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Element Symbol Atomic Number Atomic Weight 1 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

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 a

Mass number, often represented by the symbol A, the total number of nucleons (neutrons and protons) in the nucleus of an atom . All atoms of a chemical element have the same atomic number (number of protons in the nucleus) but may have different mass numbers (from having different numbers of neutrons in the nucleus). Atoms of an element with the same mass number make up an isotope of the element. Different isotopes of the same element cannot have the same mass number, but isotopes of different elements often do have the same mass number, e.g., carbon-14 (6 protons and 8 neutrons) and nitrogen-14 (7 protons and 7 neutrons).

Atomic mass unit or amu, in chemistry and physics, unit defined as exactly 1-12 the mass of an atom of carbon-12, the isotope of carbon with six protons and six neutrons in its nucleus. One amu is equal to approximately  1.66 × 10 −24  grams.

Atomic mass, the mass of a single atom , usually expressed in atomic mass units (amu). Most of the mass of an atom is concentrated in the protons and neutrons contained in the nucleus. Each proton or neutron weighs about 1 amu, and thus the atomic mass is always very close to the mass number (total number of protons and neutrons in the nucleus). Atoms of an isotope of an element all have the same atomic mass. Atomic masses are usually determined by mass spectrography (see mass spectrograph ). They have been determined with great relative accuracy, but their absolute value is less certain.

Atomic weight, mean (weighted average) of the masses of all the naturally occurring isotopes of a chemical element , as contrasted with atomic mass , which is the mass of any individual isotope. Although the first atomic weights were calculated at the beginning of the 19th cent., it was not until the discovery of isotopes by F. Soddy (c.1913) that the atomic mass of many individual isotopes was determined, leading eventually to the adoption of the atomic mass unit as the standard unit of atomic weight. Effect of Isotopes in Calculating Atomic Weight Most naturally occurring elements have one principal isotope and only insignificant amounts of other isotopes. Therefore, since the atomic mass of any isotope is very nearly a whole number, most atomic weights are nearly whole numbers, e.g., hydrogen has atomic weight 1.00797 and nitrogen has atomic weight 14.007. However, some elements have more than one principal isotope, and the atomic weight for such an element—since it is a

Isotope, in chemistry and physics, one of two or more atoms having the same atomic number but differing in atomic weight and mass number. The concept of isotope was introduced by F. Soddy in explaining aspects of radioactivity; the first stable isotope (of neon) was discovered by J. J. Thomson. The nuclei of isotopes contain identical numbers of protons, equal to the atomic number of the atom, and thus represent the same chemical element, but do not have the same number of neutrons. Thus isotopes of a given element have identical chemical properties but slightly different physical properties and very different half-lives, if they are radioactive (see half-life). For most elements, both stable and radioactive isotopes are known. Radioactive isotopes of many common elements, such as carbon and phosphorus, are used as tracers in medical, biological, and industrial research. Their radioactive nature makes it possible to follow the substances in their paths through a plant or animal body