Isotopes atomic mass/number

Ernest O. Lawrence, inventor of the cyclotron) This member of the 5f transition elements (actinide series) was discovered in March 1961 by A. Ghiorso, T. Sikkeland, A.E. Larsh, and R.M. Latimer. A 3-Mg californium target, consisting of a mixture of isotopes of mass number 249, 250, 251, and 252, was bombarded with either lOB or IIB. The electrically charged transmutation nuclei recoiled with an atmosphere of helium and were collected on a thin copper conveyor tape which was then moved to place collected atoms in front of a series of solid-state detectors. The isotope of element 103 produced in this way decayed by emitting an 8.6 MeV alpha particle with a half-life of 8 s. 

Atoms with the same number of protons but a different number of neutrons are called isotopes. To identify an isotope we use the symbol E, where E is the element s atomic symbol, Z is the element s atomic number (which is the number of protons), and A is the element s atomic mass number (which is the sum of the number of protons and neutrons). Although isotopes of a given element have the same chemical properties, their nuclear properties are different. The most important difference between isotopes is their stability. The nuclear configuration of a stable isotope remains constant with time. Unstable isotopes, however, spontaneously disintegrate, emitting radioactive particles as they transform into a more stable form. 

The most important types of radioactive particles are alpha particles, beta particles, gamma rays, and X-rays. An alpha particle, which is symbolized as a, is equivalent to a helium nucleus, fHe. Thus, emission of an alpha particle results in a new isotope whose atomic number and atomic mass number are, respectively, 2 and 4 less than that for the unstable parent isotope. 

Lead, atomic number 82, is a member of Group 14 (IVA) of the Periodic Table. Ordinary lead is bluish grey and is a mixture of isotopes of mass number 204 (15%), 206 (23.6%), 207 (22.6%), and 208 (52.3%). The average atomic weight of lead from different origins may vary as much as 0.04 units. The stable isotopes are products of decay of three naturally radioactive elements (see Radioactivity, natural) comes from the uranium series (see Uraniumand... 

Magnesium [7439-95-4] atomic number 12, is in Group 2 (IIA) of the Periodic Table between beryllium and calcium. It has an electronic configuration of 1T2T2 3T and a valence of two. The element occurs as three isotopes with mass numbers 24, 25, and 26 existing in the relative frequencies of 77, 11.5, and 11.1%, respectively. 

Proton capture processes by heavy nuclei have already been briefly mentioned in several of the preceding sections. The (p,y) reaction can also be invoked to explain the presence of a number of proton-rich isotopes of lower abundance than those of nearby normal and neutron-rich isotopes (Fig. 1.5). Such isotopes would also result from expulsion of a neutron by a y-ray, i.e. (y,n). Such processes may again be associated with supernovae activity on a very short time scale. With the exceptions of " ln and " Sn, all of the 36 isotopes thought to be produced in this way have even atomic mass numbers the lightest is Se... 

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. 

Element has no stable nuclides the value given in parentheses is the atomic mass number of the isotope of longest known half-life. However, three such elements (Th, Pa and U) do have a characteristic terrestrial isotopic composition, and for these an atomic weight is tabulated. 

Numbers in parentheses are atomic mass numbers of most stable isotopes. 

The substances we call elements are composed of atoms. Atoms in turn are made up of neutrons, protons and electrons neutrons and protons in the nucleus and electrons in a cloud of orbits around the nucleus. Nuclide is the general term referring to any nucleus along with its orbital electrons. The nuclide is characterized by the composition of its nucleus and hence by the number of protons and neutrons in the nucleus. All atoms of an element have the same number of protons (this is given by the atomic number) but may have different numbers of neutrons (this is reflected by the atomic mass numbers or atomic weight of the element). Atoms with different atomic mass but the same atomic numbers are referred to as isotopes of an element. 

Kind of Atomic Mass Number Isotope Number of Number of Number of... 

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). 

