Isotopes

12 Key excerpts on "Isotopes"

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  eBook - ePub

  The Britannica Guide to the Atom

  • Britannica Educational Publishing, Erik Gregersen(Authors)
  • 2010(Publication Date)
  • Britannica Educational Publishing

    (Publisher)

  CHAPTER 3 ISOTOPES

  A n isotope is a species of atoms of a chemical element with the same atomic number and position in the periodic table and nearly identical chemical behaviour but with different atomic masses and physical properties. Every chemical element has one or more Isotopes.

  An atom is first identified and labeled according to the number of protons in its nucleus. This atomic number is ordinarily given the symbol Z . The great importance of the atomic number derives from the observation that all atoms with the same atomic number have nearly, if not precisely, identical chemical properties. A large collection of atoms with the same atomic number constitutes a sample of an element. A bar of pure uranium, for instance, would consist entirely of atoms with atomic number 92. The periodic table of the elements assigns one place to every atomic number, and each of these places is labeled with the common name of the element, as, for example, calcium, radon, or uranium.

  Not all the atoms of an element need have the same number of neutrons in their nuclei. In fact, it is precisely the variation in the number of neutrons in the nuclei of atoms that gives rise to Isotopes. Hydrogen is a case in point. It has the atomic number 1. Three nuclei with one proton are known that contain 0, 1, and 2 neutrons, respectively. The three share the place in the periodic table assigned to atomic number 1 and hence are called Isotopes (from the Greek isos , meaning “same,” and topos , signifying “place”) of hydrogen.

  Many important properties of an isotope depend on its mass. The total number of neutrons and protons (symbol A ), or mass number, of the nucleus gives approximately the mass measured on the so-called atomic-mass-unit (amu) scale. The numerical difference between the actual measured mass of an isotope and A

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  Radiogenic Isotope Geochemistry

  A Guide for Industry Professionals

  • Bruce F. Schaefer(Author)
  • 2016(Publication Date)
  • Cambridge University Press

    (Publisher)

  1.2 Elements and Isotopes Economic geology deals with chemical processes, usually centred upon the behaviour of a single key chemical element, say Pb, Au or Cu. However, many elements contain Isotopes, which are in essence chemically identical with each other, but have a different mass due to their nuclear structure. This difference in nuclear configuration permits two fundamental types of behaviour that geoche- mists can use – namely fractionation of masses (Isotopes) of the same element, and radioactive decay from one element to another. In order to understand these, we need first to clarify what they are and distinguish between some important terms when describing atomic structure. 1.2.1 Atoms, ions and Isotopes An atom is the smallest component of a chemical element that retains the chemical properties of that element. Atoms contain three fundamental particles: protons and neutrons, which occur in the nucleus and are hence termed nucleons; and electrons, which orbit the dense nucleus. Protons and neutrons have approximately the same mass, whereas the electron is ~2000 times lighter than the nucleons. For the purposes of geological nuclear chemistry it is usually regarded as having negligible mass. Protons and neutrons weigh ~1 atomic mass unit (amu). An atom is electrically neutral, since the number of orbiting negatively charged electrons balances the number of positively charged protons in the nucleus. Neutrons have no electrical charge. Elements and Isotopes 5 The number of protons defines the atomic number (Z), and therefore which element it is. In other words, each element has a unique atomic number (e.g. Cu = 29; Au = 79), which governs the corresponding number of electrons and hence the chemical properties of the element. The atomic mass number (A) of an element is the total number of nucleons it contains, i.e. the sum of the protons (Z) and the neutrons (N). Therefore A = N + Z.

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  Cosmochemistry

  • Harry McSween, Jr, Gary Huss(Authors)
  • 2022(Publication Date)
  • Cambridge University Press

    (Publisher)

