Mass, electronic nuclidic, 17 table

Our present views on the electronic structure of atoms are based on a variety of experimental results and theoretical models which are fully discussed in many elementary texts. In summary, an atom comprises a central, massive, positively charged nucleus surrounded by a more tenuous envelope of negative electrons. The nucleus is composed of neutrons ( n) and protons ([p, i.e. H ) of approximately equal mass tightly bound by the force field of mesons. The number of protons (2) is called the atomic number and this, together with the number of neutrons (A ), gives the atomic mass number of the nuclide (A = N + Z). An element consists of atoms all of which have the same number of protons (2) and this number determines the position of the element in the periodic table (H. G. J. Moseley, 191.3). Isotopes of an element all have the same value of 2 but differ in the number of neutrons in their nuclei. The charge on the electron (e ) is equal in size but opposite in sign to that of the proton and the ratio of their masses is 1/1836.1527. 

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). 

Elements are described as atoms possessing the same number of protons. However, not all atoms of an element possess the same number of neutrons. These different varieties of an element that contain the same number of protons but different numbers of neutrons are known as isotopes. Isotopes will possess the same chemical properties, since they are determined by the number of electrons and protons, but they will have a different mass. In Chapter 5, we will look at how the number of neutrons affects the stability of the nucleus. All atoms of a specific isotope are known as nuclides. Isotopes are represented using a variety of symbols. Three of these are shown in Table 4.2. 

Nuclides with an excess of neutrons experience P decay. In the nucleus a neutron is converted into a proton, an electron and an electron antineutrino, as indicated in Table 5.1. The atomic number increases by one unit, whereas the mass number does not change (second displacement law of Soddy and Fajans). The energy of the decay process can again be calculated by comparison of the masses according to Einstein ... 

Nuclides with an excess of protons exhibit P decay. A proton in the nucleus is converted into a neutron, a positron and an electron neutrino, as indicated in Table 5.1. The atomic number decreases by one unit, and the mass number remains unchanged. As in the case of P decay, the energy of the decay process is obtained by eq. (5.10). But because now Z2 — Zi — 1, it follows that ... 

Radioactivity is characterized by the emission of energy (electromagnetic or in the form of a particle) from the nucleus of an atom, usually with associated elemental conversion. There are four basic types of radioactive decay (Table 5.4), of which alpha (a) and beta (p ) decay are most common in nature. Alpha emission is the only type of decay that causes a net mass change in the parent nuclide by loss of two protons plus two neutrons. Because two essentially weightless orbiting electrons are also lost when the equivalent of a helium nucleus is emitted, the parent nuclide transmutes into a daughter element two positions to the left on the periodic table. Thus decays by ot... 

Each nucleus is characterized by a definite atomic number Z and mass number A for clarity, we use the symbol M to denote the atomic mass in kinematic equations. The atomic number Z is the number of protons, and hence the number of electrons, in the neutral atom it reflects the atomic properties of the atom. The mass number gives the number of nucleons (protons and neutrons) isotopes are nuclei (often called nuclides) with the same Z and different A. The current practice is to represent each nucleus by the chemical name with the mass number as a superscript, e.g., 12C. The chemical atomic weight (or atomic mass) of elements as listed in the periodic table gives the average mass, i.e., the average of the stable isotopes weighted by their abundance. Carbon, for example, has an atomic weight of 12.011, which reflects the 1.1% abundance of 13C. 

A neutral atom consists of a small, dense central nucleus, about 10 cm in diameter, surrounded by a diffuse cloud of electrons whose outside diameter is around 10" cm. The nucleus contains most of the mass of the atom and carries a positive electric charge that equals a whole number times the electronic charge, 1.602101 X 10" C. This whole number is called the atomic number Z of the atom. It is identical with the serial number of the element in the periodic table. Each nucleus is made up of Z protons and a definite number N of neutrons. The total number of particles in the nucleus, N- Z, is called the mass number and is denoted by A. The mass number turns out to be the whole number nearest to the atomic weight of the nuclide. 

The various decay processes are listed in Table 1. Radioactive nuclides emit either nucleons (alpha particles, very rarely protons or neutrons) or electrons (negatrons, positrons). As an alternative to the emission of a positron, a proton may capture an electron of the K-shell (K-capture). By the emission of an alpha particle the mass number and the atomic number are reduced by the emission of electrons either the number of neutrons (jS -decay, negatron emission) or the number of protons ()S -decay, positron emission) is reduced. By K-capture also the number of protons is reduced. Due to the missing electron in the K-shell, characteristic X-rays of the newly produced atomic species are emitted. 

The first atoms of seaborgium (Sg) were identified in 1974. The longest-lived isotope of Sg has a mass number of 266. (a) How many protons, electrons, and neutrons are in a g nuclide (b) Atoms of Sg are very unstable, and it is therefore difficult to study this element s properties. Based on the position of Sg in the periodic table, what element should it most closely resemble in its chemical properties ... 

The notations for various decay modes used in this book are a for alpha decay, for / decay, P for positron decay, EC for electron capture, IT for isomeric transition, and SF for spontaneous fission. The letter m after a mass number denotes an isomer. Isomers with a half-life of less than 1 s and fission isomers are omitted from the tables. Energies are given only for the most abundant a groups and y rays for P particles the maximum energies p , are tabulated. In the last column, only the convenient methods for the production of nuclides are given nature denotes that the nuclide occurs in nature and multiple neutron capture means that this nuclide is produced by long irradiation in a high-flux reactor. 

Alternatively you can see from the periodic table that the atomic mass of magnesium is 24.31. A nuclide with a mass number of 22 is too light the N/Z ratio is too low. Therefore, Mg-22 undergoes positron emission, resulting in the conversion of a proton to a neutron. (Electron capture would accomplish the same thing as positron emission, but in Mg-22, positron emission is the only decay mode observed.)... 

Use the electron mass from Table 2.1 and the measured mass of the nuclide 18.998403 u, to determine the binding energy per nucleon (in megaelectronvolts) of this atom. 

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