Chemical element

Chemical element
The periodic table of the chemical elements

A chemical element is a pure chemical substance consisting of one type of atom distinguished by its atomic number, which is the number of protons in its nucleus.[1] Familiar examples of elements include carbon, oxygen, aluminum, iron, copper, gold, mercury, and lead.

As of November 2011, 118 elements have been identified, the latest being ununoctium in 2002.[2] Of the 118 known elements, only the first 94 are known to occur naturally on Earth (88 in non-trace amounts). Of these, 80 are stable or essentially so, while the others are radioactive, decaying into lighter elements over various timescales from fractions of a second to billions of years. Additional elements, of higher atomic numbers than those naturally occurring, have been produced artificially.

Hydrogen and helium are by far the most abundant elements in the universe. However, iron is the most abundant element (by mass) making up the Earth, and oxygen is the most common element in the Earth's crust.[3] Although all known chemical matter is composed of these elements, chemical matter itself constitutes only about 15% of the matter in the universe. The remainder is dark matter, a mysterious substance which is not composed of chemical elements since it lacks protons, neutrons or electrons.

The chemical elements are thought to have been produced by various cosmic processes, including hydrogen, helium (and smaller amounts of lithium, beryllium and boron) created during the Big Bang and cosmic-ray spallation. Production of heavier elements, from carbon to the very heaviest elements, proceeds by stellar nucleosynthesis, and these were made available for later solar system and planetary formation by supernovae, which blast these elements into space.[4] The high abundance of oxygen, silicon, and iron on Earth reflect their common production in such stars, after the lighter gaseous elements and their compounds have been subtracted. While most elements are generally viewed as stable, a small amount of natural transformation of one element to another also occurs in the present time, through decay of radioactive elements as well as other natural nuclear processes.

Relatively pure samples of isolated elements are uncommon in nature. While all of the 94 naturally occurring elements have been identified in mineral samples from the Earth's crust, only a small minority of elements are found as recognizable, relative pure minerals. Among the more common of such "native elements" are copper, silver, gold, carbon (as coal, graphite, or diamonds), sulfur, and mercury. All but a few of the most inert elements, such as noble gases and noble metals, are usually found on Earth in chemically combined form, as chemical compounds. While about 32 of the chemical elements occur on Earth in native uncombined form, most of these occur as mixtures. For example, atmospheric air is primarily a mixture of nitrogen, oxygen, and argon, and native solid elements occur in alloys, such as that of iron and nickel.

When two distinct elements are chemically combined, with the atoms held together by chemical bonds, the result is termed a chemical compound. Two thirds of the chemical elements occur on Earth only as compounds, and in the remaining third, often the compound forms of the element are most common. Chemical compounds may be composed of elements combined in exact whole-number ratios of atoms, as in water, table salt, and minerals as quartz, calcite, and some ores. However, chemical bonding of many types of elements results in crystalline solids and metallic alloys for which exact chemical formulas do not exist. Most of the solid substance of the Earth is of this latter type. The substances that form Earth's crust, mantle, and core do not have precise chemical empirical formulas.

The history of discovery and use of the elements began with primitive human societies that found native elements like copper and gold, and extracted (smelted) iron and a few other metals from their ores. Alchemists and chemists subsequently identified many more, with nearly all of the naturally-occurring elements known by 1900. The properties of the chemical elements are often summarized using the periodic table that organizes the elements by increasing atomic number into rows ("periods") in which the columns ("groups") share recurring ("periodic") physical and chemical properties. Either in its pure forms, or in various chemical compounds or mixtures, almost every element has at least one important human use. Save for short half-lived radioactive elements, all of the elements are available industrially, most to high degrees of purity.

Around two dozen of the elements are essential to various kinds of biological life. Most rare elements on Earth are not needed by life (exceptions being selenium and iodine), while a few quite common ones (aluminum and titanium) are not used. Most organisms share element needs, with a few differences. For example, ocean algae use bromine but land plants and animals seem to need none, and all animals require sodium, but some plants do not. Just six elements—carbon, hydrogen, nitrogen, oxygen, calcium, and phosphorus—make up almost 99% of the mass of a human body (see composition of the human body for a complete list). In addition to the six major elements that compose most of the human body, humans require consumption of at least a dozen more elements in the form of certain chemical compounds.



The lightest of the chemical elements are hydrogen and helium, both created by Big Bang nucleosynthesis during the first 20 minutes of the universe[5] in a ratio of around 3:1 by mass (approximately 12:1 by number of atoms). Almost all other elements found in nature, including some further hydrogen and helium created since then, were made by various natural or (at times) artificial methods of nucleosynthesis. On Earth, small amounts of new atoms are naturally produced in nucleogenic reactions, or in cosmogenic processes, such as cosmic ray spallation. New atoms are also naturally produced on Earth as radiogenic daughter isotopes of ongoing radioactive decay processes such as alpha decay, beta decay, spontaneous fission, cluster decay, and other rarer modes of decay.

Of the 94 naturally occurring elements, those with atomic numbers 1 through 40 are all considered to be stable isotopes. Elements with atomic numbers 41 through 82 are apparently stable (except technetium, element 43 and promethium, element 61) but theoretically unstable, and thus possibly mildly radioactive. The half-lives of elements 41 through 82 are so long however that their radioactive decay has yet to be detected by experiment. These "theoretical radionuclides" have half-lives at least 100 million times longer than the estimated age of the universe. Elements with atomic numbers 83 through 94 are unstable to the point that their radioactive decay can be detected. Some of these elements, notably thorium (atomic number 90) and uranium (atomic number 92), have one or more isotopes with half-lives long enough to survive as remnants of the explosive stellar nucleosynthesis that produced the heavy elements before the formation of our solar system. For example, at over 1.9×1019 years, over a billion times longer than the current estimated age of the universe, bismuth-209 (atomic number 83) has the longest known alpha decay half-life of any naturally occurring element.[6][7] The very heaviest elements (those beyond plutonium, atomic number 94) undergo radioactive decay with half-lives so short that they have only been observed as the result of experimental observation.