  Each element consists of between one and 10 stable Isotopes and many unstable (radioactive) ones, most with very short half-lives (the time it takes for half of the atoms present to decay). Isotopes of an element have the same number of protons, which define the element, but differ in the number of neutrons in the nucleus. For example, oxygen has eight protons and either eight, nine, or 10 neutrons. The Isotopes of oxygen are 16 O, 17 O, and 18 O; the numbers refer to the total number of protons þ neu- trons in the nucleus. The atomic number, Z, which is unique to each element, is the number of protons in the nucleus (equal to the number of electrons in the neutral atom). The mass number, A, is the number of protons plus the number of neutrons, N, in the nucleus (A ¼ Z þ N). Isotopes can either be stable or radioactive, and the iso- topes resulting from radioactive decay are said to be radiogenic. Natural samples contain 266 stable Isotopes comprising 81 elements, along with 65 radioactive iso- topes of these and nine additional elements. (When dis- cussing the Isotopes of more than one element, the term nuclide is typically used.) In addition, more than 1650 nuclides and 22 elements not found on Earth have been created in the laboratory with nuclear reactors and particle accelerators. Most of the man-made Isotopes are very short-lived. Most probably exist briefly in special natural settings, such as in an exploding supernova, but we typically have no way of detecting them. Radioactive technetium, which does not exist naturally in the solar system, has been identified in the spectra of highly evolved low-mass stars. The notation used for Isotopes, ions, and chemical compounds is illustrated by this example: 16 O 0 2 . The number to the upper left of the element symbol is the mass number (A).

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  Isotopic Tracers in Biology

  An Introduction to Tracer Methodology

  • Martin D. Kamen, Louis F. Fieser, Mary Fieser(Authors)
  • 2013(Publication Date)
  • Academic Press

    (Publisher)

  G E N E R A L P R O P E R T I E S OF N U C L E I 3 C. I S O T O P E S Because the chemical properties of the atom are determined by the value of the nuclear charge or atomic number, addition of neutrons to any nuclear complex of protons and neutrons changes the mass by an integral amount but does not change the nuclear charge. Since the nuclear charge determines the number of extranuclear electrons, which, in turn, deter-mines the chemistry of the atom, no change occurs in the chemical be-havior of the atom when neutrons are added to the atomic nucleus. Conse-quently there are nuclei, and hence atoms, which vary in nuclear mass but not in chemical nature. These are called Isotopes. Some elements have only one stable isotope each GBe 9 , 9F 19 , nNa 2 3 , 1 5 P 3 1 , etc.) ; others are mixtures of two or more stable Isotopes. Sulfur may be cited as an example. Four Isotopes of sulfur with mass numbers 32, 33, 34, and 36 are known. In the nomenclature described above these would be written i 6 S 3 2 , i 6 S 3 3 , leS 5 4 , and i 6 S 3 6 . Each of these nuclei contains A = 16 protons, and (A — Z) — 16, 17, 18, and 20 neutrons, respectively. The ratio of the number of neutrons to the number of protons for stable nuclei is very close to unity. It increases with increasing values of Ζ until at saBi 2 09 there is a ratio of 126/83 or 1.5. It is also possible for nuclei with the same mass number but different atomic number (isobars) to exist. Examples are 4sCd 113 and 4 9 I 1 1 1 1 3 , îsA 40 and 2oCa 40 . Finally, it is also possible that nuclei of identical charge and mass number may exist in slightly different configurations or energy states. Such nuclei are called isomers (see p. 15). The terms isotope, isobar, and isomer refer to particular species of atomic nuclei.

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  The Basics of Physics

  • Richard L. Myers(Author)
  • 2005(Publication Date)
  • Greenwood

    (Publisher)

  Nuclear medicine is used extensively for both diagnostic and therapeutic procedures. Radio- metric dating is an invaluable tool to scientists who use this method for applications such as dating relics, determining the age of the Earth, and studying climate change. In this chapter we will extend our examination of the nucleus started in chapter 10 and examine in detail the scientific principles behind human use of the atom. Atomic Mass, Atomic Numbers, and Isotopes The most basic unit of a chemical ele- ment that can undergo chemical change is an atom. Atoms of any element are identi- fied by the number of protons and neutrons 206 Nuclear Physics in the nucleus. The number of protons in the nucleus of an element is given by the atomic number. Hydrogen has one proton in its nucleus, so its atomic number is 1. The atomic number of carbon is 6 because each carbon atom contains six protons in its nucleus. Besides protons, the nucleus con- tains neutrons. The number of protons plus the number of neutrons is the mass number of an element. A standard method of sym- bolizing an element is to write the element with the mass number as a superscript and the atomic number as a subscript. Carbon-12 is written as An element's atomic number is constant, but most elements have varying mass numbers. For example, all hydrogen atoms have an atomic number of 1, and almost all hydrogen atoms have a mass number of 1. This means most hydrogen atoms have no neutrons. Some hydrogen atoms have mass numbers of 2 or 3, with one and two neutrons, respec- tively. Different forms of the same element that have different mass numbers are known as Isotopes. Three Isotopes of hydrogen are symbolized as H :H Deuterium H Hydrogen Tritium Isotopes of elements are identified by their mass numbers. Hence, the isotope C-14 is the form of carbon that contains eight neutrons.

