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[link]

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.

Notation[link]

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. 3
2He, 4
2He, 12
6C, 14
6C, 235
92U, and 239
92U, 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. 3
He, 4
He, 12
C, 14
C, 235
U, and 239
U, 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 180m
73Ta (tantalum-180m).

Radioactive, primordial, and stable isotopes[link]

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, 14
C is a radioactive form of carbon while 12
C and 13
C 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 33 nuclides with very long half-lives (over 80 million years) and 255 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 has yet been observed. The half-lives for these processes often exceed a million times the estimated age of the universe, and in fact there are 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.

History[link]

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.

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 between uranium and lead. Since the periodic table only allowed for 11 elements from uranium to lead, Soddy proposed that several types of atoms (differing in radioactive properties) could occupy the same place in the table.^[5]

The term “isotope”, Greek for “at the same place”, was suggested to Soddy in 1914 by Margaret Todd, a Scottish physician to whom he was distantly related by marriage, during a conversation in which he explained his ideas to her.^[6]

Confirmation was provided by the observation of isotopes differing in mass for a stable (non-radioactive) element by J. J. Thomson in 1913. As part of his exploration into the composition of canal 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.

Variation in properties between isotopes[link]

Chemical and molecular properties[link]

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 (1
H) and deuterium (2
H), 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.

Nuclear properties and stability[link]

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 3
2He is 1:2, the neutron:proton ratio of 238
92U is greater than 3:2. A number of lighter elements have stable nuclides with the ratio 1:1 (Z = N). The nuclide 40
20Ca (calcium-40) is the heaviest stable nuclide with the same number of neutrons and protons; all heavier stable nuclides contain more neutrons than protons.

Numbers of isotopes per element[link]

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 180m
73Ta as stable), and 26 elements have only a single stable isotope (of these, 19 are so-called mononuclidic 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).^[7] 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.

Even and odd nucleon numbers[link]

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.

Even mass number[link]

Even-mass-number nuclides, about = 154/255 = ~ 60% of all stable nuclides, are bosons, i.e. they have integer spin. Almost all are even-proton, even-neutron (EE) nuclides, which necessarily have spin 0 because of pairing; only 5 are odd-proton, odd-neutron nuclides, which have nonzero integer spin.

Pairing effects[link]

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 (lower binding 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 116
Cd 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).

Even proton-even neutron[link]

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.

Odd proton-odd neutron[link]

Only five stable nuclides contain both an odd number of protons and an odd number of neutrons: the first four odd-odd nuclides, where changing a proton to a neutron or vice versa would lead to a very lopsided proton-neutron ratio (2
1H, 6
3Li, 10
5B, and 14
7N; spins 1, 1, 3, 1) and 180m
73Ta (spin 9), the only primordial nuclear isomer, which has not yet been observed to decay despite experimental attempts.^[8] Also, four long-lived radioactive odd-odd nuclides (40
19K, 50
23V,138
57La,176
71Lu; spins 4, 6, 5, 7) occur naturally.

Of these 9 primordial odd-odd nuclides, only 14
7N is the most common isotope of a common element, because it is a part of the CNO cycle; 6
3Li and 10
5B are minority isotopes of elements that are rare compared to other light elements, while the other six isotopes make up only a tiny percentage of their elements.

None of the primordial odd-odd nuclides have spin 0 in the ground state.

Odd mass number[link]

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, giving a total of 101 stable nuclides with odd mass numbers.

Odd-mass-number nuclides are fermions, i.e. have half-integer spin. 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.^[9]

Odd proton-even neutron[link]

These 48 stable nuclides form most of the stable isotopes of the odd-numbered elements; the few odd-odds are 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, 87
37Rb, 115
49In, 151
63Eu, 187
75Re, and 209
83Bi which was recently found to decay.

Even proton-odd neutron[link]

53 stable and 3 primordial long-lived nuclides (including the fissile 235
92U) have an even number of protons and an odd number of neutrons. They are isotopes of even-Z elements, where they are a minority in comparison to the even-even isotopes which are about 3 times as numerous. Only 195
78Pt and 9
4Be are the most naturally abundant isotopes in their element, the former only by a small margin, and the latter only because the expected beryllium-8 has slightly lower binding energy than two alpha particles and therefore alpha decays.

