Protons relative masses

Nuclei that have too many protons relative to their number of neutrons correct this situation in either of two ways. They either capture one of their Is electrons or they emit a positron (a positively charged particle with the same mass as an electron). Either process effectively changes a proton to a neutron within the nucleus. 

Ans. (a) The nucleus is a distinct part of the atom. Neutrons are subatomic particles which, along with protons, are located in the nucleus. (b) Mass number refers to individual isotopes. It is the sum of the numbers of protons and neutrons. Atomic weight refers to the naturally occurring mixture of isotopes, and is the relative mass of the average atom compared to l2C. (/) Atomic mass is the same as atomic weight [see (b)]. Atomic mass unit is the unit of atomic weight. 

The structures of P-CD-modified ligands were routinely elucidated by H-NMR and ESI mass spectroscopy. Formation of Ru chloro complex 97 from 49 and 47 was demonstrated to be quantitative by H-NMR (Fig. 28). Si ificant downfield shifts (0.36-0.56 ppm) in 97 of several protons relative to free ligand 49 were observed. 

The structure of the atom is very important and gives the element its properties. An atom is arranged as a central nucleus surrounded by outer electrons. The nucleus is very small but very dense, being responsible for nearly all the mass (weight) of the atom. It is made up of two particles protons, which carry a positive electric charge (+1) and are given a relative (arbitrary) mass unit of 1 and neutrons, which have no electric charge but have the same relative mass as the proton (a relative mass unit of 1). The nucleus is only 1/100000 of the diameter of the whole atom. 

Atoms are characterized by their atomic number, which corresponds to the number of protons in the nucleus (and to the number of electrons outside it, since these are balanced for electrical neutrality), and by their mass number, which corresponds to the number of protons plus the number of neutrons in the nucleus and gives the relative mass (weight) of the atom, since the electrons contribute hardly anything to the total mass of an atom. All atoms of a given element have the same atomic number and atomic mass. When the information is useful, the atomic mass can be added to the chemical symbol, written as a small superscript to the left of the symbol. Similarly, the atomic number can be added as a small subscript to the left of the symbol. 

Electron—a negatively charged particle located outside of the atom s nucleus (we will look at this more closely later in this chapter). The electrical charge of an electron is 1.6 X 10 19 C, or a relative charge of 1.0 (charge relative to the proton). The mass of an electron is 5.486 X 10 4 amu, about 1,836 times less than a proton. 

In our study of atomic structure, we look first at the fundamental particles. These are the basic building blocks of all atoms. Atoms, and hence all matter, consist principally of three fundamental particles electrons, protons, and neutrons. Knowledge of the nature and functions of these particles is essential to understanding chemical interactions. The relative masses and charges of the three fundamental particles are shown in Table 5-1. The... 

Give the abbreviations, charges, and relative masses of protons, neutrons, and electrons. 

It is important not to confuse the mass number of an element with its atomic mass. Mass numbers are whole numbers that equal the number of protons plus neutrons in an atom. Atomic masses compare the relative masses of atoms. Atomic masses will always be decimal numbers. 

The relative mass and charge of a particle—a nucleon, another kind of elementary particle, or a nuclide—is described by the notation X, where X is the symbol for the particle, A is the mass number, or the total number of nucleons, and Z is the charge of the particle for nuclides, A is the sum of protons and neutrons and Z is the number of protons (atomic number). In this notation, the three subatomic elementary particles are... 

The nucleus contains protons, which have a positive charge equal in magnitude to the electrons negative charge, and neutrons, which have almost the same mass as protons but no charge. The neutrons function in the nucleus is not obvious. They may help hold the protons (which repel each other) together to form the nucleus, but we will not be concerned with that here. The relative masses and charges of the electron, proton, and neutron are shown in Table 3.4. 

The atomic weights listed in the periodic table are relative numbers ( C = 12.0000. ..) based upon the weighted average of naturally occurring isotopes (e.g., the atomic mass of chlorine is 35.45 reflecting the roughly 3 1 ratio of Cl to Cl). The isotope Cl has 17 protons (atomic number = 17) and 18 neutrons in its nucleus C1 has 17 protons and 20 neutrons. A more precise analysis combines the relative abundances and precise relative masses of the two stable nuclides of chlorine ( Cl 75.78 percent 34.968853. Cl 24.22 percent 36.965903) as follows Relative Atomic Mass of Chlorine 0.7578 (34.968853) + 0.2422 (36.965903) = 35.45 It is noteworthy that on rare occasions, lUPAC may introduce a very slight modification to the atomic mass provided for an element in the periodic table. The relative masses of the nuclides are known to... 

Atoms are made up of numerous smaller particles of which the most important to chemical studies are the proton, neutron, and electron. Positively chaiged protons and neutral neutrons have a relative mass of 1 u each and The atomic weights measured for elements are avera wdghts that depend on the percentages and masses of the isotopes in the naturally occuiring element. If the isotope percent abundances and isotope masses are known for an... 

TABLE 5.2 Electrical Charge and Relative Mass of Electrons, Protons, and Neutrons... 

Since most elements occur as mixtures of isotopes with different masses, the atomic mass determined for an element represents the average relative mass of all the nat-uralfy occurring isotopes of that element. The atomic masses of the individual isotopes are approximately whole numbers, because the relative masses of the protons and neutrons are approximately 1.0 amu each. Yet we find that the atomic masses given for many of the elements deviate considerably from whole numbers. 

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