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  Isotope Geology

  • Claude J. Allègre, Christopher Sutcliffe(Authors)
  • 0(Publication Date)
  • Cambridge University Press

    (Publisher)

  But in 11 Isotopy each position there are several Isotopes which di¡er by the number of neutrons N they con- tain, that is, by their mass.These Isotopes are created during nuclear processes which are collectively referred to as nucleosynthesis and which have been taking place in the stars throughoutthehistoryofthe Universe (see Chapter 4). The isotopic composition of a chemical element is expressed either as a percentage or more convenientlyas a ratio. A reference isotope is chosen relative to which the quantities of other Isotopes are expressed. Isotope ratios are expressed in terms of numbers of atoms and notofmass. For example, to study variations in the isotopic composition ofthe element strontium brought about by the radioactive decay of the isotope 87 Rb, we choose the 87 Sr/ 86 Sr isotope ratio. To study the isotopic variations of lead, we consider the 206 Pb/ 204 Pb, 207 Pb/ 204 Pb, and 208 Pb/ 204 Pb ratios. 1.3.1 The chart of the nuclides The isotopic composition of all the naturally occurring chemical elements has been deter- mined, that is, the numberof Isotopes and their proportions have been identi¢ed.The¢nd- ings have been plotted as a ( Z, N) graph, that is, the number of protons against the number ofneutrons. Figure 1.3 , details ofwhich are given in the Appendix, prompts a few remarks. Firstofall,thestableIsotopesfallintoaclearlyde¢nedzone,knownasthe valleyofstability because it corresponds to the minimum energy levels of nuclides. Initially this energy valley follows the diagonal Z ¼ N .Then, after N ¼ 20, the valley falls away from the diagonal on the side of a surplus of neutrons. It is as if, as Z increases, an even greater number of neutrons is neededtopreventthe electricallychargedprotons from repelling each otherandbreaking the nucleus apart.

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  Soil and Environmental Analysis

  Modern Instrumental Techniques

  • Keith A. Smith, Malcolm S. Cresser, Keith A. Smith, Malcolm S. Cresser(Authors)
  • 2003(Publication Date)
  • CRC Press

    (Publisher)

  8 Measurement of RadioIsotopes and Ionizing Radiation Olivia J. Marsden and Francis R. Livens The University of Manchester ; Manchester ; England I. INTRODUCTION Approximately 1700 different Isotopes are known, of which around 275 are stable. The remainder are radioactive; that is, their nuclear configurations are unstable and can change to more stable forms by nuclear transformations that are collectively known as radioactive decay. These radioactive decay processes are accompanied by the emission of particles and/or photons from the nucleus. Isotopes (or nuclides) are distinguished by the number of protons and neutrons (collectively known as nucleons) they contain and are commonly designated using mass number (A: number of protons + neutrons) and atomic number (Z: number of protons). For example, l64C is an isotope of carbon in which the nucleus contains 14 nucleons, of which six are protons. The proton number defines the chemical identity of the atom, since the proton charge must be balanced by the appropriate number of electrons, but it also duplicates the information provided by the chemical symbol and, in practice, is often omitted, hence 1 4 C. Differences in the neutron number may control the stability, or otherwise, of a nucleus but have only subtle effects on chemistry, although these can be exploited in studies of stable isotope fractionation in natural systems, for example 2 H / 1 H, 1 3 C/ 1 2 C, 1 5 N / 1 4 N, 1 7 0 / 1 6 0, 3 4 S/32S (see Chap. 9). Only a minority of the unstable Isotopes are formed in nature. Most are man-made, and the majority of these are available only in such small amounts, or are so short-lived or both, that they are unlikely to be 345 346 Marsden and Livens Figure 1 Isotopes produced by the decay of 2 3 8 U. encountered in the environment or to be of any use as radiotracers. The naturally occurring radioIsotopes fall into three groups: 1.

