MATTER & ENERGY

CONTENTS: Atoms Changes in Matter Chemical Elements Chemical Elements, alphabetic list Electricity Electromagnetic Radiation Elementary Particles Elements Periodic Table Energy Geothermal Energy Gravitation Mass Matter Metals Molecules Nuclear Energy Periodic Law Radiation Solar Energy States of Matter Wind Energy

Order in Matter Matter is composed of atoms or groups of atoms called molecules. The arrangement of particles in a material depends on the physical state of the substance. In a solid, particles form a compact structure that resists flow. Particles in a liquid have more energy than those in a solid. They can flow past one another, but they remain close. Particles in a gas have the most energy. They move rapidly and are separated from one another by relatively large distances.

Starch

Starch, common name applied to a white, granular or powdery, odorless, tasteless, complex carbohydrate, (C₆H₁₀O₅)_x, abundant in the seeds of cereal plants and in bulbs and tubers. Molecules of starch are made of hundreds or thousands of atoms, corresponding to values of x, as given in the formula above, that range from about 50 to many thousands.

Starch molecules are of two kinds. In the first kind, amylose, which constitutes about 20 percent of ordinary starch, the C₆H₁₀O₅ groups are arranged in a continuous but curled chain somewhat like a coil of rope; in the second kind, amylopectin, considerable side-branching of the molecule occurs.

Starch is manufactured by green plants during the process of photosynthesis. It forms part of the cell walls in plants, constitutes part of rigid plant fibers, and serves as a kind of energy storage for plants, because its oxidation to carbon dioxide and water releases energy. The granules of starch present in any plant have size, shape, and markings characteristic of the species of plant in which the starch is made.

Starch is almost insoluble in cold water and in alcohol, but with boiling water it gives a colloidal suspension that may form a jelly on cooling. Hot water changes starch slowly into smaller molecules called dextrins (see Dextrin). This reaction, an example of hydrolysis, is catalyzed by acids and by some enzymes. Dextrins, like starch, react with water, giving still simpler molecules, the ultimate products being maltose, C₁₂H₂₂O₁₁, a disaccharide, and glucose, C₆H ₁₂O₆, a monosaccharide (see Sugar).

The digestion of starch in the human body follows this course: The hydrolysis begins in the mouth under the action of salivary ptyalin, but is completed in the small intestine. The body does not use immediately all the glucose absorbed from the digestion of starch, but converts much of it to glycogen, which is stored in the liver. (Glycogen, called animal starch, has a structure nearly identical with that of amylopectin.) As the body requires glucose, hydrolysis of glycogen releases it into the bloodstream. Glycogen provides, therefore, an energy reserve for animals in the same way that ordinary starch does for plants.

See Carbohydrate; Human Nutrition.

Dextrin

Dextrin, amorphous, soluble carbohydrate, (C₆H₁₀O₅)_n, produced by the action on starch paste of acids, heat, or enzymes such as diatase. The first product formed in this reaction is soluble starch, which in turn hydrolyzes to form dextrin. Dextrin is prepared commercially by moistening potato starch with weak nitric acid and then drying and heating the mass at 110° C (230° F). Dextrin is used in the manufacture of beer and as a substitute for gum arabic in printing cotton fabrics. It is also used commercially as an adhesive.

Hydrolysis

Hydrolysis, type of chemical reaction in which a molecule of water, formula HOH, reacts with a molecule of a substance AB, in which A and B represent either atoms or groups of atoms. In the reaction the water molecule breaks into the fragments H⁺ and OH^-; and the molecule AB breaks into A⁺ and B^-; the fragments then join to give the final products AOH and HB. This kind of reaction is called a double decomposition or an exchange.

The Quantum Explanation of Spectral Lines

The explanation for exact spectral lines for each substance was provided by the quantum theory. In his 1913 model of the hydrogen atom Niels Bohr showed that the observed series of lines could be explained by assuming that electrons are restricted to atomic orbits in which their orbital angular momentum is an integral multiple of the quantity  h/2π, where h is Planck's constant. The integer multiple (e.g., 1, 2, 3 …) of h/2π is usually called the quantum number and represented by the symbol n.