As of 2010, there are 118 known elements (in this context, "known" means observed well-enough, even from just a few decay products, to have been differentiated from any other element).[8][9] Of these 118 elements, 94 occur naturally on Earth. Six of these occur in extreme trace quantities: technetium, atomic number 43; promethium, number 61; astatine, number 85; francium, number 87; neptunium, number 93; and plutonium, number 94. These 94 elements have been detected in the universe at large, in the spectra of stars and also supernovae, where short-lived radioactive elements are newly being made. The first 94 elements have been detected directly on Earth as primordial nuclides present from the formation of the solar system, or as naturally-occurring fission or transmutation products of uranium and thorium.

The remaining 24 heavier elements, not found today either on Earth or in astronomical spectra, have been derived artificially. All of the heavy elements that are derived solely through artificial means are radioactive, with very short half-lives; if any atoms of these elements were present at the formation of Earth, they are extremely likely to have already decayed, and if present in novae, have been in quantities too small to have been noted. Technetium was the first purportedly non-naturally occurring element to be synthesized, in 1937, although trace amounts of technetium have since been found in nature (and also the element may have been discovered naturally in 1925). This pattern of artificial production and later natural discovery has been repeated with several other radioactive naturally-occurring rare elements.

Lists of the elements are available by name, by symbol, by atomic number, by density, by melting point, and by boiling point as well as Ionization energies of the elements. The nuclides of stable and radioactive elements are also available as a list of nuclides, sorted by length of half-life for those that are unstable. One of the most convenient, and certainly the most traditional presentation of the elements, is in form of periodic table, which groups elements with similar chemical properties (and usually also similar electronic structures) together.

Atomic number

The atomic number of an element is equal to the number of protons that defines the element. For example, all carbon atoms contain 6 protons in their nucleus; so the atomic number of carbon is 6. Carbon atoms may have different numbers of neutrons; atoms of the same element having different numbers of neutrons are known as isotopes of the element.

The number of protons in the atomic nucleus also determines its electric charge, which in turn determines the number of electrons of the atom in its non-ionized state. The electrons are placed into atomic orbitals which determine the atom's various chemical properties. The number of neutrons in a nucleus usually has very little effect on an elements' chemical properties (except in the case of hydrogen and deuterium). Thus, all carbon isotopes have nearly identical chemical properties because they all have six protons and six electrons, even though carbon atoms may differ in number of neutrons. It is for this reason that atomic number rather than mass number or atomic weight is considered the identifying characteristic of a chemical element. The symbol for atomic number is Z.

Atomic mass and atomic weight

The mass number of an element, A, is the number of nucleons (protons and neutrons) in the atomic nucleus. Different isotopes of a given element are distinguished by their mass numbers, which are conventionally written as a super-index on the left hand side of the atomic symbol (e.g., 238U). The mass number is always a simple whole number and has units of "nucleons." An example of use of a mass number is "magnesium-24," which has 24 nucleons (12 protons and 12 neutrons).

Whereas the mass number simply counts the total number of neutrons and protons and is thus a natural (or whole) number, the atomic mass of a single isotope is a real number. In general, it differs in value when expressed in u for a given nuclide (or isotope) slightly from the mass number, since the mass of the protons and neutrons is not exactly 1 u, the electrons contribute a lesser share to the atomic mass as neutron number exceeds proton number, and (finally) because of the nuclear binding energy. For example, the atomic weight of chlorine-35 to five significant digits is 34.969 u and that of chlorine-37 is 36.966 u. However, the atomic mass in u of pure isotope atoms is quite close (always within 1%) to its simple mass number. The only exception to the atomic mass of an isotope atom not being a natural number is 12C, which has a mass of exactly 12 by definition, because u is defined as 1/12 of the mass of a free neutral carbon-12 atom in the ground state.

The relative atomic mass (historically and commonly also called "atomic weight") of an element is the average of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance, relative to the atomic mass unit (u). This number may be a fraction that is not close to a whole number, due to the averaging process. For example, the relative atomic mass of chlorine is 35.453 u, which differs greatly from a whole number due to being made of an average of 76% chlorine-35 and 24% chlorine-37. Whenever a relative atomic mass value differs by more than 1% from a whole number, it is due to this averaging effect resulting from significant amounts of more than one isotope being naturally present in the sample of the element in question.


Isotopes are atoms of the same element (that is, with the same number of protons in their atomic nucleus), but having different numbers of neutrons. Most (66 of 94) naturally occurring elements have more than one stable isotope. Thus, for example, there are three main isotopes of carbon. All carbon atoms have 6 protons in the nucleus, but they can have either 6, 7, or 8 neutrons. Since the mass numbers of these are 12, 13 and 14 respectively, the three isotopes of carbon are known as carbon-12, carbon-13, and carbon-14, often abbreviated to 12C, 13C, and 14C. Carbon in everyday life and in chemistry is a mixture of 12C, 13C, and (a very small fraction of) 14C atoms.

Except in the case of the isotopes of hydrogen (which differ greatly from each other in relative mass—enough to cause chemical effects), the isotopes of the various elements are typically chemically nearly indistinguishable from each other. For example, the three naturally-occurring isotopes of carbon have essentially the same chemical properties, but different nuclear properties. In this example, carbon-12 and carbon-13 are stable atoms, but carbon-14 is unstable; it is radioactive, undergoing beta decay into nitrogen-14.

All of the elements have some isotopes that are radioactive (radioisotopes), although not all of these radioisotopes occur naturally. The radioisotopes typically decay into other elements upon radiating an alpha or beta particle. If an element has isotopes that are not radioactive, they are termed "stable." All of the known stable isotopes occur naturally (see primordial isotope). The many radioisotopes that are not found in nature have been characterized from being artificially made. Certain elements have no stable isotopes and are composed only of radioactive isotopes: specifically the elements without any stable isotopes are technetium (atomic number 43), promethium (atomic number 61), and all observed elements with atomic numbers greater than 82.

Of the 80 elements with at least one stable isotope, 26 have only one stable isotope, and the mean number of stable isotopes for the 80 stable elements is 3.1 stable isotopes per element. The largest number of stable isotopes that occur for an element is 10 (for tin, element 50).