Odd neutron number[link]

Actinides with odd neutron number are generally fissile (with thermal neutrons), while those with even neutron number are generally not, though they are fissionable with fast neutrons. Only 195
78Pt, 9
4Be and 14
7N have odd neutron number and are the most naturally abundant isotope of their element.

Occurrence in nature[link]

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.^[10] 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. (See nucleosynthesis 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.

Atomic mass of isotopes[link]

The atomic mass (m_r) 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:

Failed to parse (Missing texvc executable; please see math/README to configure.): M = m_1 x_1+m_2 x_2+...+m_Nx_N

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

Applications of isotopes[link]

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.

Use of chemical and biological properties[link]

• 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^[11] 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.^[12]
• 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, they 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).
• 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.
• Isotopic substitution can be used to determine the mechanism of a reaction via the kinetic isotope effect.

Use of nuclear properties[link]

• Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes. 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.

See also[link]

• Abundance of the chemical elements
• Atom
• Table of nuclides
• Table of nuclides (complete)
• List of isotopes
• List of isotopes by half-life
• List of elements by stability of isotopes
• Isotones
• Isobars
• Radionuclide (or radioisotope)
• Nuclear medicine (includes medical isotopes)
• Isotopomer
• List of particles

Notes[link]

• Isotopes are nuclides having the same number of protons; compare:
  • Isotones are nuclides having the same number of neutrons. Failed to parse (Missing texvc executable; please see math/README to configure.): N = A - Z

• • Isobars are nuclides having the same mass number, i.e. sum of protons plus neutrons. Failed to parse (Missing texvc executable; please see math/README to configure.): A

• • Nuclear isomers are different excited states of the same type of nucleus. A transition from one isomer to another is accompanied by emission or absorption of a gamma ray, or the process of internal conversion. Isomers are by definition both isotopic and isobaric. (Not to be confused with chemical isomers.)
  • Isodiaspheres are nuclides having the same neutron excess, i.e. number of neutrons minus number of protons. Failed to parse (Missing texvc executable; please see math/README to configure.): D = N - Z

• Bainbridge mass spectrometer

References[link]

External links[link]

• National Nuclear Data Center Portal to large repository of free data and analysis programs from NNDC
• National Isotope Development Center Coordination and management of the production, availability, and distribution of isotopes, and reference information for the isotope community
• Isotope Development & Production for Research and Applications (IDPRA) U.S. Department of Energy program for isotope production and production research and development
• International Atomic Energy Agency Homepage of International Atomic Energy Agency (IAEA), an Agency of the United Nations (UN)
• Atomic Weights and Isotopic Compositions for All Elements Static table, from NIST (National Institute of Standards and Technology)
• Atomgewichte, Zerfallsenergien und Halbwertszeiten aller Isotope
• Chart of the Nuclides produced by the Knolls Atomic Power Laboratory $25
• Exploring the Table of the Isotopes at the LBNL
• Current isotope research and information isotope.info
• Emergency Preparedness and Response: Radioactive Isotopes by the CDC (Centers for Disease Control and Prevention)
• Chart of Nuclides Interactive Chart of Nuclides (National Nuclear Data Center)
• Interactive Chart of the nuclides, isotopes and Periodic Table
• The LIVEChart of Nuclides - IAEA with isotope data, in Java or HTML
• [1] Annotated bibliography for isotopes from the Alsos Digital Library for Nuclear Issuesvep:Izotop

Joe Henderson
Joe Henderson with Neil Swainson
Background information
Born	(1937-04-24)April 24, 1937 Lima, Ohio, United States
Died	June 30, 2001(2001-06-30) (aged 64) San Francisco, California, United States
Genres	Soul-jazz Mainstream jazz Hard bop Post-bop Jazz fusion
Years active	1960–1997
Labels	Milestone, Verve, Blue Note

Joe Henderson (April 24, 1937 – June 30, 2001) was an American jazz tenor saxophonist. In a career spanning more than forty years Henderson played with many of the leading American players of his day and recorded for several prominent labels, including Blue Note.

