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The mass number (A), also called atomic mass number or nucleon number, is the total number of protons and neutrons (together known as nucleons) in an atomic nucleus. Because protons and neutrons both are baryons, the mass number A is identical with the baryon number B as of the nucleus as of the whole atom or ion. The mass number is different for each different isotope of a chemical element. This is not the same as the atomic number (Z) which denotes the number of protons in a nucleus, and thus uniquely identifies an element. Hence, the difference between the mass number and the atomic number gives the number of neutrons (N) in a given nucleus: N=Aâˆ’Z.
The mass number is written either after the element name or as a superscript to the left of an element's symbol. For example, the most common isotope of carbon is carbon-12, or , which has 6 protons and 6 neutrons. The full isotope symbol would also have the atomic number (Z) as a subscript to the left of the element symbol directly below the mass number: . This is technically redundant, as each element is defined by its atomic number, so it is often omitted.
Mass number changes in radioactive decay
Different types of radioactive decay are characterized by their changes in mass number as well as atomic number, according to the radioactive displacement law of Fajans and Soddy. For example, uranium-238 usually decays by alpha decay, where the nucleus loses two neutrons and two protons in the form of an alpha particle. Thus both the atomic number and the number of neutrons decrease by 2 (Z: 92â†’90, n: 146â†’144), which decreases the mass number by 4 (A = 238â†’234); the result is an atom of thorium-234 and an alpha particle ():
On the other hand, carbon-14 naturally decays by radioactive beta decay, whereby one neutron is transmuted into a proton with the emission of an electron and an anti-neutrino. Thus the atomic number increases by 1 (Z: 6â†’7) and the mass number remains the same (A = 14), while the number of neutrons decreases by 1 (n: 8â†’7). The resulting atom is nitrogen-14, with seven protons and seven neutrons:
Another type of radioactive decay without change in mass number is emission of a gamma ray from a nuclear isomer or metastable excited state of an atomic nucleus. Since all the protons and neutrons remain in the nucleus unchanged in this process, the mass number is also unchanged.
Mass number and isotopic mass
The mass number gives an estimate of the isotopic mass measured in atomic mass units (u). For 12C the isotopic mass is exactly 12, since the atomic mass unit is defined as 1/12 of the mass of 12C. For other isotopes, the isotopic mass is usually within 0.1 u of the mass number. For example, 35Cl has a mass number of 35 and an isotopic mass of 34.96885.
There are two reasons for the difference between mass number and isotopic mass, known as the mass defect:
- The neutron is slightly heavier than the proton. This increases the mass of nuclei with more neutrons than protons relative to the atomic mass unit scale based on 12C with equal numbers of protons and neutrons.
- The nuclear binding energy varies between nuclei. A nucleus with greater binding energy has a lower total energy, and therefore a lower mass according to Einstein's mass-energy equivalence relation E = mc2. For 35Cl the isotopic mass is less than 35 so this must be the dominant factor.
Atomic mass of an element
The mass number should also not be confused with the relative atomic mass (also called atomic weight) of an element, which is the ratio of the average atomic mass of the different isotopes of that element (weighted by abundance) to the unified atomic mass unit. This weighted average can be quite different from the near-integer values for individual isotopic masses.
For instance, there are two main isotopes of chlorine: chlorine-35 and chlorine-37. In any given sample of chlorine that has not been subject to mass separation there will be roughly 75% of chlorine atoms which are chlorine-35 and only 25% of chlorine atoms which are chlorine-37. This gives chlorine a relative atomic mass of 35.5 (actually 35.4527 g/mol).
In chemistry and physics, the atomic number (also known as the proton number) is the number of protons found in the nucleus of an atom and therefore identical to the charge number of the nucleus. It is conventionally represented by the symbol Z. The atomic number uniquely identifies a chemical element. In an atom of neutral charge, the atomic number is also equal to the number of electrons.
The atomic number, Z, should not be confused with the mass number, A, which is the total number of protons and neutrons in the nucleus of an atom. The number of neutrons, N, is known as the neutron number of the atom; thus, A = Z + N. Since protons and neutrons have approximately the same mass (and the mass of the electrons is negligible for many purposes), the atomic mass of an atom is roughly equal to A.
Atoms having the same atomic number Z but different neutron number N, and hence different atomic mass, are known as isotopes. Most naturally occurring elements exist as a mixture of isotopes, and the average atomic mass of this mixture determines the element's atomic weight.
Loosely speaking, the existence of a periodic table creates an ordering for the elements. Such an ordering is not necessarily a numbering, but can be used to construct a numbering by fiat. Dmitri Mendeleev claimed he arranged his tables in order of atomic weight ("Atomgewicht") However, in deference to the observed chemical properties, he violated his own rule and placed tellurium (atomic weight 127.6) ahead of iodine (atomic weight 126.9). This placement is consistent with the modern practice of ordering the elements by proton number, Z, but this number was not known or suspected at the time.