Atoms of pure elements may bond to each other chemically in more than one way, allowing the pure element to exist in multiple structures (spacial arrangements of atoms), known as allotropes, which differ in their properties. For example, carbon can be found as diamond, which has a tetrahedral structure around each carbon atom; graphite, which has layers of carbon atoms with a hexagonal structure stacked on top of each other; graphene, which is a single layer of graphite that is incredibly strong; fullerenes, which have nearly spherical shapes; and carbon nanotubes, which are tubes with a hexagonal structure (even these may differ from each other in electrical properties). The ability for an element to exist in one of many structural forms is known as 'allotropy'.

The standard state, or reference state, of an element is defined as its thermodynamically most stable state at 1 bar at a given temperature (typically at 298.15 K). In thermochemistry, an element is defined to have an enthalpy of formation of zero in its standard state. For example, the reference state for carbon is graphite, because it is more stable than the other allotropes.


Several kinds of descriptive categorizations can be applied broadly to the elements, including consideration of their general physical and chemical properties, their states of matter under familiar conditions, their melting and boiling points, their densities, their crystal structures as solids, and their origins.

General properties

Several terms are commonly used to characterize the general physical and chemical properties of the chemical elements. A first distinction is between the metals, which readily conduct electricity, and the nonmetals, which do not, with a small group (the metalloids) having intermediate properties, often behaving as semiconductors.

A more refined classification is often shown in colored presentations of the periodic table; this system restricts the terms "metal" and "nonmetal" to only certain of the more broadly defined metals and nonmetals, adding additional terms for certain sets of the more broadly viewed metals and nonmetals. The version of this classification used in the periodic tables presented here includes: actinides, alkali metals, alkaline earth metals, halogens, lanthanides, metals (or "other metals"), metalloids, noble gases, nonmetals (or "other nonmetals"), and transition metals. In this system, the alkali metals, alkaline earth metals, and transition metals, as well as the lanthanides and the actinides, are special groups of the metals viewed in a broader sense. Similarly, the halogens and the noble gases are nonmetals, viewed in the broader sense. In some presentations, the halogens are not distinguished, with astatine identified as a metalloid and the others identified as nonmetals.

States of matter

Another commonly used basic distinction among the elements is their state of matter (phase), solid, liquid, or gas, at a selected standard temperature and pressure (STP). Most of the elements are solids at conventional temperatures and atmospheric pressure, while several are gases. Only bromine and mercury are liquids at 0 degrees Celsius (32 degrees Fahrenheit) and normal atmospheric pressure; caesium and gallium are solids at that temperature, but melt at 28.4 °C (83.2 °F) and 29.8 °C (85.6 °F), respectively.

Melting and boiling points

Melting and boiling points, typically expressed in degrees Celsius at a pressure of one atmosphere, are commonly used in characterizing the various elements. While known for most elements, either or both of these measurements is still undetermined for some of the radioactive elements available in only tiny quantities. Since helium remains a liquid even at absolute zero at atmospheric pressure, it has only a boiling point, and not a melting point, in conventional presentations.


The density at a selected standard temperature and pressure (STP) is frequently used in characterizing the elements. Density is often expressed in grams per cubic centimeter (g/cm3). Since several elements are gases at commonly encountered temperatures, their densities are usually stated for their gaseous forms; when liquefied or solidified, the gaseous elements have densities similar to those of the other elements.

When an element has allotropes with different densities, one representative allotrope is typically selected in summary presentations, while densities for each allotrope can be stated where more detail is provided. For example, the three familiar allotropes of carbon (amorphous carbon, graphite, and diamond) have densities of 1.8–2.1, 2.267, and 3.515 g/cm3 respectively.

Crystal structures

The elements studied to date as solid samples have eight kinds of crystal structures: cubic, body-centered cubic, face-centered cubic, hexagonal, monoclinic, orthorhombic, rhombohedral, and tetragonal. For some of the synthetically produced transuranic elements, available samples have been too small to determine crystal structures.

Occurrence and origin on Earth

The elements may also be categorized by their origin on Earth, with the first 94 considered naturally occurring, and those with atomic numbers beyond 94 being synthetic (produced technologically, but not known to occur naturally).

Of the naturally occurring elements, 84 are considered primordial, either stable or long-persisting, and the remaining 10 are transient elements with half lives too short to have been present at the beginning of the solar system. These are produced either "recurrently" (often found) or "incidentally" (very rare) as decay products or through other natural nuclear reactionprocesses. The 91 regularly occurring natural elements include the 80 stable or essentially stable, primordial elements (from hydrogen through lead, omitting technetium and promethium); bismuth, thorium, uranium, and plutonium (radioactive but still remaining from primordial times); and the 7 transiently existing but recurrently produced decay products of thorium, uranium, and plutonium (polonium, astatine, radon, francium, radium, actinium, and protactinium).

Three additional naturally occurring elements, technetium, promethium, and neptunium, are only incidentally occurring, present in natural materials only as transiently existing atoms produced from uranium or other heavy elements by rare nuclear processes. Note that helium is recurrently produced naturally from radioactive decay, but little if any primordial helium still exists at the Earth's surface, since this light gas readily escapes from the atmosphere into outer space.

The periodic table

The properties of the chemical elements are often summarized using the periodic table, which powerfully and elegantly organizes the elements by increasing atomic number into rows ("periods") in which the columns ("groups") share recurring ("periodic") physical and chemical properties. The current standard table contains 118 confirmed elements as of April 10, 2010.

Although earlier precursors to this presentation exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869, who intended the table to illustrate recurring trends in the properties of the elements. The layout of the table has been refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain chemical behavior.

Use of the periodic table is now ubiquitous within the academic discipline of chemistry, providing an extremely useful framework to classify, systematize and compare all the many different forms of chemical behavior. The table has also found wide application in physics, geology, biology, materials science, engineering, agriculture, medicine, nutrition, environmental health, and astronomy. Its principles are especially important in chemical engineering.

Nomenclature and symbols

The various chemical elements are formally identified by their unique atomic numbers, by their accepted names, and by their symbols.

Atomic numbers

The known elements have atomic numbers from 1 through 118, conventionally presented as Arabic numerals.