When an electron changes from an orbit of higher energy (higher angular momentum) to one of lower energy, a photon of light energy is emitted whose frequency ν is related to the energy difference ΔE by the equation ν=ΔE/h.For hydrogen, the frequencies of the spectral lines are given by ν=cR (1/n_f²−1/n_i²) where c is the speed of light, Ris the Rydberg constant, and n_f and n_i are the final and initial quantum numbers of the electron orbits (n_i is always greater than n_f). The series of spectral lines for which n_f=1 is known as the Lyman series; that for n_f=2 is the Balmer series; that for n_f=3 is the Paschen series; that for n_f=4 is the Brackett series; and that for n_f=5 is the Pfund series. The Bohr theory was not as successful in explaining the spectra of other substances, but later developments of the quantum theory showed that all aspects of atomic and molecular spectra can be explained quantitatively in terms of energy transitions between different allowed quantum states.

Spectrum

Spectrum, arrangement or display of light or other form of radiation separated according to wavelength, frequency, energy, or some other property. Beams of charged particles can be separated into a spectrum according to mass in a mass spectrometer (see mass spectrograph). Physicists often find it useful to separate a beam of particles into a spectrum according to their energy.

Continuous and Line Spectra

Dispersion, the separation of visible light into a spectrum, may be accomplished by means of a prism or a diffraction grating. Each different wavelength or frequency of visible light corresponds to a different color, so that the spectrum appears as a band of colors ranging from violet at the short-wavelength (high-frequency) end of the spectrum through indigo, blue, green, yellow, and orange, to red at the long-wavelength (low-frequency) end of the spectrum. In addition to visible light, other types of electromagnetic radiation may be spread into a spectrum according to frequency or wavelength.

The spectrum formed from white light contains all colors, or frequencies, and is known as a continuous spectrum. Continuous spectra are produced by all incandescent solids and liquids and by gases under high pressure. A gas under low pressure does not produce a continuous spectrum but instead produces a line spectrum, i.e., one composed of individual lines at specific frequencies characteristic of the gas, rather than a continuous band of all frequencies. If the gas is made incandescent by heat or an electric discharge, the resulting spectrum is a bright-line, or emission, spectrum, consisting of a series of bright lines against a dark background. A dark-line, or absorption, spectrum is the reverse of a bright-line spectrum; it is produced when white light containing all frequencies passes through a gas not hot enough to be incandescent. It consists of a series of dark lines superimposed on a continuous spectrum, each line corresponding to a frequency where a bright line would appear if the gas were incandescent. The Fraunhofer lines appearing in the spectrum of the sun are an example of a dark-line spectrum; they are caused by the absorption of certain frequencies of light by the cooler, outer layers of the solar atmosphere. Line spectra of either type are useful in chemical analysis, since they reveal the presence of particular elements. The instrument used for studying line spectra is the spectroscope.

learn more: The Quantum Explanation of Spectral Lines

Conservation Laws

Conservation laws, in physics, basic laws that together determine which processes can or cannot occur in nature; each law maintains that the total value of the quantity governed by that law, e.g., mass or energy, remains unchanged during physical processes. Conservation laws have the broadest possible application of all laws in physics and are thus considered by many scientists to be the most fundamental laws in nature.

Conservation of Classical Processes

Most conservation laws are exact, or absolute, i.e., they apply to all possible processes; a few conservation laws are only partial, holding for some types of processes but not for others. By the beginning of the 20th cent. physics had established conservation laws governing the following quantities: energy, mass (or matter), linear momentum, angular momentum, and electric charge. When the theory of relativity showed (1905) that mass was a form of energy, the two laws governing these quantities were combined into a single law conserving the total of mass and energy.

Conservation of Elementary Particle Properties

With the rapid development of the physics of elementary particles during the 1950s, new conservation laws were discovered that have meaning only on this subatomic level. Laws relating to the creation or annihilation of particles belonging to the baryon and lepton classes of particles have been put forward. According to these conservation laws, particles of a given group cannot be created or destroyed except in pairs, where one of the pair is an ordinary particle and the other is an antiparticle belonging to the same group. Recent work has raised the possibility that the proton, which is a type of baryon, may in fact be unstable and decay into lighter products; the postulated methods of decay would violate the conservation of baryon number. To date, however, no such decay has been observed, and it has been determined that the proton has a lifetime of at least 10³¹ years. Two partial conservation laws, governing the quantities known as strangeness and isotopic spin, have been discovered for elementary particles. Strangeness is conserved during the so-called strong interactions and the electromagnetic interactions, but not during the weak interactions associated with particle decay; isotopic spin is conserved only during the strong interactions.