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  Radioactive Tracers in Biology

  An Introduction to Tracer Methodology

  • Martin D. Kamen, Louis F. Fieser, Mary Fieser(Authors)
  • 2013(Publication Date)
  • Academic Press

    (Publisher)

  The atomic number is usually written as a left subscrip to the chemical symbol, i.e., iH, nNa, i 5 P. The nucleus is completely identified if both the mass number A and the atomic number Z are indi cated; thus iH 1 , nNa 2 3 , and i B P 31 refer to certain kinds of nuclei for th< elements hydrogen, sodium, and phosphorus, respectively. In fact, th< chemical symbol is superfluous in this nomenclature but is retained fo] convenience in writing nuclear reactions. 1 2 RADIOACTIVE TRACERS IN BIOLOGY Isotopes. According to modern concepts, the nuclei of atoms are built up by combinations of protons and neutrons. The neutron has a mass number of 1 but is uncharged. It may be considered the uncharged analog of the proton. Because the chemical properties of the atom are determined by the value of the nuclear charge or atomic number, addition of neutrons to any nuclear complex of protons and neutrons changes the mass by an integral amount but does not change the nuclear charge. Since the nuclear charge determines the number of extranuclear electrons, which, in turn, determines the chemistry of the atom, no change occurs in chemical behavior of the atom when neutrons are added to the atomic nucleus. Consequently, there are nuclei, and hence atoms, which vary in nuclear mass but not in chemical nature. These are called Isotopes. Some elements have only one stable isotope each ( 4 Be 9 , 9F 1 9 , nNa 2 3 , 15P 3 1 , etc.); others are mixtures of two or more stable Isotopes. On this basis the mass number A is the total number of particles in the nucleus. The difference A — Z is the number of neutrons. Sulfur may be cited as an example. Four Isotopes of sulfur with mass numbers 32, 33, 34, and 36 are known. In the nomenclature discussed above these would be written ieS 32 , uS 3 3 , 1 6 S 3 4 , and 1 6 S 3 6 . Each of these nuclei contains A = 16 protons, and (A — Z) = 16, 17, 18, and 20 neutrons, respectively.

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  Isotopes in Biology

  • George Wolf(Author)
  • 2013(Publication Date)
  • Academic Press

    (Publisher)

  IV Principles and Conditions for the Use of Isotopes in Biology Section I. Identical Chemical and Biochemical Behavior of Isotopic Compounds A. Cases of Nonidentical Chemical Behavior 1. T H E ISOTOPE EFFECT The tacit assumption is made by the isotope chemist that the isotopic compound behaves exactly like the normal compound. In fact, identical chemical behavior is the first condition for the use of Isotopes in chemical or biological experiments. The aim is to trace the normal molecules through reactions and transformations. If the isotopic molecules are to be effective tracers, they must obviously behave in the same way as the normal ones. The radioactive molecule expresses its distinctiveness only at the moment of disintegration of the isotopic atom within it. Chemical reactions between atoms depend on the num-ber and arrangement of electrons orbiting around the nucleus, and these are identical for normal and isotopic atoms. However, the rates of chemical reactions are to some extent influenced by the mass of the reacting atoms, as well as by the number of electrons, and Isotopes are, of course, different in mass from normal atoms. Therefore, differences 4 5 46 IV. PRINCIPLES AND CONDITIONS in behavior between isotopic and normal compounds will be observed, if one looks at reaction velocities and equi-libria of reactions, or the consequences thereof. Molecules containing the lighter isotope react faster than those which carry the heavier one. This difference in behavior is called the Isotope Effect. It can generally be assumed to be negligible for the Isotopes of carbon, since the differences between reaction rates depend on the ratio of the atomic weights, which for C 12 and C 14 is not far from unity (14/12). But it becomes a factor of some importance for the hy-drogen Isotopes, H 1 , H 2 , and H 3 , since the ratio of atomic weights is now 2/1 or 3/1, and maximum rate differences can be 73$ between H 3 and H 1 .

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  Visualizing Everyday Chemistry

  • Douglas P. Heller, Carl H. Snyder(Authors)
  • 2015(Publication Date)
  • Wiley

    (Publisher)

  It appears as a superscript immediately preceding an element’s symbol. A neutron re- sembles a proton except that the neutron carries no electri- cal charge. At least one neutron resides in all nuclei except in the nuclei of the simplest and most abundant isotope of hydrogen, symbolized as shown. Mass number ⎯→ Atomic number ⎯→ 1 1 H 3 Isotopes and Atomic Mass 40 • What are Isotopes? Isotopes are atoms of a given element that differ in the number of neutrons they contain and therefore in their mass Figure 2.14 • The Isotopes of hydrogen Figure 2.20 • A version of the modern periodic table The term periodic table reflects the periodicity, or repeat- ing pattern, of the elements as they appear in the rows and columns of this chart. For example, for helium, neon, argon, krypton and xenon, a series of gases with little or no reactiv- ity, the difference in atomic numbers—8, 8, 18, 18—is mir- rored by the difference in atomic numbers of lithium, sodium, potassium, rubidium and cesium, a series of metals that react similarly with water.

