next, we apply those ideas to the simplest atom of all, the hydrogen atom. Finally, we extend these concepts to the atoms of all the elements of the periodic table and see the origin of the periodicity of the elements. 

To explain his observations, Rutherford proposed a new h /pothesis for atomic structure. He suggested that every atom has a tiny central core, called the nucleus, within which all the positive charge and most of the mass is concentrated. Electrons surround this central core, as shown schematically in Figure 2-16. Electrons occupy a volume that is huge compared with the size of the nucleus, but each electron has such a small mass that alpha particles are not deflected by the electrons. Consequently, an alpha particle is deflected only when it passes very near a nucleus, and it bounces back only when it collides head-on with a nucleus. Because most of the volume of an atom contains only electrons, most projectiles pass through the foil without being affected. 

Ernest Rutherfords proposed atomic structure added to the problems posed to nineteenth century physics by the ultraviolet catastrophe and the photoelectric effect. Rutherfords atom had a negatively charged electron circling a positively charged nucleus. The physics of the day predicted that the atom would emit radiation, causing the electron to lose energy and spiral down into the nucleus. Theory predicted that Rutherfords atom could not exist. Clearly, science needed new ideas to explain these three anomalies. 

Fig. 7.78 Linear relation of the quadmpole splitting A q = ( jl)eqQ (1 + j /3)l/2 and the isomer shift b for aurous (a) and auric (b) compounds. Also included is a correlation with the relative change in electron density at the gold nucleus, Ali/r(o)P, as derived from Dirac-Fock atomic structure calculations for several electron configurations of gold. An approximate scale of the EFG (in the principal axes system) is given on the right-hand ordinate (from [341])...

The carbon atom has, outside its nucleus, six electrons which, on the Bohr theory of atomic structure, were believed to be arranged in orbits at increasing distance from the nucleus. These orbits corres-... 

In the early part of the twentieth century, then, a simple model of atomic structure became accepted, now known as the Rutherford nuclear model of the atom, or, subsequently, the Bohr-Rutherford model. This supposed that most of the mass of the atom is concentrated in the nucleus, which consists of protons (positively charged particles) and neutrons (electrically neutral particles, of approximately the same mass). The number of protons in the nucleus is called the atomic number, which essentially defines the nature of... 

The hydrogen atom has the simplest atomic structure of all elements and consists of a nucleus and one electron. A neutral H atom can join a second electron, which forms the negative ion, H. Atomic hydrogen is formed as a result of different chemical reactions, but its lifetime is extremely short, as the atoms join each other to form a... 

Only a few relevant points about the atomic structures are summarized in the following. Table 4.1 collects basic data about the fundamental physical constants of the atomic constituents. Neutrons (Jn) and protons (ip), tightly bound in the nucleus, have nearly equal masses. The number of protons, that is the atomic number (Z), defines the electric charge of the nucleus. The number of neutrons (N), together with that of protons (A = N + Z) represents the atomic mass number of the species (of the nuclide). An element consists of all the atoms having the same value of Z, that is, the same position in the Periodic Table (Moseley 1913). The different isotopes of an element have the same value of Z but differ in the number of neutrons in their nuclei and therefore in their atomic masses. In a neutral atom the electronic envelope contains Z electrons. The charge of an electron (e ) is equal in size but of opposite sign to that of a proton (the mass ratio, mfmp) is about 1/1836.1527). 

The exploration of atomic structure began in 1911, when Ernest Rutherford, a New Zealander who worked in Canada and England, discovered that atoms had a dense central nucleus that contained positively charged particles, which he named protons. (See Table 3-1.) it was soon established that each chemical element was characterized by a specific number of protons in each atom. A hydrogen atom has 1 proton, helium has 2, lithium has 3, and so forth through the periodic table. The atomic number is the number of protons for each element. 

Mass spectrometry is more than 100 years old and has yielded basic results and profound insights for the development of atomic physics. The rapid development of nuclear physics, in particular, would be unthinkable without the application of mass spectrometric methods. Mass spectrometry has contributed to conclusive evidence for the hypothesis of the atomic structure of matter. So far mass spectrometry has supplied specific results on the structure of the nucleus of atoms. Nobel prizes have been awarded to a number of scientists (Thomson, Wien, Aston, Paul, Fenn and Tanaka) associated with the birth and development of mass spectrometry, or in which mass spectrometry has aided an important discovery (e.g., for the discovery of fullerenes by Curl, Kroto and Smalley). 

During the past several years it has been shown that both molybdenum and tungsten in oxidation states of +4 or thereabouts, have marked predilection to form trinuclear duster species. For the tungsten cluster spedes, there are three different structural types with respect to the ligand arrangements (Figure 24 a, b, c) which are based on an equilateral triangle of M—M bonded atoms. Structure (a) contains th [W3(j 3-X)(ju-Y)3] nucleus, where X = 0, Cl,... 

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