Conservation of Natural Symmetries

One very important discovery has been the link between conservation laws and basic symmetries in nature. For example, empty space possesses the symmetries that it is the same at every location (homogeneity) and in every direction (isotropy); these symmetries in turn lead to the invariance principles that the laws of physics should be the same regardless of changes of position or of orientation in space. The first invariance principle implies the law of conservation of linear momentum, while the second implies conservation of angular momentum. The symmetry known as the homogeneity of time leads to the invariance principle that the laws of physics remain the same at all times, which in turn implies the law of conservation of energy. The symmetries and invariance principles underlying the other conservation laws are more complex, and some are not yet understood.

Three special conservation laws have been defined with respect to symmetries and invariance principles associated with inversion or reversal of space, time, and charge. Space inversion yields a mirror-image world where the "handedness" of particles and processes is reversed; the conserved quantity corresponding to this symmetry is called space parity, or simply parity, P. Similarly, the symmetries leading to invariance with respect to time reversal and charge conjugation (changing particles into their antiparticles) result in conservation of time parity, T, and charge parity, C. Although these three conservation laws do not hold individually for all possible processes, the combination of all three is thought to be an absolute conservation law, known as the CPT theorem, according to which if a given process occurs, then a corresponding process must also be possible in which particles are replaced by their antiparticles, the handedness of each particle is reversed, and the process proceeds in the opposite direction in time. Thus, conservation laws provide one of the keys to our understanding of the universe and its material basis.

Momentum

Momentum, in mechanics, the quantity of motion of a body, specifically the product of the mass of the body and its velocity. Momentum is a vector quantity; i.e., it has both a magnitude and a direction, the direction being the same as that of the velocity vector. When an external force acts upon a body or a system of bodies in motion, it causes a change in the momentum of the body. The impulse of a force acting on a body is the product of the force and the duration of time in which it acts and is equal to the change in momentum of the body. When no external force acts upon a body in motion or a system of bodies there is no change in the total momentum even though, as in the case of a system of bodies, there may be an internal disturbance of the system resulting in changes in the momenta of individual bodies. This conclusion is commonly known as the principle of the conservation of momentum (see conservation laws, in physics). The momentum of a body should not be confused with its kinetic energy. The distinction between them can be seen in the action of a pile driver. The distance to which the pile is driven depends upon its kinetic energy; the length of time required for the action to cease, upon its momentum. In addition to the momentum a body has because of its linear motion, the body may also have angular momentum because of rotation. The angular momentum of a particle rotating about a point is equal to the product of the mass of the particle, its angular velocity, and the square of its distance from the axis of rotation. More simply, the angular momentum is the product of the instantaneous linear momentum and the distance. Angular momentum is a vector quantity directed perpendicular to the plane of motion.

Official Symbols and Names for the Elements

Each element is assigned an official symbol by the International Union of Pure and Applied Chemistry (IUPAC). For example, the symbol for carbon is C, and the symbol for silver is Ag [Lat. argentum = silver]. There are several ways of designating an isotope. One designation consists of the name or symbol of the element followed by a hyphen and the mass number of the isotope; thus the isotope of carbon with mass number 12 can be designated carbon-12 or C-12. The mass number is often written as a superscript, e.g., C¹²; sometimes the atomic number is written as a subscript preceding the symbol, e.g., ₆C¹². The IUPAC rules for nomenclature of inorganic chemistry state that the subscript atomic number and superscript mass number should both precede the symbol, e.g., _¹²⁶C.

Many isotopes were given special names and symbols when they were first discovered in natural radioactive decay series (e.g., uranium-235 was called actinouranium and represented by the symbol AcU). This practice is discouraged in the modern nomenclature except in the case of hydrogen. The isotopes hydrogen-2 and hydrogen-3 are usually called deuterium and tritium, respectively. Hydrogen-1, the most abundant isotope, has the name protium but is usually simply called hydrogen. Newly discovered elements that have been synthesized by one laboratory and not yet confirmed by a second are given a provisional name based on Greek and Latin roots; when the discovery is confirmed, the laboratory that first made it may suggest a name for the element.

