Nuclear ChemistryNuclear chemistry is a branch of chemistry in which the nuclear chemists frequently span several areas such as organic, analytical, inorganic and physical chemistry. Nuclear analytical techniques are an important part of the arsenal of the modern analytical chemist. 

The study of the actinides and trans-actinide elements has involved the joint efforts of nuclear and inorganic chemists in extending knowledge of the periodic table. Nuclear chemistry is concerned with the changes happening in the nucleus of the atom.

The emission of radiations from a radioactive material comes from the fact that a radioactive isotope is unstable and it converts to a stable isotope by emitting radiations like α, ß positron and Gamma -rays, etc. What is Nuclear Chemistry?Back to Top

1. A nuclear reaction is different from a chemical reaction.2. In a chemical reaction, atoms of the reactants combine by a rearrangement of extra nuclear electrons but the nuclei of the atoms remain unchanged.3. In a nuclear reaction, however, it is the nucleus of the atom which is involved.4. The number of protons or neutrons in the nucleus changes to form a new element itself.

"A study of the nuclear changes in atoms is termed as Nuclear chemistry". So, nuclear chemistry is the study of phenomenon involving nuclear reactions, like radioactivity. Nuclear

chemistry also deals with the energy released from nuclear reactions, and its usage.History of Nuclear ChemistryBack to Top

The history of Nuclear chemistry dates back to 1895, with the discovery of X- rays by William Roentgen. In early 1896, Henri Becquerel was carrying out a series of experiments on fluorescence. He had used photographic film between two pieces of paper. 

When he developed the photographic film, he found that it had the same appearance as if it had been exposed to light. And after this, by accident, he developed the photographic plates, which was kept in the same drawer, as Uranium. To his surprise, the plate had been blackened. He thought that it was a new type of fluorescence. But, actually, he had come across a phenomenon of radioactivity. So, accidentally, radioactivity was discovered by Henri Becquerel.

The name radioactivity, was invented some time later by Marie Curie. She won Nobel prize for her discovery in 1903 with Henri Becquerel and Pierre Curie. Thereby evolved the branch of chemistry, called Nuclear chemistry. 

The discovery of radioactivity also brought into account many other processes, such as fission and fusion, which again were used as source of energy in many reactors. And also, with the discovery of radioactivity and other phenomenon related to radioactive elements, many new elements were brought into light.

Nuclear Symbol ChemistryBack to Top

There are various types of radiations involved in Nuclear chemistry. They have their own representations. Let us look into these radiations and their symbols.

Types of radiations and nuclear symbols

The radioactive radiations are of three types. They were sorted out by Rutherford in 1902, by passing them between two oppositely charged plates. The ones bending towards negative plate carried positive charge and were named as alpha rays. Those bending towards the positive plate and carrying negative charge were called as beta rays. The third type of radiation, being uncharged, passed straight through the electric field and were named as gamma rays.

Symbols of these rays

1. Alpha rays

Represented as α. These are positively charged rays. Since the alpha rays has a mass of 4 amu and charge _2, they are actually helium nuclei. So, they are also represented as

4 2He or 4

 2 α

2. Beta Rays

Negatively charged rays. They are represented as: β. Since they have a mass similar to electron, they are also represented as e-. They have a unit negative charge.

0-1e or 0

-1 β 

3. Gamma rays

These rays are neutral, with no charge. They are simply represented by the symbol : γ

RadioactivityBack to Top

A number of elements such as Uranium, and radium are unstable. Their atomic nucleus breaks on its own accord to form a smaller atomic nucleus of another element.

The protons and neutrons, in the unstable nucleus, regroup to give the new nucleus. This causes the release of excess particles and energy from original nucleus, which we know as radiation. The elements whose atomic nucleus emits radiation are said to be radioactive.

Radioactive decay can be defined as: " The spontaneous breaking down of the unstable atoms is termed radioactive disintegration or radioactive decay."

" The disintegration or decay of unstable atoms accompanied by emission of radiation is called radioactivity". The radioactive radiations can be detected and measured by a number of

methods. Some important methods are:

Cloud chamber method Geiger-Muller counter Ionization chamber method Scintillation counter method

Three types of radioactive decay occurs

1. Alpha decay .2. beta decay.3. Gamma decay .

→   Read   More Nuclear Waste FactsBack to Top

The major concern, on using the nuclear materials for any purpose, is their disposal. It is a known fact that the nuclear reactions, or rather the radioactivity, does not stop anytime. It is a spontaneous process, therefore, continues even after the material has been discarded. Thus, the disposal of the nuclear waste has to be done with utmost care.

The products of fission, like Ba-139 and Kr-92 are themselves radioactive. They emit dangerous radiations for several hundred years. The waste is usually packed in concrete barrels which are buried deep in the earth or dumped in sea. But the fear is that any leakage and corrosion of the storage vessels may eventually contaminate the water supplies.

Uses of Nuclear ChemistryBack to Top

Nuclear chemistry finds its use in many fields. Though there are disadvantages in the form of atom bombs, nuclear reactivity for war fare, etc, they find many useful applications too. Some of them are

1. Light -water nuclear power plant

Most commercial power plants today are light-water reactors. U-235 is used as a fuel here. The Uranium -235 rods are submerged into it. A lot of energy is produced from the reaction in the nuclear reactor. A reactor, once started, can supply power for a lot of generations. About 15% of consumable electricity in U.S.A is provided by Light-water reactors.

2. Breeder reactorThis is again another reactor which taps energy from nuclear reactions. Here too, U-235 is used for the production of electricity.

3. Radioactive dating

This is a very important use of radioactivity. The age of an old piece of wood can be determined using radioactive dating technique. A plant, while alive, takes up both normal carbon, C-12 and radioactive carbon, C-14. When the plant dies, uptake of carbon from atmosphere stops. Though, the C-12 does not show any change, the decay of C-14 starts with the release of Beta radiation. This helps in detecting of the age of a wood piece.

4. Medicine

Many radionuclide are used in medicine to detect cancerous cells, any other defect in organs, etc. These radioactive elements are combined with other compounds, or elements and administered orally. And, after some time, the path traveled by these radio isotopes are detected using a detector. The complete length traveled by it can be seen and the problems can be easily detected.

Themes   >   Science   >   Chemistry   >   Nuclear Chemistry   >   Isotopes   > What are Natural Isotopes?

Isotopes are atoms of the same element with different masses. They get these different masses by having different numbers of neutrons in their nucleii. They are the same type of atom, however, because their nucleii have the same number of protons in them.Isotopes of atoms that occur in nature come in two flavors: stable and unstable (radioactive). Some of the unstable isotopes are only moderately unstable and can therefore still persist in nature today. The isotope 238-U is a good example. It is radioactive but it's half life is 4.43 billion years. The Earth itself is 4.55 billion years old, so we now have roughly half of the 238-U on Earth that we had when Earth was formed. When an unstable isotope decays, it makes a new atom of a different element. Stable, isotopes, on the other hand, do not decay.What determines whether an isotope is stable or unstable, and if it is unstable, how unstable it is (i.e., how short it's half life is) depends on the energy of the configuration of that particular nucleus. This is a subject of nuclear physics and too detailed to go into here.Just so you know, there are also non-Natural (man made) isotopes. These are all radioactive.Themes > Science > Chemistry > Nuclear Chemistry > Isotopes > Isotope Properties and Neutron Cross Sections

This is subsection lists most isotopes of interest for direct, indirect, or theoretical nuclear weapon applications. Basic isotopes characteristics, and summary tables of significant neutron reaction cross sections are included.The neutronic data is derived from the authoritative ENDF-VI evaluated nuclear data base compiled and maintained by the National Nuclear Data Center (NNDC) at the Brookhaven National Laboratory (BNL). The data listings below were prepared with the assistance of the Japan Atomic Energy Research Insititue (JAERI).Notes:

The Maxwellian average cross sections are for a peak neutron energy distribution at 0.0253 eV (a room temperature thermal distribution).

Molar volume is the minimum volume per mole for the densest phase at standard pressure and temperature (STP)

SF = spontaneous fission Critical mass estimates for the fissile isotopes are given for bare spheres at the

densest STP phase. Where available these estimates are from other sources and are based either on actual experimental measurement, or reasonably sophisticated numerical computations. All fissile isotopes also include for comparison critical mass estimates made by me, using an exact criticality solution with a one-group representation of neutronic properties. The one-group parameters are fission spectrum averages calculated from the ENDF-VI evaluated nuclear data base.

As can be seen where outside critical mass estimates are also available, this one-group calculation method consistently underestimates the true critical mass - primarily because it does not take into account the effects of inelastic scattering in softening the neutron spectrum. The one-group calculated critical mass estimates are thus lower bounds on the true value. Comparison between the one-group calculations and the actual values for the highly fissile isotopes for which good experimental data is available (U-233, U-235, Pu-239, and Pu-241) shows a consistent underestimate of 70-75% of the true value. For less fissile isotopes, where critical mass estimates have been offered by others (these are mostly calculated estimates also, but with more sophisticated models), the underestimates are more severe (at worst 22-29% of the 'true' value for Pu-242). This too is to be expected because the effects of inelastic scattering is relatively greater in less fissile materials. On the other hand, the estimates for extremely fissile transuranics like the californium isotopes should be quite good.

Isotope

From Wikipedia, the free encyclopedia

This article is about the atomic variants of chemical elements. For the British jazz fusion band, see Isotope (band).

"Isotopes" redirects here. For the minor league baseball team, see Albuquerque Isotopes.

Isotopes are variants of a particular chemical element. While all isotopes of a given element share the same number of protons, each isotope differs from the others in its number of neutrons. The

term isotope is formed from the Greek roots isos (ἴσος "equal") and topos (τόπος "place"). Hence: "the same place," meaning that different isotopes of a single element occupy the same position

on the periodic table. The number of protons within the atom's nucleus uniquely identifies an element, but a given element may in principle have any number of neutrons. The number

of nucleons(protons and neutrons) in the nucleus is the mass number, and each isotope of a given element has a different mass number.

For example, carbon-12, carbon-13 and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13 and 14 respectively. The atomic number of carbon is 6 which means that

every carbon atom has 6 protons, so that the neutron numbers of these isotopes are 6, 7 and 8 respectively.

Isotope vs. nuclide

A nuclide is an atom with a specific number of protons and neutrons in the nucleus, for example carbon-13 with 6 protons and 7 neutrons. The nuclide concept (referring to individual nuclear species) emphasizes nuclear properties over chemical properties, while the isotope concept (grouping all atoms of each element) emphasizes chemical over nuclear. The neutron number has drastic effects on nuclear properties, but its effect on chemical properties is negligible in most elements, and still quite small in the case of the very lightest elements, although it does matter in some circumstances (for hydrogen, the lightest of all elements, the isotope effect is large enough to strongly affect biology). Since isotope is the older term, it is better known than nuclide, and is still sometimes used in contexts where nuclide might be more appropriate, such as nuclear technology and nuclear medicine.

[edit]Notation

An isotope and/or nuclide is specified by the name of the particular element (this indicates the atomic number implicitly) followed by a hyphen and the mass number (e.g. helium-3, helium-4,carbon-12, carbon-14, uranium-235 and uranium-239).[1] When a chemical symbol is used, e.g., "C" for carbon, standard notation (now known as "AZE notation" because A is the mass number, Z the atomic number, and E for element) is to indicate the number of nucleons with a superscript at the upper left of the chemical symbol and to indicate the atomic number with a subscript at the lower left (e.g. 32He, 42He, 126C, 146C, 23592U, and 23992U, respectively).[2] Since the atomic number is implied by the element symbol, it is common to state only the mass number in the superscript and leave out the atomic number subscript (e.g. 3He, 4He, 12C, 14C, 235U, and 239U, respectively). The letter m is sometimes appended after the mass number to indicate a nuclear isomer, a metastable or energetically-excited nuclear state (rather than the lowest-energy ground state), for example 180m73Ta (tantalum-180m).[edit]Radioactive, primordial, and stable isotopes

Some isotopes are radioactive, and are therefore described as radioisotopes or radionuclides, while others have never been observed to undergo radioactive decay and are described as stable isotopes. For example, 14C is a radioactive form of carbon while 12C and 13C are stable isotopes. There are about 339 naturally occurring nuclides on Earth,[3] of which 288 are primordial nuclides, meaning that they have existed since the solar system's formation. These include 34 nuclides with very long half-lives (over 80 million years) and 254 which are formally considered as "stable isotopes",[3] since they have not been observed to decay.

Many apparently "stable" isotopes are predicted by theory to be radioactive, with extremely long half-lives (this does not count the possibility of proton decay, which would make all nuclides ultimately unstable). Of the 255 nuclides never observed to decay, only 90 of these (all from the first 40 elements) are stable in theory to all known forms of decay. Element 41 (niobium) is theoretically unstable via spontaneous fission, but this has never been detected. Many other stable nuclides are in theory energetically susceptible to other known forms of decay, such as alpha

decay or double beta decay, but no decay products have yet been observed. The predicted half-lives for these nuclides often greatly exceed the estimated age of the universe, and in fact there are also 27 known radionuclides (see primordial nuclide) with half-lives longer than the age of the universe.

Adding in the radioactive nuclides that have been created artificially, there are more than 3100 currently known nuclides.[4] These include 905 nuclides which are either stable, or have half-lives longer than 60 minutes. See list of nuclides for details.

[edit]History

[edit]Radioactive isotopes

The existence of isotopes was first suggested in 1912 by the radiochemist Frederick Soddy, based on studies of radioactive decay chains which indicated about 40 different species described asradioelements (i.e. radioactive elements) between uranium and lead, although the periodic table only allowed for 11 elements from uranium to lead[5].

Several attempts to separate these new radioelements chemically had failed[6]. For example, Soddy had shown in 1910 that mesothorium (later shown to be Ra-228), radium (Ra-226, the longest-lived isotope), and thorium X (Ra-224) are impossible to separate[7]. Attempts to place the radioelements in the periodic table led Soddy and Kazimierz Fajans independently to propose their radioactive displacement law in 1913, to the effect that alpha decay produced an element two places to the left in the periodic table, while beta decay emission produced an element one place to the right.[8] Soddy recognized that emission of an alpha particle followed by two beta particles led to the formation of an element chemically identical to the initial element but with a mass four units lighter and with different radioactive properties.

Soddy proposed that several types of atoms (differing in radioactive properties) could occupy the same place in the table. For example, the alpha-decay of uranium-235 forms thorium-231, while the beta decay of actinium-230 forms thorium-230[6] The term “isotope”, Greek for “at the same place”, was suggested to Soddy by Margaret Todd, a Scottish physician and family friend, during a conversation in which he explained his ideas to her.[9][10][7][11][12]

In the bottom right corner of JJ Thomson's photographic plate are the separate impact marks for the two isotopes of neon: neon-20 and neon-22.

In 1914 T.W. Richards found variations between the atomic weight of lead from different mineral sources, attributable to variations in isotopic composition due to different radioactive origins[6][13].

[edit]Stable isotopes

The first evidence for isotopes of a stable (non-radioactive) element was found by J. J. Thomson in 1913 as part of his exploration into the composition ofcanal rays (positive ions). Thomson channeled streams of neon ions through a magnetic and an electric field and measured their deflection by placing a photographic plate in their path. Each stream created a glowing patch on the plate at the point it struck. Thomson observed two separate patches of light on the photographic plate (see image), which suggested two different parabolas of deflection. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest.

F.W. Aston subsequently discovered different stable isotopes for numerous elements using a mass spectrograph. In 1919 Aston studied neon with sufficient resolution to show that the two isotopic masses are very close to the integers 20 and 22, and that neither is equal to the known molar mass (20.2) of neon gas. This is an example of Aston’s whole number rule for isotopic masses, which states that large deviations of elemental molar masses from integers are primarily due to the fact that the element is a mixture of isotopes. Aston similarly showed that the molar mass of chlorine (35.45) is a weighted average of the almost integral masses for the two isotopes Cl-35 and Cl-37.[14]

[edit]Variation in properties between isotopes

[edit]Chemical and molecular properties

A neutral atom has the same number of electrons as protons. Thus, different isotopes of a given element all have the same number of protons and share a similar electronic structure. Because the chemical behavior of an atom is largely determined by its electronic structure, different isotopes exhibit nearly identical chemical behavior. The main exception to this is the kinetic isotope effect: due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. This is most pronounced for protium (1H) and deuterium (2H), because deuterium has twice the mass of protium. The mass effect between deuterium and the relatively light protium also affects the behavior of their respective chemical bonds, by means of changing the center of gravity (reduced mass) of the atomic systems. However, for heavier elements, which have more neutrons than lighter elements, the ratio of the nuclear mass to the collective electronic mass is far greater, and the relative mass difference between isotopes is much less. For these two reasons, the mass-difference effects on chemistry are usually negligible.

Isotope half-lives. Note that the plot for stable isotopes diverges from the line, protons Z = neutrons N as the element number Z becomes larger

In similar manner, two molecules that differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structure and therefore almost indistinguishable physical and chemical properties (again with deuterium providing the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. As a consequence, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range.

[edit]Nuclear properties and stability

See also: Stable isotope, List of nuclides, and List of elements by stability of isotopes

Atomic nuclei consist of protons and neutrons bound together by the residual strong force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, stabilize the nucleus in two ways. Their copresence pushes protons slightly apart, reducing the electrostatic repulsion between the protons, and they exert the attractive nuclear force on each other and on protons. For this reason, one or more neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, so does the ratio of neutrons to

protons necessary to ensure a stable nucleus (see graph at right). For example, although the neutron:proton ratio of 32He is 1:2, the neutron:proton ratio of 23892U is greater than 3:2. A number of lighter elements have stable nuclides with the ratio 1:1 (Z = N). The nuclide 4020Ca (calcium-40) is the heaviest stable nuclide with the same number of neutrons and protons; (theoretically, it's S-32) all heavier stable nuclides contain more neutrons than protons.[edit]Numbers of isotopes per element

Of the 80 elements with a stable isotope, the largest number of stable isotopes observed for any element is ten (for the element tin). Xenon is the only element that has nine stable isotopes. No element has eight stable isotopes. Four elements have seven stable isotopes, nine have six stable isotopes, nine have five stable isotopes, nine have four stable isotopes, five have three stable isotopes, 16 have two stable isotopes (counting 180m73Ta as stable), and 26 elements have only a single stable isotope (of these, 19 are so-calledmononuclidic elements, having a single primordial stable isotope that dominates and fixes the atomic weight of the natural element to high precision; 3 radioactive mononuclidic elements occur as well).[15] In total, there are 255 nuclides that have not been observed to decay. For the 80 elements that have one or more stable isotopes, the average number of stable isotopes is 255/80 = 3.2 isotopes per element.[edit]Even and odd nucleon numbers

Even vs. odd mass number (A). Hydrogen-1 not included

Even Odd Total

Stable 152 102 254

Long-lived 25 8 33

All primordial 177 110 287

The proton:neutron ratio is not the only factor affecting nuclear stability. Adding neutrons to isotopes can vary their nuclear spins and nuclear shapes, causing differences in neutron capture cross-sections and gamma spectroscopy and nuclear magnetic resonance properties.

[edit]Even mass number

Even-mass-number nuclides, which comprise 152/255 = ~ 60% of all stable nuclides, are bosons, i.e. they have integer spin. Almost all (148 of the 152) are even-proton, even-neutron (EE) nuclides, which necessarily have spin 0 because of pairing. The remainder of the stable bosonic nuclides are 5 odd-proton, odd-neutron stable nuclides (see below, these are: 21H, 63Li, 105B, 14

7N and 180m73Ta). All have nonzero integer spin.[edit]Pairing effects

Even/odd Z, N (Hydrogen-1 not included)

p,n EE OO EO OE Total

Stable 147 5 53 49 254

Long-lived 21 4 3 5 33

All primordial 168 9 56 54 287

Beta decay of an even-even nucleus produces an odd-odd nucleus, and vice versa. An even number of protons or of neutrons are more stable (lowerbinding energy) because of pairing effects, so even-even nuclei are much more stable than odd-odd. One effect is that there are few stable odd-odd nuclides, but another effect is to prevent beta decay of many even-even nuclei into another even-even nucleus of the same mass number but lower energy, because decay proceeding one step at a time would have to pass through an odd-odd nucleus of higher energy. Double beta decay directly from even-even to even-even skipping over an odd-odd nuclide is only occasionally possible, and even then with a half-life greater than a billion times the age of the universe. For example, the double beta emitter 116Cd has a half-life of 2.9×1019 years. This makes for a larger number of stable even-even nuclides, up to three for some mass numbers, and up to seven for some atomic (proton) numbers.

For example, the extreme stability of helium-4 due to a double pairing of 2 protons and 2 neutrons prevents any nuclides containing five or eight nucleons from existing for long enough to serve as platforms for the buildup of heavier elements via nuclear fusion in stars (see triple alpha process).

[edit]Even proton-even neutron

There are 148 stable even-even nuclides, forming 58% of the 255 stable nuclides. There are also 21 primordial long-lived even-even nuclides. As a result, many of the 41 even-numbered elements from 2 to 82 have many primordial isotopes. Half of these even-numbered elements have six or more stable isotopes.

All even-even nuclides have spin 0 in their ground state.

[edit]Odd proton-odd neutron

Only five stable nuclides contain both an odd number of protons and an odd number of neutrons. The first four "odd-odd" nuclides occur in low mass nuclides, for which changing a proton to a neutron or vice versa would lead to a very lopsided proton-neutron ratio (21H, 6

3Li, 105B, and 147N; spins 1, 1, 3, 1). The only other entirely "stable" odd-odd nuclide is 180m73Ta (spin 9), the only primordial nuclear isomer, which has not yet been observed to decay despite experimental attempts.[16] Also, four long-lived radioactive odd-odd nuclides (4019K, 5023V,13857La,17671Lu; spins 4, 6, 5, 7) occur naturally. As in the case of 180m73Ta decay of high spin nuclides by beta decay (including electron capture), gamma decay, or internal conversion is greatly inhibited if the only decay possible between isobar nuclides (or in the case of 180m73Ta between nuclear isomers of the same nuclide) involves high multiples of a change in spin of 1 unit, the "preferred" change of spin that is associated with rapid decay. This high-spin inhibition of decay is the cause of the five heavy stable or long-lived odd-proton, odd-neutron nuclides discussed above. For an example of this effect where the spin effect is subtracted, tantalum-180, the odd-odd low-spin (theoretical) decay product of primordial tantalum-180m, itself has a half life of only about 8 hours.

Many odd-odd radionuclides (like tantalum-180) with comparatively short half lives are known. Almost invariably, these decay by positive or negative beta decay, in order to produce stable even-even isotopes which have paired protons and paired neutrons. In some odd-odd radionuclides where the ratio of neither protons or neutrons is "excessive" (i.e., falls too far from the ratio of maximal stability), this decay can procede in either direction, turning a proton into a neutron, or vice versa. An example is 6429Cu, which can decay either by positron emission to 6428Ni, or by electron emission to 6430Zn.

Of the nine primordial odd-odd nuclides (five stable and four radioactive with long half lives), only 147N is the most common isotope of a common element. This is the case because it is a part of the CNO cycle. The nuclides 63Li and 105B are minority isotopes of elements that are themselves rare compared to other light elements, while the other six isotopes make up only a tiny percentage of the natural abundnace of their elements. For example, 180m73Ta is thought to be the rarest of the 253 stable isotopes.

None of the primordial (i.e., stable or nearly stable) odd-odd nuclides have spin 0 in the ground state. This is because the single unpaired neutron and unpaired proton have a larger nuclear forceattraction to each other if their spins are alligned (producing a total spin of at least 1 unit), instead of anti-alligned. See deuterium for the simplest case of this nuclear behavior.

[edit]Odd mass number

For a given odd mass number, there can be only a single beta-stable nuclide, since there is not a difference in binding energy between even-odd and odd-even comparable to that between even-even and odd-odd, leaving other nuclides of the same mass number (isobars) free to beta decay towards the lowest-mass one. For 5, 147, 151, and 209+, the beta-stable isobar of that mass number can alpha decay. (In theory, mass number 143 to 155, 160 to 162, and 165+ can also alpha decay). This gives a total of 101 stable nuclides with odd mass numbers. There are another 18 radioactive primordial nuclides (which by definition all have relatively long half lives, greater than 80 million years).

Odd-mass-number nuclides are fermions, i.e. have half-integer spin. Generally speaking, since odd-mass-number nuclides always have an even number of either neutrons or protons, the even-numbered particles usually form part of a "core" in the nucleus with a spin of zero. The nucleon with the odd number (whether protons or neutrons) then form a second core with nucleons paired-off, with most of the nuclear spin due to the orbital angular momentum and spin angular momentum of the last remaining nucleon. 29 of the 117 primordial odd-mass nuclides have spin 1/2, 30 have spin 3/2, 24 have spin 5/2, 17 have spin 7/2, and 9 have spin 9/2.[17]

The odd-mass number stable nuclides are divided (roughly evenly) into odd-proton-even-neutron, and odd-neutron-even-proton nuclides, which are more thoroughly discussed below.