Biography[link]

Early life[link]

From a very large family with five sisters and nine brothers, Henderson was born in Lima, Ohio, and was encouraged by his parents and older brother James T. to study music. He dedicated his first album to them "for being so understanding and tolerant" during his formative years. Early musical interests included drums, piano, saxophone and composition. According to Kenny Dorham, two local piano teachers who went to school with Henderson's brothers and sisters, Richard Patterson and Don Hurless, gave him a knowledge of the piano.^[1] He was particularly enamored of his brother's record collection. It seems that a hometown drummer, John Jarette, advised Henderson to listen to musicians like Lester Young, Stan Getz, Dexter Gordon and Charlie Parker.^[1] He also liked Flip Phillips, Lee Konitz and the Jazz at the Philharmonic recordings. However, Parker became his greatest inspiration. His first approach to the saxophone was under the tutelage of Herbert Murphy in high school. In this period of time, he wrote several scores for the school band and rock groups.

By eighteen, Henderson was active on the Detroit jazz scene of the mid-'50s, playing in jam sessions with visiting New York stars. While attending classes of flute and bass at Wayne State University, he further developed his saxophone and compositional skills under the guidance of renowned teacher Larry Teal at the Teal School of Music. In late 1959, he formed his first group.^[1] By the time he arrived at Wayne State University, he had transcribed and memorized so many Lester Young solos that his professors believed he had perfect pitch. Classmates Yusef Lateef, Barry Harris and Donald Byrd undoubtedly provided additional inspiration.^[2] He also studied music at Kentucky State College.

Shortly prior to his army induction in 1960, Henderson was commissioned by UNAC to write some arrangements for the suite "Swings and Strings", which was later performed by a ten-member orchestra and the local dance band of Jimmy Wilkins.^[1]

Early career[link]

He spent two years (1960–1962) in the U.S. Army: firstly in Fort Benning, where he even competed in the army talent show and won the first place, then in Fort Belvoir, where he was chosen for a world tour, with a show to entertain soldiers. While in Paris, he met Kenny Drew and Kenny Clarke. Then he was sent to Maryland to conclude his draft. In 1962, he was finally discharged and promptly moved to New York. He first met trumpeter Kenny Dorham, an invaluable guidance for him, at saxophonist Junior Cook's place. That very evening, they went see Dexter Gordon playing at the Birdland. Henderson was asked by Gordon himself to play something with his rhythm section; needless to say, he happily accepted.^[1]

Although Henderson's earliest recordings were marked by a strong hard-bop influence, his playing encompassed not only the bebop tradition, but R&B, Latin and avant-garde as well. He soon joined Horace Silver's band and provided a seminal solo on the jukebox hit "Song for My Father". After leaving Silver's band in 1966, Henderson resumed freelancing and also co-led a big band with Kenny Dorham. His arrangements for the band went unrecorded until the release of Joe Henderson Big Band (Verve) in 1996.

Blue Note[link]

From 1963 to 1968, Joe appeared on nearly thirty albums for Blue Note, including five released under his name. The recordings ranged from relatively conservative hard-bop sessions (Page One, 1963) to more explorative sessions (Inner Urge and Mode for Joe, 1966). He played a prominent role in many landmark albums under other leaders for the label including most of Horace Silver's swinging and soulful Song For My Father, Herbie Hancock's dark and densely orchestrated The Prisoner, Lee Morgan's hit album The Sidewinder and 'out' albums with pianist Andrew Hill (Black Fire 1963 and Point of Departure, 1964) and drummer Pete La Roca (Basra, 1965).

In 1967, there was a notable, but brief, association with Miles Davis's quintet featuring Herbie Hancock, Wayne Shorter, Ron Carter and Tony Williams, although the band was never recorded. Henderson's adaptability and eclecticism would become even more apparent in the years to follow.