Since the elements can be uniquely sequenced by atomic number, conventionally from lowest to hightest (as in a periodic table), sets of elements are sometimes specified by such notation as "through", "beyond", or "from ... through", as in "through iron", "beyond uranium", or "from lanthanum through lutetium". The terms "light" and "heavy" are sometimes also used informally to indicate relative atomic numbers (not densities!), as in "lighter than carbon" or "heavier than lead", although technically the weight or mass of atoms of an element (their atomic weights or atomic masses) do not always increase monotonically with their atomic numbers.

Element names

The naming of various substances now known as elements precedes the atomic theory of matter, as names were given locally by various cultures to various minerals, metals, compounds, alloys, mixtures, and other materials, although at the time it was not known which chemicals were elements and which compounds. As they were identified as elements, the existing names for anciently-known elements (e.g., gold, mercury, iron) were kept in most countries. National differences emerged over the names of elements either for convenience, linguistic niceties, or nationalism. For a few illustrative examples: German speakers use "Wasserstoff" (water substance) for "hydrogen", "Sauerstoff" (acid substance) for "oxygen" and "Stickstoff" (smothering substance) for "nitrogen", while English and some romance languages use "sodium" for "natrium" and "potassium" for "kalium", and the French, Italians, Greeks, Portuguese and Poles prefer "azote/azot/azoto" (from roots meaning "no life") for "nitrogen".

For purposes of international communication and trade, the official names of the chemical elements both ancient and more recently recognized are decided by the International Union of Pure and Applied Chemistry (IUPAC), which has decided on a sort of international English language, drawing on traditional English names even when an element's chemical symbol is based on a Latin or other traditional word, for example adopting "gold" rather than "aurum" as the name for the 79th element (Au). IUPAC prefers the British spellings "aluminium" and "caesium" over the U.S. spellings "aluminum" and "cesium", and the U.S. "sulfur" over the British "sulphur". However, elements that are practical to sell in bulk in many countries often still have locally used national names, and countries whose national language does not use the Latin alphabet are likely to use the IUPAC element names.

According to IUPAC, chemical elements are not proper nouns in English; consequently, the full name of an element is not routinely capitalized in English, even if derived from a proper noun, as in californium and einsteinium. Isotope names of chemical elements are also uncapitalized if written out, e.g., carbon-12 or uranium-235. Chemical element symbols are always capitalized (see below).

In the second half of the twentieth century, physics laboratories became able to produce nuclei of chemical elements with half-lives too short for an appreciable amount of them to exist at any time. These are also named by IUPAC, which generally adopts the name chosen by the discoverer. This practice can lead to the controversial question of which research group actually discovered an element, a question that has delayed naming of elements with atomic number of 104 and higher for a considerable time. (See element naming controversy).

Precursors of such controversies involved the nationalistic namings of elements in the late 19th century. For example, lutetium was named in reference to Paris, France. The Germans were reluctant to relinquish naming rights to the French, often calling it cassiopeium. Similarly, the British discoverer of niobium originally named it columbium, in reference to the New World. It was used extensively as such by American publications prior to international standardization.

Chemical symbols

Specific chemical elements

Before chemistry became a science, alchemists had designed arcane symbols for both metals and common compounds. These were however used as abbreviations in diagrams or procedures; there was no concept of atoms combining to form molecules. With his advances in the atomic theory of matter, John Dalton devised his own simpler symbols, based on circles, which were to be used to depict molecules.

The current system of chemical notation was invented by Berzelius. In this typographical system chemical symbols are not used as mere abbreviations – though each consists of letters of the Latin alphabet – they are symbols intended to be used by peoples of all languages and alphabets. The first of these symbols were intended to be fully universal; since Latin was the common language of science at that time, they were abbreviations based on the Latin names of metals – Cu comes from Cuprum, Fe comes from Ferrum, Ag from Argentum. The symbols were not followed by a period (full stop) as abbreviations were. Later chemical elements were also assigned unique chemical symbols, based on the name of the element, but not necessarily in English. For example, sodium has the chemical symbol 'Na' after the Latin natrium. The same applies to "W" (wolfram) for tungsten, "Fe" (ferrum) for iron, "Hg" (hydrargyrum) for mercury, "Sn" (stannum) for tin, "K" (kalium) for potassium, "Au" (aurum) for gold, "Ag" (argentum) for silver, "Pb" (plumbum) for lead, "Cu" (Cuprum) for copper, and "Sb" (stibium) for antimony.

Chemical symbols are understood internationally when element names might need to be translated. There are sometimes differences; for example, the Germans have used "J" instead of "I" for iodine, so the character would not be confused with a Roman numeral.

The first letter of a chemical symbol is always capitalized, as in the preceding examples, and the subsequent letters, if any, are always lower case (small letters). Thus, the symbols for californium or einsteinium are Cf and Es.

General chemical symbols

There are also symbols for series of chemical elements, for comparative formulas. These are one capital letter in length, and the letters are reserved so they are not permitted to be given for the names of specific elements. For example, an "X" is used to indicate a variable group amongst a class of compounds (though usually a halogen), while "R" is used for a radical, meaning a compound structure such as a hydrocarbon chain. The letter "Q" is reserved for "heat" in a chemical reaction. "Y" is also often used as a general chemical symbol, although it is also the symbol of yttrium. "Z" is also frequently used as a general variable group. "L" is used to represent a general ligand in inorganic and organometallic chemistry. "M" is also often used in place of a general metal.

At least one additional, two-letter generic chemical symbol is also in informal usage, "Ln" for any lanthanide element.

Isotope symbols

Isotopes are distinguished by the atomic mass number (total protons and neutrons) for a particular isotope of an element, with this number combined with the pertinent element's symbol. IUPAC prefers that isotope symbols be written in superscript notation when practical, for example 12C and 235U. However, other notations, such as carbon-12 and uranium-235, or C-12 and U-235, are also used.

As a special case, the three naturally occurring isotopes of the element hydrogen are often specified as H for 1H (protium), D for 2H (deuterium), and T for 3H (tritium). This convention is easier to use in chemical equations, replacing the need to write out the mass number for each atom. For example, the formula for heavy water may be written D2O instead of 2H2O.