[edit]Odd proton-even neutron

These 48 stable nuclides, stabalized by their even numbers of paired neutrons, form most of the stable isotopes of the odd-numbered elements; the very few odd-odd nuclides comprise the others. There are 41 odd-numbered elements with Z = 1 through 81, of which 32 have one stable odd-even isotope, the elements technetium (43Tc) and promethium (61Pm) have no stable isotopes, and chlorine (17Cl), potassium (19K), copper (29Cu), gallium (31Ga), bromine (35Br), silver (47Ag), antimony (51Sb), iridium (Ir), and thallium (81Tl), have two each, making a total of 48 stable odd-even isotopes. There are also five primordial long-lived radioactive odd-even isotopes, 8737Rb, 11549In, 15163Eu, 18775Re, and 20983Bi which was recently found to decay.[edit]Even proton-odd neutron

Even-odd long-lived

Decay Half-life

11348Cd

beta7.7×1015 a

14762Sm

beta 1.06×1011 a

23592U

alpha 7.04×108 a

These 53 stable (and also 3 primordial long-lived nuclides, including the fissile 23592U) have an even number of protons and an odd number of neutrons. By definition, they are all isotopes of even-Z elements, where they are a minority in comparison to the even-even isotopes which are about 3 times as numerous.

Because of their odd neutron numbers, these nuclides tend to have large neutron capture cross sections, due to the energy that results from neutron-pairing effects.

These stable even-proton odd-neutron nuclides tend to be uncommon by abundance in nature, generally because in order to form and be enter into primordial abundance, they must have escaped capturing neutrons to form yet other stable even-even isotopes, during both the s-process and r-process of neutron capture, during nucleosynthesis in stars. For this reason, only 19578Pt and 94Be are the most naturally abundant isotopes of their element, the former only by a small margin, and the latter only because the expected beryllium-8 has lower binding energy than two alpha particles and therefore alpha decays.[edit]Odd neutron number

Neutron number parity(H-1 with no neutrons excluded)

N Even Odd

Stable 196 58

Long-lived 26 7

All primordial 222 65

Actinides with odd neutron number are generally fissile (with thermal neutrons), while those with even neutron number are generally not, though they are fissionablewith fast neutrons. Only 19578Pt, 94Be and 147N have odd neutron number and are the most naturally abundant isotope of their element.[edit]Occurrence in nature

See also: Abundance of the chemical elements

Elements are composed of one or more naturally occurring isotopes. The unstable (radioactive) isotopes are either primordial or postprimordial. Primordial isotopes were a product of stellar nucleosynthesis or another type of nucleosynthesis such as cosmic ray spallation, and have persisted down to the present because their rate of decay is so slow (e.g., uranium-238 and potassium-40). Postprimordial isotopes were created by cosmic ray bombardment as cosmogenic nuclides (e.g.,tritium, carbon-14), or by the decay of a radioactive primordial isotope to a

radioactive radiogenic nuclide daughter (e.g., uranium to radium). A few isotopes also continue to be naturally synthesized as nucleogenic nuclides, by some other natural nuclear reaction, such as when neutrons from natural nuclear fission are absorbed by another atom.

As discussed above, only 80 elements have any stable isotopes, and 26 of these have only one stable isotope. Thus, about two thirds of stable elements occur naturally on Earth in multiple stable isotopes, with the largest number of stable isotopes for an element being ten, for tin (50Sn). There are about 94 elements found naturally on Earth (up to plutonium inclusive), though some are detected only in very tiny amounts, such as plutonium-244. Scientists estimate that the elements that occur naturally on Earth (some only as radioisotopes) occur as 339 isotopes (nuclides) in total.[18] Only 255 of these naturally occurring isotopes are stable in the sense of never having been observed to decay as of the present time An additional 33 primordial nuclides (to a total of 288 primordial nuclides), are radioactive with known half-lives, but have half-lives longer than 80 million years, allowing them to exist from the beginning of the solar system. See list of nuclides for details.

All the known stable isotopes occur naturally on Earth; the other naturally occurring-isotopes are radioactive but occur on Earth due to their relatively long half-lives, or else due to other means of ongoing natural production. These include the afore-mentioned cosmogenic nuclides, the nucleogenic nuclides, and any radiogenic radioisotopes formed by ongoing decay of a primordial radioactive isotope, such as radon and radium from uranium.

An additional ~3000 radioactive isotopes not found in nature have been created in nuclear reactors and in particle accelerators. Many short-lived isotopes not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae. An example is aluminum-26, which is not naturally found on Earth, but which is found in abundance on an astronomical scale.

The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses. Before the discovery of isotopes, empirically determined noninteger values of atomic mass confounded scientists. For example, a sample of chlorine contains 75.8% chlorine-35 and 24.2% chlorine-37, giving an average atomic mass of 35.5 atomic mass units.

According to generally accepted cosmology theory, only isotopes of hydrogen and helium, traces of some isotopes of lithium and beryllium, and perhaps some boron, were created at the Big Bang, while all other isotopes were synthesized later, in stars and supernovae, and in interactions between energetic particles such as cosmic rays, and previously produced isotopes. (Seenucleosynthesis for details of the various processes thought to be responsible for isotope production.) The respective abundances of isotopes on Earth result from the quantities formed by these processes, their spread through the galaxy, and the rates of decay for isotopes that are unstable. After the initial coalescence of the solar system, isotopes were redistributed according to mass, and the isotopic composition of elements varies slightly from planet to planet. This sometimes makes it possible to trace the origin of meteorites.

[edit]Atomic mass of isotopes

The atomic mass (mr) of an isotope is determined mainly by its mass number (i.e. number of nucleons in its nucleus). Small corrections are due to the binding energy of the nucleus (see mass defect), the slight difference in mass between proton and neutron, and the mass of the electrons associated with the atom, the latter because the electron:nucleon ratio differs among isotopes.

The mass number is a dimensionless quantity. The atomic mass, on the other hand, is measured using the atomic mass unit based on the mass of the carbon-12 atom. It is denoted with symbols "u" (for unit) or "Da" (for Dalton).

The atomic masses of naturally occurring isotopes of an element determine the atomic mass of the element. When the element contains N isotopes, the equation below is applied for the atomic mass M:

where m1, m2, ..., mN are the atomic masses of each individual isotope, and x1, ..., xN are the relative abundances of these isotopes.

[edit]Applications of isotopes

[edit]Purification

Main article: isotope separation

Several applications exist that capitalize on properties of the various isotopes of a given element. Isotope separation is a significant technological challenge, particularly with heavy elements such as uranium or plutonium. Lighter elements such as lithium, carbon, nitrogen, and oxygen are commonly separated by gas diffusion of their compounds such as CO and NO. The separation of hydrogen and deuterium is unusual since it is based on chemical rather than physical properties, for example in the Girdler sulfide process. Uranium isotopes have been separated in bulk by gas diffusion, gas centrifugation, laser ionization separation, and (in the Manhattan Project) by a type of production mass spectrometry.

[edit]Use of chemical and biological properties

Main articles: isotope geochemistry, cosmochemistry, and paleoclimatology

Isotope analysis  is the determination of isotopic signature, the relative abundances of isotopes of a given element in a particular sample. For biogenic substances in particular, significant variations of isotopes of C, N and O can occur. Analysis of such variations has a wide range of applications, such as the detection of adulteration of food products[19] or the geographic origins of products using isoscapes. The identification of certain meteorites as having originated on Mars is based in part upon the isotopic signature of trace gases contained in them.[20]

Isotopic substitution can be used to determine the mechanism of a chemical reaction via the kinetic isotope effect. Another common application is isotopic labeling, the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each

other. However, by using isotopes of different masses, even different nonradioactive stable isotopes can be distinguished by mass spectrometry or infrared spectroscopy. For example, in 'stable isotope labeling with amino acids in cell culture (SILAC)' stable isotopes are used to quantify proteins. If radioactive isotopes are used, they can be detected by the radiation they emit (this is called radioisotopic labeling).

[edit]Use of nuclear properties

A technique similar to radioisotopic labeling is radiometric dating: using the known half-life of an unstable element, one can calculate the amount of time that has elapsed since a known level of isotope existed. The most widely known example is radiocarbon dating used to determine the age of carbonaceous materials.

Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes, both radioactive and stable. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are 1H, 2D,15N, 13C, and 31P.

Mössbauer spectroscopy  also relies on the nuclear transitions of specific isotopes, such as 57Fe. Radionuclides  also have important uses. Nuclear power and nuclear weapons development require relatively large quantities of specific isotopes. Nuclear medicine and radiation

oncologyutilize radioisotopes respectively for medical diagnosis and treatment.

List of isotopes

From Wikipedia, the free encyclopedia

A table of chemical elements ordered by atomic number and color coded according to type of element. Given is each element's name, element symbol, atomic mass (or most stable isotope),

and in some cases a list of isotopes with their own article.

The name of the element links to an article about isotopes of that element, and the symbol links to the main article about the element.

Some element categories in the periodic table

Metals

Metalloids

NonmetalsUnknownchemicalproperties

Alkalimetals

Alkalineearth metals

Inner transition metalsTransitionmetals

Post-transitionmetals

Othernonmetals

HalogensNoblegasesLanthanides Actinides

Z Name Sym Mass (g/mol) Selected isotopes

1 Hydrogen H 1.00794(7)[1] [2] [3] Hydrogen-1, Deuterium (Hydrogen-2), Tritium (Hydrogen-3), Hydrogen-4, Hydrogen-5

2 Helium He 4.002602(2)[1] [3] Diproton (hypothetical Helium-2 nucleus), Helium-3, Helium-4

3 Lithium Li 6.941(2)[1] [2] [3] [4] Lithium-6 , Lithium-7

4 Beryllium Be 9.012182(3) Beryllium-8 , Beryllium-10

5 Boron B 10.811(7)[1] [2] [3]

6 Carbon C 12.0107(8)[1] [3] Carbon-11, Carbon-12, Carbon-13, Carbon-14

7 Nitrogen N 14.0067(2)[1] [3] Nitrogen-13, Nitrogen-14, Nitrogen-15

8 Oxygen O 15.9994(3)[1] [3] Oxygen-13, Oxygen-15, Oxygen-16, Oxygen-17, Oxygen-18

Z Name Sym Mass (g/mol) Selected isotopes

9 Fluorine F 18.9984032(5) Fluorine-17, Fluorine-18, Fluorine-19

10 Neon Ne 20.1797(6)[1] [2] Neon-20, Neon-21, Neon-22

11 Sodium Na 22.98976928(2) Sodium-22

12 Magnesium Mg 24.3050(6) Magnesium-23, Magnesium-24, Magnesium-25, Magnesium-26

13 Aluminium Al 26.9815386(8) Aluminium-26

14 Silicon Si 28.0855(3)[3] Silicon-28, Silicon-29, Silicon-30

15 Phosphorus P 30.973762(2) Phosphorus-30, Phosphorus-31, Phosphorus-32, Phosphorus-33

16 Sulfur S 32.065(5)[1] [3]

17 Chlorine Cl 35.453(2)[1] [2] [3] Chlorine-35, Chlorine-37

Z Name Sym Mass (g/mol) Selected isotopes

18 Argon Ar 39.948(1)[1] [3] Argon-36, Argon-38, Argon-40

19 Potassium K 39.0983(1)

20 Calcium Ca 40.078(4)[1] Calcium-48

21 Scandium Sc 44.955912(6)

22 Titanium Ti 47.867(1) Titanium-44

23 Vanadium V 50.9415(1) Vanadium-49, Vanadium-50

24 Chromium Cr 51.9961(6)

25 Manganese Mn 54.938045(5)

26 Iron Fe 55.845(2) Iron-52, Iron-54, Iron-55, Iron-56, Iron-57, Iron-58, Iron-60

27 Cobalt Co 58.933195(5) Cobalt-56, Cobalt-57, Cobalt-59, Cobalt-60

Z Name Sym Mass (g/mol) Selected isotopes

28 Nickel Ni 58.6934(2) Nickel-62

29 Copper Cu 63.546(3)[3] Copper-64

30 Zinc Zn 65.409(4)

31 Gallium Ga 69.723(1)

32 Germanium Ge 72.64(1)

33 Arsenic As 74.92160(2)

34 Selenium Se 78.96(3)[3] Selenium-72, Selenium-76, Selenium-78, Selenium-79, Selenium-82,

35 Bromine Br 79.904(1)

36 Krypton Kr 83.798(2)[1] [2] Krypton-85

37 Rubidium Rb 85.4678(3)[1]

Z Name Sym Mass (g/mol) Selected isotopes

38 Strontium Sr 87.62(1)[1] [3] Strontium-90

39 Yttrium Y 88.90585(2)

40 Zirconium Zr 91.224(2)[1] Zirconium-93, Zirconium-96

41 Niobium Nb 92.906 38(2)

42 Molybdenum Mo 95.94(2)[1]

43 Technetium Tc [98.9063][5] Technetium-99, Technetium-99m

44 Ruthenium Ru 101.07(2)[1] Ruthenium-106

45 Rhodium Rh 102.90550(2)

46 Palladium Pd 106.42(1)[1] Palladium-102, Palladium-103, Palladium-107

47 Silver Ag 107.8682(2)[1]

Z Name Sym Mass (g/mol) Selected isotopes

48 Cadmium Cd 112.411(8)[1] Cadmium-113m

49 Indium In 114.818(3)

50 Tin Sn 118.710(7)[1] Tin-121m, Tin-126

51 Antimony Sb 121.760(1)[1]

52 Tellurium Te 127.60(3)[1] Tellurium-124

53 Iodine I 126.90447(3) Iodine-123, Iodine-124, Iodine-125, Iodine-129, Iodine-131, Iodine-135

54 Xenon Xe 131.293(6)[1] [2] Xenon-124, Xenon-125, Xenon-133, Xenon-135, Xenon-136

55 Caesium Cs 132.9054519(2) Caesium-134, Caesium-135, Caesium-137

56 Barium Ba 137.327(7)

57 Lanthanum La 138.90547(7)[1]

Z Name Sym Mass (g/mol) Selected isotopes

58 Cerium Ce 140.116(1)[1]

59 Praseodymium Pr 140.90765(2)

60 Neodymium Nd 144.242(3)[1]

61 Promethium Pm [146.9151][5] Promethium-147

62 Samarium Sm 150.36(2)[1] Samarium-149, Samarium-151

63 Europium Eu 151.964(1)[1] Europium-155

64 Gadolinium Gd 157.25(3)[1]

65 Terbium Tb 158.92535(2)

66 Dysprosium Dy 162.500(1)[1]

67 Holmium Ho 164.93032(2)

Z Name Sym Mass (g/mol) Selected isotopes

68 Erbium Er 167.259(3)[1]

69 Thulium Tm 168.93421(2)

70 Ytterbium Yb 173.04(3)[1]

71 Lutetium Lu 174.967(1)[1]

72 Hafnium Hf 178.49(2)

73 Tantalum Ta 180.9479(1)

74 Tungsten W 183.84(1)

75 Rhenium Re 186.207(1)

76 Osmium Os 190.23(3)[1]

77 Iridium Ir 192.217(3) Iridium-192

Z Name Sym Mass (g/mol) Selected isotopes

78 Platinum Pt 195.084(9)

79 Gold Au 196.966569(4)

80 Mercury Hg 200.59(2)

81 Thallium Tl 204.3833(2) Thallium-205, Thallium-208

82 Lead Pb 207.2(1)[1] [3] Lead-206, Lead-208, Lead-209, Lead-214,

83 Bismuth Bi 208.98040(1) Bismuth-209

84 Polonium Po [208.9824][5]

85 Astatine At [209.9871][5]

86 Radon Rn [222.0176][5]

87 Francium Fr [223.0197][5]

Z Name Sym Mass (g/mol) Selected isotopes

88 Radium Ra [226.0254][5]

89 Actinium Ac [227.0278][5]

90 Thorium Th 232.03806(2)[5] [1] Thorium-228, Thorium-229, Thorium-230, Thorium-231, Thorium-232, Thorium-233, Thorium-234

91 Protactinium Pa 231.03588(2)[5] Protactinium-230

92 Uranium U 238.02891(3)[5] [1] [2] Uranium-230, Uranium-231, Uranium-232, Uranium-233, Uranium-234, Uranium-235, Uranium-236, Uranium-238, Uranium-239

93 Neptunium Np [237.0482][5] Neptunium-235, Neptunium-236

94 Plutonium Pu [244.0642][5] Plutonium-238, Plutonium-239, Plutonium-240, Plutonium-241, Plutonium-242, Plutonium-244

95 Americium Am [243.0614][5] Americium-241, Americium-242, Americium-243

96 Curium Cm [247.0703][5]

97 Berkelium Bk [247.0703][5]

Z Name Sym Mass (g/mol) Selected isotopes

98 Californium Cf [251.0796][5]

99 Einsteinium Es [252.0829][5]

100 Fermium Fm [257.0951][5]

101 Mendelevium Md [258.0986][5]

102 Nobelium No [259.1009][5]

103 Lawrencium Lr [262.11][5]

104 Rutherfordium Rf [265.12][5]

105 Dubnium Db [268.13][5]

106 Seaborgium Sg [271.13][5]

107 Bohrium Bh [270][5]

Z Name Sym Mass (g/mol) Selected isotopes

108 Hassium Hs [277.15][5]

109 Meitnerium Mt [276.15][5]

110 Darmstadtium Ds [281.16][5]

111 Roentgenium Rg [280.16][5]

112 Copernicium Cn [285.17][5]

113 Ununtrium Uut [284.18][5]

114 Flerovium Fl [289.19][5]

115 Ununpentium Uup [288.19][5]

116 Livermorium Lv [293][5]

117 Ununseptium Uus [294][5]

Z Name Sym Mass (g/mol) Selected isotopes

118 Ununoctium Uuo [294][5]

Some element categories in the periodic table

Metals

Metalloids

NonmetalsUnknownchemicalproperties

Alkalimetals

Alkalineearth metals

Inner transition metalsTransitionmetals

Post-transitionmetals

Othernonmetals

HalogensNoblegasesLanthanides Actinides

An isotope is a form of a chemical element whose atomic nucleus contains a specific number of neutron s, in addition to the number of proton s that uniquely defines the element. The nuclei of most atom s contain neutrons as well as protons. (An exception is the common form of hydrogen, whose nucleus consists of a lone proton.) Every chemical element has more than one isotope. For any element, one of the isotopes is more abundant in nature than any of the others, although often multiple isotopes of a single element are mixed.

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atomic mass unit (AMU or amu)

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The isotope of an element is defined by the nucleon number, which is the sum of the number of protons and the number of neutrons in the atomic nucleus. The nucleon number is customarily written as a superscript preceding the chemical symbol for the element. For example, 16 O represents oxygen-16, which has 8 protons and 8 neutrons, while 12 C represents carbon-12, with 6 protons and 6 neutrons. These are the most common naturally occurring isotopes of oxygen and carbon, respectively. Some carbon-14 is found in nature. An atom of carbon-14 contains 6 protons and 8 neutrons and is denoted 14 C. Over time, 14 C decays into 12 C.

Sometimes the isotope of an element is denoted by writing the nucleon number after the chemical symbol, and not as a superscript. Thus, some texts will denote carbon-14 as C-14 or C14 instead of 14 C.

Certain isotopes of elements are unstable, giving off ionizing radiation, also known as radioactivity. Such an isotope is called a radioisotope. Carbon-14 is a radioisotope of carbon. In the case of an element that is radioactive in all its known forms, such as uranium (U), certain isotopes are more radioactive than others, and/or give off different proportions of the various types of ionizing radiation.

What Are Isotopes?Isotopes are atoms of the same element that have different numbers of neutrons in the nucleus and therefore different atomic weights (shown as a superscript numeral in front of the element's symbol)

Isotopes may be stable or radioactive. Most elements have two or more isotopes. The isotopes most commonly used in geochemistry include: carbon (12C,13C,14C), hydrogen (1H, 2H, 3H), oxygen (16O,18O), sulfur (32S, 34S) and nitrogen (14N, 15N).

The elements listed above are important constituents in biological systems and are also involved in many geochemical reactions. Small differences in the concentration of isotopes exist in chemically identical compounds because of differences in the origin or certain processes that have occurred after the compounds was produced. These characteristics make isotope analyses very useful for determining the source of certain compounds in the environment and/or determining the geochemical reactions that have affected the concentration of the compounds or materials of interest. The radioactive isotopes are often used to determine the age of different types of materials.

IsotopesIsotope is the word we use to talk about two different forms of the same element that have different numbers of neutrons in the nucleus. A given element is identified by the numer of protons in its nucleus; that's its atomic number. And two different isotopes of a given element have the same number of protons (of course), but each has a different number of neutrons in its nucleus.

In chemistry, elements differ from one another by the number of protons in their nucleus. In their neutral state, elements also have the same number of electrons as protons. But the number of neutrons can vary from one atom of a given element to the next atom of that same element. Said another way, one element may exists in a number of different forms called isotopes. Each element has a number of different isotopes, and this is determined by the number of neutrons in the nucleus. Let's look at an example.

ExamplesIf we consider the case of oxygen, all atoms of this element have 8 protons. But the isotope oxygen-15 has 7 neutrons, while the isotope oxygen-16 (the most common isotope of oxygen) has 8 neutrons. The isotope oxygen-17 has 9 neutrons. You probably already noticed that the number of the isotope equals the atomic number of the element plus the number of neutrons in the nucleus of that isotope. This number is also called the mass number, the sum of the atomic number and number of neutrons. Different isotopes of the one element are almost physically and chemically identical. Let's look at one more.

When we investigate the metal uranium, we sometimes recall that there are a pair of isotopes that are important. They are uranium-235 and uranium-238. Sure, there are other isotopes, but U-238 is the most common (over 99% of natural abundance). We also know of U-235 because it is the isotope that is fissile, and the one we use to fuel most nuclear reactors. Both isotopes are uranium, and both have 92 protons in their nucleus. But there are 146 neutrons in an atom of U-238, and there are 143 neutrons in an atom of U-235.

Now when you see something like O-16 or U-235 or any other chemical symbol followed by a number, you'll know exactly what is being specified by this form of scientific expression. Additionally, mass number may be indicated by the superscript before the symbol, such as 16O.

IsotopesAtoms of the same element can have different numbers of neutrons; the different possible versions of each element are called isotopes. For example, the most common isotope of hydrogen has no neutrons at all; there's also a hydrogen isotope called deuterium, with one neutron, and another, tritium, with two neutrons.

Hydrogen Deuterium Tritium

If you want to refer to a certain isotope, you write it like this: AXZ. Here X is the chemical symbol for the element, Z is the atomic number, and A is the number of neutrons and protons combined, called the mass number. For instance, ordinary hydrogen is written 1H1, deuterium is2H1, and tritium is 3H1. 

No; there are "preferred" combinations of neutrons and protons, at which the forces holding nuclei together seem to balance best. Light elements tend to have about as many neutrons as protons; heavy elements apparently need more neutrons than protons in order to stick together. Atoms with a few too many neutrons, or not quite enough, can sometimes exist for a while, but they'reunstable. 

I'm not sure what you mean by "unstable." Do atoms just fall apart if they don't have the right number of neutrons? 

Well, yes, in a way. Unstable atoms areradioactive: their nuclei change or decay by spitting out radiation, in the form of particles or electromagnetic waves. 

Beta Decay

Does it just kick out one of the neutrons? 

No, it can't do that; the neutrons are stuck too firmly where they are. What it can do...well, I'll let you see for yourself. In the applet, click on the button labeled H3 (for hydrogen 3, or tritium). 

The neutron turns into a proton! 3H1 becomes 3He2. 

Right. An unstable isotope of hydrogen has converted itself into a stable isotope of helium. You'll notice that 3H1 and 3He2 have the same mass number, which is good, because mass has to be conserved.

There is a problem, though. Electric charge also has to be conserved.

Hydrogen has only one proton, and helium has two, so you'd end up with twice as much positive charge as you started with. How do you get around that? 

When 3H metamorphoses into helium 3, it also gives off an electron--which has hardly any mass, and is endowed with a negative charge that exactly cancels one proton. This process is known as beta decay, and the electron is called a beta particle in this context.

You can write out the nuclear reaction involved in the beta decay of tritium by giving the electron a "mass number" of 0 and an "atomic number" of -1:

3H1 => 3He2 + 0e-1

Notice that the mass numbers on each side add up to the same total (3 = 3 + 0), and so do the charges (1 = 2 + -1). This must always be true in any nuclear reaction.

Positrons, Alpha Particles, and Gamma RaysWhat happens when an atom doesn't have enoughneutrons to be stable?

That's the case with beryllium 7, 7Be4. Click on it in the applet and see what happens. 

It decays to lithium 7--so a proton turns into a neutron. That makes sense...but how do you deal with the electric charge problem now? Going from Be to Li, you lose charge; emitting an electron would just make things worse. 

No, there are other possibilities. Some heavy isotopes decay by spitting out alpha particles. These are actually helium 4 nuclei--clumps of two neutrons and two protons each. A typical alpha decay looks like this:

238U92 => 234Th90 + 4He2

There's also a third type of radioactive emission. After alpha or beta decay, a nucleus is often left in an excited state--that is, with some extra energy. It then "calms down" by releasing this energy in the form of a very high-frequency photon, or electromagnetic wave, known as a gamma ray.Click on the advanced button for more information about why this happens.

HalflifeThe applet lists a "halflife" for each radioactive isotope. What does that mean?

The halflife is the amount of time it takes for half of the atoms in a sample to decay. The halflife for a given isotope is always the same   ; it doesn't depend on how many atoms you have or on how long they've been sitting around.

For example, the applet will tell you that the halflife of beryllium 11 is 13.81 seconds. Let's say you start with, oh, 16 grams of 11Be. Wait 13.81 seconds, and you'll have 8 grams left; the rest will have decayed to boron 11. Another 13.81 seconds go by, and you're left with 4 grams of 11Be; 13.81 seconds more, and you have 2 grams...you get the idea. 