Nonmetal

Nonmetal, chemical element possessing certain properties by which it is distinguished from a metal. In general, this distinction is drawn on the basis that a nonmetal tends to accept electrons and form negative ions and that its oxide is acidic. Nonmetals are poor conductors of heat and electricity and do not have the luster of metals. Arsenic, antimony, selenium, and tellurium exhibit both nonmetallic and metallic properties and are called metalloids. Unlike the metals, which are all solids (with the exception of mercury) under ordinary conditions of temperature and pressure, the nonmetals appear in all three states. Argon, chlorine, fluorine, helium, hydrogen, krypton, neon, nitrogen, oxygen, and xenon are normally gases. Bromine is a liquid. Boron, carbon, iodine, phosphorus, silicon, and sulfur are solids. Certain of them, e.g., boron, carbon, iodine, silicon, and sulfur, form crystals, as do the metals. In hardness they vary considerably. Carbon in its allotropic form, the diamond, is the hardest element known. With the exception of carbon, sulfur, nitrogen, oxygen, and the inert gases—argon, helium, krypton, neon, and xenon—the nonmetals do not occur uncombined in nature, but exist in numerous relatively abundant compounds, among which are the oxides, halides (binary halogen compounds), sulfides, carbonates, nitrates, phosphates, silicates, and sulfates. With a few exceptions, the nonmetallic elements are important chiefly for their compounds. For the properties and uses of specific nonmetals, see the separate articles on these elements.

Properties of the Elements

Properties of an element are sometimes classed as either chemical or physical. Chemical properties are usually observed in the course of a chemical reaction, while physical properties are observed by examining a sample of the pure element. The chemical properties of an element are due to the distribution of electrons around the atom's nucleus, particularly the outer, or valence, electrons; it is these electrons that are involved in chemical reactions. A chemical reaction does not affect the atomic nucleus; the atomic number therefore remains unchanged in a chemical reaction.

Some properties of an element can be observed only in a collection of atoms or molecules of the element. These properties include color, density, melting point, boiling point, and thermal and electrical conductivity. While some of these properties are due chiefly to the electronic structure of the element, others are more closely related to properties of the nucleus, e.g., mass number.

The elements are sometimes grouped according to their properties. One major classification of the elements is as metals, nonmetals, and metalloids. Elements with very similar chemical properties are often referred to as families; some families of elements include the halogens, the inert gases, and the alkali metals. In the periodic table the elements are arranged in order of increasing atomic weight in such a way that the elements in any column have similar properties.

Compound

Compound, in chemistry, a substance composed of atoms of two or more elements in chemical combination, occurring in a fixed, definite proportion and arranged in a fixed, definite structure. A compound is often represented by its chemical formula. The formula for water is H₂O, and for sodium chloride, NaCl. The formula weight of a compound can be determined from its formula. The molecular weight of a molecular compound can be determined from its molecular formula. Two or more distinct compounds that have the same molecular formula but different properties are called isomers.

Formation and Decomposition of Compounds

Compounds are formed from simpler substances by chemical reaction. Some compounds can be formed directly from their constituent elements, e.g., water from hydrogen and oxygen: 2H₂ + O₂ → 2H₂O. Other compounds are formed by reaction of an element with another compound; e.g., sodium hydroxide (NaOH) is formed (and hydrogen gas released) by the reaction of sodium metal with water: 2Na + 2H₂O → 2NaOH + H₂↑. Compounds are also made by reaction of other compounds; e.g., sodium hydroxide reacts with hydrogen chloride (HCl) to form sodium chloride and water: HCl + NaOH → NaCl + H₂O. Complex molecules such as proteins are formed by a series of reactions involving elements and simple compounds.

Compounds can be decomposed by chemical means into elements or simpler compounds. Water is broken down into hydrogen and oxygen by electrolysis. Candle wax, a mixture of hydrocarbons, is changed in the candle flame by combustion (with oxygen) to a mixture of the simpler compounds carbon dioxide (CO₂) and water. Life is based on numerous reactions in which energy is stored and released as compounds are produced and decomposed.