Origin of the elements

Estimated distribution of dark matter and dark energy in the universe. Only the fraction of the mass and energy in the universe labeled "atoms" is composed of chemical elements.

Only about 4% of the total mass of the universe is made of atoms or ions, and thus represented by chemical elements. This fraction is about 15% of the total matter, with the remainder of the matter (85%) being dark matter. The nature of dark matter is unknown, but it is not composed of atoms of chemical elements because it contains no protons, neutrons, or electrons. (The remaining non-matter part of the mass of the universe is composed of the even more mysterious dark energy).

The universe's 94 naturally occurring chemical elements are thought to have been produced by at least four cosmic processes. Most of the hydrogen and helium in the universe was produced primordially in the first few minutes of the Big Bang. Three recurrently occurring later processes are thought to have produced the remaining elements. Stellar nucleosynthesis, an ongoing process, produces all elements from carbon through iron in atomic number, but little lithium, beryllium, or boron. Elements heavier in atomic number than iron, as heavy as uranium and plutonium, are produced by explosive nucleosynthesis in supernovas and other cataclysmic cosmic events. Cosmic ray spallation (fragmentation) of carbon, nitrogen, and oxygen is important to the production of lithium, beryllium and boron.

During the early phases of the Big Bang, nucleosynthesis of hydrogen nuclei resulted in the production of hydrogen-1 (protonium, 1H) and helium-4 (4He), as well as a smaller amount of deuterium (2H) and very minuscule amounts (on the order of 10−10) of lithium and beryllium. Even smaller amounts of boron may have been produced in the Big Bang, since it has been observed in some very old stars, while carbon has not.[10] It is generally agreed that no heavier elements than boron were produced in the Big Bang. As a result, the primordial abundance of atoms (or ions) consisted of roughly 75% 1H, 25% 4He, and 0.01% deuterium, with only tiny traces of lithium, beryllium, and perhaps boron.[11] Subsequent enrichment of galactic halos occurred due to stellar nucleosynthesis and supernova nucleosynthesis.[12] However, the element abundance in intergalactic space can still closely resemble primordial conditions, unless it has been enriched by some means.

On Earth (and elsewhere), trace amounts of various elements continue to be produced from other elements as products of natural transmutation processes. These include some produced by cosmic rays or other nuclear reactions (see cosmogenic and nucleogenic nuclides), and others produced as decay products of long-lived primordial nuclides.[13] For example, trace (but detectable) amounts of carbon-14 (14C) are continually produced in the atmosphere by cosmic rays impacting nitrogen atoms, and argon-40 (40Ar) is continually produced by the decay of primordially occurring but unstable potassium-40 (40K). Also, three primordially occurring but radioactive actinides, thorium, uranium, and plutonium, decay through a series of recurrently produced but unstable radioactive elements such as radium and radon, which are transiently present in any sample of these metals or their ores or compounds. Three other radioactive elements, technetium, promethium, and neptunium, occur only incidentally in natural materials, produced as individual atoms by natural fission of the nuclei of various heavy elements or in other rare nuclear processses.

Human technology has produced various additional elements beyond these first 94, with those through atomic number 118 now known.


The following graph (note log scale) shows abundance of elements in our solar system. The table shows the twelve most common elements in our galaxy (estimated spectroscopically), as measured in parts per million, by mass.[14] Nearby galaxies that have evolved along similar lines have a corresponding enrichment of elements heavier than hydrogen and helium. The more distant galaxies are being viewed as they appeared in the past, so their abundances of elements appear closer to the primordial mixture. As physical laws and processes appear common throughout the visible universe, however, it is expected that these galaxies will likewise have evolved similar abundances of elements.

The abundance of the chemical elements on Earth varies from air to crust to ocean, and in various types of life. The abundance of elements in Earth's crust differs from those in the universe mainly in selective loss of the very lightest elements (hydrogen and helium) and also volatile neon, carbon, nitrogen and sulfur, as a result of solar heating in the early formation of the solar system. Aluminum is also far more common in the Earth and Earth's crust than the universe and solar system. The composition of the human body, by contrast, more closely follows the composition of seawater, save that the human body has additional stores of carbon and nitrogen which are necessary to form the proteins and nucleic acids that are characteristic of living organisms. Certain kinds of organisms require particular additional elements, for example the magnesium in chlorophyll in green plants, the calcium in mollusc shells, or the iron in the hemoglobin in vertebrate animals' red blood cells.

Abundances of the chemical elements in the Solar system. Hydrogen and helium are most common, from the Big Bang. The next three elements (Li, Be, B) are rare because they are poorly synthesized in the Big Bang and also in stars. The two general trends in the remaining stellar-produced elements are: (1) an alternation of abundance in elements as they have even or odd atomic numbers, and (2) a general decrease in abundance, as elements become heavier.
Elements in our galaxy Parts per million
by mass
Hydrogen 739,000
Helium 240,000
Oxygen 10,400
Carbon 4,600
Neon 1,340
Iron 1,090
Nitrogen 960
Silicon 650
Magnesium 580
Sulfur 440
Potassium 210
Nickel 100

Periodic table highlighting dietary elements[15]

H   He
Li Be   B C N O F Ne
Na Mg   Al Si P S Cl Ar
K Ca Sc   Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y   Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
Cs Ba La * Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac ** Rf Db Sg Bh Hs Mt Ds Rg
  * Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
  ** Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
The four organic basic elements Quantity elements Essential trace elements Suggested function from biochemistry and handling but no identified biological function in humans


Mendeleev's 1869 periodic table

Evolving definitions

The concept of an "element" as an undivisible substance has developed through three major historical phases: Classical definitions (such as those of the ancient Greeks), chemical definitions, and atomic definitions.

Classical definitions

Ancient philosophy posited a set of classical elements to explain observed patterns in nature. These elements originally referred to earth, water, air and fire rather than the chemical elements of modern science.