Hmmm...so a lot of decays happen really fast when there are lots of atoms, and then things slow down when there aren't so many. The halflife is always the same, but the half gets smaller and smaller. 

Notice how the decays are fast and furious at the beginning and slow down over time; you can see this both from the color changes in the top window and from the graph.

You'll also notice that the pattern of atoms in the top picture is random-looking, and different each time you run the applet, but the graph below always has the same shape. It's impossible to predict when a specific atom is going to decay, but you can predict the number of atoms that will decay in a certain time period. 

Radioisotopes: What Are They and How Are They Made?

What are isotopes?The isotopes of an element are all the atoms that have in their nucleus the number of protons (atomic number) corresponding to the chemical behavior of that element. However, the isotopes of a single element vary in the number of neutrons in their nuclei. Since they still have the same number of protons, all these isotopes of an element have identical chemical behavior. But since they have different numbers of neutrons, these isotopes of the same element may have different radioactivity. An isotope that is radioactive is called a radioisotope or radionuclide. Two examples may help clarify this.

The most stable isotope of uranium, U-238, has an atomic number of 92 (protons) and an atomic weight of 238 (92 protons plus 146 neutrons). The isotope of uranium of greatest importance in atomic bombs, U-235, though, has three fewer neutrons. Thus, it also has an atomic number of 92 (since the number of protons has not changed) but an atomic weight of 235 (92 protons plus only 143 neutrons). The chemical behavior of U-235 is identical to all other forms of uranium, but its nucleus is less stable, giving it higher radioactivity and greater susceptibility to the chain reactions that power both atomic bombs and nuclear fission reactors.

Another example is iodine, an element essential for health; insufficient iodine in one's diet can lead to a goiter. Iodine also is one of the earliest elements whose radioisotopes were used in what is now called nuclear medicine. The most common, stable form of iodine has an atomic number of 53 (protons) and an atomic weight of 127 (53 protons plus 74 neutrons). Because its nucleus has the "correct" number of neutrons, it is stable and is not radioactive. A less stable form of iodine also has 53 protons (this is what makes it behave chemically as iodine) but four extra neutrons, for a total atomic weight of 131 (53 protons and 78 neutrons). With "too many" neutrons in its nucleus, it is unstable and radioactive, with a half-life of eight days. Because it behaves chemically as iodine, it travels throughout the body and localizes in the thyroid gland just like the stable form of iodine. But, because it is radioactive, its presence can be detected. Iodine 131 thus became one of the earliest radioactive tracers.

How can different isotopes of an element be produced?How can isotopes be produced--especially radioisotopes, which can serve many useful purposes? There are two basic methods: separation and synthesis.

Some isotopes occur in nature. If radioactive, these usually are radioisotopes with very long half-lives. Uranium 235, for example, makes up about 0.7 percent of the naturally occurring uranium on the earth.[89]The challenge is to separate this very small amount from the much larger bulk of other forms of uranium. The difficulty is that all these forms of uranium, because they all have the same number of electrons, will have identical chemical behavior: they will bind in identical fashion to other atoms. Chemical separation, developing a chemical reaction that will bind only uranium atoms, will separate out uranium atoms, but not distinguish among different isotopes of uranium. The only difference among the uranium isotopes is their atomic weight. A method had to be developed that would sort atoms according to weight.

One initial proposal was to use a centrifuge. The basic idea is simple: spin the uranium atoms as if they were on a very fast merry-go-round. The heavier ones will drift toward the outside faster and can be drawn off. In practice the technique was an enormous challenge: the goal was to draw off that very small portion of uranium atoms that were lighter than their brethren. The difficulties were so enormous the plan was abandoned in 1942.[90] Instead, the technique of gaseous diffusion was developed. Again, the basic idea was very simple: the rate at which gas passed (diffused) through a filter depended on the weight of the gas molecules: lighter molecules diffused more quickly. Gas molecules that contained U-235 would diffuse slightly faster than gas molecules containing the more common but also heavier U-238. This method also presented formidable technical challenges, but was eventually implemented in the gigantic gas diffusion plant at Oak Ridge, Tennessee. In this process, the uranium was chemically combined with fluorine to form a hexafluoride gas prior to separation by diffusion. This is not a practical method for extracting radioisotopes for scientific and medical use. It was extremely expensive and could only supply naturally occurring isotopes.

A more efficient approach is to artificially manufacture radioisotopes. This can be done by firing high-speed particles into the nucleus of an atom. When struck, the nucleus may absorb the particle or become unstable and emit a particle. In either case, the number of particles in the nucleus would be altered, creating an isotope. One source of high-speed particles could be a cyclotron. A cyclotron accelerates particles around a circular race track with periodic pushes of an electric field. The particles gather speed with each push, just as a child swings higher with each push on a swing. When traveling fast enough, the particles are directed off the race track and into the target.

A cyclotron works only with charged particles, however. Another source of bullets are the neutrons already shooting about inside a nuclear reactor. The neutrons normally strike the nuclei of the fuel, making them unstable and causing the nuclei to split (fission) into two large fragments and two to three "free" neutrons. These free neutrons in turn make additional nuclei unstable, causing further fission. The result is a chain reaction. Too many neutrons can lead to an uncontrolled chain reaction, releasing too much heat and perhaps causing a "meltdown." Therefore, "surplus" neutrons are usually absorbed by "control rods." However, these surplus neutrons can also be absorbed by targets of carefully selected material placed in the reactor. In this way the surplus neutrons are used to create radioactive isotopes of the materials placed in the targets.

With practice, scientists using both cyclotrons and reactors have learned the proper mix of target atoms and shooting particles to "cook up" a wide variety of useful radioisotopes.

RADIOACTIVE HALF-LIFE (CONTINUED)

After this reading this section you will be able to do the following:

Describe carbon dating and how half-life information is used. Explain how a radiographer uses half-life information.

As we have mentioned before each radioactive isotope has its own decay pattern. Not only does it decay by giving off energy and matter, but it also decays at a rate that is characteristic to itself. The rate at which a radioactive isotope decays is measured in half-life. The term half-life is defined as the time it takes for one-half of the atoms of a radioactive material to disintegrate. Half-lives for various radioisotopes can range from a few microseconds to billions of years. See the table below for a list of radioisotopes and each of unique their half-lives.

Radioisotope Half-life

Polonium-215 0.0018 seconds

Bismuth-212 60.5 seconds

Sodium-24 15 hours

Iodine-131 8.07 days

Cobalt-60 5.26 years

Radium-226 1600 years

Uranium-238 4.5 billion years

How does the half-life affect an isotope?

Let's look closely at how the half-life affects an isotope. Suppose you have 10 grams of Barium-139. It has a half-life of 86 minutes. After 86 minutes, half of the atoms in the sample would have decayed into another element, Lanthanum-139. Therefore, after one half-life, you would have 5 grams of Barium-139, and 5 grams of Lanthanum-139. After another 86 minutes, half of the 5 grams of Barium-139 would decay into Lanthanum-139; you would now have 2.5 grams of Barium-139 and 7.5 grams of Lanthanum-139.

How is half-life information used in carbon dating?

The half-lives of certain types of radioisotopes are very useful to know. They allow us to determine the ages of very old artifacts. Scientists can use the half-life of Carbon-14 to determine the approximate age of organic objects less than 40,000 years old. By determining how much of the carbon-14 has transmutated, scientist can calculate and estimate the age of a substance. This technique is known as Carbon dating. Isotopes with longer half-lives such as Uranium-238 can be used to date even older objects.

You will learn more about carbon dating in the next sub-unit.

Uses of the half-life in NDT

In the field of nondestructive testing radiographers (people who produce radiographs to inspect objects) also use half-life information. A radiographer who works with radioisotopes needs to know the specific half-life to properly determine how much radiation the source in the camera is producing so that the film can be exposed properly. After one half-life of a given radioisotope, only one half as much of the original number of atoms remains active. Another way to look at this is that if the radiation intensity is cut in half; the source will have only half as many curies as it originally had. It is important to recognize that the intensity or amount of radiation is decreasing due to age but not the penetrating energy of the radiation. The energy of the radiation for a given isotope is considered to be constant for the life of the isotope.

Review:

1. The half-life of radioisotopes varies from seconds to billions of years.2. Carbon-dating uses the half-life of Carbon-14 to find the approximate age of an object that is 40,000 years old or younger.3. Radiographers use half-life information to make adjustments in the film exposure time due to the changes in radiation intensity that occurs as radioisotopes

degrade.

Half-life

From Wikipedia, the free encyclopedia

This article is about the scientific and mathematical term. For other uses, see Half-life (disambiguation).

This article needs additional citations for verification. Please help improve this article by adding citations to reliable sources. Unsourced material may be challenged and removed. (July 2009)

Half-life (t½) is the time required for a quantity to fall to half its value as measured at the beginning of the time period. In physics, it is typically used to describe

a property of radioactive decay, but may be used to describe any quantity which follows an exponential decay.

The original term, dating to 1907, was "half-life period", which was shortened to "half-life" in the early 1950s.[1]

Half-life is used to describe a quantity undergoing exponential decay, and is constant over the lifetime of the decaying quantity. It is a characteristic unitfor the

exponential decay equation. The term "half-life" may generically be used to refer to any period of time in which a quantity falls by half, even if the decay is not

exponential. For a general introduction and description of exponential decay, see exponential decay. For a general introduction and description of non-

exponential decay, see rate law.

The converse of half-life is doubling time.

The table on the right shows the reduction of a quantity in terms of the number of half-lives elapsed.

Number ofhalf-liveselapsed

Fractionremaining

Percentageremaining

0 1/1 100

1 1/2 50

2 1/4 25

3 1/8 12.5

4 1/16 6.25

5 1/32 3.125

6 1/64 1.563

7 1/128 0.781

... ... ...

n 2-n 100/(2n)

Contents

[hide]

1 Probabilistic nature of half-life

2 Formulas for half-life in exponential decay

o 2.1 Decay by two or more processes

o 2.2 Examples

3 Half-life in non-exponential decay

4 Half-life in biology and pharmacology

5 See also

6 References

7 External links

[edit]Probabilistic nature of half-life

Simulation of many identical atoms undergoing radioactive decay, starting with either 4 atoms per box (left) or 400 (right). The number at the top is how many half-lives have elapsed. Note thelaw

of large numbers: With more atoms, the overall decay is more regular and more predictable.

A half-life usually describes the decay of discrete entities, such as radioactive atoms, which have unstable nuclei. In that case, it does not work to use the definition "half-life is the time required for

exactly half of the entities to decay". For example, if there is just one radioactive atom with a half-life of one second, there will not be "one-half of an atom" left after one second. There will be either

zero atoms left or one atom left, depending on whether or not that atom happened to decay.

Instead, the half-life is defined in terms of probability. It is the time when the expected value of the number of entities that have decayed is equal to half the original number. For example, one can

start with a single radioactive atom, wait its half-life, and then check whether or not it has decayed. Perhaps it did, but perhaps it did not. But if this experiment is repeated again and again, it will be

seen that - on average - it decays within the half-life 50% of the time.

In some experiments (such as the synthesis of a superheavy element), there is in fact only one radioactive atom produced at a time, with its lifetime individually measured. In this case, statistical

analysis is required to infer the half-life. In other cases, a very large number of identical radioactive atoms decay in the measured time range. In this case, the law of large numbers ensures that

the number of atoms that actually decay is approximately equal to the number of atoms that are expected to decay. In other words, with a large enough number of decaying atoms, the probabilistic

aspects of the process could be neglected.

There are various simple exercises that demonstrate probabilistic decay, for example involving flipping coins or running a statistical computer program.[2][3][4] For example, the image on the right is a

simulation of many identical atoms undergoing radioactive decay. Note that after one half-life there are not exactly one-half of the atoms remaining, only approximately, because of the random

variation in the process. However, with more atoms (right boxes), the overall decay is smoother and less random-looking than with fewer atoms (left boxes), in accordance with the law of large

numbers.

[edit]Formulas for half-life in exponential decay

Main article: Exponential decay

An exponential decay process can be described by any of the following three equivalent formulas:

where

N0 is the initial quantity of the substance that will decay (this quantity may be measured in grams, moles, number of atoms, etc.),

N(t) is the quantity that still remains and has not yet decayed after a time t,

t1/2 is the half-life of the decaying quantity,

τ is a positive number called the mean lifetime of the decaying quantity,

λ is a positive number called the decay constant of the decaying quantity.

The three parameters  ,  , and λ are all directly related in the following way:

where ln(2) is the natural logarithm of 2 (approximately 0.693).

[show]Click "show" to see a detailed derivation of the relationship between half-life, decay time, and decay constant.

By plugging in and manipulating these relationships, we get all of the following equivalent descriptions of exponential decay, in terms of the half-life:

Regardless of how it's written, we can plug into the formula to get

 as expected (this is the definition of "initial quantity")

 as expected (this is the definition of half-life)

, i.e. amount approaches zero as t approaches infinity as expected (the longer we wait, the less remains).

[edit]Decay by two or more processes

Some quantities decay by two exponential-decay processes simultaneously. In this case, the actual half-life T1/2 can be related to the half-lives t1 and t2 that the

quantity would have if each of the decay processes acted in isolation:

For three or more processes, the analogous formula is:

For a proof of these formulas, see Decay by two or more processes.

[edit]Examples

Main article: Exponential decay--Applications and examples

There is a half-life describing any exponential-decay process. For example:

The current flowing through an RC circuit or RL circuit decays with a half-life of   or  , respectively. For this example,

the term half time might be used instead of "half life", but they mean the same thing.

In a first-order chemical reaction, the half-life of the reactant is  , where λ is the reaction rate constant.

In radioactive decay, the half-life is the length of time after which there is a 50% chance that an atom will have undergone nuclear decay. It

varies depending on the atom type and isotope, and is usually determined experimentally. See List of nuclides.

[edit]Half-life in non-exponential decay

Main article: Rate equation

The decay of many physical quantities is not exponential—for example, the evaporation of water from a puddle, or (often) the chemical reaction of a

molecule. In such cases, the half-life is defined the same way as before: as the time elapsed before half of the original quantity has decayed.

However, unlike in an exponential decay, the half-life depends on the initial quantity, and the prospective half-life will change over time as the

quantity decays.

As an example, the radioactive decay of carbon-14 is exponential with a half-life of 5730 years. A quantity of carbon-14 will decay to half of its

original amount (on average) after 5730 years, regardless of how big or small the original quantity was. After another 5730 years, one-quarter of the

original will remain. On the other hand, the time it will take a puddle to half-evaporate depends on how deep the puddle is. Perhaps a puddle of a

certain size will evaporate down to half its original volume in one day. But on the second day, there is no reason to expect that one-quarter of the

puddle will remain; in fact, it will probably be much less than that. This is an example where the half-life reduces as time goes on. (In other non-

exponential decays, it can increase instead.)

The decay of a mixture of two or more materials which each decay exponentially, but with different half-lives, is not exponential. Mathematically, the

sum of two exponential functions is not a single exponential function. A common example of such a situation is the waste of nuclear power stations,

which is a mix of substances with vastly different half-lives. Consider a sample containing a rapidly decaying element A, with a half-life of 1 second,

and a slowly decaying element B, with a half-life of one year. After a few seconds, almost all atoms of the element A have decayed after repeated

halving of the initial total number of atoms; but very few of the atoms of element B will have decayed yet as only a tiny fraction of a half-life has

elapsed. Thus, the mixture taken as a whole does not decay by halves.

[edit]Half-life in biology and pharmacology

Main article: Biological half-life

A biological half-life or elimination half-life is the time it takes for a substance (drug, radioactive nuclide, or other) to lose one-half of its

pharmacologic, physiologic, or radiological activity. In a medical context, the half-life may also describe the time that it takes for the concentration

in blood plasma of a substance to reach one-half of its steady-state value (the "plasma half-life").

The relationship between the biological and plasma half-lives of a substance can be complex, due to factors including accumulation in tissues,

active metabolites, and receptor interactions.[5]

While a radioactive isotope decays almost perfectly according to so-called "first order kinetics" where the rate constant is a fixed number, the

elimination of a substance from a living organism usually follows more complex chemical kinetics.

For example, the biological half-life of water in a human being is about seven to 14 days, though this can be altered by his/her behavior. The

biological half-life of cesium in human beings is between one and four months. This can be shortened by feeding the person prussian blue, which

acts as a solid ion exchanger that absorbs the cesium while releasing potassium ions in their place.

Radioactive decay

From Wikipedia, the free encyclopedia

For particle decay in a more general context, see Particle decay. For more information on hazards of various kinds of radiation from decay, see Ionizing radiation.

"Radioactive" redirects here. For other uses, see Radioactive (disambiguation).

This article needs additional citations for verification. Please help improve this article by adding citations to reliable sources. Unsourced material may be challenged and removed. (April 2010)

Alpha decay is one example type ofradioactive decay, in which an atomic nucleus emits an alpha particle, and thereby transforms (or 'decays') into an atom with amass number 4 less and atomic

number 2 less. Many other types of decays are possible.

Nuclear physics

Radioactive decay

Nuclear fission

Nuclear fusion

Classical decays[show]

Advanced decays[show]

Emission processes[show]

Capturing[show]

High energy processes[show]

Nucleosynthesis [show]

Scientists[show]

V

T

E

Radioactive decay is the process by which an atomic nucleus of an unstable atom loses energy by emitting ionizing particles (ionizing radiation). There are many different types of radioactive

decay (see table below). A decay, or loss of energy, results when an atom with one type of nucleus, called the parent radionuclide, transforms to an atom with a nucleus in a different state, or to a

different nucleus containing different numbers of protonsand neutrons. Either of these products is named the daughter nuclide. In some decays the parent and daughter are different chemical

elements, and thus the decay process results in nuclear transmutation (creation of an atom of a new element).

The first decay processes to be discovered were alpha decay, beta decay, and gamma decay. Alpha decay occurs when the nucleus ejects an alpha particle (helium nucleus). This is the most

common process of emitting nucleons, but in rarer types of decays, nuclei can eject protons, or specific nuclei of other elements (in the process called cluster decay). Beta decay occurs when the

nucleus emits an electron or positron and a type ofneutrino, in a process that changes a proton to a neutron or the other way around. The nucleus may capture an orbiting electron, converting a

proton into an neutron (electron capture). All of these processes result in nuclear transmutation.

By contrast, there exist radioactive decay processes that do not result in transmutation. The energy of an excited nucleus may be emitted as agamma ray in gamma decay, or used to eject an

orbital electron by interaction with the excited nucleus in a process called internal conversion. Radioisotopes occasionally emit neutrons, and this results in a change in an element from

one isotope to another.

One type of radioactive decay results in products which are not defined, but appear in a range of "pieces" of the original nucleus. This decay is called spontaneous fission. This decay happens

when a large unstable nucleus spontaneously splits into two (and occasionally three) smaller daughter nuclei, and usually emits gamma rays, neutrons, or other particles as a consequence.

Radioactive decay is a stochastic (i.e., random) process at the level of single atoms, in that, according to quantum theory, it is impossible to predict when a particular atom will decay.[1] However,

the chance that a given atom will decay is constant over time. For a large number of atoms, the decay rate for the collection is computable from the measured decay constants of the nuclides (or

equivalently from the half-lifes).

Contents

[hide]

1 Natural origin

2 Decay phenomena

3 Discovery

4 Danger of radioactive substances

5 Types of decay

6 Decay modes in table form

7 Decay chains and multiple modes

8 Occurrence and applications

9 Radioactive decay rates

o 9.1 Units of radioactivity measurements

10 Mathematics of radioactive decay

o 10.1 Universal law of radioactive decay

10.1.1 One-decay process

10.1.2 Chain-decay processes

10.1.3 Alternative decay modes

o 10.2 Corollaries of the decay laws

o 10.3 Decay timing: definitions and relations

10.3.1 Time constant and mean-life

10.3.2 Half-life

o 10.4 Example

11 Changing decay rates

12 See also

13 Notes

14 References

15 External links

[edit]Natural origin

Primordial nuclides found in the earth are residues from ancient supernova explosions which occurred before the formation of the solar system. They are the long-lived fraction of radionuclides

surviving in the primordial solar nebula through planet accretion until the present. The naturally occurring short-lived radiogenic radionuclides found in rocks are the daughters of these

radiactiveprimordial nuclides. Another minor source of naturally occurring radioactive nuclides are cosmogenic nuclides, formed by cosmic ray bombardment of material in the

Earth's atmosphere or crust. For a summary table showing the number of stable nuclides and of radioactive nuclides in each category, see radionuclide. Radionuclides can also be produced

artificially e.g. using particle accelerators or nuclear reactors.

[edit]Decay phenomena

The trefoil symbol is used to indicate radioactive material.

The new radioactive danger symbol is used to indicate dangerous ionizing radioactive material.

The neutrons and protons that constitute nuclei, as well as other particles that approach close enough to them, are governed by several interactions. The strong nuclear force, not observed at the

familiar macroscopic scale, is the most powerful force over subatomic distances. The electrostatic force is almost always significant, and, in the case of beta decay, the weak nuclear force is also

involved.

The interplay of these forces produces a number of different phenomena in which energy may be released by rearrangement of particles in the nucleus, or else the change of one type of particle

into others. These rearrangements and transformations may be hindered energetically, so that they do not occur immediately. In certain cases, random quantum vacuum fluctuations are theorized

to promote relaxation to a lower energy state (the "decay") in a phenomenon known as quantum tunneling. Radioactive decay half-life of nuclides has been measured over timescales of 55 orders

of magnitude, from 2.3 x 10-23 second (for hydrogen-7) to 6.9 x 1031 seconds (for tellurium-128)[2]. The limits of these timescales are set by the sensitivity of instrumentation only, and there are no

known natural limits to how brief or long a decay half life for radioactive decay of a radionuclide may be.

The decay process, like all hindered energy transformations, may be analogized by a snowfield on a mountain. While friction between the ice crystals may be supporting the snow's weight, the

system is inherently unstable with regard to a state of lower potential energy. A disturbance would thus facilitate the path to a state of greater entropy: The system will move towards the ground

state, producing heat, and the total energy will be distributable over a larger number of quantum states. Thus, an avalanche results. The total energy does not change in this process, but, because

of the law of entropy, avalanches happen only in one direction and that is toward the "ground state" — the state with the largest number of ways in which the available energy could be distributed.

Such a collapse (a decay event) requires a specific activation energy. For a snow avalanche, this energy comes as a disturbance from outside the system, although such disturbances can be

arbitrarily small. In the case of an excited atomic nucleus, the arbitrarily small disturbance comes from quantum vacuum fluctuations. A radioactive nucleus (or any excited system in quantum

mechanics) is unstable, and can, thus, spontaneously stabilize to a less-excited system. The resulting transformation alters the structure of the nucleus and results in the emission of either a

photon or a high-velocity particle that has mass (such as an electron, alpha particle, or other type).

[edit]Discovery

Radioactivity was discovered in 1896 by the French scientist Henri Becquerel, while working on phosphorescent materials. These materials glow in the dark after exposure to light, and he

suspected that the glow produced in cathode ray tubes by X-rays might be associated with phosphorescence. He wrapped a photographic plate in black paper and placed various

phosphorescent salts on it. All results were negative until he used uranium salts. The result with these compounds was a blackening of the plate. These radiations were called Becquerel Rays.

Pierre and Marie Curie in their Paris laboratory, before 1907

It soon became clear that the blackening of the plate had nothing to do with phosphorescence, because the plate blackened when the mineral was in the dark. Non-phosphorescent salts of

uranium and metallic uranium also blackened the plate. It was clear that there is a form of radiation that could pass through paper that was causing the plate to become black.

At first it seemed that the new radiation was similar to the then recently discovered X-rays. Further research by Becquerel, Ernest Rutherford, Paul Villard,Pierre Curie, Marie Curie, and others

discovered that this form of radioactivity was significantly more complicated. Different types of decay can occur, producing very different types of radiation. Rutherford was the first to realize that

they all occur with the same mathematical exponential formula (see below), and Rutherford and his student Frederick Soddy were first to realize that many decay processes resulted in

the transmutation of one element to another. Subsequently, the radioactive displacement law of Fajans and Soddy was formulated to describe the products of alpha and beta decay.

The early researchers also discovered that many other chemical elements besides uranium have radioactive isotopes. A systematic search for the total radioactivity in uranium ores also

guided Marie Curie to isolate a new element polonium and to separate a new element radium from barium. The two elements' chemical similarity would otherwise have made them difficult to

distinguish.

[edit]Danger of radioactive substances

Main article: Ionizing radiation

The danger classificationsign of radioactive materials

Alpha particles may be completely stopped by a sheet of paper, beta particlesby aluminum shielding. Gamma rays can only be reduced by much more substantial mass, such as a very thick layer

oflead.

Different types of decay of a radionuclide. Vertical: atomic number Z, Horizontal: neutron number N

The dangers of radioactivity and radiation were not immediately recognized. Acute effects of radiation were first observed in the use of X-rays when electrical engineer and physicist Nikola

Tesla intentionally subjected his fingers to X-rays in 1896.[3] He published his observations concerning the burns that developed, though he attributed them to ozone rather than to X-rays. His

injuries later healed.

The genetic effects of radiation, including the effect of cancer risk, were recognized much later. In 1927, Hermann Joseph Muller published research showing genetic effects, and in 1946 was

awarded the Nobel prize for his findings.

Before the biological effects of radiation were known, many physicians and corporations began marketing radioactive substances as patent medicine, glow-in-the-dark pigments. Examples were

radium enema treatments, and radium-containing waters to be drunk as tonics. Marie Curie protested this sort of treatment, warning that the effects of radiation on the human body were not well

understood. Curie later died from aplastic anemia, likely caused by exposure to ionizing radiation. By the 1930s, after a number of cases of bone necrosis and death of enthusiasts, radium-

containing medicinal products had been largely removed from the market (radioactive quackery).