Properties of Compounds

A compound has unique properties that are distinct from the properties of its elemental constituents. One familiar chemical compound is water, a liquid that is nonflammable and does not support combustion. It is composed of two elements: hydrogen, an extremely flammable gas, and oxygen, a gas that supports combustion. A compound differs from a mixture in that the components of a mixture retain their own properties and may be present in many different proportions. The components of a mixture are not chemically combined; they can be separated by physical means. A mixture of hydrogen and oxygen gases is still a gas and can be separated by physical methods. If the mixture is ignited, however, the two gases undergo a rapid chemical combination to form water. Although the hydrogen and oxygen can occur in any proportion in a mixture of gases, they are always combined in the exact proportion of two atoms of hydrogen to one atom of oxygen when combined in the compound water. Another familiar compound is sodium chloride (common salt). It is composed of the silvery metal sodium and the greenish poisonous gas chlorine combined in the proportion of one atom of sodium to one atom of chlorine.

Molecular and Ionic Compounds

Water is a molecular compound; it is made up of electrically neutral molecules, each containing a fixed number of atoms. Sodium chloride is an ionic compound; it is made up of electrically charged ions that are present in fixed proportions and are arranged in a regular, geometric pattern (called crystalline structure) but are not grouped into molecules. The atoms in a compound are held together by chemical bonding.

Synthetic Elements / Transactinide Elements

Synthetic elements, in chemistry, radioactive elements that were not discovered occurring in nature but as artificially produced isotopes. They are technetium (at. no. 43), which was the first element to be synthesized, promethium (at. no. 61), astatine (at. no. 85), francium (at. no. 87), and the transuranium elements (at. no. 93 and beyond in the periodic table). Some of these elements have since been shown to exist in minute amounts in nature, usually as short-lived members of natural radioactive decay series (see radioactivity).

The synthetic elements through at. no. 100 (fermium) are created by bombarding a heavy element, such as uranium or plutonium, with neutrons or alpha particles. The synthesis of the transfermium elements (elements with at. no. 101 or greater) is accomplished by the fusion of the nuclei of two lighter elements. Elements 101 through 106 were first produced by fusing the nuclei of slightly lighter elements, such as californium, with those of light elements, such as carbon. Elements 107 through 112 were first produced by fusing the nuclei of medium-weight elements, such as bismuth or lead, with those of other medium-weight elements, such as iron, nickel, or zinc. Element 114 was first produced by fusing the nuclei of plutonium and calcium and subsequently by fusing the nuclei of lead and krypton, as was element 116. Element 115 was produced by bombarding americium with calcium, and element 113 resulted from the radioactive decay of element 115. The claim by Lawrence Berkeley National Laboratory to have created element 118 has been retracted.)

The transfermium elements are produced in very small quantities (one atom at a time), and identification is therefore very difficult because of half-lives ranging from minutes to milliseconds and the need to identify the products by methods other than known chemical separations. This has led to controversy over reported discoveries and over the naming of the elements. It has been predicted that one isotope of element 114—containing 114 protons and 184 neutrons—would be very stable because its nucleus would have a full complement of protons and neutrons. Termed an "island of stability," its half-life might be measured in years. However, none of the three isotopes of element 114 synthesized as yet have as many as 184 neutrons, and their half-lives are still in the millisecond range.

Transactinide elements (chemistry), in the periodic table, elements with atomic numbers higher than 103.

Ununhexium

Ununhexium, artificially produced radioactive chemical element; symbol Uuh; at. no. 116; mass number of most stable isotope 292; m.p., b.p., sp. gr., and valence unknown. Situated in Group 16 of the periodic table, it is expected to have properties similar to those of polonium and tellurium.

In 1999 a research team at the Lawrence Berkeley National Laboratory in Calif. bombarded lead-208 atoms with high-energy krypton-86 ions to create, apparently, ununoctium (element 118) atoms. The Uuo-293 isotope that they synthesized emitted an alpha particle to decay into Uuh-289, which has a life-life of about 0.6 millisecond, which then emitted an alpha particle to decay into ununquadium (element 114). Although the Berkeley laboratory retracted its claim for creating ununoctium in 2001, other research teams have since created ununhexium directly. No name has yet been adopted for element 116, which is therefore called ununhexium, from the Latin roots un for one and hex for six, under a convention for neutral temporary names proposed by the International Union of Pure and Applied Chemistry (IUPAC) in 1980.