The term 'elements' (stoicheia) was first used by the Greek philosopher Plato in about 360 BCE, in his dialogue Timaeus, which includes a discussion of the composition of inorganic and organic bodies and is a speculative treatise on chemistry. Plato believed the elements introduced a century earlier by Empedocles were composed of small polyhedral forms: tetrahedron (fire), octahedron (air), icosahedron (water), and cube (earth).[16][17]

Aristotle, c. 350 BCE, also used the term stoicheia and added a fifth element called aether, which formed the heavens. Aristotle defined an element as:

Element – one of those bodies into which other bodies can decompose, and that itself is not capable of being divided into other.[18]

Chemical definitions

In 1661, chemist Robert Boyle showed that there were more than just the four classical elements that the ancients had assumed.[19] The first modern list of chemical elements was given in Antoine Lavoisier's 1789 Elements of Chemistry, which contained thirty-three elements, including light and caloric.[20] By 1818, Jöns Jakob Berzelius had determined atomic weights for forty-five of the forty-nine then-accepted accepted elements. Dmitri Mendeleev had sixty-six elements in his periodic table of 1869.

From Boyle until the early 20th century, an element was defined as a pure substance that could not be decomposed into any simpler substance.[19] Put another way, a chemical element cannot be transformed into other chemical elements by chemical processes. Elements during this time were generally distinguished by their atomic weights, a property measurable with fair accuracy by available analytical techniques.

Atomic definitions

The 1913 discovery by Henry Moseley that the nuclear charge is the physical basis for an atom's atomic number, further refined when the nature of protons and neutrons became appreciated, eventually led to the current definition of an element, based on atomic number (number of protons per atomic nucleus). The use of atomic numbers, rather than atomic weights, to distinguish elements has greater predictive value (since these numbers are integers), and also resolves some ambiguities in the chemistry-based view due to varying properties of isotopes and allotropes within the same element. Currently IUPAC defines an element to exist if it has isotopes with a lifetime longer than the 10−14 seconds which takes the nucleus to form an electronic cloud.[21]

By 1914, seventy-two elements were known, all naturally occurring.[22] The remaining naturally occurring elements were discovered or isolated is subsequent decades, and various additional elements have also been produced synthetically, with much of that work pioneered by Glenn T. Seaborg. In 1955, element 101 was discovered and named mendelevium in honor of D.I. Mendeleev, the first to arrange the elements in a periodic manner. Most recently, the synthesis of element 118 was reported in October 2006, and the synthesis of element 117 was reported in April 2010.[23]

Discovery and recognition of various elements

Ten materials familiar to various prehistoric cultures are now known to be chemical elements: Carbon, copper, gold, iron, lead, mercury, silver, sulfur, tin, and zinc. Three additional materials now accepted as elements, arsenic, antimony, and bismuth, were recognized as distinct substances prior to 1500 AD. Phosphorus, cobalt, and platinum were isolated before 1750.

Most of the remaining naturally occurring chemical elements were identified and characterized by 1900, including:

Elements isolated or produced since 1900 include:

Recently discovered elements

The first transuranium element (element with atomic number greater than 92) discovered was neptunium in 1940. As of February 2010, only the elements up to 112, copernicium, have been confirmed as discovered by IUPAC, while claims have been made for synthesis of elements 113, 114, 115, 116, 117[24] and 118. The discovery of element 112 was acknowledged in 2009, and the name 'copernicium' and the atomic symbol 'Cn' were suggested for it.[25] The name and symbol were officially endorsed by IUPAC on February 19, 2010.[26] The heaviest element that is believed to have been synthesized to date is element 118, ununoctium, on October 9, 2006, by the Flerov Laboratory of Nuclear Reactions in Dubna, Russia.[9][27] Element 117 was the latest element claimed to be discovered, in 2009.[24] IUPAC officially recognized ununquadium and ununhexium, elements 114 and 116, in June 2011.[28]

List of the 118 known chemical elements

The following sortable table includes the 118 known chemical elements, with the names linking to the Wikipedia articles on each.