[edit]Types of decay

Types of radioactive decay related to N and Z numbers

As for types of radioactive radiation, it was found that an electric or magnetic field could split such emissions into three types of beams. The rays were given thealphabetic names alpha, beta,

and gamma, in order of their ability to penetrate matter. While alpha decay was seen only in heavier elements (atomic number 52,tellurium, and greater), the other two types of decay were seen in

all of the elements. Spontaneous decay is evident in elements of atomic number ninety or greater.

In analyzing the nature of the decay products, it was obvious from the direction of electromagnetic forces induced upon the radiations by external magnetic and electric fields that alpha

particles carried a positive charge, beta particles carried a negative charge, and gamma rays were neutral. From the magnitude of deflection, it was clear that alpha particles were much more

massive than beta particles. Passing alpha particles through a very thin glass window and trapping them in a discharge tube allowed researchers to study the emission spectrum of the resulting

gas, and ultimately prove that alpha particles are helium nuclei. Other experiments showed the similarity between classical beta radiation and cathode rays: They are both streams of electrons.

Likewise gamma radiation and X-rays were found to be similar high-energy electromagnetic radiation.

The relationship between the types of decays also began to be examined: For example, gamma decay was almost always found associated with other types of decay, and occurred at about the

same time, or afterward. Gamma decay as a separate phenomenon (with its own half-life, now termed isomeric transition), was found in natural radioactivity to be a result of the gamma decay of

excited metastable nuclear isomers, which were in turn created from other types of decay.

Although alpha, beta, and gamma radiations were found most commonly, other types of decay were eventually discovered. Shortly after the discovery of thepositron in cosmic ray products, it was

realized that the same process that operates in classical beta decay can also produce positrons (positron emission). In an analogous process, instead of emitting positrons and neutrinos, some

proton-rich nuclides were found to capture their own atomic electrons (electron capture), and emit only a neutrino (and usually also a gamma ray). Each of these types of decay involves the

capture or emission of nuclear electrons or positrons, and acts to move a nucleus toward the ratio of neutrons to protons that has the least energy for a given total number of nucleons (neutrons

plus protons).

A theoretical process of positron capture (analogous to electron capture) is possible in antimatter atoms, but has not been observed since the complex antimatter atoms are not available.[4] This

would required antimatter atoms at least as complex as beryllium-7, which is the lightest known isotope of normal matter to undergo decay by electron capture.

Shortly after the discovery of the neutron in 1932, Enrico Fermi realized that certain rare decay reactions yield neutrons as a decay particle (neutron emission). Isolated proton emission was

eventually observed in some elements. It was also found that some heavy elements may undergo spontaneous fission into products that vary in composition. In a phenomenon called cluster

decay, specific combinations of neutrons and protons other than alpha particles (helium nuclei) were found to be spontaneously emitted from atoms.

Other types of radioactive decay that emit previously seen particles were found, but by different mechanisms. An example is internal conversion, which results in electron and sometimes high-

energy photon emission, even though it involves neither beta nor gamma decay. This type of decay (like isomeric transition gamma decay) does not transmute one element to another.

Rare events that involve a combination of two beta-decay type events happening simultaneously (see below) are known. Any decay process that does not violate conservation of energy or

momentum laws (and perhaps other particle conservation laws) is permitted to happen, although not all have been detected. An interesting example (discussed in a final section) is bound state

beta decay of rhenium-187. In this process, an inverse of electron capture, beta electron-decay of the parent nuclide is not accompanied by beta electron emission, because the beta particle has

been captured into the K-shell of the emitting atom. An antineutrino, however, is emitted.

[edit]Decay modes in table form

Radionuclides can undergo a number of different reactions. These are summarized in the following table. A nucleus with mass number A and atomic number Z is represented as (A, Z). The

column "Daughter nucleus" indicates the difference between the new nucleus and the original nucleus. Thus, (A − 1, Z) means that the mass number is one less than before, but the atomic

number is the same as before.

Mode of decay Participating particles Daughter nucleus

Decays with emission of nucleons:

Alpha decay An alpha particle (A = 4, Z = 2) emitted from nucleus (A − 4, Z − 2)

Proton emission A proton ejected from nucleus (A − 1, Z − 1)

Neutron emission A neutron ejected from nucleus (A − 1, Z)

Double proton emission Two protons ejected from nucleus simultaneously (A − 2, Z − 2)

Spontaneous fission Nucleus disintegrates into two or more smaller nuclei and other particles —

Cluster decay Nucleus emits a specific type of smaller nucleus (A1, Z1) smaller than, or larger than, an alpha particle(A − A1, Z − Z1) + (A1, Z1)

Different modes of beta decay:

β − decay A nucleus emits an electron and an electron antineutrino (A, Z + 1)

Positron emission (β + decay ) A nucleus emits a positron and an electron neutrino (A, Z − 1)

Electron capture A nucleus captures an orbiting electron and emits a neutrino; the daughter nucleus is left in an excited unstable state (A, Z − 1)

Bound state beta decayA nucleus beta decays to electron and antineutrino, but the electron is not emitted, as it is captured into an empty K-shell; the daughter nucleus is left in an excited and unstable state. This process is suppressed except in ionized atoms that have K-shell vacancies.

(A, Z + 1)

Double beta decay A nucleus emits two electrons and two antineutrinos (A, Z + 2)

Double electron capture A nucleus absorbs two orbital electrons and emits two neutrinos – the daughter nucleus is left in an excited and unstable state (A, Z − 2)

Electron capturewith positron emission

A nucleus absorbs one orbital electron, emits one positron and two neutrinos (A, Z − 2)

Double positron emission A nucleus emits two positrons and two neutrinos (A, Z − 2)

Transitions between states of the same nucleus:

Isomeric transition Excited nucleus releases a high-energy photon (gamma ray) (A, Z)

Internal conversion Excited nucleus transfers energy to an orbital electron, which is subsequently ejected from the atom (A, Z)

Radioactive decay results in a reduction of summed rest mass, once the released energy (the disintegration energy) has escaped in some way (for example, the products might be captured and

cooled, and the heat allowed to escape). Although decay energy is sometimes defined as associated with the difference between the mass of the parent nuclide products and the mass of the

decay products, this is true only of rest mass measurements, where some energy has been removed from the product system. This is true because the decay energy must always carry mass with

it, wherever it appears (see mass in special relativity) according to the formula E = mc2. The decay energy is initially released as the energy of emitted photons plus the kinetic energy of massive

emitted particles (that is, particles that have rest mass). If these particles come to thermal equilibrium with their surroundings and photons are absorbed, then the decay energy is transformed to

thermal energy, which retains its mass.

Decay energy therefore remains associated with a certain measure of mass of the decay system invariant mass. The energy of photons, kinetic energy of emitted particles, and, later, the thermal

energy of the surrounding matter, all contribute to calculations of invariant mass of systems. Thus, while the sum of rest masses of particles is not conserved in radioactive decay, the systemmass

and system invariant mass (and also the system total energy) is conserved throughout any decay process.

[edit]Decay chains and multiple modes

The daughter nuclide of a decay event may also be unstable (radioactive). In this case, it will also decay, producing radiation. The resulting second daughter nuclide may also be radioactive. This

can lead to a sequence of several decay events. Eventually, a stable nuclide is produced. This is called a decay chain (see this article for specific details of important natural decay chains).

Gamma-ray energy spectrum of uranium ore (inset). Gamma-rays are emitted by decayingnuclides, and the gamma-ray energy can be used to characterize the decay (which nuclide is decaying

to which). Here, using the gamma-ray spectrum, several nuclides that are typical of the decay chain of 238U have been identified: 226Ra,214Pb, 214Bi.

An example is the natural decay chain of 238U, which is as follows:

decays, through alpha-emission, with a half-life of 4.5 billion years to thorium-234

which decays, through beta-emission, with a half-life of 24 days to protactinium-234

which decays, through beta-emission, with a half-life of 1.2 minutes to uranium-234

which decays, through alpha-emission, with a half-life of 240 thousand years to thorium-230

which decays, through alpha-emission, with a half-life of 77 thousand years to radium-226

which decays, through alpha-emission, with a half-life of 1.6 thousand years to radon-222

which decays, through alpha-emission, with a half-life of 3.8 days to polonium-218

which decays, through alpha-emission, with a half-life of 3.1 minutes to lead-214

which decays, through beta-emission, with a half-life of 27 minutes to bismuth-214

which decays, through beta-emission, with a half-life of 20 minutes to polonium-214

which decays, through alpha-emission, with a half-life of 160 microseconds to lead-210

which decays, through beta-emission, with a half-life of 22 years to bismuth-210

which decays, through beta-emission, with a half-life of 5 days to polonium-210

which decays, through alpha-emission, with a half-life of 140 days to lead-206, which is a stable nuclide.

Some radionuclides may have several different paths of decay. For example, approximately 36% of bismuth-212 decays, through alpha-emission, to thallium-208 while approximately 64%

of bismuth-212 decays, through beta-emission, to polonium-212. Both the thallium-208 and thepolonium-212 are radioactive daughter products of bismuth-212, and both decay directly to

stable lead-208.

[edit]Occurrence and applications

According to the Big Bang theory, stable isotopes of the lightest five elements (H, He, and traces of Li, Be, and B) were produced very shortly after the emergence of the universe, in a process

called Big Bang nucleosynthesis. These lightest stable nuclides (including deuterium) survive to today, but any radioactive isotopes of the light elements produced in the Big Bang (such astritium)

have long since decayed. Isotopes of elements heavier than boron were not produced at all in the Big Bang, and these first five elements do not have any long-lived radioisotopes. Thus, all

radioactive nuclei are, therefore, relatively young with respect to the birth of the universe, having formed later in various other types of nucleosynthesis in stars (in particular, supernovae), and also

during ongoing interactions between stable isotopes and energetic particles. For example, carbon-14, a radioactive nuclide with a half-life of only 5730 years, is constantly produced in Earth's

upper atmosphere due to interactions between cosmic rays and nitrogen.

Nuclides that are produced by radioactive decay are called radiogenic nuclides, whether they themselves are stable or not. There exist stable radiogenic nuclides that were formed from short-

livedextinct radionuclides in the early solar system.[5][6] The extra presence of these stable radiogenic nuclides (such as Xe-129 from primordial I-129) against the background of primordial stable

nuclides can be inferred by various means.

Radioactive decay has been put to use in the technique of radioisotopic labeling, which is used to track the passage of a chemical substance through a complex system (such as a livingorganism).

A sample of the substance is synthesized with a high concentration of unstable atoms. The presence of the substance in one or another part of the system is determined by detecting the locations

of decay events.

On the premise that radioactive decay is truly random (rather than merely chaotic), it has been used in hardware random-number generators. Because the process is not thought to vary

significantly in mechanism over time, it is also a valuable tool in estimating the absolute ages of certain materials. For geological materials, the radioisotopes and some of their decay products

become trapped when a rock solidifies, and can then later be used (subject to many well-known qualifications) to estimate the date of the solidification. These include checking the results of

several simultaneous processes and their products against each other, within the same sample. In a similar fashion, and also subject to qualification, the rate of formation of carbon-14 in various

eras, the date of formation of organic matter within a certain period related to the isotope's half-life may be estimated, because the carbon-14 becomes trapped when the organic matter grows and

incorporates the new carbon-14 from the air. Thereafter, the amount of carbon-14 in organic matter decreases according to decay processes that may also be independently cross-checked by

other means (such as checking the carbon-14 in individual tree rings, for example).

[edit]Radioactive decay rates

The decay rate, or activity, of a radioactive substance are characterized by:

Constant quantities:

The half-life—t1/2, is the time taken for the activity of a given amount of a radioactive substance to decay to half of its initial value; see List of nuclides.

The mean lifetime— τ, "tau" the average lifetime of a radioactive particle before decay.

The decay constant— λ, "lambda" the inverse of the mean lifetime.

Although these are constants, they are associated with statistically random behavior of populations of atoms. In consequence predictions using these constants are less accurate for small number

of atoms.

In principle the reciprocal of any number greater than one— a half-life, a third-life, or even a (1/√2)-life—can be used in exactly the same way as half-life; but the half-life t1/2 is adopted as the

standard time associated with exponential decay.

Time-variable quantities:

Total activity— A, is number of decays per unit time of a radioactive sample.

Number of particles—N, is the total number of particles in the sample.

Specific activity—SA, number of decays per unit time per amount of substance of the sample at time set to zero (t = 0). "Amount of substance" can be the mass, volume or moles of the initial

sample.

These are related as follows:

where N0 is the initial amount of active substance — substance that has the same percentage of unstable particles as when the substance was formed.

[edit]Units of radioactivity measurements

The SI unit of radioactive activity is the becquerel (Bq), in honor of the scientist Henri Becquerel. One Bq is defined as one transformation (or decay or disintegration) per second.

Since sensible sizes of radioactive material contains many atoms, a Bq is a tiny measure of activity; amounts giving activities on the order of GBq (gigabecquerel, 1 x 109 decays per

second) or TBq (terabecquerel, 1 x 1012 decays per second) are commonly used.

Another unit of radioactivity is the curie, Ci, which was originally defined as the amount of radium emanation (radon-222) in equilibrium with one gram of pure radium, isotope Ra-226.

At present it is equal, by definition, to the activity of any radionuclide decaying with a disintegration rate of 3.7 × 1010 Bq, so that 1 curie (Ci) = 3.7 × 1010 Bq. The use of Ci is currently

discouraged by the SI. Low activities are also measured in disintegrations per minute (dpm).

[edit]Mathematics of radioactive decay

For the mathematical details of exponential decay in general context, see exponential decay.

For related derivations with some further details, see half-life.

For the analogous mathematics in 1st order chemical reactions, see Consecutive reactions.

[edit]Universal law of radioactive decay

Radioactivity is one very frequent example of exponential decay. The law describes the statistical behavior of a large number of nuclides, rather than individual ones. In the following

formalism, the number of nuclides or nuclide population N, is of course a discrete variable (a natural number) - but for any physical sample N is so large (amounts of L =

1023, avagadro's constant) that it can be treated as a continuous variable. Differential calculus is needed to set up differential equations for modelling the behaviour of the nuclear

decay.

[edit]One-decay process

Consider the case of a nuclide A decaying into another B by some process A → B (emission of other particles, like electron neutrinos ν

e and electrons e– in beta decay, are irrelevant in what follows). The decay of an unstable nucleus is entirely random and it is impossible to predict when a particular atom will decay.[1] However, it is equally likely to decay at any time. Therefore, given a sample of a particular radioisotope, the number of decay events −dN expected to occur in a small interval of

time dt is proportional to the number of atoms present N, that is

Particular radionuclides decay at different rates, so each has its own decay constant λ. The expected decay −dN/N is proportional to an increment of time, dt:

The negative sign indicates that N decreases as time increases, as each decay event follows one after another. The solution to this first-order differential equation is

the function:

where N0 is the value of N at time t = 0.

This equation is of particular interest; the behaviour of numerous important quantities can be found from it (see below). Although the parent decay distribution follows an

exponential, observations of decay times will be limited by a finite integer number of N atoms and follow Poisson statistics as a consequence of the random nature of the

process.

We have for all time t:

where Ntotal is the constant number of particles throughout the decay process, clearly equal to the initial number of A nuclides since this is the initial substance.

If the number of non-decayed A nuclei is:

then the number of nuclei of B, i.e. number of decayed A nuclei, is

[edit]Chain-decay processes

Chain of two decays

Now consider the case of a chain of two decays: one nuclide A decaying into another B by one process, then B decaying into another C by a second

process, i.e. A → B → C. The previous equation cannot be applied to a decay chain, but can be generalized as follows. The decay rate of B is proportional to

the number of nuclides of B present, so again we have:

but care must be taken. Since A decays into B, then B decays into C, the activity of A adds to the total number of B nuclides in the present

sample, before those B nuclides decay and reduce the number of nuclides leading to the later sample. In other words, the number of second generation

nuclei B increases as a result of the first generation nuclei decay of A, and decreases as a result of its own decay into the third generation nuclei C.

[7] The proportionality becomes an equation:

adding the increasing (and correcting) term obtains the law for a decay chain for two nuclides:

The equation is not

since this implies the number of atoms of B is only decreasing as time increases, which is not the case. The rate of change of NB, that

is dNB/dt, is related to the changes in the amounts of Aand B, NB can increase as B is produced from A and decrease as B produces C.

Re-writing using the previous results:

The subscripts simply refer to the respective nuclides, i.e. NA is the number of nuclides of type A, NA0 is the initial number of nuclides of

type A, λA is the decay constant for A - and similarly for nuclide B. Solving this equation for NB gives:

Naturally this equation reduces to the previous solution, in the case B is a stable nuclide (λB = 0):

as shown above for one decay. The solution can be found by the integration factor method, where the integrating factor is eλB

t.

This case is perhaps the most useful, since it can derive both the one-decay equation (above) and the equation for multi-

decay chains (below) more directly.

Chain of any number of decays

For the general case of any number of consecutive decays in a decay chain, i.e. A1 → A2 ··· → Ai ··· → AD, where D is the

number of decays and i is a dummy index (i = 1, 2, 3, ...D), each nuclide population can be found in terms of the previous

population. In this case N2 = 0, N3 = 0,..., ND = 0. Using the above result in a recursive form:

The general solution to the recursive problem are given by Bateman's equations[8]:

[edit]Alternative decay modes

In all of the above examples, the initial nuclide decays into only one product. Consider the case of one initial nuclide

which can decay into two products, that is A → B + C. We have for all timet:

in which,

so the relations follow in parallel:

indicating that the total decay constant is that of A, given by:

Solving this equation for NA:

When measuring the production of one nuclide, one can only observe the total decay

constant λA. The decay constants λB and λC determine the probability for the decay to result in

products Bor C as follows:

These perhaps seemingly disjionted results are consistent:

[edit]Corollaries of the decay laws

The solutions to the above differential equations are sometimes written using

quantities related to the number of nuclide particles N in a sample,

where L is Avogadro's constant,6.023×1023, andAr is the relative atomic mass number,

and the amount of the substance is in moles.

The activity: A = λN.

The amount of substance: n = N/L.

The mass: M = Arn = ArN/L.

Collecting these results together for convenience: N = A/λ = Ln = LM/Ar.

Equivalent ways to write the decay solutions, then, are as follows:

One-decay processes

The solution

can be written:

Notice how we can simply replace each quantity (on both sides of the

equation), since they are directly proportional to N and so the

constants cancel (constant at least for a particular nuclide).

Chain-decay processes

For the two-decay chain,

its almost as simple:

[edit]Decay timing: definitions and relations

[edit]Time constant and mean-life

For the one-decay solution A → B:

the equation indicates that the decay

constant λ has units of t-1, and can thus also be

represented as 1/τ, where τ is a characteristic

time of the process called the time constant.

In a radioactive decay process, this time

constant is also the mean lifetime for decaying

atoms. Each atom "lives" for a finite amount of

time before it decays, and it may be shown that

this mean lifetime is the arithmetic mean of all

the atoms' lifetimes, and that it is τ, which again

is related to the decay constant as follows:

This form is also true for two-decay

processes simultaneously A → B + C,

inserting the equivalent values of decay

constants (as given above)

into the decay solution leads to:

Simulation of many identical

atoms undergoing radioactive

decay, starting with either 4 atoms

(left) or 400 (right). The number at

the top indicates how many half-

liveshave elapsed. Note the law of

large numbers: with more atoms,

the overall decay is less random.

[edit]Half-life

A more commonly used parameter

is the half-life. Given a sample of a

particular radionuclide, the half-life

is the time taken for half the

radionuclide's atoms to decay. For

the case of one-decay nuclear

reactions:

the half-life is related to the

decay constant as follows:

set N = N0/2 and t = T1/2 to

obtain

This relationship between

the half-life and the decay

constant shows that

highly radioactive

substances are quickly

spent, while those that

radiate weakly endure

longer. Half-lives of

known radionuclides vary

widely, from more

than 10 19   years , such as

for the very nearly stable

nuclide 209Bi, to

10−23 seconds for highly

unstable ones.

The factor of ln(2) in the

above relations results

from the fact that concept

of "half-life" is merely a

way of selecting a

different base other than

the natural base e for the

lifetime expression. The

time constant τ is the e -

1 -life, the time until only

1/e remains, about

36.8%, rather than the

50% in the half-life of a

radionuclide. Thus, τ is

longer than t1/2. The

following equation can be

shown to be valid:

Since radioactive

decay is exponential

with a constant

probability, each

process could as

easily be described

with a different

constant time period

that (for example)

gave its "(1/3)-life"

(how long until only

1/3 is left) or "(1/10)-

life" (a time period

until only 10% is left),

and so on. Thus, the

choice of τ and t1/2 for

marker-times, are

only for convenience,

and from convention.

They reflect a

fundamental principle

only in so much as

they show that

the same

proportion of a given

radioactive

substance will decay,

during any time-

period that one

chooses.

Mathematically,

the nth life for the

above situation

would be found in the

same way as above

—by setting N = N0/n,

and substituting into

the decay solution to

obtain

[edit]Example

A sample of 14C,

whose half-life is

5730 years, has

a decay rate of

14 disintegration

per minute (dpm)

per gram of

natural carbon.

An artefact is

found to have

radioactivity of 4

dpm per gram of

its present C,

how old is the

artefact?

Using the above

equation, we

have:

where: 

 years,

 years.

[

edit]

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Half LifeRadioactive stuff and what is meant by it having a "half life"

If you're wondering what a HALF LIFE is, and why scientists talk about atomic elements having a half life rather than a whole life, here are some answers.

STUFF

Stuff is made of atoms. Different stuff is made of different atoms. Most types of stuff (gold, oxygen, iron, helium, etc) have atoms which are stable and will continue to remain the same even if they are kept a very long time. In contrast, radioactive elements (such as plutonium, uranium, radium, etc) have atoms which are unstable.

SPLITTING THE ATOM

Radioactive elements' unstable atoms will SPLIT. They will do this at random, on their own, without any persuasion. It's as if the radioactive atoms are too big for comfort and just have to split. Splitting of the atom occurs spontaneously, and randomly. In general a few small bits come off. The small bits are the radioactivity and what's remaining is an atom which is a different element. So, what you see is one kind of stuff changing into a different kind of stuff, and giving off particles at the same time. The particles of radioactivity coming off can be detected using a Geiger-counter which makes that familiar clicking sound.

LEAD TURNS TO GOLD?

It would be handy if lead turned to gold, but it doesn't usually go that way, as lead isn't unstable. What usually happens is that a heavy radioactive element turns into another radioactive element which then turns into lead.

THE TIMESCALE

Each different radioactive element has its own characteristic and will turn into another material at a particular rate. The timescale at which this happens is known as "the half life".

Why is it HALF LIFE?

Supposing there is a radioactive element Frankensteinium which has a characteristic of changing into lead over a period of years and it is described by scientists as "having a half life of one year". That would mean that after one year, half of the Frankensteinium has turned into lead. So, if there was one kilogramme of Frankensteinium to start with, then a year later there would be only 500 grammes of Frankensteinium, together with 500 grammes of lead.

The misconception which is commonly believed by folk is that after two years it would all have turned to lead and there would be no radioactive material remaining. It's not so. What happens in the second year is very similar to what happened in the first year. Half of the Frankensteinium turns to lead. So after two years there is 750 grammes of lead and 250 grammes of Frankensteinium yet to turn. After another year, half of the 250 grammes remaining will turn into lead leaving only 125 grammes. This goes on indefinitely, and for each year (each half-life), half of the remaining material will turn into lead.

The amount of radioactivity given off reduces, as 125g of Frankensteinium gives off less split off material than the 1000g which was there to start with.

After ten half lives (ten years in the case of the hypothetical element Frankensteinium), only one gramme of the radioactive element remains, the other 999g having turned to lead.

After twenty half lives, only one millionth of the original material still remains.

Real atomic elements

For actual atomic elements the situation is more complicated. Materials such as plutonium, uranium, and thorium change into other elements which then turn into other elements, in a sequence. Also, there's more than one kind of each element. These

different types (isotopes) have different radioactive characteristics. Also, the different materials have widely different half-lives. Some are several minutes, some several years, and some have a half life of millions of years, or a tiny fraction of a second. See extended periodic table

How long before it's all gone?

So, if you are wondering how long a heap of radioactive waste is going to be around, the answer is generally "far too long", as the stuff will be a continuously changing mixture of stuff which will go on atomically splitting and changing (transmuting) for a very long time before it eventually turns to lead. It's an ecological problem.

Although in theory the radioactive material will never have "all gone", there is a practical way of thinking about it based on what an "empty bottle" is. If you have a bottle of drink, and you pour it all out into glasses, so no more comes out, then is that an "empty" bottle? If you think it is practically empty, even if there's a drop left in it, then it's possible to consider some point at which radioactive material has ALMOST all gone.

However, if the half life of the radioactive material is "70,000 years" (a guess placed on modern radioactive waste being produced from nuclear power stations), then it's going to be around for a very long time whichever way you think about it. Certainly much longer than the lifespan of a nuclear industry that can't be sure of balancing its accounts in 50 years time.

Half-lifeFrom Wikipedia, the free encyclopedia

This article is about the scientific and mathematical term. For other uses, see Half-life (disambiguation).

