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.
Showing posts with label element. Show all posts
Showing posts with label element. Show all posts

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.

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.

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:

Popular Posts