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  Nuclear Engineering Fundamentals

  A Practical Perspective

  • Robert E. Masterson(Author)
  • 2017(Publication Date)
  • CRC Press

    (Publisher)

  Some heavy elements of particular importance to the nuclear power industry (such as uranium and plutonium) have a very large number of Isotopes, but only a few of them exist in measurable amounts. For example, uranium has a total of nine Isotopes, but only three of them are considered to be stable. Plutonium has a total of 20 Isotopes, but none of them are considered to be stable on a geological time frame. Elements of an intermediate atomic weight such as barium, cesium, iodine, xenon, krypton, rubidium, and strontium also have Isotopes that have important nuclear and medical applications.

  Elements that are not stable decay primarily through alpha or beta decay. This decay is sometimes accompanied by the emission of a gamma ray (an energetic photon) that helps to return the product nucleus to a stable state.

  Student Exercise 9.3 If there were more than 120 elements in the periodic table, what would you expect the next orbital in the sixth shell to be?

  TABLE 9.6 List of the Elements and Their Properties

  9.7 Meaning of the Fractional Atomic Weights

  Hydrogen is the lightest naturally occurring element with one proton and one electron. However, it has two additional Isotopes called deuterium and tritium that have one and two neutrons, respectively. Another important point to keep in mind is that many of the elements in the periodic table have atomic weights that are not whole numbers such as 1, 2, 4, 6, or 12 . For example, the atomic weight of carbon, which is usually made out of six protons and six neutrons, happens to be 12.011. The reader may wonder why the atomic weight is not a whole number. Did someone make a mistake when the periodic table was built or is there something more complex happening here?

  The answer to this question lies in the fact that about two-thirds of the naturally occurring have more than one isotope,* and the atomic weights in the periodic table simply represent the measured average of the atomic weights of these Isotopes . For example, carbon has three naturally occurring Isotopes—Carbon-12, Carbon-13, and Carbon-14, and each of these Isotopes occurs in nature in slightly different amounts.

  Hence, the weight for the element carbon that appears in the periodic table is the average of the atomic weight of these three Isotopes—adjusted for the relative amounts of each isotope that are present

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  Groundwater Geochemistry and Isotopes

  • Ian Clark(Author)
  • 2015(Publication Date)
  • CRC Press

    (Publisher)

  Rare are configurations such as HD 18 O (atomic mass 21) or heavy water D 2 O (atomic mass 20) that is concentrated and used as a neutron moderator in CanDU nuclear reactors. Although oxygen can exist with more than 10 or fewer than 8 neutrons, these radionuclides are not stable and readily decay to other elements. Similarly, other ele-ments can have a greater or lesser complement of neutrons, creating a host of stable and unstable or radioactive nuclides (atoms with specific number of neutrons and protons). The 92 naturally occurring elements have some 250 stable Isotopes. The stable and unstable Isotopes for the light elements up to atomic number 8 are shown on a neutron–proton chart in Figure 4.1. The law of mass action applies not only for solutes in groundwaters but also for Isotopes. Stable Isotopes of an element behave the same chemically, and so partici-pate in all the reactions discussed in Chapter 3. However, their slight mass difference affects bond strength such that they accumulate preferentially on one side or the other of a reaction. It is this mass-dependent partitioning or fractionation of iso-topes that makes them good tracers of the origin and reactions of water and solutes. RadioIsotopes also fractionated by their mass, but it is their decay that brings the important element of time to hydrogeology. 98 Groundwater Geochemistry and Isotopes STABLE ISOTOPE FRACTIONATION AND DISTILLATION Isotopes are a useful tool to trace the origin of water and solutes. Like geochemical components of a system, the partitioning of Isotopes in the environment is facilitated by the thermodynamics of equilibrium and kinetic reactions. For most systems, their distribution is controlled by interplay of two processes, the fractionation of Isotopes during any physical or chemical reaction, and the distillation of Isotopes from reac-tant reservoir as the reaction proceeds. These two processes work together to parti-tion Isotopes into different reservoirs.

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