This article needs additional citations for verification. Please help improve this article by adding citations to reliable sources. Unsourced material may be challenged and removed. (July 2009)

Half-life (t½) is the time required for a quantity to fall to half its value as measured at the beginning of the time period. In physics, it is typically used to describe a property

of radioactive decay, but may be used to describe any quantity which follows an exponential decay.

The original term, dating to 1907, was "half-life period", which was shortened to "half-life" in the early 1950s.[1]

Half-life is used to describe a quantity undergoing exponential decay, and is constant over the lifetime of the decaying quantity. It is a characteristic unitfor the exponential

decay equation. The term "half-life" may generically be used to refer to any period of time in which a quantity falls by half, even if the decay is not exponential. For a general

introduction and description of exponential decay, see exponential decay. For a general introduction and description of non-exponential decay, see rate law.

The converse of half-life is doubling time.

The table on the right shows the reduction of a quantity in terms of the number of half-lives elapsed.

Number ofhalf-liveselapsed

Fractionremainin

g

Percentageremaining

0 1/1 100

1 1/2 50

2 1/4 25

3 1/8 12.5

4 1/16 6.25

5 1/32 3.125

6 1/64 1.563

7 1/128 0.781

... ... ...

n 2-n 100/(2n)

Contents

[hide]

1 Probabilistic nature of half-life

2 Formulas for half-life in exponential decay

o 2.1 Decay by two or more processes

o 2.2 Examples

3 Half-life in non-exponential decay

4 Half-life in biology and pharmacology

5 See also

6 References

7 External links

[edit]Probabilistic nature of half-life

Simulation of many identical atoms undergoing radioactive decay, starting with either 4 atoms per box (left) or 400 (right). The number at the top is how many half-lives have elapsed. Note thelaw of large numbers: With more atoms, the overall decay is more regular and more predictable.

A half-life usually describes the decay of discrete entities, such as radioactive atoms, which have unstable nuclei. In that case, it does not work to use the definition "half-life is the time required for exactly

half of the entities to decay". For example, if there is just one radioactive atom with a half-life of one second, there will not be "one-half of an atom" left after one second. There will be either zero atoms left

or one atom left, depending on whether or not that atom happened to decay.

Instead, the half-life is defined in terms of probability. It is the time when the expected value of the number of entities that have decayed is equal to half the original number. For example, one can start with a

single radioactive atom, wait its half-life, and then check whether or not it has decayed. Perhaps it did, but perhaps it did not. But if this experiment is repeated again and again, it will be seen that - on

average - it decays within the half-life 50% of the time.

In some experiments (such as the synthesis of a superheavy element), there is in fact only one radioactive atom produced at a time, with its lifetime individually measured. In this case, statistical analysis is

required to infer the half-life. In other cases, a very large number of identical radioactive atoms decay in the measured time range. In this case, the law of large numbers ensures that the number of atoms

that actually decay is approximately equal to the number of atoms that are expected to decay. In other words, with a large enough number of decaying atoms, the probabilistic aspects of the process could

be neglected.

There are various simple exercises that demonstrate probabilistic decay, for example involving flipping coins or running a statistical computer program.[2][3][4] For example, the image on the right is a

simulation of many identical atoms undergoing radioactive decay. Note that after one half-life there are not exactly one-half of the atoms remaining, only approximately, because of the random variation in

the process. However, with more atoms (right boxes), the overall decay is smoother and less random-looking than with fewer atoms (left boxes), in accordance with the law of large numbers.

[edit]Formulas for half-life in exponential decay

Main article: Exponential decay

An exponential decay process can be described by any of the following three equivalent formulas:

where

N0 is the initial quantity of the substance that will decay (this quantity may be measured in grams, moles, number of atoms, etc.),

N(t) is the quantity that still remains and has not yet decayed after a time t,

t1/2 is the half-life of the decaying quantity,

τ is a positive number called the mean lifetime of the decaying quantity,

λ is a positive number called the decay constant of the decaying quantity.

The three parameters  ,  , and λ are all directly related in the following way:

where ln(2) is the natural logarithm of 2 (approximately 0.693).

[show]Click "show" to see a detailed derivation of the relationship between half-life, decay time, and decay constant.

By plugging in and manipulating these relationships, we get all of the following equivalent descriptions of exponential decay, in terms of the half-life:

Regardless of how it's written, we can plug into the formula to get

 as expected (this is the definition of "initial quantity")

 as expected (this is the definition of half-life)

, i.e. amount approaches zero as t approaches infinity as expected (the longer we wait, the less remains).

[edit]Decay by two or more processes

Some quantities decay by two exponential-decay processes simultaneously. In this case, the actual half-life T1/2 can be related to the half-lives t1 and t2 that the quantity

would have if each of the decay processes acted in isolation:

For three or more processes, the analogous formula is:

For a proof of these formulas, see Decay by two or more processes.

[edit]Examples

Main article: Exponential decay--Applications and examples

There is a half-life describing any exponential-decay process. For example:

The current flowing through an RC circuit or RL circuit decays with a half-life of   or  , respectively. For this example, the

term half time might be used instead of "half life", but they mean the same thing.

In a first-order chemical reaction, the half-life of the reactant is  , where λ is the reaction rate constant.

In radioactive decay, the half-life is the length of time after which there is a 50% chance that an atom will have undergone nuclear decay. It varies

depending on the atom type and isotope, and is usually determined experimentally. See List of nuclides.

[edit]Half-life in non-exponential decay

Main article: Rate equation

The decay of many physical quantities is not exponential—for example, the evaporation of water from a puddle, or (often) the chemical reaction of a

molecule. In such cases, the half-life is defined the same way as before: as the time elapsed before half of the original quantity has decayed. However,

unlike in an exponential decay, the half-life depends on the initial quantity, and the prospective half-life will change over time as the quantity decays.

As an example, the radioactive decay of carbon-14 is exponential with a half-life of 5730 years. A quantity of carbon-14 will decay to half of its original

amount (on average) after 5730 years, regardless of how big or small the original quantity was. After another 5730 years, one-quarter of the original will

remain. On the other hand, the time it will take a puddle to half-evaporate depends on how deep the puddle is. Perhaps a puddle of a certain size will

evaporate down to half its original volume in one day. But on the second day, there is no reason to expect that one-quarter of the puddle will remain; in fact,

it will probably be much less than that. This is an example where the half-life reduces as time goes on. (In other non-exponential decays, it can increase

instead.)

The decay of a mixture of two or more materials which each decay exponentially, but with different half-lives, is not exponential. Mathematically, the sum of

two exponential functions is not a single exponential function. A common example of such a situation is the waste of nuclear power stations, which is a mix

of substances with vastly different half-lives. Consider a sample containing a rapidly decaying element A, with a half-life of 1 second, and a slowly decaying

element B, with a half-life of one year. After a few seconds, almost all atoms of the element A have decayed after repeated halving of the initial total

number of atoms; but very few of the atoms of element B will have decayed yet as only a tiny fraction of a half-life has elapsed. Thus, the mixture taken as

a whole does not decay by halves.

[edit]Half-life in biology and pharmacology

Main article: Biological half-life

A biological half-life or elimination half-life is the time it takes for a substance (drug, radioactive nuclide, or other) to lose one-half of its pharmacologic,

physiologic, or radiological activity. In a medical context, the half-life may also describe the time that it takes for the concentration in blood plasma of a

substance to reach one-half of its steady-state value (the "plasma half-life").

The relationship between the biological and plasma half-lives of a substance can be complex, due to factors including accumulation in tissues,

active metabolites, and receptor interactions.[5]

While a radioactive isotope decays almost perfectly according to so-called "first order kinetics" where the rate constant is a fixed number, the elimination of

a substance from a living organism usually follows more complex chemical kinetics.

For example, the biological half-life of water in a human being is about seven to 14 days, though this can be altered by his/her behavior. The biological half-

life of cesium in human beings is between one and four months. This can be shortened by feeding the person prussian blue, which acts as a solid ion

exchanger that absorbs the cesium while releasing potassium ions in their place.

List of radioactive isotopes by half-lifeFrom Wikipedia, the free encyclopedia

  (Redirected from List of isotopes by half-life)

This list is incomplete; you can help by expanding it.

This is a list of radioactive isotopes ordered by half-life from the shortest to the longest.

Contents

[hide]

1 10 -18 seconds and less

2 10 -12 seconds

3 10 -9 seconds

4 10 -6 seconds

5 10 -3 seconds

6 10 0 seconds

7 10 3 seconds

8 10 6 seconds

9 10 9 seconds

10 10 12 seconds

11 10 15 seconds

12 10 18 seconds

13 10 21 seconds

14 10 24 seconds and beyond

15 See also

16 Note

[edit]10-18 seconds and less

isotopehalf-life

10−24 seconds

hydrogen- 23

7

lithium-4 75.7

hydrogen-5

80

hydrogen-4

100

hydrogen-6

290

lithium-5 300

boron-7 350

helium-5 760

half-life10−21 seconds

helium-10 1.52

lithium-10 2

carbon-8 2

helium-7 3.04

beryllium-6 5

helium-9 7

boron-9 800

half-life10−18 seconds

beryllium-8 81.9

[edit]10-12 seconds

isotopehalf-life

10−12 seconds

boron-16 190

beryllium-13

500

[edit]10-9 seconds

isotopehalf-life

10−9 seconds

lithium-12 10

boron-18 26

carbon-21 30

beryllium-15 200

beryllium-16 200

[edit]10-6 seconds

isotopehalf-life

10−6 seconds

Copernicium-277 240

[edit]10-3 seconds

isotopehalf-life

10−3 seconds

hassium-265 2

boron-19 2.92

meitnerium-266

3.4

boron-17 5.08

carbon-22 6.2

lithium-11 8.59

boron-15 9.87

boron-14 12.5

carbon-20 16

boron-13 17.33

boron-12 20.2

beryllium-12 21.49

carbon-19 46.2

carbon-18 92

bohrium-262 102

helium-8 119

carbon-9 126.5

lithium-9 178.3

carbon-17 193

carbon-16 747

boron-8 770

helium-6 806.7

lithium-8 839.9

[edit]100 seconds

isotopehalf-lifeseconds

carbon-15 2.449

flerovium-289 2.6

beryllium-14 4.84

beryllium-11 13.81

carbon-10 19.29

dubnium-261 27

seaborgium-266 30

dubnium-262 34

rutherfordium-261

61

nobelium-253 102

[edit]103 seconds

isotope

half-life

minutes 103 seconds

carbon-11 20.334 1.2200

nobelium-259 58 3.5

hours 103 seconds

fluorine-18 1.8295 6.586

mendelevium-257

5.52 19.9

erbium-165 10.36 37.3

fermium-252 25.39 91.4

erbium-160 28.58 102.9

days 103 seconds

fermium-253 3 260

manganese-52 5.591 483.1

thulium-167 9.25 799

[edit]106 seconds

isotope

half-life

days 106 seconds

vanadium-48 15.9735 1.38011

californium-253 17.81 1.539

chromium-51 27.7025 2.39350

mendelevium-258 51.5 4.45

beryllium-7 53.12 4.590

californium-254 60.5 5.23

cobalt-56 77.27 6.676

scandium-46 83.79 7.239

sulfur-35 87.32 7.544

thulium-168 93.1 8.04

fermium-257 100.5 8.68

thulium-170 128.6 11.11

polonium-210 138 11.9

cobalt-57 271.79 23.483

vanadium-49 330 29

californium-248 333.5 28.81

ruthenium-106 373.59 32.278

neptunium-235 396.1 34.22

cadmium-109 462.6 39.97

years 106 seconds

thulium-171 1.92 61

sodium-22 2.602 82.1

rhodium-101 3.3 100

cobalt-60 5.2714 166.35

hydrogen-3 12.32 389

californium-250 13.08 413

niobiummeta state Nb-93m

16.13 509

curium-243 29.1 920

[edit]109 seconds

isotope

half-life

years 109 seconds

titanium-44 63 2.0

uranium-232 68.9 2.17

plutonium-238 87.7 2.77

nickel-63 100.1 3.16

silicon-32 170 5.4

argon-39 269 8.5

californium-249

351 11.1

silver-108 418 13.2

americium-241 432.2 13.64

niobium-91 680 21

californium-251

898 28.3

carbon-14[1] 5,715 180.4

americium-243 7,370 233

curium-245 8,500 270

niobium-94 20,300 640

plutonium-239 24,110 761

[edit]1012 seconds

isotope

half-life

103 years 1012 seconds

nickel-59 76 2.4

neptunium-236 154 4.9

uranium-233 159.2 5.02

technetium-99 211.1 6.66

uranium-234 245.5 7.75

chlorine-36 301 9.5

curium-248 340 11

aluminium-26 717 22.6

106 years 1012 seconds

beryllium-10 1.36 43

zirconium-93 1.53 48

technetium-97 2.6 82

manganese-53 3.74 118

technetium-98 4.2 130

palladium-107 6.5 210

curium-247 15.6 490

uranium-236 23.42 739

[edit]1015 seconds

isotope

half-life

106 years 1015 seconds

niobium-92 34.7 1.10

plutonium-244

80 2.5

samarium-146 103.5 3.27

uranium-235 703.8 22.21

109 years 1015 seconds

potassium-40 1.277 40.3

uranium-238 4.468 141.0

thorium-232 14.056 443.6

[edit]1018 seconds

isotope

half-life

109 years 1018 seconds

lutetium-176 38.5 1.21

rhenium-187 41.2 ± .2 1.30 ± 0.0063

rubidium-87 49.23 1.554

lanthanum-138 102 3.2

samarium-147 106 3.3

platinum-190 650 ± 30 21 ± 0.95

[edit]1021 seconds

isotope half-life

1012 years 1021 seconds

barium-130 70 2.2

gadolinium-152 108 3.4

indium-115 441 13.9

tellurium-123 599 18.9

1015 years 1021 seconds

hafnium-174 2.0 63

osmium-186 2.0 63

neodymium-144

2.29 72

samarium-148 7 220

cadmium-113 7.7 240

[edit]1024 seconds and beyond

isotope

half-life

1018 years 1024 seconds

vanadium-50 0.14 ± 0.04 4.4 ± 1.3

tungsten-180 1.8 ± .2 57 ± 6.3

europium-151 5 160

molybdenum-100

8.5 270

bismuth-209 19 ± 2 600 ± 63

zirconium-96 20 630

cadmium-116 31 ± 4 980 ± 130

1021 years 1027 seconds

selenium-82 0.097 3.1

tellurium-130 0.79 25

germanium-76 1.8 0.057

xenon-136 2.11 ± 0.25 0.067 ± 0.0079

neodymium-150 6.7 0.21

calcium-48 40 1.3

1024 years 1030 seconds

tellurium-128 2.2 ± 0.3 69 ± 9.5

List of elements by stability of isotopesFrom Wikipedia, the free encyclopedia

This is a list of the chemical elements and their isotopes, listed in terms of stability.

See also: List of nuclides, Stable isotope, Primordial nuclide, and Isotope#Nuclear properties and stability

Atomic nuclei consist of protons and neutrons, which attract each other through the nuclear force, while protons repel each other via the electric force due to their positive charge. These two forces compete, leading to some combinations of neutrons and

protons being more stable than others. Neutrons stabilize the nucleus, because they attract each other and protons equally by the strong nuclear force, which helps offset the electrical repulsion between protons. As a result, as the number of protons

increases, an increasing ratio of neutrons to protons is needed to form a stable nucleus.

However, if too many or too few neutrons are present with regard to the optimum ratio, the nucleus becomes unstable and subject to certain types of nuclear decay. Unstable isotopes decay through various radioactive decay pathways, most

commonly alpha decay, beta decay, or electron capture. Many other rare types of decay, such as spontaneous fission or cluster decay are known. (See radioactive decay for details.)

Contents

  [hide] 

1   Overview

2   Elements by number of primordial isotopes

3   Tables

4   Elements with no primordial isotopes

5   See also

6   Footnotes

7   References

[edit]Overview

Isotope half-lives. Note that the darker more stable isotope region departs from the line of protons (Z) = neutrons (N), as the element number Z becomes larger

Of the first 82 elements in the periodic table, 80 have isotopes considered to be stable.[1] Technetium, promethium (atomic numbers 43 and 61, respectively[a]) and all the elements with an atomic number over 82 have only isotopes that are known to

decompose throughradioactive decay. They are not expected to have any stable, undiscovered ones; therefore lead is considered the heaviest stable element. However, it is possible that some isotopes that are presently considered stable will be revealed

to decay with extremely long half-lives (as was the case in 2003 with bismuth-209 which had been previously considered to be stable).[2][3] This list depicts what is agreed upon by the consensus of the scientific community as of 2008. [1]

For each of the 80 stable elements, the number of the stable isotopes is given. Only 90 isotopes are expected to be perfectly stable, and an additional 163 are energetically unstable, but have never been observed to decay. Thus, 253 isotopes ( nuclides)

are stable by definition (including Ta-180m, for which no decay has yet been observed). Those that are found in the future to be radioactive are expected to have half-lives usually longer than 10 22 years (for example, xenon-134).

Of the chemical elements, only one element (tin) has 10 such stable isotopes, one (xenon) has eight isotopes, four have seven isotopes, eight have six isotopes, ten have five isotopes, nine have four isotopes, five have three stable isotopes, 16 have two

stable isotopes, and 26 have a single stable isotope.[1]

Additionally, about 29 nuclides of the 94 naturally-occurring elements have unstable isotopes with a half-life larger than the age of theSolar System (~109 years or more).[b] An additional 6 nuclides have half-lives longer than 80 million years, which is far

less than the age of the solar system, but long enough for some of them to have survived. These 35 radioactive naturally occurring nuclides comprise theradioactive primordial nuclides. The total number of primordial nuclides is then 253 (the stable

nuclides) plus the 35 radioactive primordial nuclides, for a total of 288 primordial nuclides. This number is subject to change if new shorter-lived primordials are identified on Earth.

One of the primordial nuclides is Ta-180m which is predicted to have a half-life in excess of 1015 years, but has never been observed to decay. The even longer half-life of 7.7 x 1024 years of tellurium-128 was measured by a unique method of detecting

radiogenic daughterxenon-128 and is presently the longest known experimentally measured half-life.[4] Another notable example is the only naturally-occurring isotope of bismuth, which has been predicted to be unstable with a very long half-life, but has

only recently been observed to decay. Because of their long half-lives, such isotopes are still found on Earth in various abundances, and together with the stable isotopes they are called primordial isotopes. All the primordial isotopes are given in order of

their decreasing abundance on Earth.[c]. For a list of primordial nuclides in order of half-life, see list of nuclides.

There are 80 elements with at least one stable isotope, but 114 to 118 chemical elements are known, depending on official confirmation (118 are given in this table). All elements to element 98 are found in nature, and the remainder of discovered

elements are artificially produced, with isotopes all known to be highly radioactive with relatively short half-lives (see below). The elements in this list are ordered according to the lifetime of their most stable isotope. [1] Of these, four elements

(bismuth, thorium, uranium and plutonium) are primordial because they have long enough half-lives to still be found on Earth,[d] while all the others are produced either by radioactive decay or are synthesized in laboratories and nuclear reactors. Only 13 of

the 38 known-but-unstable elements (assuming the total number of elements is 118) have isotopes with a half-life of at least 100 years. Every known isotope of the remaining 25 elements is highly radioactive; they are used in academic research and

sometimes in industry and medicine.[e] Some of the heavier elements in the periodic table may be revealed to have yet-undiscovered isotopes with longer lifetimes than those listed here. [f]

About 339 nuclides are found in nature, on Earth. These comprise 253 stable isotopes, and with the addition of the 35 long-lived radioisotopes with half-lives longer than 80 million years, a total of 288 primordial nuclides, as noted above. The nuclides

found naturally comprise not only the 288 primordials, but also include about 51 more short-lived isotopes (defined by a half-life less than 80 million years, too short to have survived from the formation of the Earth) that are daughters of primordial isotopes

(such as radium from uranium); or else are made by energetic natural processes, such as carbon-14 made from atmospheric nitrogen by bombardment from cosmic rays.

[edit]Elements by number of primordial isotopes

An even number of protons or of neutrons are more stable (lower binding energy) because of pairing effects, so even-even nuclides are much more stable than odd-odd. One effect is that there are few stable odd-odd nuclides: in fact only five are stable,

with another four having half-lives longer than a billion years.

Another effect is to prevent beta decay of many even-even nuclides into another even-even nuclide of the same mass number but lower energy, because decay proceeding one step at a time would have to pass through an odd-odd nuclide of higher

energy. (Double beta decay directly from even-even to even-even, skipping over an odd-odd nuclide, is only occasionally possible, and is a process so strongly hindered that it has a half-life greater than a billion times the age of the universe.) This makes

for a larger number of stable even-even nuclides, up to three for some mass numbers, and up to seven for some atomic (proton) numbers and at least four for all stable even-Z elements beyond iron except for strontium.

Since a nucleus with an odd number of protons is relatively less stable, odd-numbered elements tend to have fewer stable isotopes. Of the 26 "monoisotopic" elements that have only a single stable isotope, all but one have an odd atomic number — the

single exception being beryllium.

[edit]Tables

The following tables give the elements with primordial nuclides, which means the element may still be identified on Earth from natural sources, having been present since the Earth was formed out of the solar nebula. Thus, none are shorter-lived daughters

of longer-lived parental primordials, such as radon.

The tables of elements are sorted in order of decreasing number of nuclides associated with each element. (For a list sorted entirely in terms of half-lives of nuclides, with mixing of elements, seeList of nuclides.) Stable and unstable (marked decays)

nuclides are given, with symbols for unstable (radioactive) nuclides in italics. Note that the sorting does not quite give the elements purely in order of stable nuclides, since some elements have a larger number of long-lived unstable nuclides, which place

them ahead of elements with a larger number of stable nuclides. By convention, nuclides are counted as "stable" if they have never been observed to decay by experiment or from observation of decay products (extremely long lived nuclides unstable only

in theory, such as tantalum-180m, are counted as stable).

The first table is for even-atomic numbered elements, which tend to have far more primordial nuclides, due to stability conferred by proton-proton pairing. A second separate table is given for odd-atomic numbered elements, which tend to have far fewer

stable and long-lived (primordial) unstable nuclides.