  • Atomic number, name, and symbol all serve independently as unique identifiers.
  • Names are those accepted by IUPAC; provisional names for recently produced elements not yet formally named are in parentheses.
  • Group, period, and block refer to an element's position in the periodic table.
  • State of matter (solid, liquid, or gas) applies at standard temperature and pressure conditions (STP).[citation needed]
  • Occurrence distinguishes naturally occurring elements, categorized as either primordial or transient (from decay), and additional synthetic elements that have been produced technologically, but are not known to occur naturally.
  • Description summarizes an element's properties using the broad categories commonly presented in periodic tables: Actinide, alkali metal, alkaline earth metal, halogen, lanthanide, metal, metalloid, noble gas, non-metal, and transition metal.
List of elements
Name Symbol Group Period Block State at
Occurrence Description
1 Hydrogen H 1 1 s Gas Primordial Non-metal
2 Helium He 18 1 s Gas Primordial Noble gas
3 Lithium Li 1 2 s Solid Primordial Alkali metal
4 Beryllium Be 2 2 s Solid Primordial Alkaline earth metal
5 Boron B 13 2 p Solid Primordial Metalloid
6 Carbon C 14 2 p Solid Primordial Non-metal
7 Nitrogen N 15 2 p Gas Primordial Non-metal
8 Oxygen O 16 2 p Gas Primordial Non-metal
9 Fluorine F 17 2 p Gas Primordial Halogen
10 Neon Ne 18 2 p Gas Primordial Noble gas
11 Sodium Na 1 3 s Solid Primordial Alkali metal
12 Magnesium Mg 2 3 s Solid Primordial Alkaline earth metal
13 Aluminium Al 13 3 p Solid Primordial Metal
14 Silicon Si 14 3 p Solid Primordial Metalloid
15 Phosphorus P 15 3 p Solid Primordial Non-metal
16 Sulfur S 16 3 p Solid Primordial Non-metal
17 Chlorine Cl 17 3 p Gas Primordial Halogen
18 Argon Ar 18 3 p Gas Primordial Noble gas
19 Potassium K 1 4 s Solid Primordial Alkali metal
20 Calcium Ca 2 4 s Solid Primordial Alkaline earth metal
21 Scandium Sc 3 4 d Solid Primordial Transition metal
22 Titanium Ti 4 4 d Solid Primordial Transition metal
23 Vanadium V 5 4 d Solid Primordial Transition metal
24 Chromium Cr 6 4 d Solid Primordial Transition metal
25 Manganese Mn 7 4 d Solid Primordial Transition metal
26 Iron Fe 8 4 d Solid Primordial Transition metal
27 Cobalt Co 9 4 d Solid Primordial Transition metal
28 Nickel Ni 10 4 d Solid Primordial Transition metal
29 Copper Cu 11 4 d Solid Primordial Transition metal
30 Zinc Zn 12 4 d Solid Primordial Transition metal
31 Gallium Ga 13 4 p Solid Primordial Metal
32 Germanium Ge 14 4 p Solid Primordial Metalloid
33 Arsenic As 15 4 p Solid Primordial Metalloid
34 Selenium Se 16 4 p Solid Primordial Non-metal
35 Bromine Br 17 4 p Liquid Primordial Halogen
36 Krypton Kr 18 4 p Gas Primordial Noble gas
37 Rubidium Rb 1 5 s Solid Primordial Alkali metal
38 Strontium Sr 2 5 s Solid Primordial Alkaline earth metal
39 Yttrium Y 3 5 d Solid Primordial Transition metal
40 Zirconium Zr 4 5 d Solid Primordial Transition metal
41 Niobium Nb 5 5 d Solid Primordial Transition metal
42 Molybdenum Mo 6 5 d Solid Primordial Transition metal
43 Technetium Tc 7 5 d Solid Transient Transition metal
44 Ruthenium Ru 8 5 d Solid Primordial Transition metal
45 Rhodium Rh 9 5 d Solid Primordial Transition metal
46 Palladium Pd 10 5 d Solid Primordial Transition metal
47 Silver Ag 11 5 d Solid Primordial Transition metal
48 Cadmium Cd 12 5 d Solid Primordial Transition metal
49 Indium In 13 5 p Solid Primordial Metal
50 Tin Sn 14 5 p Solid Primordial Metal
51 Antimony Sb 15 5 p Solid Primordial Metalloid
52 Tellurium Te 16 5 p Solid Primordial Metalloid
53 Iodine I 17 5 p Solid Primordial Halogen
54 Xenon Xe 18 5 p Gas Primordial Noble gas
55 Caesium Cs 1 6 s Solid Primordial Alkali metal
56 Barium Ba 2 6 s Solid Primordial Alkaline earth metal
57 Lanthanum La 3 6 f Solid Primordial Lanthanide
58 Cerium Ce 3 6 f Solid Primordial Lanthanide
59 Praseodymium Pr 3 6 f Solid Primordial Lanthanide
60 Neodymium Nd 3 6 f Solid Primordial Lanthanide
61 Promethium Pm 3 6 f Solid Transient Lanthanide
62 Samarium Sm 3 6 f Solid Primordial Lanthanide
63 Europium Eu 3 6 f Solid Primordial Lanthanide
64 Gadolinium Gd 3 6 f Solid Primordial Lanthanide
65 Terbium Tb 3 6 f Solid Primordial Lanthanide
66 Dysprosium Dy 3 6 f Solid Primordial Lanthanide
67 Holmium Ho 3 6 f Solid Primordial Lanthanide
68 Erbium Er 3 6 f Solid Primordial Lanthanide
69 Thulium Tm 3 6 f Solid Primordial Lanthanide
70 Ytterbium Yb 3 6 f Solid Primordial Lanthanide
71 Lutetium Lu 3 6 d Solid Primordial Lanthanide
72 Hafnium Hf 4 6 d Solid Primordial Transition metal
73 Tantalum Ta 5 6 d Solid Primordial Transition metal
74 Tungsten W 6 6 d Solid Primordial Transition metal
75 Rhenium Re 7 6 d Solid Primordial Transition metal
76 Osmium Os 8 6 d Solid Primordial Transition metal
77 Iridium Ir 9 6 d Solid Primordial Transition metal
78 Platinum Pt 10 6 d Solid Primordial Transition metal
79 Gold Au 11 6 d Solid Primordial Transition metal
80 Mercury Hg 12 6 d Liquid Primordial Transition metal
81 Thallium Tl 13 6 p Solid Primordial Metal
82 Lead Pb 14 6 p Solid Primordial Metal
83 Bismuth Bi 15 6 p Solid Primordial Metal
84 Polonium Po 16 6 p Solid Transient Metalloid
85 Astatine At 17 6 p Solid Transient Halogen
86 Radon Rn 18 6 p Gas Transient Noble gas
87 Francium Fr 1 7 s Solid Transient Alkali metal
88 Radium Ra 2 7 s Solid Transient Alkaline earth metal
89 Actinium Ac 3 7 f Solid Transient Actinide
90 Thorium Th 3 7 f Solid Primordial Actinide
91 Protactinium Pa 3 7 f Solid Transient Actinide
92 Uranium U 3 7 f Solid Primordial Actinide
93 Neptunium Np 3 7 f Solid Transient Actinide
94 Plutonium Pu 3 7 f Solid Primordial Actinide
95 Americium Am 3 7 f Solid Synthetic Actinide
96 Curium Cm 3 7 f Solid Synthetic Actinide
97 Berkelium Bk 3 7 f Solid Synthetic Actinide
98 Californium Cf 3 7 f Solid Synthetic Actinide
99 Einsteinium Es 3 7 f Solid Synthetic Actinide
100 Fermium Fm 3 7 f Solid Synthetic Actinide
101 Mendelevium Md 3 7 f Solid Synthetic Actinide
102 Nobelium No 3 7 f Solid Synthetic Actinide
103 Lawrencium Lr 3 7 d Solid Synthetic Actinide
104 Rutherfordium Rf 4 7 d Synthetic Transition metal
105 Dubnium Db 5 7 d Synthetic Transition metal
106 Seaborgium Sg 6 7 d Synthetic Transition metal
107 Bohrium Bh 7 7 d Synthetic Transition metal
108 Hassium Hs 8 7 d Synthetic Transition metal
109 Meitnerium Mt 9 7 d Synthetic
110 Darmstadtium Ds 10 7 d Synthetic
111 Roentgenium Rg 11 7 d Synthetic
112 Copernicium Cn 12 7 d Synthetic Transition metal
113 (Ununtrium) Uut 13 7 p Synthetic
114 (Ununquadium) Uuq 14 7 p Synthetic
115 (Ununpentium) Uup 15 7 p Synthetic
116 (Ununhexium) Uuh 16 7 p Synthetic
117 (Ununseptium) Uus 17 7 p Synthetic
118 (Ununoctium) Uuo 18 7 p Synthetic