Milestone[link]

Signing with Orrin Keepnews's fledgling Milestone label in 1967 marked a new phase in Henderson’s career. He co-led the Jazz Communicators with Freddie Hubbard from 1967-1968. Henderson was also featured on Hancock's Fat Albert Rotunda for Warner Bros. It was during this time that Henderson began to experiment with jazz-funk fusion, studio overdubbing, and other electronic effects. Song and album titles like Power To the People, In Pursuit of Blackness, and Black Narcissus reflected his growing political awareness and social consciousness, although the last album was named after the Powell and Pressburger film of 1947.

After a brief association with Blood, Sweat & Tears in 1971, Henderson moved to San Francisco and added teaching to his résumé. He continued to record and perform as always, but seemed to be taken for granted by jazz audiences.

Later career and death[link]

Though he occasionally worked with Echoes of an Era, the Griffith Park Band and Chick Corea, Henderson remained primarily a leader throughout the 1980s. An accomplished and prolific composer, he began to focus more on reinterpreting standards and his own earlier compositions. Blue Note attempted to position the artist at the forefront of a resurgent jazz scene in 1986 with the release of the two-volume State of the Tenor recorded at the Village Vanguard in New York City. The albums (with Ron Carter on bass and Al Foster on drums) revisited the tenor trio form used by Sonny Rollins in 1957 on his own live Vanguard albums for the same label. Henderson established his basic repertoire for the next seven or eight years, with Monk's "Ask Me Now" becoming a signature ballad feature.

It was only after the release of An Evening with Joe Henderson, a live trio set (featuring Charlie Haden and Al Foster) for the Italian independent label Red Records that Henderson underwent a major career change: Verve took notice of him and in the early 1990s signed him. That label adopted a 'songbook' approach to recording him, coupling it with a considerable marketing and publicity campaign, which more successfully positioned Henderson at the forefront of the contemporary jazz scene. His 1992 'comeback' album Lush Life: The Music of Billy Strayhorn was a commercial and critical success and followed by tribute albums to Miles Davis, Antonio Carlos Jobim and a rendition of the George Gershwin opera Porgy and Bess.

On June 30, 2001, Joe Henderson died due to heart failure after a long battle with emphysema.^[3]

Discography[link]

As leader[link]

Artist rendering of Joe Henderson

Linoleum Print of the late Joe Henderson by Stuart Armstrong

Blue Note Records

• 1963: Page One
• 1963: Our Thing
• 1964: In 'n Out
• 1964: Inner Urge
• 1966: Mode for Joe
• 1985: The State of the Tenor: Live at the Village Vanguard, Vols. 1 & 2

Milestone Records

• 1967: The Kicker
• 1968: Tetragon
• 1969: Power to the People
• 1970: If You're Not Part of the Solution, You're Part of the Problem
• 1971: In Pursuit of Blackness
• 1971: Joe Henderson in Japan
• 1972: Black Is the Color
• 1973: Multiple
• 1973: The Elements
• 1975: Canyon Lady
• 1976: Black Miracle
• 1976: Black Narcissus

Verve Records

• 1968: Four
• 1968: Straight, No Chaser
• 1992: Lush Life: The Music of Billy Strayhorn
• 1992: So Near, So Far (Musings for Miles)
• 1994: Double Rainbow: The Music of Antonio Carlos Jobim
• 1996: Big Band
• 1997: Porgy & Bess

Red Records

• 1987: Evening with Joe Henderson - with Charlie Haden, Al Foster
• 1991: The Standard Joe - with Rufus Reid, Al Foster
• 2009: More from an Evening with Joe Henderson

Jazz Door

• 1973: 6tet/4tet - with Kenny Barron, Cedar Walton
• 1994: Live - with Bheki Mseleku, George Mraz, Al Foster
• 2001: Sunrise in Tokyo: Live in 1971- with Terumasa Hino, Masabumi Kikuchi