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

50 tin 10 — 120Sn 118Sn 116Sn 119Sn 117Sn 124Sn 122Sn 112Sn 114Sn 115Sn

54 xenon 8 1 132Xe 129Xe 131Xe 134Xe 136Xe 130Xe 128Xe 124Xe 126Xe

48 cadmium 6 2 114Cd 112Cd 111Cd 110Cd 113Cd 116Cd 106Cd 108Cd

52 tellurium 5 3 130Te 128Te 126Te 125Te 124 Te 122Te 123Te 120Te

62 samarium 5 3 152Sm 154Sm 147Sm 149 Sm 148Sm 150Sm 144Sm14

6Sm

44 ruthenium 7 — 102Ru 104Ru 101Ru 99Ru 100Ru 96Ru 98Ru

66 dysprosium 7 — 164Dy 162Dy 163Dy 161Dy 160Dy 158Dy 156Dy

70 ytterbium 7 — 174Yb 172Yb 173Yb 171Yb 176Yb 170Yb 168Yb

80 mercury 7 — 202Hg 200Hg 199Hg 201Hg 198Hg 204Hg 196Hg

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

42 molybdenum 6 1 98Mo 96Mo 95Mo 92Mo 100Mo 97Mo 94Mo

56 barium 6 1 138Ba 137Ba 136Ba 135Ba 134Ba 132Ba 130Ba

64 gadolinium 6 1 158Gd 160Gd 156Gd 157Gd 155Gd 154Gd 152Gd

76 osmium 6 1 192Os 190Os 189Os 188Os 187Os 186Os 184Os

60 neodymium 5 2 142Nd 144Nd 146Nd 143Nd 145Nd 148Nd 150Nd

36 krypton 6 — 84Kr 86Kr 82Kr 83Kr 80Kr 78Kr

46 palladium 6 — 106Pd 108Pd 105Pd 110Pd 104Pd 102Pd

68 erbium 6 — 166Er 168Er 167Er 170Er 164Er 162Er

20 calcium 5 1 40Ca 44Ca 42Ca 48 Ca 43Ca 46Ca

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

34 selenium 5 1 80Se 78Se 76Se 82 Se 77Se 74Se

72 hafnium 5 1 180Hf 178Hf 177Hf 179Hf 176Hf 174Hf

78 platinum 5 1 195Pt 194Pt 196Pt 198Pt 192Pt 190Pt

22 titanium 5 — 48Ti 46Ti 47Ti 49Ti 50Ti

28 nickel 5 — 58Ni 60Ni 62Ni 61Ni 64Ni

30 zinc 5 — 64Zn 66Zn 68Zn 67Zn 70Zn

32 germanium 4 1 74Ge 72Ge 70Ge 73Ge 76Ge

40 zirconium 4 1 90Zr 94Zr 92Zr 91Zr 96 Zr

74 tungsten 4 1 184W 186W 182W 183W 180W

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

16 sulfur 4 — 32S 34S 33S 36S

24 chromium 4 — 52Cr 53Cr 50Cr 54Cr

26 iron 4 — 56 Fe 54Fe 57Fe 58Fe

38 strontium 4 — 88Sr 86Sr 87Sr 84Sr

58 cerium 4 — 140Ce 142Ce 138Ce 136Ce

82 lead 4 — 208Pb 206Pb 207Pb 204Pb

8 oxygen 3 — 16O 18 O 17O

10 neon 3 — 20Ne 22Ne 21Ne

12 magnesium 3 — 24Mg 26Mg 25Mg

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

14 silicon 3 — 28Si 29Si 30Si

18 argon 3 — 40Ar 36Ar 38Ar

2 helium 2 — 4 He 3 He

6 carbon 2 — 12 C 13 C

92 uranium 0 2 238 U [d] 235 U

4 beryllium 1 — 9Be

90 thorium 0 1 232 Th

94 plutonium 0 1 244 Pu

Primordial isotopes of odd-Z elements

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

Z Element Stab Decunstable: italics

odd N on pink

19 potassium 2 1 39K 41K 40K

1 hydrogen 2 — 1 H 2 H

3 lithium 2 — 7Li 6Li

5 boron 2 — 11B 10B

7 nitrogen 2 — 14 N 15 N

17 chlorine 2 — 35Cl 37Cl

29 copper 2 — 63Cu 65Cu

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

31 gallium 2 — 69Ga 71Ga

35 bromine 2 — 79Br 81Br

47 silver 2 — 107Ag 109Ag

51 antimony 2 — 121Sb 123Sb

73 tantalum 2 — 181Ta 180mTa

77 iridium 2 — 193Ir 191Ir

81 thallium 2 — 205Tl 203Tl

23 vanadium 1 1 51V 50V

37 rubidium 1 1 85Rb 87Rb

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

49 indium 1 1 115In 113In

57 lanthanum 1 1 139La 138La

63 europium 1 1 153Eu 151Eu

71 lutetium 1 1 175Lu 176Lu

75 rhenium 1 1 187Re 185Re

9 fluorine 1 — 19F

11 sodium 1 — 23Na

13 aluminium 1 — 27Al

15 phosphorus 1 — 31P

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

21 scandium 1 — 45Sc

25 manganese 1 — 55Mn

27 cobalt 1 — 59Co

33 arsenic 1 — 75As

39 yttrium 1 — 89Y

41 niobium 1 — 93Nb

45 rhodium 1 — 103Rh

53 iodine 1 — 127I

55 caesium 1 — 133Cs

Primordial isotopes (in order of decreasing abundance on Earth[c]) of even-Z elements

Z ElementStabli

[1]

Decays[b][1]

unstable in italics[b]

odd neutron number on pink

59praseodymiu

m1 — 141Pr

65 terbium 1 — 159Tb

67 holmium 1 — 165Ho

69 thulium 1 — 169Tm

79 gold 1 — 197Au

83 bismuth 0 1 209 Bi

[edit]Elements with no primordial isotopes

No primordial isotopesLongest lived isotope in years/days

Z Element t 1/2 of[g][1]

Longestlived

isotope

96 curium 1.56×107 a 247Cm

43 technetium 4.2×106 a 98Tc[a]

93 neptunium 2.144×106 a 237Np

91 protactinium 32,760 a 231Pa

95 americium 7,370 a 243 Am

88 radium 1,602 a 226Ra

97 berkelium 1,380 a 247Bk

98 californium 898 a 251Cf

No primordial isotopesLongest lived isotope in years/days

Z Element t 1/2 of[g][1]

Longestlived

isotope

84 polonium 103 a 209Po

89 actinium 21.77 a 227Ac

61 promethium 17.7 a 145Pm[a]

99 einsteinium 1.29 a 252Es[f]

100 fermium 100.5 d 257Fm[f]

101 mendelevium 51.5 d 258Md[f]

86 radon 3.82 d 222Rn

105 dubnium 1.3 d 268Db[f]

No primordial isotopes

No primordial isotopesLongest lived isotope in years/days

Z Element t 1/2 of[g][1]

Longestlived

isotope

Longest lived isotope in hour/min/sec

Z Element t 1/2 of[g][1]

Longestlived

isotope

103 lawrencium 10 h[h]264Lr[f]

85 astatine 8.1 h 210At

107 bohrium 1.5 h[h] 273Bh[f]

104 rutherfordium 1.3 h 265Rf[f]

106 seaborgium 1 h[h] 272Sg[f]

108 hassium 1 h[h] 276Hs[f]

No primordial isotopesLongest lived isotope in years/days

Z Element t 1/2 of[g][1]

Longestlived

isotope

102 nobelium 58 min 259No[f]

87 francium 22.0 min 223Fr[f]

113 ununtrium [i] 20 min[h] 287Uut[f]

111 roentgenium 10 min[h] 283Rg[f]

109 meitnerium 6 min[h] 279Mt[f]

115 ununpentium [i] 1 min[h] 291Uup[f]

112 copernicium 34 s 285Cn[f]

110 darmstadtium 10 s 278Ds[f]

114 flerovium 2.7 s 289Fl[f]

No primordial isotopesLongest lived isotope in years/days

Z Element t 1/2 of[g][1]

Longestlived

isotope

116 livermorium 5.3×10−2 s 293Lv[f]

117 ununseptium [i] 7.8×10−2 s 294Uus[f]

118 ununoctium [i] 8.9×10−4 s 294Uuo[f]

Periodic table colored according to the number of stable isotopes. Elements with odd atomic numbers have at most one or two stable isotopes, while elements up to lead with even atomic numbers all have three or more stable isotopes, except for the first three: helium, beryllium, and carbon.

Periodic table with elements colored according to the half-life of their most stable isotope.

  Elements which contain at least one stable isotope;

  Radioactive elements: the most stable isotope is very long-lived, with half-life of over four million years;

  Radioactive elements: the most stable isotope has half-life between 800 and 34,000 years;

  Radioactive elements: the most stable isotope has half-life between one day and 103 years;

  Highly radioactive elements: the most stable isotope has half-life between one minute and one day;

  Extremely radioactive elements: the most stable isotope has half-life less than a minute. Very little is known about these elements due to their extreme instability and radioactivity.

Radioactive decayFrom Wikipedia, the free encyclopedia

For particle decay in a more general context, see Particle decay. For more information on hazards of various kinds of radiation from decay, see Ionizing radiation.

"Radioactive" redirects here. For other uses, see Radioactive (disambiguation).

This article needs additional citations for verification. Please help improve this article by adding citations to reliable sources. Unsourced material may be challenged and removed. (April 2010)

Alpha decay is one example type ofradioactive decay, in which an atomic nucleus emits an alpha particle, and thereby transforms (or 'decays') into an atom with amass number 4 less and atomic number 2 less. Many other types of decays are possible.

Nuclear physics

Radioactive decay

Nuclear fission

Nuclear fusion

Classical decays[show]

Advanced decays[show]

Emission processes[show]

Capturing[show]

High energy processes[show]

Nucleosynthesis [show]

Scientists[show]

V

T

E

Radioactive decay is the process by which an atomic nucleus of an unstable atom loses energy by emitting ionizing particles (ionizing radiation). There are many different types of radioactive decay (see

table below). A decay, or loss of energy, results when an atom with one type of nucleus, called the parent radionuclide, transforms to an atom with a nucleus in a different state, or to a different nucleus

containing different numbers of protonsand neutrons. Either of these products is named the daughter nuclide. In some decays the parent and daughter are different chemical elements, and thus the decay

process results in nuclear transmutation (creation of an atom of a new element).

The first decay processes to be discovered were alpha decay, beta decay, and gamma decay. Alpha decay occurs when the nucleus ejects an alpha particle (helium nucleus). This is the most common

process of emitting nucleons, but in rarer types of decays, nuclei can eject protons, or specific nuclei of other elements (in the process called cluster decay). Beta decay occurs when the nucleus emits

an electron or positron and a type ofneutrino, in a process that changes a proton to a neutron or the other way around. The nucleus may capture an orbiting electron, converting a proton into an neutron

(electron capture). All of these processes result in nuclear transmutation.

By contrast, there exist radioactive decay processes that do not result in transmutation. The energy of an excited nucleus may be emitted as agamma ray in gamma decay, or used to eject an orbital

electron by interaction with the excited nucleus in a process called internal conversion. Radioisotopes occasionally emit neutrons, and this results in a change in an element from one isotope to another.

One type of radioactive decay results in products which are not defined, but appear in a range of "pieces" of the original nucleus. This decay is called spontaneous fission. This decay happens when a large

unstable nucleus spontaneously splits into two (and occasionally three) smaller daughter nuclei, and usually emits gamma rays, neutrons, or other particles as a consequence.

Radioactive decay is a stochastic (i.e., random) process at the level of single atoms, in that, according to quantum theory, it is impossible to predict when a particular atom will decay.[1] However, the chance

that a given atom will decay is constant over time. For a large number of atoms, the decay rate for the collection is computable from the measured decay constants of the nuclides (or equivalently from

the half-lifes).

Contents

[hide]

1 Natural origin

2 Decay phenomena

3 Discovery

4 Danger of radioactive substances

5 Types of decay

6 Decay modes in table form

7 Decay chains and multiple modes

8 Occurrence and applications

9 Radioactive decay rates

o 9.1 Units of radioactivity measurements

10 Mathematics of radioactive decay

o 10.1 Universal law of radioactive decay

10.1.1 One-decay process

10.1.2 Chain-decay processes

10.1.3 Alternative decay modes

o 10.2 Corollaries of the decay laws

o 10.3 Decay timing: definitions and relations

10.3.1 Time constant and mean-life

10.3.2 Half-life

o 10.4 Example

11 Changing decay rates

12 See also

13 Notes

14 References

15 External links

[edit]Natural origin

Primordial nuclides found in the earth are residues from ancient supernova explosions which occurred before the formation of the solar system. They are the long-lived fraction of radionuclides surviving in

the primordial solar nebula through planet accretion until the present. The naturally occurring short-lived radiogenic radionuclides found in rocks are the daughters of these radiactiveprimordial nuclides.

Another minor source of naturally occurring radioactive nuclides are cosmogenic nuclides, formed by cosmic ray bombardment of material in the Earth's atmosphere or crust. For a summary table showing

the number of stable nuclides and of radioactive nuclides in each category, see radionuclide. Radionuclides can also be produced artificially e.g. using particle accelerators or nuclear reactors.

[edit]Decay phenomena

The trefoil symbol is used to indicate radioactive material.

The new radioactive danger symbol is used to indicate dangerous ionizing radioactive material.

The neutrons and protons that constitute nuclei, as well as other particles that approach close enough to them, are governed by several interactions. The strong nuclear force, not observed at the

familiar macroscopic scale, is the most powerful force over subatomic distances. The electrostatic force is almost always significant, and, in the case of beta decay, the weak nuclear force is also involved.

The interplay of these forces produces a number of different phenomena in which energy may be released by rearrangement of particles in the nucleus, or else the change of one type of particle into others.

These rearrangements and transformations may be hindered energetically, so that they do not occur immediately. In certain cases, random quantum vacuum fluctuations are theorized to promote relaxation

to a lower energy state (the "decay") in a phenomenon known as quantum tunneling. Radioactive decay half-life of nuclides has been measured over timescales of 55 orders of magnitude, from 2.3 x 10-

23 second (for hydrogen-7) to 6.9 x 1031 seconds (for tellurium-128)[2]. The limits of these timescales are set by the sensitivity of instrumentation only, and there are no known natural limits to how brief or long

a decay half life for radioactive decay of a radionuclide may be.

The decay process, like all hindered energy transformations, may be analogized by a snowfield on a mountain. While friction between the ice crystals may be supporting the snow's weight, the system is

inherently unstable with regard to a state of lower potential energy. A disturbance would thus facilitate the path to a state of greater entropy: The system will move towards the ground state, producing heat,

and the total energy will be distributable over a larger number of quantum states. Thus, an avalanche results. The total energy does not change in this process, but, because of the law of entropy,

avalanches happen only in one direction and that is toward the "ground state" — the state with the largest number of ways in which the available energy could be distributed.

Such a collapse (a decay event) requires a specific activation energy. For a snow avalanche, this energy comes as a disturbance from outside the system, although such disturbances can be arbitrarily

small. In the case of an excited atomic nucleus, the arbitrarily small disturbance comes from quantum vacuum fluctuations. A radioactive nucleus (or any excited system in quantum mechanics) is unstable,

and can, thus, spontaneously stabilize to a less-excited system. The resulting transformation alters the structure of the nucleus and results in the emission of either a photon or a high-velocity particle that

has mass (such as an electron, alpha particle, or other type).

[edit]Discovery

Radioactivity was discovered in 1896 by the French scientist Henri Becquerel, while working on phosphorescent materials. These materials glow in the dark after exposure to light, and he suspected that the

glow produced in cathode ray tubes by X-rays might be associated with phosphorescence. He wrapped a photographic plate in black paper and placed various phosphorescent salts on it. All results were

negative until he used uranium salts. The result with these compounds was a blackening of the plate. These radiations were called Becquerel Rays.

Pierre and Marie Curie in their Paris laboratory, before 1907

It soon became clear that the blackening of the plate had nothing to do with phosphorescence, because the plate blackened when the mineral was in the dark. Non-phosphorescent salts of uranium and

metallic uranium also blackened the plate. It was clear that there is a form of radiation that could pass through paper that was causing the plate to become black.

At first it seemed that the new radiation was similar to the then recently discovered X-rays. Further research by Becquerel, Ernest Rutherford, Paul Villard,Pierre Curie, Marie Curie, and others discovered

that this form of radioactivity was significantly more complicated. Different types of decay can occur, producing very different types of radiation. Rutherford was the first to realize that they all occur with the

same mathematical exponential formula (see below), and Rutherford and his student Frederick Soddy were first to realize that many decay processes resulted in the transmutation of one element to

another. Subsequently, the radioactive displacement law of Fajans and Soddy was formulated to describe the products of alpha and beta decay.

The early researchers also discovered that many other chemical elements besides uranium have radioactive isotopes. A systematic search for the total radioactivity in uranium ores also guided Marie

Curie to isolate a new element polonium and to separate a new element radium from barium. The two elements' chemical similarity would otherwise have made them difficult to distinguish.

[edit]Danger of radioactive substances

Main article: Ionizing radiation

The danger classificationsign of radioactive materials

Alpha particles may be completely stopped by a sheet of paper, beta particlesby aluminum shielding. Gamma rays can only be reduced by much more substantial mass, such as a very thick layer oflead.

Different types of decay of a radionuclide. Vertical: atomic number Z, Horizontal: neutron number N

The dangers of radioactivity and radiation were not immediately recognized. Acute effects of radiation were first observed in the use of X-rays when electrical engineer and physicist Nikola Tesla intentionally

subjected his fingers to X-rays in 1896.[3] He published his observations concerning the burns that developed, though he attributed them to ozone rather than to X-rays. His injuries later healed.

The genetic effects of radiation, including the effect of cancer risk, were recognized much later. In 1927, Hermann Joseph Muller published research showing genetic effects, and in 1946 was awarded

the Nobel prize for his findings.

Before the biological effects of radiation were known, many physicians and corporations began marketing radioactive substances as patent medicine, glow-in-the-dark pigments. Examples were

radium enema treatments, and radium-containing waters to be drunk as tonics. Marie Curie protested this sort of treatment, warning that the effects of radiation on the human body were not well understood.

Curie later died from aplastic anemia, likely caused by exposure to ionizing radiation. By the 1930s, after a number of cases of bone necrosis and death of enthusiasts, radium-containing medicinal products

had been largely removed from the market (radioactive quackery).

[edit]Types of decay

Types of radioactive decay related to N and Z numbers

As for types of radioactive radiation, it was found that an electric or magnetic field could split such emissions into three types of beams. The rays were given thealphabetic names alpha, beta, and gamma, in

order of their ability to penetrate matter. While alpha decay was seen only in heavier elements (atomic number 52,tellurium, and greater), the other two types of decay were seen in all of the elements.

Spontaneous decay is evident in elements of atomic number ninety or greater.

In analyzing the nature of the decay products, it was obvious from the direction of electromagnetic forces induced upon the radiations by external magnetic and electric fields that alpha particles carried a

positive charge, beta particles carried a negative charge, and gamma rays were neutral. From the magnitude of deflection, it was clear that alpha particles were much more massive than beta particles.

Passing alpha particles through a very thin glass window and trapping them in a discharge tube allowed researchers to study the emission spectrum of the resulting gas, and ultimately prove that alpha

particles are helium nuclei. Other experiments showed the similarity between classical beta radiation and cathode rays: They are both streams of electrons. Likewise gamma radiation and X-rays were found

to be similar high-energy electromagnetic radiation.

The relationship between the types of decays also began to be examined: For example, gamma decay was almost always found associated with other types of decay, and occurred at about the same time,

or afterward. Gamma decay as a separate phenomenon (with its own half-life, now termed isomeric transition), was found in natural radioactivity to be a result of the gamma decay of excited

metastable nuclear isomers, which were in turn created from other types of decay.

Although alpha, beta, and gamma radiations were found most commonly, other types of decay were eventually discovered. Shortly after the discovery of thepositron in cosmic ray products, it was realized

that the same process that operates in classical beta decay can also produce positrons (positron emission). In an analogous process, instead of emitting positrons and neutrinos, some proton-rich nuclides

were found to capture their own atomic electrons (electron capture), and emit only a neutrino (and usually also a gamma ray). Each of these types of decay involves the capture or emission of nuclear

electrons or positrons, and acts to move a nucleus toward the ratio of neutrons to protons that has the least energy for a given total number of nucleons (neutrons plus protons).

A theoretical process of positron capture (analogous to electron capture) is possible in antimatter atoms, but has not been observed since the complex antimatter atoms are not available. [4] This would

required antimatter atoms at least as complex as beryllium-7, which is the lightest known isotope of normal matter to undergo decay by electron capture.

Shortly after the discovery of the neutron in 1932, Enrico Fermi realized that certain rare decay reactions yield neutrons as a decay particle (neutron emission). Isolated proton emission was eventually

observed in some elements. It was also found that some heavy elements may undergo spontaneous fission into products that vary in composition. In a phenomenon called cluster decay, specific

combinations of neutrons and protons other than alpha particles (helium nuclei) were found to be spontaneously emitted from atoms.

Other types of radioactive decay that emit previously seen particles were found, but by different mechanisms. An example is internal conversion, which results in electron and sometimes high-energy photon

emission, even though it involves neither beta nor gamma decay. This type of decay (like isomeric transition gamma decay) does not transmute one element to another.

Rare events that involve a combination of two beta-decay type events happening simultaneously (see below) are known. Any decay process that does not violate conservation of energy or momentum laws

(and perhaps other particle conservation laws) is permitted to happen, although not all have been detected. An interesting example (discussed in a final section) is bound state beta decay of rhenium-187. In

this process, an inverse of electron capture, beta electron-decay of the parent nuclide is not accompanied by beta electron emission, because the beta particle has been captured into the K-shell of the

emitting atom. An antineutrino, however, is emitted.

[edit]Decay modes in table form

Radionuclides can undergo a number of different reactions. These are summarized in the following table. A nucleus with mass number A and atomic number Z is represented as (A, Z). The column

"Daughter nucleus" indicates the difference between the new nucleus and the original nucleus. Thus, (A − 1, Z) means that the mass number is one less than before, but the atomic number is the same as

before.

Mode of decay Participating particles Daughter nucleus

Decays with emission of nucleons:

Alpha decay An alpha particle (A = 4, Z = 2) emitted from nucleus (A − 4, Z − 2)

Proton emission A proton ejected from nucleus (A − 1, Z − 1)

Neutron emission A neutron ejected from nucleus (A − 1, Z)

Double proton emission Two protons ejected from nucleus simultaneously (A − 2, Z − 2)

Spontaneous fission Nucleus disintegrates into two or more smaller nuclei and other particles —

Cluster decay Nucleus emits a specific type of smaller nucleus (A1, Z1) smaller than, or larger than, an alpha particle(A − A1, Z − Z1) + (A1, Z1)

Different modes of beta decay:

β − decay A nucleus emits an electron and an electron antineutrino (A, Z + 1)

Positron emission (β + decay ) A nucleus emits a positron and an electron neutrino (A, Z − 1)

Electron capture A nucleus captures an orbiting electron and emits a neutrino; the daughter nucleus is left in an excited unstable state (A, Z − 1)

Bound state beta decayA nucleus beta decays to electron and antineutrino, but the electron is not emitted, as it is captured into an empty K-shell; the daughter nucleus is left in an excited and unstable state. This process is suppressed except in ionized atoms that have K-shell vacancies.

(A, Z + 1)

Double beta decay A nucleus emits two electrons and two antineutrinos (A, Z + 2)

Double electron capture A nucleus absorbs two orbital electrons and emits two neutrinos – the daughter nucleus is left in an excited and unstable state (A, Z − 2)

Electron capturewith positron emission

A nucleus absorbs one orbital electron, emits one positron and two neutrinos (A, Z − 2)

Double positron emission A nucleus emits two positrons and two neutrinos (A, Z − 2)

Transitions between states of the same nucleus:

Isomeric transition Excited nucleus releases a high-energy photon (gamma ray) (A, Z)

Internal conversion Excited nucleus transfers energy to an orbital electron, which is subsequently ejected from the atom (A, Z)

Radioactive decay results in a reduction of summed rest mass, once the released energy (the disintegration energy) has escaped in some way (for example, the products might be captured and cooled, and

the heat allowed to escape). Although decay energy is sometimes defined as associated with the difference between the mass of the parent nuclide products and the mass of the decay products, this is true

only of rest mass measurements, where some energy has been removed from the product system. This is true because the decay energy must always carry mass with it, wherever it appears (see mass in

special relativity) according to the formula E = mc2. The decay energy is initially released as the energy of emitted photons plus the kinetic energy of massive emitted particles (that is, particles that have rest

mass). If these particles come to thermal equilibrium with their surroundings and photons are absorbed, then the decay energy is transformed to thermal energy, which retains its mass.

Decay energy therefore remains associated with a certain measure of mass of the decay system invariant mass. The energy of photons, kinetic energy of emitted particles, and, later, the thermal energy of

the surrounding matter, all contribute to calculations of invariant mass of systems. Thus, while the sum of rest masses of particles is not conserved in radioactive decay, the systemmass and

system invariant mass (and also the system total energy) is conserved throughout any decay process.

[edit]Decay chains and multiple modes

The daughter nuclide of a decay event may also be unstable (radioactive). In this case, it will also decay, producing radiation. The resulting second daughter nuclide may also be radioactive. This can lead to

a sequence of several decay events. Eventually, a stable nuclide is produced. This is called a decay chain (see this article for specific details of important natural decay chains).

Gamma-ray energy spectrum of uranium ore (inset). Gamma-rays are emitted by decayingnuclides, and the gamma-ray energy can be used to characterize the decay (which nuclide is decaying to which). Here, using the gamma-ray spectrum, several nuclides that are typical of the decay chain of  238U

have been identified: 226Ra,214Pb, 214Bi.

An example is the natural decay chain of 238U, which is as follows:

decays, through alpha-emission, with a half-life of 4.5 billion years to thorium-234

which decays, through beta-emission, with a half-life of 24 days to protactinium-234

which decays, through beta-emission, with a half-life of 1.2 minutes to uranium-234

which decays, through alpha-emission, with a half-life of 240 thousand years to thorium-230

which decays, through alpha-emission, with a half-life of 77 thousand years to radium-226

which decays, through alpha-emission, with a half-life of 1.6 thousand years to radon-222

which decays, through alpha-emission, with a half-life of 3.8 days to polonium-218

which decays, through alpha-emission, with a half-life of 3.1 minutes to lead-214

which decays, through beta-emission, with a half-life of 27 minutes to bismuth-214

which decays, through beta-emission, with a half-life of 20 minutes to polonium-214

which decays, through alpha-emission, with a half-life of 160 microseconds to lead-210

which decays, through beta-emission, with a half-life of 22 years to bismuth-210

which decays, through beta-emission, with a half-life of 5 days to polonium-210

which decays, through alpha-emission, with a half-life of 140 days to lead-206, which is a stable nuclide.

Some radionuclides may have several different paths of decay. For example, approximately 36% of bismuth-212 decays, through alpha-emission, to thallium-208 while approximately 64% of bismuth-

212 decays, through beta-emission, to polonium-212. Both the thallium-208 and thepolonium-212 are radioactive daughter products of bismuth-212, and both decay directly to stable lead-208.

[edit]Occurrence and applications

According to the Big Bang theory, stable isotopes of the lightest five elements (H, He, and traces of Li, Be, and B) were produced very shortly after the emergence of the universe, in a process called Big

Bang nucleosynthesis. These lightest stable nuclides (including deuterium) survive to today, but any radioactive isotopes of the light elements produced in the Big Bang (such astritium) have long since

decayed. Isotopes of elements heavier than boron were not produced at all in the Big Bang, and these first five elements do not have any long-lived radioisotopes. Thus, all radioactive nuclei are, therefore,

relatively young with respect to the birth of the universe, having formed later in various other types of nucleosynthesis in stars (in particular, supernovae), and also during ongoing interactions between stable

isotopes and energetic particles. For example, carbon-14, a radioactive nuclide with a half-life of only 5730 years, is constantly produced in Earth's upper atmosphere due to interactions between cosmic

rays and nitrogen.

Nuclides that are produced by radioactive decay are called radiogenic nuclides, whether they themselves are stable or not. There exist stable radiogenic nuclides that were formed from short-livedextinct

radionuclides in the early solar system.[5][6] The extra presence of these stable radiogenic nuclides (such as Xe-129 from primordial I-129) against the background of primordial stable nuclides can be inferred

by various means.

Radioactive decay has been put to use in the technique of radioisotopic labeling, which is used to track the passage of a chemical substance through a complex system (such as a livingorganism). A sample

of the substance is synthesized with a high concentration of unstable atoms. The presence of the substance in one or another part of the system is determined by detecting the locations of decay events.