Ununoctium

Ununoctium (y'nənŏk`tēəm), artificially produced radioactive chemical element; symbol Uuo; at. no. 118. Scientists from the Joint Institute for Nuclear Research in Dubna, Russia, and Lawrence Livermore National Laboratory in California collaborated in the discovery of ununoctium in experiments conducted in 2002 and 2005. They bombarded atoms of californium-249 with ions of calcium-48. Among the products of the bombardments were three atoms of ununoctium-294 (one atom in 2002 and two in 2005), each of which decayed in 0.9 milliseconds into an atom of ununhexium by emitting an alpha particle. No name has yet been adopted for element 118, which is therefore called ununoctium, from the Latin roots un for one and oct for eight, under a convention for neutral temporary names proposed by the International Union of Pure and Applied Chemistry (IUPAC) in 1980.

In 1999 a research team at the Lawrence Berkeley National Laboratory in Calif. bombarded lead-208 atoms with high-energy krypton-86 ions to create what an analysis showed to be three atoms of element 118 with mass number 293 and a half-life of less than a millisecond. In 2001, however, the team retracted its claim to have produced ununoctium after other laboratories failed to reproduce their results and after a reanalysis of the original data did not show the production of element 118. A subsequent investigation suggested that the original finding was the result of fraud on the part of one of the team scientists.

See also synthetic elements; transactinide elements; transuranium elements.