See also


  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "chemical element".
  2. ^ Oganessian, YT (2007). "Heaviest nuclei from 48Ca-induced reactions". Journal of Physics G: Nuclear and Particle Physics 34 (4): R165–R242. Bibcode 2007JPhG...34..165O. doi:10.1088/0954-3899/34/4/R01. Retrieved 2011-05-07. 
  3. ^ Los Alamos National Laboratory (2011). "Periodic Table of Elements: Oxygen". Los Alamos, New Mexico: Los Alamos National Security, LLC. Retrieved 2011-05-07. 
  4. ^ E. M. Burbidge, G. R. Burbidge, W. A. Fowler, F. Hoyle (1957). "Synthesis of the Elements in Stars". Reviews of Modern Physics 29 (4): 547–650. Bibcode 1957RvMP...29..547B. doi:10.1103/RevModPhys.29.547. 
  5. ^ See the timeline on p.10 of Gaitskell, R.; et al. (2006). "Evidence for Dark Matter". Physical Review C 74 (4): 044602. Bibcode 2006PhRvC..74d4602O. doi:10.1103/PhysRevC.74.044602. 
  6. ^ Dumé, Belle (2003-04-23). "Bismuth breaks half-life record for alpha decay". (Bristol, England: Institute of Physics). Retrieved 2011-05-07. 
  7. ^ Marcillac, Pierre de; Noël Coron, Gérard Dambier, Jacques Leblanc, and Jean-Pierre Moalic (2003). "Experimental detection of alpha-particles from the radioactive decay of natural bismuth". Nature 422 (6934): 876–8. Bibcode 2003Natur.422..876D. doi:10.1038/nature01541. PMID 12712201. 
  8. ^ Sanderson, K. (17 October 2006). "Heaviest element made – again". Nature News. doi:10.1038/news061016-4. 
  9. ^ a b Schewe, P.; Stein, B. (17 October 2006). "Elements 116 and 118 Are Discovered". Physics News Update. American Institute of Physics. Retrieved 2006-10-19. 
  10. ^ Wilford, J.N. (14 January 1992). "Hubble Observations Bring Some Surprises". New York Times. 
  11. ^ Wright, E.L. (12 September 2004). "Big Bang Nucleosynthesis". UCLA, Division of Astronomy. Retrieved 2007-02-22. 
  12. ^ Wallerstein, G.; et al. (1999). "Synthesis of the elements in stars: forty years of progress". Reviews of Modern Physics 69 (4): 995–1084. Bibcode 1997RvMP...69..995W. doi:10.1103/RevModPhys.69.995. 
  13. ^ Earnshaw, A; Greenwood, N (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. 
  14. ^ Croswell, K. (1996). Alchemy of the Heavens. Anchor. ISBN 0-385-47214-5. 
  15. ^ Ultratrace minerals. Authors: Nielsen, Forrest H. USDA, ARS Source: Modern nutrition in health and disease / editors, Maurice E. Shils ... et al.. Baltimore : Williams & Wilkins, c1999., p. 283-303. Issue Date: 1999 URI: [1]
  16. ^ Plato (2008) [c. 360 BC]. Timaeus. Forgotten Books. p. 45. ISBN 978-1606200186. 
  17. ^ Hillar, M. (2004). "The Problem of the Soul in Aristotle's De anima". NASA/WMAP. Retrieved 2006-08-10. 
  18. ^ Partington, J.R. (1937). A Short History of Chemistry. New York: Dover Publications. ISBN 0486659771. 
  19. ^ a b Boyle, R. (1661). The Sceptical Chymist. London. ISBN 0922802904. 
  20. ^ Lavoisier, A.L. (1790). Elements of chemistry translated by Robert Kerr. Edinburgh. pp. 175–176. ISBN 9780415179140. 
  21. ^ Transactinide-2.
  22. ^ Carey, G.W. (1914). The Chemistry of Human Life. Los Angeles. ISBN 0766128407. 
  23. ^ Glanz, J. (6 April 2010). "Scientists Discover Heavy New Element". New York Times. 
  24. ^ a b W., Greiner. "Recommendations". 31st meeting, PAC for Nuclear Physics. Joint Institute for Nuclear Research. 
  25. ^ "IUPAC Announces Start of the Name Approval Process for the Element of Atomic Number 112". IUPAC. 20 July 2009. Retrieved 2009-08-27. 
  26. ^ "IUPAC (International Union of Pure and Applied Chemistry): Element 112 is Named Copernicium". IUPAC. 20 February 2010. 
  27. ^ Oganessian, Yu. Ts.; et al. (2006). "Synthesis of the isotopes of elements 118 and 116 in the 249Cf and 245Cm+48Ca fusion reactions". Physical Review C 74 (4): 044602. Bibcode 2006PhRvC..74d4602O. doi:10.1103/PhysRevC.74.044602. 
  28. ^ "Two ultra-heavy elements added to the periodic table". 6 June 2011. 

Further reading

  • Ball, Philip (2004). The Elements: A Very Short Introduction. Oxford University Press. ISBN 0192840991. 
  • Emsley, John (2003). Nature's Building Blocks: An A-Z Guide to the Elements. Oxford University Press. ISBN 0198503407. 
  • Gray, Theodore (2009). The Elements: A Visual Exploration of Every Known Atom in the Universe. Black Dog & Leventhal Publishers Inc. ISBN 1579128149. 
  • Scerri, E.R. (2007). The Periodic Table, Its Story and Its Significance. Oxford University Press. 
  • Strathern, Paul (2000). Mendeleyev's Dream: The Quest for the Elements. Hamish Hamilton Ltd. ISBN 024114065X. 

External links

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