On the premise that radioactive decay is truly random (rather than merely chaotic), it has been used in hardware random-number generators. Because the process is not thought to vary significantly in

mechanism over time, it is also a valuable tool in estimating the absolute ages of certain materials. For geological materials, the radioisotopes and some of their decay products become trapped when a rock

solidifies, and can then later be used (subject to many well-known qualifications) to estimate the date of the solidification. These include checking the results of several simultaneous processes and their

products against each other, within the same sample. In a similar fashion, and also subject to qualification, the rate of formation of carbon-14 in various eras, the date of formation of organic matter within a

certain period related to the isotope's half-life may be estimated, because the carbon-14 becomes trapped when the organic matter grows and incorporates the new carbon-14 from the air. Thereafter, the

amount of carbon-14 in organic matter decreases according to decay processes that may also be independently cross-checked by other means (such as checking the carbon-14 in individual tree rings, for

example).

[edit]Radioactive decay rates

The decay rate, or activity, of a radioactive substance are characterized by:

Constant quantities:

The half-life—t1/2, is the time taken for the activity of a given amount of a radioactive substance to decay to half of its initial value; see List of nuclides.

The mean lifetime— τ, "tau" the average lifetime of a radioactive particle before decay.

The decay constant— λ, "lambda" the inverse of the mean lifetime.

Although these are constants, they are associated with statistically random behavior of populations of atoms. In consequence predictions using these constants are less accurate for small number of atoms.

In principle the reciprocal of any number greater than one— a half-life, a third-life, or even a (1/√2)-life—can be used in exactly the same way as half-life; but the half-life t1/2 is adopted as the standard time

associated with exponential decay.

Time-variable quantities:

Total activity— A, is number of decays per unit time of a radioactive sample.

Number of particles—N, is the total number of particles in the sample.

Specific activity—SA, number of decays per unit time per amount of substance of the sample at time set to zero (t = 0). "Amount of substance" can be the mass, volume or moles of the initial sample.

These are related as follows:

where N0 is the initial amount of active substance — substance that has the same percentage of unstable particles as when the substance was formed.

[edit]Units of radioactivity measurements

The SI unit of radioactive activity is the becquerel (Bq), in honor of the scientist Henri Becquerel. One Bq is defined as one transformation (or decay or disintegration) per second. Since

sensible sizes of radioactive material contains many atoms, a Bq is a tiny measure of activity; amounts giving activities on the order of GBq (gigabecquerel, 1 x 109 decays per second) or TBq

(terabecquerel, 1 x 1012 decays per second) are commonly used.

Another unit of radioactivity is the curie, Ci, which was originally defined as the amount of radium emanation (radon-222) in equilibrium with one gram of pure radium, isotope Ra-226. At

present it is equal, by definition, to the activity of any radionuclide decaying with a disintegration rate of 3.7 × 1010 Bq, so that 1 curie (Ci) = 3.7 × 1010 Bq. The use of Ci is currently discouraged

by the SI. Low activities are also measured in disintegrations per minute (dpm).

[edit]Mathematics of radioactive decay

For the mathematical details of exponential decay in general context, see exponential decay.

For related derivations with some further details, see half-life.

For the analogous mathematics in 1st order chemical reactions, see Consecutive reactions.

[edit]Universal law of radioactive decay

Radioactivity is one very frequent example of exponential decay. The law describes the statistical behavior of a large number of nuclides, rather than individual ones. In the following formalism,

the number of nuclides or nuclide population N, is of course a discrete variable (a natural number) - but for any physical sample N is so large (amounts of L = 1023, avagadro's constant) that it

can be treated as a continuous variable. Differential calculus is needed to set up differential equations for modelling the behaviour of the nuclear decay.

[edit]One-decay process

Consider the case of a nuclide A decaying into another B by some process A → B (emission of other particles, like electron neutrinos ν

e and electrons e– in beta decay, are irrelevant in what follows). The decay of an unstable nucleus is entirely random and it is impossible to predict when a particular atom will decay. [1] However,

it is equally likely to decay at any time. Therefore, given a sample of a particular radioisotope, the number of decay events −dN expected to occur in a small interval of time dt is proportional to

the number of atoms present N, that is

Particular radionuclides decay at different rates, so each has its own decay constant λ. The expected decay −dN/N is proportional to an increment of time, dt:

The negative sign indicates that N decreases as time increases, as each decay event follows one after another. The solution to this first-order differential equation is the function:

where N0 is the value of N at time t = 0.

This equation is of particular interest; the behaviour of numerous important quantities can be found from it (see below). Although the parent decay distribution follows an

exponential, observations of decay times will be limited by a finite integer number of N atoms and follow Poisson statistics as a consequence of the random nature of the process.

We have for all time t:

where Ntotal is the constant number of particles throughout the decay process, clearly equal to the initial number of A nuclides since this is the initial substance.

If the number of non-decayed A nuclei is:

then the number of nuclei of B, i.e. number of decayed A nuclei, is

[edit]Chain-decay processes

Chain of two decays

Now consider the case of a chain of two decays: one nuclide A decaying into another B by one process, then B decaying into another C by a second process, i.e. A →

B → C. The previous equation cannot be applied to a decay chain, but can be generalized as follows. The decay rate of B is proportional to the number of nuclides

of B present, so again we have:

but care must be taken. Since A decays into B, then B decays into C, the activity of A adds to the total number of B nuclides in the present

sample, before those B nuclides decay and reduce the number of nuclides leading to the later sample. In other words, the number of second generation

nuclei B increases as a result of the first generation nuclei decay of A, and decreases as a result of its own decay into the third generation nuclei C.[7] The

proportionality becomes an equation:

adding the increasing (and correcting) term obtains the law for a decay chain for two nuclides:

The equation is not

since this implies the number of atoms of B is only decreasing as time increases, which is not the case. The rate of change of NB, that is dNB/dt, is

related to the changes in the amounts of Aand B, NB can increase as B is produced from A and decrease as B produces C.

Re-writing using the previous results:

The subscripts simply refer to the respective nuclides, i.e. NA is the number of nuclides of type A, NA0 is the initial number of nuclides of

type A, λA is the decay constant for A - and similarly for nuclide B. Solving this equation for NB gives:

Naturally this equation reduces to the previous solution, in the case B is a stable nuclide (λB = 0):

as shown above for one decay. The solution can be found by the integration factor method, where the integrating factor is eλB

t. This

case is perhaps the most useful, since it can derive both the one-decay equation (above) and the equation for multi-decay chains

(below) more directly.

Chain of any number of decays

For the general case of any number of consecutive decays in a decay chain, i.e. A1 → A2 ··· → Ai ··· → AD, where D is the number of

decays and i is a dummy index (i = 1, 2, 3, ...D), each nuclide population can be found in terms of the previous population. In this case N2 =

0, N3 = 0,..., ND = 0. Using the above result in a recursive form:

The general solution to the recursive problem are given by Bateman's equations[8]:

[edit]Alternative decay modes

In all of the above examples, the initial nuclide decays into only one product. Consider the case of one initial nuclide which

can decay into two products, that is A → B + C. We have for all timet:

in which,

so the relations follow in parallel:

indicating that the total decay constant is that of A, given by:

Solving this equation for NA:

When measuring the production of one nuclide, one can only observe the total decay constant λA.

The decay constants λB and λC determine the probability for the decay to result in products Bor C as

follows:

These perhaps seemingly disjionted results are consistent:

[edit]Corollaries of the decay laws

The solutions to the above differential equations are sometimes written using quantities

related to the number of nuclide particles N in a sample, where L is Avogadro's

constant,6.023×1023, andAr is the relative atomic mass number, and the amount of the

substance is in moles.

The activity: A = λN.

The amount of substance: n = N/L.

The mass: M = Arn = ArN/L.

Collecting these results together for convenience: N = A/λ = Ln = LM/Ar.

Equivalent ways to write the decay solutions, then, are as follows:

One-decay processes

The solution

can be written:

Notice how we can simply replace each quantity (on both sides of the

equation), since they are directly proportional to N and so the constants

cancel (constant at least for a particular nuclide).

Chain-decay processes

For the two-decay chain,

its almost as simple:

[edit]Decay timing: definitions and relations

[edit]Time constant and mean-life

For the one-decay solution A → B:

the equation indicates that the decay

constant λ has units of t-1, and can thus also be

represented as 1/τ, where τ is a characteristic

time of the process called the time constant.

In a radioactive decay process, this time constant

is also the mean lifetime for decaying atoms.

Each atom "lives" for a finite amount of time

before it decays, and it may be shown that this

mean lifetime is the arithmetic mean of all the

atoms' lifetimes, and that it is τ, which again is

related to the decay constant as follows:

This form is also true for two-decay

processes simultaneously A → B + C, inserting

the equivalent values of decay constants (as

given above)

into the decay solution leads to:

Simulation of many identical atoms undergoing

radioactive decay, starting with either 4 atoms (left)

or 400 (right). The number at the top indicates how

many half-liveshave elapsed. Note the law of large

numbers: with more atoms, the overall decay is less

random.

[edit]Half-life

A more commonly used parameter

is the half-life. Given a sample of a

particular radionuclide, the half-life

is the time taken for half the

radionuclide's atoms to decay. For

the case of one-decay nuclear

reactions:

the half-life is related to the

decay constant as follows:

set N = N0/2 and t = T1/2 to obtain

This relationship between

the half-life and the decay

constant shows that highly

radioactive substances are

quickly spent, while those

that radiate weakly endure

longer. Half-lives of known

radionuclides vary widely,

from more than 10 19   years ,

such as for the very nearly

stable nuclide 209Bi, to

10−23 seconds for highly

unstable ones.

The factor of ln(2) in the

above relations results

from the fact that concept

of "half-life" is merely a

way of selecting a different

base other than the natural

base e for the lifetime

expression. The time

constant τ is the e -1 -life,

the time until only

1/e remains, about 36.8%,

rather than the 50% in the

half-life of a radionuclide.

Thus, τ is longer than t1/2.

The following equation can

be shown to be valid:

Since radioactive

decay is exponential

with a constant

probability, each

process could as

easily be described

with a different

constant time period

that (for example)

gave its "(1/3)-life"

(how long until only

1/3 is left) or "(1/10)-

life" (a time period

until only 10% is left),

and so on. Thus, the

choice of τ and t1/2 for

marker-times, are only

for convenience, and

from convention. They

reflect a fundamental

principle only in so

much as they show

that the same

proportion of a given

radioactive substance

will decay, during any

time-period that one

chooses.

Mathematically,

the nth life for the

above situation would

be found in the same

way as above—by

setting N = N0/n, and

substituting into the

decay solution to

obtain

[edit]Example

A sample of 14C,

whose half-life is

5730 years, has a

decay rate of 14

disintegration per

minute (dpm) per

gram of

natural carbon.

An artefact is

found to have

radioactivity of 4

dpm per gram of

its present C, how

old is the

artefact?

Using the above

equation, we

have:

where: 

 years,

 years.

[

edit]

Ch

ang

ing

dec

ay

rate

s

The

radi

oact

ive

dec

ay

mod

es

of el

ectr

on

capt

ure 

and 

inter

nal

con

vers

ion 

are

kno

wn

to

be

slig

htly

sen

sitiv

e to

che

mic

al

and

envi

ron

men

tal

effe

cts

whic

h

cha

nge

the

elec

troni

c

stru

ctur

e of

the

ato

m,

whic

h in

turn

affe

cts

the

pres

enc

e

of 1

s an

d 2s 

elec

tron

s

that

parti

cipa

te in

the

dec

ay

proc

ess.

A

sma

ll

num

ber

of

mos

tly

light

nucl

ides

are

affe

cted

.

For

exa

mpl

e, c

hem

ical

bon

ds c

an

affe

ct

the

rate

of

elec

tron

capt

ure

to a

sma

ll

degr

ee

(in

gen

eral,

less

than

1%)

dep

endi

ng

on

the

prox

imit

y of

elec

tron

s to

the

nucl

eus

in

bery

lliu

m.

In 

7Be,

a

diffe

renc

e of

0.9

%

has

bee

n

obs

erve

d

bet

wee

n

half-

lives

in

met

allic

and

insu

latin

g

envi

ron

men

ts.[9] 

This

relat

ively

larg

e

effe

ct is

bec

aus

e

bery

lliu

m is

a

sma

ll

ato

m

who

se

vale

nce

elec

tron

s

are

in 2

s at

omi

c

orbit

als,

whic

h

are

subj

ect

to

elec

tron

capt

ure

in 

7Be

bec

aus

e

(like

all s 

ato

mic

orbit

als

in

all

ato

ms)

they

natu

rally

pen

etrat

e

into

the

nucl

eus.

Rhe

niu

m-

187 

is a

mor

e

spe

ctac

ular

exa

mpl

e. 18

7Re

nor

mall

y be

ta

dec

ays 

to 18

7Os

with

a ha

lf-

life 

of

41.6

×

109 

y,[10] 

but

stud

ies

usin

g

fully

ioni

sed 

187R

e at

oms

(bar

e

nucl

ei)

hav

e

foun

d

that

this

can

decr

eas

e to

only

33

y.

This

is

attri

bute

d to

"

bou

nd-

stat

e

β -   d

eca

y" of

the

fully

ioni

sed

ato

m

—

the

elec

tron

is

emit

ted

into

the

"K-

shel

l"

(1s 

ato

mic

orbit

al),

whic

h

can

not

occ

ur

for

neut

ral

ato

ms

in

whic

h all

low-

lyin

g

bou

nd

stat

es

are

occ

upie

d.[11]

A

num

ber

of

exp

erim

ents

hav

e

foun

d

that

dec

ay

rate

s of

othe

r

mod

es

of

artifi

cial

and

natu

rally

occ

urrin

g

radi

oiso

tope

s

are,

to a

high

degr

ee

of

prec

isio

n,

unaf

fect

ed

by

exte

rnal

con

ditio

ns

suc

h as

tem

pera

ture,

pres

sure

, the

che

mic

al

envi

ron

men

t,

and

elec

tric,

mag

neti

c, or

grav

itati

onal

field

s.[12] 

Co

mpa

riso

n of

labo

rato

ry

exp

erim

ents

over

the

last

cent

ury,

stud

ies

of

the

Okl

o na

tural

nucl

ear

reac

tor (

whic

h

exe

mpli

fied

the

effe

cts

of

ther

mal

neut

rons

on

nucl

ear

dec

ay),

and

astr

oph

ysic

al

obs

erva

tion

s of

the

lumi

nosi

ty

dec

ays

of

dist

ant

sup

erno

vae

(whi

ch

occ

urre

d far

awa

y so

the

light

has

take

n a

grea

t

deal

of

time

to

reac

h

us),

for

exa

mpl

e,

stro

ngly

indi

cate

that

dec

ay

rate

s

hav

e

bee

n

con

stan

t (at

leas

t to

withi

n

the

limit

atio

ns

of

sma

ll

exp

erim

enta

l

erro

rs)

as a

func

tion

of

time

as

well.

Rec

ent

resu

lts

sug

gest

the

pos

sibili

ty

that

dec

ay

rate

s

mig

ht

hav

e a

wea

k

dep

end

enc

e

(0.5

%

or

less

) on

envi

ron

men

tal

fact

ors.

It

has

bee

n

sug

gest

ed

that

mea

sure

men

ts of

dec

ay

rate

s

o

fsilic

on-

32, 

man

gan

ese-

54,

and 

radi

um-

226 

exhi

bit

sma

ll

sea

son

al

vari

atio

ns

(of

the

orde

r of

0.1

%),

prop

ose

d to

be

relat

ed

to

eith

er

sola

r

flare

acti

vity

or

dist

anc

e

from

the

sun.

[13][14]

[15] H

owe

ver,

suc

h

mea

sure

men

ts

are

high

ly

sus

cept

ible

to

syst

ema

tic

erro

rs,

and

a

sub

seq

uent

pap

er[16] 

has

foun

d no

evid

enc

e for

suc

h

corr

elati

ons

in

six

othe

r

isot

ope

s,

and

sets

upp

er

limit

s on

the

size

of

any

suc

h

effe

cts.

IsotopesMost atoms have several naturally occurring isotopes (click here for a list of elements that have no isotopes). An isotope is an atom which contains a different number of neutrons in its nucleus than some other atom of the same element. This means that different isotopes of an element will have different masses, since both the protons and the neutrons contribute about equally to the mass of an atom.  (Here is a great source of information about sub-atomic particles from a physics point of view .)

Not all isotopes are equally abundant in nature. For example, here are the naturally occurring isotopes of Hydrogen (Hydrogen-2 is the only common isotope which has its own name, and is generally called Deuterium). 

 

Compare these abundances of hydrogen to those for some other elements. Notice that the atomic mass is close to the sum of the number of neutrons and protons, but not exactly the same. 

 

Carbon-12 is special, since its mass is defined to be exactly 12. It is the element chosen to be the reference mass in calculating atomic masses. 

You'll also notice that the atomic masses of each element are known very accurately. It would be impossible to construct a balance that had moving parts with that kind of accuracy--the vibration of the earth's crust would never allow it to become stable. Today's atomic masses are not calculated using a balance at all. Instead, they are measured using various forms of mass spectroscopy. 

A mass spectrometer can provide extremely accurate measurement of the mass of an individual atom. Here's how it works. A moving charged particle (an ion of an atom) will bend if it is placed in a magnetic field. The amount that it bends depends on its charge (+2 charged particles would bend twice as much as +1 charged ones, and negatively charged particles would bend the opposite direction). More importantly, the amount that the particle's path curves depends on its mass. 

It is sort of like rolling a steel shot put past a magnet. Because of its heavy mass, the path of the shot put would be very straight, compared to the path for a small ball bearing. In other words, a light particle will bend more in a magnetic field. In a mass spectrometer, atoms are charged by bombarding them with electrons. These charged atoms pass through a strong magnetic field, and their path bends. The lightest atoms bend a lot. The heavier atoms bend very little. Some type of detector is used to measure where the atoms arrive. By measuring the curvature of the path, and comparing it a known atom - our standard being Carbon-12 - very accurate masses can be calculated. 

In the original mass spectrograph, a photographic film was the detector. The film would be darkened by the impacting charged particles. The more particles that arrived, the darker the film would become. Measuring the amount of this exposure allowed scientists to calculate the percentage abundance of each isotope. In a modern mass spectrometer electronic devices are used to measure the number of ions which arrive, but the principle remains the same - heavy atoms and light atoms can be separated very precisely on the basis of their mass. 

Once the precise atomic masses of each isotope, and their relative abundances are known, the average mass of the isotope as found in nature can be calculated. In fact, the atomic masses are known very precisely, as shown in the above tables, but the abundances are much less accurately known. This is partly because the abundance of isotopes can vary from one location on earth to another. Certain naturally occurring process can concentrate the abundance of one type of isotope in one location as compared to another. Lead's isotopic abundance is one of the least reproducible because various isotopes are the final products of the radioactive decay of a number of heavy elements. To find the average mass of the mixture of isotopes found in nature the following procedure is used: multiply the abundance of each naturally occurring isotope of element by its precise atomic mass and then sum them all. Here are the results for a couple of elements. 

  Atomic mass (the weighted average of all naturally occurring isotopes) is not given for those atoms which have no stable isotopes. In these cases the mass number of the most stable isotope is reported, often in brackets, for example Technetium (98). 

List of Radioactive Elements

A radioactive element is one with an unstable nucleus, which radiates alpha, beta or gamma radiation and gets converted to a stable element. This article

has a comprehensive list of radioactive elements and their properties.Ads by GoogleTag Manager   SuperTag allows you to manage all tags in a single platform easily  supertag.datalicious.com

This Buzzle article has a list of radioactive elements that abound in nature, arranged in the order of increasing atomic number, along with their decay

modes. 

Let us understand the phenomenon of radioactivity. Radioactivity arrived on the scene of world physics in the 19th century, just when people thought they

knew everything in physics. With its discovery in 1896, radioactivity opened up a Pandora's box of questions and revealed a new world, waiting to be

explored in the microcosm of the atomic nucleus. 

What is Radioactivity?

Radioactivity is a very interesting phenomenon in nature. Classical Electromagnetism cannot explain radioactivity. It's a spontaneous and random

phenomenon whereby nuclei of certain chemical elements like Uranium, radiate gamma rays (high frequency electromagnetic radiation), beta particles

(electrons or positrons) and alpha particles (Helium Nuclei). 

By the emission of these particles and radiation, the unstable nucleus gets converted into a stabler nucleus. This is called radioactive decay.

The Term 'Radioactive' - A Misnomer 

A radioactive element is a fundamental element whose atomic nuclei demonstrates the phenomenon of radioactivity. The name 'radioactive' may suggest

to you that radioactive elements radiate radio waves, but unfortunately that is not so! The name 'radioactivity' is a misnomer because these elements have

nothing to do with radio waves! The reason is that energy and frequency of a gamma ray which is emitted by a radioactive element, is far beyond that of

the radio band of electromagnetic spectrum! So, we are just stuck up with the name! 

What Makes an Element Radioactive?

To understand radioactivity, we need to explore the structure of an atomic nucleus. Every nucleus contains neutrons as well as protons. Neutrons are

neither positively charged, nor negatively charged, they are neutral particles. Protons are positively charged. As you might remember from high school

physics, like charges repel each other while unlike charges attract each other. In the nucleus, protons and neutrons are cramped together in a really very

small space. 

The protons in the nucleus, all being positively charged, repel each other! So if all the protons repel each other, how does the nucleus stay glued together

and remain stable? It is because of the 'Nuclear Force'. 

This force is more stronger than the electromagnetic force, but the range of this force is only limited to size of the nucleus, unlike electromagnetic force

whose range is infinite. This nuclear force acts between the protons and neutrons, irrespective of the charge and it's always strongly attractive. However, it

has limitations of range. So, in the nucleus, there is a constant tussle between the repelling electromagnetic coulomb force of protons and the attractive

strong nuclear force. 

In a nucleus like Uranium, which has almost 92 protons, coulomb repulsive force becomes too much for the nuclear force to contain. Subsequently, the

nucleus is very unstable and radioactive decay occurs and Uranium decays into a more stable element. Such an unstable nucleus like Uranium, when

gently tapped by a neutron, splits up into two other nuclei through nuclear fission, releasing tremendous amount of energy in the process! This is the

principle on which nuclear energy and nuclear weapons are based.

The radioactive elements listed below shows all the decay modes of Uranium. A full explanation of radioactivity can only be given, if we plunge deep into

quantum physics and elementary particle physics.

Types of Radioactive Decay

This decay may occur in any of the following three ways:

Alpha Decay : Nucleus emits a helium nucleus (called an Alpha Particle) and gets converted to another nucleus with atomic number lesser by 2 and

atomic weight lesser by 4.

Beta Decay : Beta decay could be of two types; either through emission of an electron or positron (the antiparticle of electron). Electron emission

causes an increase in the atomic number by 1, while positron emission causes a decrease in the atomic number by 1. In some cases, double beta

decay may occur, involving the emission of two beta particles.

Gamma Decay : Gamma decay just changes the energy level of the nucleus.

Electron Capture : One of the rarest decay modes is electron capture. In this phenomenon, an electron is captured or absorbed by a proton rich

nucleus. This leads to the conversion of a proton into a neutron in the nucleus, along with release of an electron neutrino. This leads to a decrease in

atomic number (transmuting the element in the process), while leaving the atomic mass number unchanged.

A radioactive element may have more than one decay mode.

Radioactive Isotopes

When two nuclei have the same atomic number, but different atomic weight or mass numbers, then they are said to be isotopes. Isotopes have the same

chemical properties but different physical properties. For example, carbon has two isotopes, 6C14 and 6C12. Both have the same atomic number, but different

number of neutrons. The one with the two extra neutrons is radioactive and undergoes radioactive decay. The radioactive isotope of carbon was used to

develop carbon dating tool, which has made the dating of various relics possible.

Half-Life of a Radioactive Element

Half-life is the amount of time required, for half quantity of radioactive element to decay. For example C14has a half life of 5730 years. That is, if you take 1

gm of C14, then half of it will have been decayed in 5730 years. In the list presented below, half-lives of all the radioactive elements are presented.

Radioactive Elements List

Here is a detailed and comprehensive list of radioactive elements along with their atomic and mass numbers, decay modes and half-lives. Here 'Beta Decay

(β-)' denotes electron emission while Beta Decay (β+) denotes positron emission.