A simple numbering based on periodic table position was never entirely satisfactory. Besides iodine and tellurium, several other pairs of elements (such as cobalt and nickel) were known to have nearly identical or reversed atomic weights, leaving their placement in the periodic table by chemical properties to be in violation of known physical properties. Another problem was that the gradual identification of more and more chemically similar and indistinguishable lanthanides, which were of an uncertain number, led to inconsistency and uncertainty in the numbering of all elements at least from lutetium (element 71) onwards (hafnium was not known at this time).
In 1911, Ernest Rutherford gave a model of the atom in which a central core held most of the atom's mass and a positive charge which, in units of the electron's charge, was to be approximately equal to half of the atom's atomic weight, expressed in numbers of hydrogen atoms. This central charge would thus be approximately half the atomic weight (though it was almost 25% off the figure for the atomic number in gold (Z = 79, A = 197), the single element from which Rutherford made his guess). Nevertheless, in spite of Rutherford's estimation that gold had a central charge of about 100 (but was element Z = 79 on the periodic table), a month after Rutherford's paper appeared, Antonius van den Broek first formally suggested that the central charge and number of electrons in an atom was exactly equal to its place in the periodic table (also known as element number, atomic number, and symbolized Z). This proved eventually to be the case.
The experimental situation improved dramatically after research by Henry Moseley in 1913. Moseley, after discussions with Bohr who was at the same lab (and who had used Van den Broek's hypothesis in his Bohr model of the atom), decided to test Van den Broek and Bohr's hypothesis directly, by seeing if spectral lines emitted from excited atoms fit the Bohr theory's demand that the frequency of the spectral lines be proportional to a measure of the square of Z.
To do this, Moseley measured the wavelengths of the innermost photon transitions (K and L lines) produced by the elements from aluminum (Z = 13) to gold (Z = 79) used as a series of movable anodic targets inside an x-ray tube. The square root of the frequency of these photons (x-rays) increased from one target to the next in a linear fashion. This led to the conclusion (Moseley's law) that the atomic number does closely correspond (with an offset of one unit for K-lines, in Moseley's work) to the calculated electric charge of the nucleus, i.e. the proton number Z. Among other things, Moseley demonstrated that the lanthanide series (from lanthanum to lutetium inclusive) must have 15 membersâ€”no fewer and no moreâ€”which was far from obvious from the chemistry at that time.
The conventional symbol Z presumably comes from the German word (atomic number).
Each element has a specific set of chemical properties as a consequence of the number of electrons present in the neutral atom, which is Z. The configuration of these electrons follows from the principles of quantum mechanics. The number of electrons in each element's ele
atom [Gr.,=uncuttable (indivisible)], basic unit of matter ; more properly, the smallest unit of a chemical element having the properties of that element. Structure of the Atom The atom consists of a central, positively charged core, the nucleus , and negatively charged particles called electrons that are found in orbits around the nucleus. The Nucleus Almost the entire mass of the atom is concentrated in the nucleus, which occupies only a tiny fraction of the atom's volume. The nucleus of an atom consists of neutrons and protons, the neutron being an uncharged particle and the proton a positively charged one. Their masses are almost equal. Atoms containing the same number of protons but different numbers of neutrons represent different forms, or isotopes , of the same element. The Electrons Surrounding the nucleus of an atom are its electrons; for a neutral atom, the number of electrons is equal to the atomic number. The outermost electrons of an atom determine its chemical and electrical properties. An atom may combine chemically with another atom in various ways, either by giving up or receiving electrons, thus setting up an electrical attraction between the atoms (see ion ), or by sharing one or more pairs of electrons (see chemical bond ). Because metals have few outermost electrons and tend to give them up easily, they are good conductors of electricity or heat (see conduction ). The electrons are often described as revolving about the nucleus as the planets revolve about the sun. This picture, however, is misleading. The quantum theory has shown that all particles in motion also have certain wave properties. For a particle the size of an electron, these properties are of considerable importance. As a result the electrons in an atom cannot be pictured as localized in space, but rather should be viewed as smeared out over the entire orbit so that they form a cloud of charge. The electron clouds around the nucleus represent regions in which the electrons are most likely to be found. The shapes of these clouds can be very complex, in marked contrast to the simple elliptical orbits of planets. Surprisingly, the sizes of all atoms are comparable, in spite of the large differences in the number of electrons they contain. Atomic Weight and Number The atomic number of an atom is simply the number of protons in its nucleus. The atomic weight of an atom is given in most cases by the mass number of the atom, equal to the total number of protons and neutrons combined. An atom may be conveniently symbolized by its chemical symbol with the atomic number and mass number written as subscript and superscript, respectively. For example, the symbol for uranium is U (atomic number 92); the isotopes of uranium with atomic weights 235 and 238 are indicated by 23592 U and 23892 U. Development of Atomic Theory Early Atomic Theory The atomic theory, which holds that matter is composed of tiny, indivisible particles in constant motion, was proposed in the 5th cent. BC by the Greek philosophers Leucippus and Democritus and was adopted by the Roman Lucretius. However, Aristotle did