Elements

Element	Symbol	Atomic Number	Atomic Weight1	Melting Point (Degrees Celsius)	Boiling Point (Degrees Celsius)
actinium	Ac	89	227.0278	1050.	3200. ±300
aluminum	Al	13	26.98154	660.37	2467.
americium	Am	95	(243)	1172.	2600.
antimony	Sb	51	121.75	630.74	1750.
argon	Ar	18	39.948	−189.2	−185.7
arsenic	As	33	74.9216	817. (at 28 atmospheres)	613. (sublimates)
astatine	At	85	(210)	302. (est.)	337. (est.)
barium	Ba	56	137.33	725.	1640.
berkelium	Bk	97	(247)	1050.	2590.
beryllium	Be	4	9.01218	1278. ±5	2970.
bismuth	Bi	83	208.9804	271.3	1560. ±5
bohrium	Bh	107	(262)	—	—
boron	B	5	10.81	2300.	2550. (sublimates)
bromine	Br	35	79.904	−7.2	58.78
cadmium	Cd	48	112.41	320.9	765.
calcium	Ca	20	40.08	839. ±2	1484.
californium	Cf	98	(251)	900.	1470.
carbon	C	6	12.011	∼3550.	4827.
cerium	Ce	58	140.12	799.	3426.
cesium	Cs	55	132.9054	28.40	669.3
chlorine	Cl	17	35.453	−100.98	−34.6
chromium	Cr	24	51.996	1857. ±20	2672.
cobalt	Co	27	58.9332	1495.	2870.
copper	Cu	29	63.546	1083.4 ±0.2	2567.
curium	Cm	96	(247)	1340. ±40	3110.
darmstadtium	Ds	110	(271)	—	—
dubnium	Db	105	(262)	—	—
dysprosium	Dy	66	162.50	1412.	2562.
einsteinium	Es	99	(252)	857.	—
erbium	Er	68	167.26	1529.	2863.
europium	Eu	63	151.96	822.	1597.
fermium	Fm	100	(257)	1527.	—
fluorine	F	9	18.998403	−219.62	−188.14
francium	Fr	87	(223)	(27) (est.)	(677) (est.)
gadolinium	Gd	64	157.25	1313. ±1	3266.
gallium	Ga	31	69.72	29.78	2403.
germanium	Ge	32	72.59	937.4	2830.
gold	Au	79	196.9665	1064.43	2808.
hafnium	Hf	72	178.49	2227. ±20	4602.
hassium	Hs	108	(265)	—	—
helium	He	2	4.0026	<−272.2	−268.934
holmium	Ho	67	164.9304	1474.	2425.
hydrogen	H	1	1.00794	−259.14	−252.87
indium	In	49	114.82	156.61	2080.
iodine	I	53	126.9045	113.5	184.35
iridium	Ir	77	192.22	2410.	4130.
iron	Fe	26	55.847	1535.	2750.
krypton	Kr	36	83.80	−156.6	−152.30 ±0.10
lanthanum	La	57	138.9055	921.	3457.
lawrencium	Lr	103	(262)	1627.	—
lead	Pb	82	207.2	327.502	1740.
lithium	Li	3	6.941	180.54	1342.
lutetium	Lu	71	174.967	1663.	3395.
magnesium	Mg	12	24.305	648.8 ±0.5	1090.
manganese	Mn	25	54.9380	1244. ±3	1962.
meitnerium	Mt	109	(266)	—	—
mendelevium	Md	101	(258)	827.	—
mercury	Hg	80	200.59	−38.842	356.58
molybdenum	Mo	42	95.94	2617.	4612.
neodymium	Nd	60	144.24	1021.	3068.
neon	Ne	10	20.179	−248.67	−246.048
neptunium	Np	93	237.0482	640. ±1	3902. (est.)
nickel	Ni	28	58.69	1453.	2732.
niobium	Nb	41	92.9064	2468. ±10	4742.
nitrogen	N	7	14.0067	−209.86	−195.8
nobelium	No	102	(259)	827.	—
osmium	Os	76	190.2	3045. ±30	5027. ±100
oxygen	O	8	15.9994	−218.4	−182.962
palladium	Pd	46	106.42	1554.	2970.
phosphorus	P	15	30.97376	44.1 (white)	280. (white)
platinum	Pt	78	195.08	1772.	3827. ±100
plutonium	Pu	94	(244)	641.	3232.
polonium	Po	84	(209)	254.	962.
potassium	K	19	39.0983	63.25	760.
praseodymium	Pr	59	140.9077	931.	3512.
promethium	Pm	61	(145)	1042	3000. (est.)
protactinium	Pa	91	231.0359	<1600.	4026.
radium	Ra	88	226.0254	700.	1140.
radon	Rn	86	(222)	−71.	−61.8
rhenium	Re	75	186.207	3180.	5627. (est.)
rhodium	Rh	45	102.9055	1966. ±3	3727. ±100
roentgenium	Rg	111	(272)	—	—
rubidium	Rb	37	85.4678	38.89	686.
ruthenium	Ru	44	101.07	2310.	3900.
rutherfordium	Rf	104	(261)	—	—
samarium	Sm	62	150.36	1072. ±5	1791.
scandium	Sc	21	44.9559	1541.	2831.
seaborgium	Sg	106	(266)	—	—
selenium	Se	34	78.96	217.	684.9 ±1.0
silicon	Si	14	28.0855	1410.	2355.
silver	Ag	47	107.8682	961.93	2212.
sodium	Na	11	22.98977	97.81 ±0.03	882.9
strontium	Sr	38	87.62	269.	1384.
sulfur	S	16	32.06	112.8	444.674
tantalum	Ta	73	180.9479	2996.	5425. ±100
technetium	Tc	43	(98)	2200.	4877.
tellurium	Te	52	127.60	449.5 ±0.3	989.8 ±3.8
terbium	Tb	65	158.9254	1356.	3123.
thallium	Tl	81	204.383	303.5	1457. ±10
thorium	Th	90	232.0381	1750.	∼4790.
thulium	Tm	69	168.9342	1545. ±15	1947.
tin	Sn	50	118.69	231.9681	2270.
titanium	Ti	22	47.88	1660. ±10	3287.
tungsten	W	74	183.85	3410. ±20	5660.
ununbium	Uub	112	(285)	—	—
ununhexium	Uuh	116	(292)	—	—
ununoctium	Uuo	118	(294)	—	—
ununpentium	Uup	115	(288)	—	—
ununquadium	Uuq	114	(289)	—	—
ununtrium	Uut	113	(284)	—	—
uranium	U	92	238.0289	1132.3 ±0.8	3818.
vanadium	V	23	50.9415	1890. ±10	3380.
xenon	Xe	54	131.29	−111.9	−107.1 ±3
ytterbium	Yb	70	173.04	819.	1194.
yttrium	Y	39	88.9059	1522. ±8	3338.
zinc	Zn	30	65.38	419.58	907.
zirconium	Zr	40	91.22	1852. ±2	4377.