Radioactive Element Atomic Number Atomic Mass Number Decay Type Half-Life

Hydrogen (H) 1 3 Beta Decay (β-) 12.32 years

Beryllium (Be) 4 7 Electron Capture (ε), Gamma Decay) 53.12 Days

Beryllium (Be) 4 8 Alpha 7 x 10-17 sec

Beryllium (Be) 4 10 Beta Decay (β-) 1,360,000 years

Carbon (C) 6 14 Beta Decay (β-) 5,730 years

Calcium (Ca) 20 41 Electron Capture (ε) 103,000 years

Calcium(Ca) 20 46 Double Beta Decay (β-β-) > 2.8 x 1015 years

Calcium(Ca) 20 48 Double Beta Decay (β-β-) > 4 x 1019

Iron (Fe) 26 54 Double Electron Capture (ε) > 3.1 x 1022 years

Iron (Fe) (Synthetic) 26 55 Electron Capture (ε) 2.73 years

Iron (Fe) (Synthetic) 26 59 Beta Decay (β-) 44.503 days

Iron (Fe) (Synthetic) 26 60 Beta Decay (β-) 2,600,000 years

Cobalt (Co) (Synthetic) 27 56 Electron Capture (ε) 77.27 days

Cobalt (Co) (Synthetic) 27 57 Electron Capture (ε) 271.79 days

Cobalt (Co) (Synthetic) 27 58 Electron Capture (ε) 70.86 days

Cobalt (Co) (Synthetic) 27 60 Beta Decay (β-), Double Gamma 5.2714 years

Nickel (Ni) 28 59 Electron Capture (ε) 76,000 years

Nickel (Ni) (Synthetic) 28 63 Beta Decay (β-) 100.1 years

Zinc (Zn) (Synthetic) 30 65 Electron Capture (ε), Gamma 243.8 days

Zinc (Zn) (Synthetic) 30 72 Beta Decay (β-) 46.5 hours

Selenium (Se) 34 79 Beta Decay (β-) 3.27 x 105 years

Selenium (Se) 34 82 Double Beta Decay (β- β-) 1.08 x 1020 years

Krypton (Kr) 36 85 Beta Decay (β-) 10.756 years

Rubidium (Rb) 37 87 Beta Decay (β-) 4.88 x 1010 years

Strontium (Sr) 38 89 Electron Capture (ε), Beta Decay (β-) 50.52 days

Strontium (Sr) 38 90 Beta Decay (β-) 28.9 years

Yttrium (Y) 39 90 Beta Decay (β-), Gamma 2.67 days

Yttrium (Y) 39 91 Beta Decay (β-), Gamma 58.5 days

Zirconium (Zr) 40 93 Beta Decay (β-) 1.53 x 106 years

Zirconium (Zr) 40 94 Double Beta Decay (β-) > 1.1 x 1017 years

Zirconium (Zr) 40 96 Double Beta Decay (β-) 2 x 1019 years

Niobium (Nb) (Metastable) 41 93 Beta Decay (β-),Gamma 16.13 years

Niobium (Nb) 41 95 Beta Decay (β-), Gamma 34.991 days

Molybdenum (Mo) 42 93 Electron Capture (ε) 4 x 103 years

Technetium (Tc) 43 99 Beta Decay (β-) 2.111 x 105 years

Ruthenium (Ru) 44 103 Beta Decay (β-), Gamma 39.26 days

Ruthenium(Ru) 44 106 Beta Decay (β-) 373.59 days

Palladium (Pd) 46 107 Beta Decay (β-), Gamma 6.5 million years

Silver (Ag) 47 111 Beta Decay (β-), Gamma 7.45 days

Tin (Sn) 50 126 Beta Decay (β-) 2.3 x 105 years

Antimony (Sb) 51 125 Beta Decay (β-) 2.7582 years

Tellurium (Te) 52 127 Beta Decay (β-), Gamma 9.35 hours

Tellurium (Te) 52 129 Beta Decay (β-) 69.6 minutes

Iodine (I) 53 123 Electron Capture (ε), Gamma 13 hours

Iodine (I) 53 129 Beta Decay (β-) 15.7 million years

Iodine (I) 53 131 Beta Decay (β-), Gamma 8.02070 days

Xenon (Xe) 54 125 Electron Capture (ε) 16.9 hours

Xenon (Xe) 54 127 Electron Capture (ε) 36.345 days

Xenon (Xe) 54 133 Beta Decay (β-) 5.247 days

Cesium (Cs) 55 134 Electron Capture (ε), Beta Decay (β-) 2.0648 years

Cesium (Cs) 55 135 Beta Decay (β-) 2.3 million years

Cesium (Cs) 55 137 Beta Decay (β-), Gamma 30.17 years

Cerium (Ce) 58 144 Beta Decay (β-) 285 days

Promethium (Pm) 61 147 Beta Decay (β-), Gamma 2.6234 years

Europium (Eu) 63 154 Beta Decay (β-), Beta Decay (β+), Gamma 16 years

Europium (Eu) 63 155 Beta Decay (β-) 2 years

Iridium (Ir) (Synthetic) 77 188 Electron Capture (ε) 1.73 days

Iridium (Ir) (Synthetic) 77 189 Electron Capture (ε) 13.2 days

Iridium (Ir) (Synthetic) 77 190 Electron Capture (ε) 11.8 days

Iridium (Ir) (Synthetic) 77 192 Beta Decay (β-), Electron Capture (ε) 73.827 days

Iridium (Ir) (Synthetic, Metastable) 77 192 Gamma Decay 241 years

Iridium (Ir) (Synthetic) 77 193 Gamma Decay 10.5 days

Iridium (Ir) (Synthetic) 77 194 Beta Decay (β-) 19.3 hours

Iridium (Ir) (Synthetic, Metastable) 77 194 Gamma Decay 171 days

Lead (Pb) 82 210 Beta Decay (β-), Alpha 21 years

Bismuth (Bi) 83 210 Alpha 3 million years

Polonium (Po) 84 210 Alpha 138 days

Radon (Rn) 86 220 Alpha, Beta Decay (β+) 1 min

Radon (Rn) 86 222 Alpha 4 days

Radium (Ra) 88 224 Alpha 4 days

Radium (Ra) 88 225 Beta Decay (β-) 15 days

Radium (Ra) 88 226 Alpha 1,622 years

Thorium (Th) 90 228 Alpha 2 years

Thorium (Th) 90 229 Alpha 7,340 years

Thorium (Th) 90 230 Alpha 80,000 years

Thorium (Th) 90 232 Alpha 14 years

Thorium (Th) 90 234 Beta Decay (β-) 24 days

Proactinium (Pa) 91 234 Beta Decay (β-) 6.75 hours

Uranium (U) 92 233 Alpha 159,200 years

Uranium (U) 92 234 Alpha 245,500 years

Uranium (U) 92 235 Alpha 7.038 x 108 years

Uranium (U) 92 236 Alpha 2.342 x 107 years

Uranium (U) 92 238 Alpha 4.468 billion years

Neptunium (Np) (Synthetic) 93 237 Alpha 2.144 million years

Plutonium (Pu) 94 238 Alpha 87.74 years

Plutonium (Pu) 94 239 Alpha 2.41 x 104 years

Plutonium (Pu) 94 240 Alpha 6.5 x 103 years

Plutonium (Pu) 94 241 Beta Decay (β-) 14 years

Plutonium (Pu) 94 242 Alpha 3.73 x 105 years

Plutonium (Pu) 94 244 Alpha 8.08 x 107 years

Americium (Am) 95 241 Alpha 432.2 years

Americium (Am) (Metastable) 95 242 Alpha, Gamma 141 years

Americium (Am) 95 243 Alpha 7,370 years

Curium (Cm) 96 242 Alpha 160 days

Curium (Cm) 96 243 Alpha 29.1 years

Curium (Cm) 96 244 Alpha 18.1 years

Curium (Cm) 96 247 Alpha 15.6 million years

These radioactive isotopes have a lot of applications today, ranging from medicine to atomic energy. Since these radioactive elements are harmful, burning

up radioactive waste or disposing it, is difficult. Every development in science and technology brings in new problems. Now, it's for us to decide, how we

intend to use the power of technology placed in our hands.

Half Lives

Radiation-related consulting and services from Integrated Environmental Management, Inc.

The half-life of a radioactive element is the time that it takes for one half of the atoms of that substance to disintegrate into another nuclear form. These can range from mere fractions of a second, to many billions of years. In addition, the half-life of a particular radionuclide is unique to that radionuclide, meaning that knowledge of the half-life leads to the identity of the radionuclide. The following is a listing of the half-lives of commonly-encountered radionuclides, with the units of each as shown.

Please note that some of the longer half-lives are written in scientific notation (i.e., 7.2E1 is equal to 7.2 x 10, or 72.)

Actinium

Ac-225 - 10.0 days Ac-227 - 21.773 years Ac-228 - 6.13 hours

Americium

Am-241 - 432.2 years Am-242 - 16.02 hours Am-242m - 152 years Am-243 - 7380 years

Antimony

Sb-124 - 60.20 days Sb-125 - 2.77 years Sb-126 - 12.4 days Sb-126m - 19.0 minutes Sb-127 - 3.85 days

Argon

Ar-41 - 1.827 hours

Astatine

At-217 - 0.0323 seconds At-218 - 2 seconds

Barium

Ba-137m - 2.552 minutes Ba-139 - 82.7 minutes Ba-140 - 12.74 days Ba-141 - 18.27 minutes Ba-142 - 10.6 minutes

Beryllium

Be-10 - 1.6E6 years Be-7 - 53.44 days

Bismuth

Bi-210 - 5.012 days Bi-211 - 2.14 minutes Bi-212 - 60.55 minutes Bi-213 - 45.65 minutes Bi-214 - 19.9 minutes

Bromine

Br-82 - 35.30 hours Br-83 - 2.39 hours Br-84 - 31.80 minutes

Cadmium

Cd-113m - 13.6 years Cd-115m - 44.6 days

Calcium

Ca-41 - 1.3E5 years Ca-47 - 4.53 days

Californium

Cf-252 - 2.638 years

Carbon

C-11 - 20.38 minutes C-14 - 5730 years

Cerium

Ce-141 - 32.50 days Ce-143 - 33.0 hours Ce-144 - 284.3 days

Cesium

Cs-134 - 2.062 years Cs-134m - 2.90 hours Cs-135 - 2.3E6 years Cs-136 - 13.1 days Cs-137 - 30.0 years Cs-138 - 32.2 minutes

Chromium

Cr-51 - 27.704 days

Cobalt

Co-56 - 78.76 days Co-57 - 270.9 days Co-58 - 70.8 days Co-60 - 5.27 years

Copper

Cu-61 - 3.408 hours Cu-64 - 12.701 hours

Curium

Cm-242 - 162.8 days Cm-243 - 28.5 years Cm-244 - 18.11 years Cm-245 - 8500 years Cm-246 - 4730 years Cm-247 - 1.56E7 years Cm-248 - 3.39E5 years

Europium

Eu-152 - 13.33 years Eu-154 - 8.8 years Eu-155 - 4.96 years Eu-156 - 15.19 days

Fluorine

F-18 - 109.74 minutes

Francium

Fr-221 - 4.8 minutes Fr-223 - 21.8 minutes

Gadolinium

Gd-152 - 1.08E14 years

Gallium

Ga-67 - 3.261 days

Gold

Au-198 - 2.696 days

Holmium

Ho-166m - 1.20E3 years

Hydrogen

H-3 - 12.35 years

Indium

In-111 - 2.83 days In-113 - 1.658 hours In-115 - 5.1E15 years

Iodine

I-123 - 13.2 hours I-125 - 60.14 days I-129 - 1.57E7 years I-130 - 12.36 hours I-131 - 8.04 days I-132 - 2.30 hours

I-133 - 20.8 hours I-134 - 52.6 minutes I-135 - 6.61 hours

Iridium

Ir-192 - 74.02 days

Iron

Fe-55 - 2.7 years Fe-59 - 44.53 days

Krypton

Kr-83m - 1.83 hours Kr-85 - 10.72 years Kr-85m - 4.48 hours Kr-87 - 76.3 minutes Kr-88 - 2.84 minutes

Lanthanum

La-140 - 40.272 hours La-141 - 3.93 hours La-142 - 92.5 minutes

Lead

Pb-209 - 3.253 hours

Pb-210 - 22.3 years Pb-211 - 36.1 minutes Pb-212 - 10.64 hours Pb-214 - 26.8 minutes Pb-214 - 26.8 minutes

Manganese

Mn-52 - 5.591 days Mn-52m - 21.1 minutes Mn-54 - 312.5 days Mn-56 - 2.579 hours Mn-57 - 36.08 hours

Mercury

Hg-197 - 64.1 hours Hg-203 - 46.60 days

Molybdenum

Mo-93 - 3.5E3 years Mo-99 - 66.0 hours

Neodymium

Nd-147 - 10.98 days

Neptunium

Np-237 - 2.14E6 years Np-238 - 2.117 days Np-239 - 2.355 days Np-240 - 65 minutes Np-240m - 7.4 minutes

Nickel

Ni-59 - 7.5E4 years Ni-63 - 96 years Ni-65 - 2.520 hours

Niobium

Nb-93m - 13.6 years Nb-95 - 35.15 days Nb-95m - 86.6 hours Nb-97 - 72.1 minutes Nb-97m - 60 seconds

Nitrogen

N-13 - 9.97 minutes N-16 - 7.13 seconds

Oxygen

O-15 - 122.24 seconds

Palladium

Pd-107 - 6.5E6 years Pd-109 0- 13.427 hours

Phosphorus

P-32 - 14.29 days

Plutonium

Pu-238 - 87.74 years Pu-239 - 24065 years Pu-240 - 6537 years Pu-241 - 14.4 years Pu-242 - 3.76E5 years Pu-243 - 4.956 hours Pu-244 - 8.26E7 years

Polonium

Po-210 - 138.38 days Po-211 - 0.516 seconds Po-212 - 0.305 microseconds Po-213 - 4.2 microseconds Po-214 - 164.3 microseconds Po-215 - 0.00178 seconds Po-216 - 0.15 seconds Po-218 - 3.05 minutes

Potassium

K-40 - 1.27E9 years K-42 - 12.36 hours K-43 - 22.6 hours

Praseodymium

Pr-143 - 13.56 days Pr-144 - 17.28 minutes Pr-144m - 7.2 minutes

Promethium

Pm-147 - 2.6234 years Pm-148 - 5.37 days Pm-148 - 41.3 days Pm-149 - 53.08 hours Pm-151 - 28.40 hours

Protactinium

Pa-231 - 3.28E4 years Pa-233 - 27.0 days Pa-234 - 6.70 hours Pa-234m - 1.17 minutes

Radium

Ra-223 - 11.434 days Ra-224 - 3.66 days

Ra-225 - 14.8 days Ra-226 - 1600 years Ra-228 - 5.75 years

Radon

Rn-219 - 3.96 seconds Rn-220 - 55.6 seconds Rn-222 - 3.824 days

Rhenium

Re-187 - 5E10 years

Rhodium

Rh-103m - 56.12 minutes Rh-105 - 35.36 hours Rh-106 - 29.9 seconds

Rubidium

Rb-86 - 18.66 days Rb-87 - 4.7E10 years Rb-88 - 17.8 minutes Rb-89 - 15.2 minutes

Ruthenium

Ru-103 - 39.28 days

Ru-105 - 4.44 hours Ru-106 - 368.2 days Ru-97 - 2.9 days

Samarium

Sm-147 - 1.06E11 years Sm-151 - 90 years Sm-153 - 46.7 hours

Scandium

Sc-44 - 3.927 hours Sc-46 - 83.83 days Sc-47 - 3.351 days Sc-48 - 43.7 hours

Selenium

Se-75 - 119.78 days Se-79 - 65000 years

Silver

Ag-110 - 24.6 seconds Ag-110m - 249.9 days Ag-111 - 7.45 days

Sodium

Na-22 - 2.602 years Na-24 - 15.00 hours

Strontium

Sr-85 - 64.84 days Sr-87m - 2.81 hours Sr-89 - 50.5 days Sr-90 - 29.12 years Sr-91 - 9.5 hours Sr-92 - 2.71 hours

Sulfur

S-35 - 87.44 days

Technetium

Tc-101 - 14.2 minutes Tc-99 - 2.13E5 years Tc-99m - 6.02 hours

Tellurium

Te-125m - 58 days Te-127 - 9.35 hours Te-127m - 109 days Te-129 - 69.6 minutes Te-129m - 33.6 days

Te-131 - 25.0 minutes Te-131m - 30 hours Te-132 - 78.2 hours Te-133 - 12.45 minutes Te-133m - 55.4 minutes Te-134 - 41.8 minutes

Terbium

Tb-160 - 72.3 days

Thallium

Tl-201 - 73.06 hours Tl-207 - 4.77 minutes Tl-208 - 3.07 minutes Tl-209 - 2.20 minutes

Thorium

Th-227 - 18.718 days Th-228 - 1.913 years Th-229 - 7340 years Th-230 - 7.7E4 years Th-231 - 25.52 hours Th-232 - 1.41E10 years Th-234 - 24.10 days

Tin

Sn-119m - 293.1 days Sn-123 - 129.2 days Sn-125 - 9.64 days Sn-126 - 1.0E5 years

Tungsten

W-181 - 121.2 days W-185 - 75.1 days W-187 - 23.9 hours

Uranium

U-232 - 72 years U-233 - 1.59E5 years U-234 - 2.445E5 years U-235 - 7.03E8 years U-236 - 2.34E7 years U-237 - 6.75 days U-238 - 4.47E9 years U-240 - 14.1 hours

Vanadium

V-48 - 16.238 days

Xenon

Xe-131m - 11.9 days

Xe-133 - 5.245 days Xe-133m - 2.188 days Xe-135 - 9.09 hours Xe-135m - 15.29 minutes Xe-138 - 14.17 minutes

Ytterbium

Yb-169 - 32.01 days

Yttrium

Y-90 - 64.0 hours Y-91 - 58.51 days Y-91m - 49.71 minutes Y-92 - 3.54 hours Y-93 - 10.1 hours

Zinc

Zn-65 - 243.9 days Zn-69 - 57 minutes

Zirconium

Zr-93 - 1.53E6 years Zr-95 - 63.98 days Zr-97 - 16.90 hours From Yahoo Answers

Question:I know there are different types of isotopes, but I just dont know the name of any. What are three different kinds, and what do they do?

Answers:Isotopes are just different 'versions' of the atom of an element. Most elements have at least two isotopes, and neither one is THE isotope, it's just that some are

more common than others. For example, the basic hydrogen atom has only one proton and an electron. An isotope of it has one proton, electron and a neutron, and

another one has one proton, electron and two neutrons. What changes is the number of neutrons, everything else remains the same. And any of these three variations is

an isotope of hydrogen.

Question:This question is to test your general knowledge of science from memory. (No fair looking up the answer. I can't stop you from cheating but please don't.) Tell

me the basic definition of each if you actually know. Don't feel bad if you do not. They are good words to know if you do not yet. Thanks for your honesty in answering with

out research. :-)

Answers:Good question and I am proud to say that I DO know the difference. The more common one, an "Isotope" is when there is a heavier form of one atom of an

element, due to it having more neutrons than normal. Normally there is the same number of protons and neutrons (like helium has 2+2) but since neutrons are neutral

there can be many more isotopes when each one has one more neutron. Basic Hydrogen has no neutron in fact and only a proton. (Called Protium) Then Deuterium has

1 proton and 1 neutron and Tritium has 2 neutrons and one proton. Those are all Isotopes of Hydrogen. An "Allotrope" is when you have different forms of the way atoms

bond together of the same element. The best example is Carbon forms both Graphite and also Diamonds but they are two different Allotropes. Diamonds have carbon

atoms bonded together in a 3D crystal structure and Graphite is actually sheets of connected carbon atoms which then slide over each other. They are both 100% carbon

but different allotropes.

Question:

Answers:An allotrope is a structurally different form of an element; "Carbon has the most number of allotropes. Scientists have already found eight of them, including

amorphous carbon allotrope (manifestations include coal and soot), carbon nanofoam, carbon nanotube, the diamond allotrope, fullerene allotrope, graphite, lonsdaleite,

and ceraphite allotrope." It is the ability of an element to exist in more than one physical state. Many elements can appear naturally in more than one form at room

temperature. The allotropes are usually different due to different crystal shape, a lack of crystal shape, or through the attachments of more than one atom of an element.

An isotope is an element that has different numbers of neutrons in its nucleus. Examples are hydrogen-1, hydrogen-2 and hydrogen-3, known as protium, deuterium, and

tritium respectively. protium or hydrogen-1 has a nucleus with 1 proton and no neutrons. Deuterium has 1 proton and 1 neutron in its nucleus, while tritium has 1 proton

and 2 neutrons in its nucleus.

A list of radioactive elementsHere is a detailed and comprehensive list of radioactive elements along with their atomic and mass numbers, decay modes and half lives. Here 'Beta Decay (-)' denotes Electron emission while Beta Decay (+) denotes Positron emission.

 

Radioactive Element

Atomic Number

Atomic Mass Number

Decay Type Half Life

Hydrogen (H) 1 3 Beta Decay (Beta-) 12 years

Beryllium (be) 4 10 Beta Decay (Beta-) 2,700,000 years

Carbon (C) 6 14 Beta Decay (Beta-) 5,730 years

Calcium (Ca) 20 41 Beta Decay (Beta+) 100,000 years

Iron (Fe) 26 59 Beta Decay (Beta-) 45 days

Cobalt (Co) 27 60 Beta Decay (Beta-) Gamma 5 years

Nickel (Ni) 28 59 Beta Decay (Beta+) 80,000 years

Zinc (Zn) 30 65 Beta Decay (Beta-) Gamma 145 days

Selenium (Se) 34 79 Beta Decay (Beta-) 70,000 years

Krypton (Kr) 36 85 Beta Decay (Beta-) Gamma 10 years

Krypton (Kr) 36 90 Beta Decay (Beta-) Gamma 33 seconds

Rubidium (Rb)37 87 Beta Decay (Beta-) 47 Billion years

Strontium (Sr) 38 89 Beta Decay (Beta-) 53 days

Strontium (Sr) 38 90 Beta Decay (Beta-) 28 years

Yttrium (Y) 39 90 Beta Decay (Beta-) Gamma 64 hrs

Yttrium (Y) 39 91 Beta Decay (Beta-) 58 days

Zirconium (Zr) 40 93 Beta Decay (Beta-) 950,000

Zirconium (Zr) 40 95 Beta Decay (Beta-) 65 days

Niobium (Nb) 41 93 Gamma 4 years

Niobium (Nb) 41 95 Beta Decay (Beta-) Gamma 35 days

Molybdenum (Mo)

42 93 Beta Decay (Beta+) 10,000 years

Technetium (Tc)

43 99 Beta Decay (Beta-) Gamma 210,000

Ruthenium (Ru)

44 103 Beta Decay (Beta-) 40 days

Ruthenium (Ru)

44 106 Beta Decay (Beta-) 1 year

Palladium 46 107 Beta Decay (Beta-) Gamma 7 million years

Silver (Ag) 47 110 Beta Decay (Beta-) Gamma 249 days

Tin (Sn) 50 126 Beta Decay (Beta-) 100,000 years

Antimony (Sb) 51 125 Beta Decay (Beta-) 2 years

Tellurium (Te) 52 127 Beta Decay (Beta-) Gamma 105 days

Tellurium (Te) 52 129 Beta Decay (Beta-) 67 minutes

Iodine (I) 53 129 Beta Decay (Beta-) Gamma 17.2 million years

Iodine (I) 53 131 Beta Decay (Beta-) Gamma 8 days

Iodine (I) 53 134 Beta Decay (Beta-) Gamma 52 minutes

Xenon (Xe) 54 133 Beta Decay (Beta-) Gamma 5 days

Xenon (Xe) 54 137 Beta Decay Beta-) Gamma 4 minutes

Xenon (Xe) 54 138 Beta Decay (Beta-) Gamma 14 minutes

Cesium (Cs) 55 134 Beta Decay (Beta-) Gamma 2 years

Cesium (Cs) 55 135 Beta Decay (Beta-) Gamma 2 million years

Cesium (Cs) 55 137 Beta Decay (Beta-) Gamma 30 years

Cesium (Ce) 58 144 Beta Decay (Beta-) 285 days

Promethium (Pm)

61 147 Beta Decay (Beta-) Gamma 2 years

Europium (Eu)63 154Beta Decay (Beta-), Beta Decay (Beta+) Gamma

16 years

Europium (Eu)63 155 Beta Decay (Beta-) 2 tears

Lead (Pb) 82 210 Beta Decay (Beta-) Alpha 21 years

Bismuth (Bi) 83 210 Alpha 3 million years

Polonium (Po) 84 210 Alpha 138 days

Radon (Rn) 86 220 Alpha, Beta Decay (Beta+) 1 minute

Radon (Rn) 86 222 Alpha 4 days

Radium (Ra) 88 224 Alpha 4 days

Radium (Ra) 88 225 Beta Decay (Beta-) 15 days

Radium (Ra) 88 226 Alpha 1,622 years

Thorium (Th) 90 228 Alpha 2 years

Thorium (Th) 90 229 Alpha 7,340 years

Thorium (Th) 90 230 Alpha 80,000 years

Thorium (Th) 90 232 Alpha 14 years

Thorium (Th) 90 234 Beta Decay (Beta-) 24 days

Proactinium (Pa)

91 226 Alpha, Beta Decay (Beta+) 2 minutes

Uranium (U) 92 233 Alpha 162,000 years

Uranium (U) 92 234 Alpha 248,000 years

Uranium (U) 92 235 Alpha 713 million years

Uranium (U) 92 236 Alpha 23.9 million years

Uranium (U) 92 238 Alpha 4.51 billion years

Neptunium (Np)

93 237 Alpha 2.2 million years

Plutonium (Pu)

94 236 Alpha 285 years

Plutonium (Pu)

94 238 Alpha 86 years

Plutonium (Pu)

94 239 Alpha 24,390 years

Plutonium (Pu)

94 240 Alpha 6,580 years

Plutonium (Pu)

94 241 Beta Decay (Beta-), Alpha 13 years

Plutonium (Pu)

94 242 Alpha 379,000 years

Plutonium (Pu)

94 243 Alpha 5 years

Plutonium (Pu)

94 244 Alpha 76 million years

Americium (Am)

95 241 Alpha 458 years

Americium (Am)

95 242Beta Decay (Beta-), Beta Decay (Beta+), Alpha, Gamma

16 hours

Americium (Am)

95 243 Alpha 7,950

Curium (Cm) 96 242 Alpha 163 days

Curium (Cm) 96 243 Alpha 35 years

Curium (Cm) 96 244 Alpha 18 years

Curium (Cm) 96 247 Alpha 40 million years