Other labels

• 1977: Barcelona (Enja Records) - with Wayne Darling, Ed Soph
• 1979: Relaxin' at Camarillo (Contemporary Records) with Chick Corea, either Tony Dumas or Richard Davis on bass, Peter Erskine or Tony Williams drums
• 1980: Mirror, Mirror (Pausa Records) with Chick Corea, Ron Carter, Billy Higgins
• 1999: Warm Valley (West Wind Records) - with Tony Martucci, Tommy Cecil, Louis Scherr

As sideman[link]

• 1963: Kenny Dorham - Una Mas (Blue Note)
• 1963: Andrew Hill - Black Fire (Blue Note)
• 1963: Johnny Coles - Little Johnny C (Blue Note)
• 1963: Grant Green - Am I Blue (Blue Note)
• 1963: Grant Green - Idle Moments (Blue Note)
• 1963: Blue Mitchell - Step Lightly (Blue Note)
• 1963: Antonio Diaz "Chocolaté" Mena - Eso Es Latin Jazz...Man!
• 1963: Bobby Hutcherson - The Kicker (Blue Note - released 1999)
• 1964: Grant Green- Solid (Blue Note)
• 1964: Lee Morgan - The Sidewinder (Blue Note)
• 1964: Kenny Dorham - Trompeta Toccata (Blue Note)
• 1964: Horace Silver - Song for My Father (Blue Note)
• 1964: Andrew Hill - Black Fire (Blue Note)
• 1964: Andrew Hill - Point Of Departure (Blue Note)
• 1964: Freddie Roach - Brown Sugar (Blue Note)
• 1964: Duke Pearson - Wahoo! (Blue Note)
• 1965: Andrew Hill - Pax (Blue Note)
• 1965: Pete La Roca - Basra (Blue Note)
• 1965: Larry Young - Unity (Blue Note)
• 1965: Horace Silver - The Cape Verdean Blues (Blue Note)
• 1966: Nat Adderley - Sayin' Somethin'
• 1966: Duke Pearson - Sweet Honey Bee (Blue Note)
• 1967: McCoy Tyner - The Real McCoy (Blue Note)
• 1969: George Benson - Tell It Like It Is (A&M/CTI)
• 1969: Herbie Hancock - The Prisoner (Blue Note)
• 1969: Herbie Hancock - Fat Albert Rotunda
• 1970: Freddie Hubbard - Red Clay
• 1970: Freddie Hubbard - Straight Life
• 1970: Alice Coltrane - Ptah, the El Daoud
• 1971: Blue Mitchell - Vital Blue
• 1972: Miroslav Vitouš - Mountain In The Clouds
• 1973: Ron Carter - All Blues (CTI)
• 1973: Flora Purim - Butterfly Dreams
• 1974: Johnny Hammond - Higher Ground
• 1974: Charles Earland - Leaving This Planet
• 1974: Patrice Rushen - Prelusion
• 1976: Coke Escovedo - Comin' At Ya!
• 1976: Roy Ayers - Daddy Bug & Friends
• 1977: Flora Purim - Encounter
• 1977: Woody Shaw - Rosewood (Columbia)
• 1978: Freddie Hubbard - Super Blue
• 1979: Art Farmer - Yama (CTI)
• 1982: Mal Waldron - One Entrance, Many Exits
• 1989: Jon Ballantyne - Sky Dance (Justin Time)
• 1991: Wynton Marsalis - Thick In The South: Soul Gestures In Southern Blue, Vol. 1
• 1991: McCoy Tyner - New York Reunion
• 1992: Kenny Garrett - Black Hope
• 1994: Kitty Margolis - Evolution
• 1994: Roy Hargrove - With the Tenors of Our Time
• 1995: Shirley Horn - The Main Ingredient
• 1999: Terence Blanchard - Jazz in Film
• 2004: Charlie Haden/Joe Henderson/Al Foster - The Montreal Tapes: Tribute to Joe Henderson (Verve, 1989)

References[link]

External links[link]

• The Joe Henderson Discography
• "Twelve Essential Joe Henderson Tracks" by S. Victor Aaron (Jazz.com)
• Joe Henderson "Lush Life" Solo: Transcription and Analysis