not accept the theory, and it was ignored for many centuries. Interest in the atomic theory was revived during the 18th cent. following work on the nature and behavior of gases (see gas laws ). From Dalton to the Periodic Table Modern atomic theory begins with the work of John Dalton, published in 1808. He held that all the atoms of an element are of exactly the same size and weight (see atomic weight ) and are in these two respects unlike the atoms of any other element. He stated that atoms of the elements unite chemically in simple numerical ratios to form compounds. The best evidence for his theory was the experimentally verified law of simple multiple proportions , which gives a relation between the weights of two elements that combine to form different compounds. Evidence for Dalton's theory also came from Michael Faraday's law of electrolysis . A major development was the periodic table , devised simultaneously by Dmitri Mendeleev and J. L. Meyer, which arranged atoms of different elements in order of increasing atomic weight so that elements with similar chemical properties fell into groups. By the end of the 19th cent. it was generally accepted that matter is composed of atoms that combine to form molecules. Discovery of the Atom's Structure In 1911, Ernest Rutherford developed the first coherent explanation of the structure of an atom. Using alpha particles emitted by radioactive atoms, he showed that the atom consists of a central, positively charged core, the nucleus , and negatively charged particles called electrons that orbit the nucleus. There was one serious obstacle to acceptance of the nuclear atom, however. According to classical theory, as the electrons orbit about the nucleus, they are continuously being accelerated (see acceleration ), and all accelerated charges radiate electromagnetic energy. Thus, they should lose their energy and spiral into the nucleus. This difficulty was solved by Niels Bohr (1913), who applied the quantum theory developed by Max Planck and Albert Einstein to the problem of atomic structure. Bohr proposed that electrons could circle a nucleus without radiating energy only in orbits for which their orbital angular momentum was an integral multiple of Planck's constant h divided by 2Ï€. The discrete spectral lines (see spectrum ) emitted by each element were produced by electrons dropping from allowed orbits of higher energy to those of lower energy, the frequency of the photon of light emitted being proportional to the energy difference between the orbits. Around the same time, experiments on x-ray spectra (see X ray ) by H. G. J. Moseley showed that each nucleus was characterized by an atomic number, equal to the number of unit positive charges associated with it. By rearranging the periodic table according to atomic number rather than atomic weight, a more systematic arrangement was obtained. The development of quantum mechanics during the 1920s resulted in a satisfactory explanation for all phenomena related to the role of electrons in atoms and all aspects of their associated spectra. With the discovery of the neutron in 1932 the modern picture of the atom was complete. Contemporary Studies of the Atom With many of the problems of individual atomic structure and behavior now solved, attention has turned to both smaller and larger scales. On a smaller scale the atomic nucleus is being studied in order to determine the details of its structure and to develop sources of energy from nuclear fission and fusion (see nuclear energy ), for the atom is not at all indivisible, as the ancient philosophers thought, but can undergo a number of possible changes. On a larger scale new discoveries about the behavior of large groups of atoms have been made (see solid-state physics ). The question of the basic nature of matter has been carried beyond the atom and now centers on the nature of and relations between the hundreds of elementary particles that have been discovered in addition to the proton, neutron, and electron. Some of these particles have been used to make new types of exotic "atoms" such as positronium (see antiparticle ) and muonium (see muon ). Bibliography See G. Gamow, The Atom and Its Nucleus (1961); H. A. Boorse and L. Motz, ed., The World of the Atom (2 vol., 1966); B. H. Bransden and C. J. Joachain, Physics of Atoms and Molecules (1986).
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Answers:atomic mass is protons plus neutrons (all in the nucleus) atomic number is the number of protons (or electrons) in an atom so mass = 18+17 = 35
Answers:OMG! I should have know this without finding a periodic chart. I used to be a biology major at BYU. I had four semesters of chemistry under my belt. Oh, well, I don't have to use this info any more, I guess. Mn
Answers:The atomic no. of Ca is 20. Atomic no. is also the mass or Proton no. and is equal to nucleon no. all of them are 20. See the periodic tables of elements.
Answers:An element's atomic number is the number of protons in its nucleus. Number of protons specifies atom type. The atomic number also specifies order on the periodic table. For an electrically neutral atom, atomic number also corresponds to count of electrons. Mass number is the number of nucleons (both protons and neutrons) in an atom's nucleus. This is roughly equal to the mass of the atom, because protons and neutrons are the most significant contributes to the atom's mass. Different neutron counts can exist among one particular atom, and atoms differing in neutron counts of the same element are called isotopes. The atomic mass is an overall average of the effective mass of billions of atoms distributed evenly based on isotopic abundance. 18. Key word: "atom". Atoms are defined as neutral. Ions are non-neutral systems resembling the atom. Thus, because the atom is neutral, 11 protons also means 11 electrons. 19. 6 protons, 8 neutrons, 6 electrons 20. "All atoms of a given element are identical." Isotopes indicate that not all atoms of the same type are identical, because the neutron counts differ. 22. Let's just say, the Bohr model is oversimplified. The Bohr model has us visualize the atom as a miniature solar system, which is an incomplete view of reality. Electrons aren't easily located, nor are they definitely in one location at a time.