Examples of isotopes are abundant. The major form of hydrogen is represented as H (or H-1), with one proton H, known as the isotope deuterium or heavy hydrogen, consists of one proton and one neutron (thus an amu of 2) and is the isotope of hydrogen called tritium with an amu of 3. Carbon-12 ( C or C-12) is the most abundant form of carbon, though carbon has several isotopes. One is the C isotope, a radioactive isotope of carbon that is used as a tracer and to determine dates of organic artifacts. Uranium-238 is the radioactive isotope (Note The atomic number is placed as a subscript prefix to the element s symbol—for example, —and the atomic mass number can be written either as a dash and number fol-... 

Titanium has five naturally occurring isotopes with mass numbers ranging from 46 to 50 their percentage isotopic abundances are given in Table 1, with the atomic masses of the individual nucleides. 

When radium (A = 88) emits an alpha particle, its atomic number reduces by 2 and becomes the new element radon (A = 86). The resulting atomic mass is reduced by 4. If th e radium was of the most common isotope, 226, then the radon isotope would have atomic mass number 222. 

The atomic number may be defined as the number of protons in an atomic nucleus, or the positive charge of the nucleus, expressed in terms of the electronic charge. Atomic number usually is denoted by the symbol Z. In the symbolic designation of individual nuclides, the atomic number sometimes is written as a subscript lo Ihe left of the chemical symbol of the atomic species, such as x160 for the oxygen isotope of mass number 16. This usage is redundant, in that the chemical symbol per se specifies the atomic number of the nuclide. 

PROBLEM 2.5 Chlorine, one of the elements in common table salt (sodium chloride), has two main isotopes, with mass numbers 35 and 37. Look up the atomic number of chlorine, tell how many neutrons each isotope contains, and give the standard symbol for each. 

Elements differ from one another according to how many protons their atoms contain, a value called the atomic number (Z) of the element. The sum of an atom s protons and neutrons is its mass number (A). Although all atoms of a specific element have the same atomic number, different atoms of an element can have different mass numbers, depending on how many neutrons they have. Atoms with identical atomic numbers but different mass numbers are called isotopes. Atomic masses are measured using the atomic mass unit (amu), defined as 1/12 the mass of a 12C atom. Because both protons and neutrons have a mass of approximately 1 amu, the mass of an atom in atomic mass units (the isotopic mass) is numerically close to the atom s mass number. The element s atomic mass is a weighted mass average for naturally occurring isotope mixtures. 

When the radioactive g2 Pb isotope emits two alpha particles, it transmutes into another atom. Find the atomic number and the atomic mass number of that atom. 

A radioactive isotope 3 K transmutes into Ar by decaying through one P+ emission. Find the atomic mass number of Ar. 

Neutron emission occurs by the removal of a neutron from a nucleus. However, it is very difficult to trace a neutron emission, since it occurs very rarely and at a high velocity. A new element is not formed as a result of this decay, but an isotope of the original element is formed. That is, after a neutron emission the atomic number of the element remains unchanged whereas the atomic mass number decreases by one. The following nuclear equation is an example for a neutron emission ... 

A value in parentheses for an element without any stable nuclides is the atomic mass number of the isotope of that element of longest known half-life. 

Element abundance data were useful not only in astrophysics and cosmology but also in the attempts to understand the structure of the atomic nucleus. [74] As mentioned, this line of reasoning was adopted by Harkins as early as 1917, of course based on a highly inadequate picture of the nucleus. It was only after 1932, with the discovery of the neutron as a nuclear component, that it was realized that not only is the atomic mass number related to isotopic abundance, but so are the proton and neutron numbers individually. Cosmochemical data played an important part in the development of the shell model, first proposed by Walter Elsasser and Kurt Guggenheimer in 1933-34 but only turned into a precise quantitative theory in the late 1940s. [75] Guggenheimer, a physical chemist, used isotopic abundance data as evidence of closed nuclear shells with nucleon numbers 50 and 82. 

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