¹ Parentheses indicate most stable isotope.

Element

Element, in chemistry, a substance that cannot be decomposed into simpler substances by chemical means. A substance such as a compound can be decomposed into its constituent elements by means of a chemical reaction, but no further simplification can be achieved. An element can, however, be decomposed into simpler substances, such as protons and neutrons or various combinations of them, by the methods of particle physics, e.g., by bombardment of the nucleus.

The Atom

The smallest unit of a chemical element that has the properties of that element is called an atom. Many elements (e.g., helium) occur as single atoms. Other elements occur as molecules made up of more than one atom. Elements that ordinarily occur as diatomic molecules include hydrogen, nitrogen, oxygen, and the halogens, but oxygen also occurs as a triatomic form called ozone. Phosphorus usually occurs as a tetratomic molecule, and crystalline sulfur occurs as molecules containing eight atoms.

Atomic Number and Mass Number

Regardless of how many atoms the element is composed of, each atom has the same number of protons in its nucleus, and this is different from the number in the nucleus of any other element. Thus this number, called the atomic number (at. no.), defines the element. For example, the element carbon consists of atoms all with at. no. 6, i.e., all having 6 protons in the nucleus; any atom with at. no. 6 is a carbon atom. By 2006, 117 elements were known, ranging from hydrogen with an at. no. of 1 to an as yet unnamed element (temporarily known as ununoctium) with an at. no. of 118. (See the table entitled Elements for an alphabetical list of all the elements, including their symbols, atomic numbers, atomic weights, and melting and boiling points.) The nuclei of most atoms also contain neutrons. The total number of protons and neutrons in the nucleus of an atom is called the mass number. For example, the mass number of a carbon atom with 6 protons and 6 neutrons in its nucleus is 12.

Isotopes

Although all atoms of an element have the same number of protons in their nuclei, they may not all have the same number of neutrons. Atoms of an element with the same mass number make up an isotope of the element. All known elements have isotopes; some have more than others. Hydrogen, for example, has only 3 isotopes, while xenon has 16. Approximately 300 naturally occurring isotopes are known, and more than 2,500 radioactive isotopes have been artificially produced (see synthetic elements). There are 13 isotopes of carbon, having from 2 to 14 neutrons in the nucleus and therefore mass numbers from 8 to 20.

Not all of the elements have stable isotopes. Some have only radioactive isotopes, which decay to form other isotopes, usually of other elements (see radioactivity). In some cases all the isotopes of an element are very unstable, and the element is therefore not found in nature. Only 94 of the elements are known to occur naturally on earth. Of these, 6 occur in minute amounts produced by the decay of other elements. These 6 extremely scarce elements and those that do not occur at all naturally were discovered when they were produced in the laboratory; they are often called the man-made, artificially produced, or synthetic elements.

Atomic Mass and Atomic Weight

Atoms are not very massive; a carbon atom weighs about 2 × 10−23 grams. Because atoms have so little mass, a unit much smaller than the gram is used. In the current system (adopted in 1960–61) the unit of atomic mass, called atomic mass unit (amu), is defined as exactly 1-12 the mass of an atom of carbon-12. The atomic weight of an element is the mean (weighted average) of the atomic masses of all the naturally occurring isotopes. Carbon has two principal naturally occurring isotopes, carbon-12 and carbon-13. Carbon-12, whose mass is defined as exactly 12 amu, constitutes 98.89% of naturally occurring carbon; carbon-13, whose mass is 13.00335 amu, constitutes 1.11%. (There are also small traces of the radioactive isotope carbon-14.) The atomic weight of the element is determined by multiplying the percent abundance of each isotope by the atomic mass of the isotope, adding these products, and dividing by 100. However, isotope abundance is often determined by the medium of the source, solid, liquid, or gas, and the average atomic weight may fluctuate. Thus, for carbon, [(98.89 × 12.000) + (1.11 × 13.00335)]/100 = 12.01115, which is the atomic weight of the element carbon in amu. Certain synthetic elements exist only momentarily in the form of a few short-lived isotopes; in such cases the concept of atomic weight cannot be applied.

learn more:

• Properties of the Elements
• Official Symbols and Names for the Elements
• The Elements through the Ages