Metal Cd

[Janita Locklin]

Ametal may be a chemical element such as iron; an alloy such as stainless steel; or a molecular compound such as polymeric sulfur nitride.[4] The general science of metals is called metallurgy, a subtopic of materials science; aspects of the electronic and thermal properties are also within the scope of condensed matter physics and solid-state chemistry, it is a multidisciplinary topic. In colloquial use materials such as steel alloys are referred to as metals, while others such as polymers, wood or ceramics are nonmetallic materials.

A metal conducts electricity at a temperature of absolute zero,[5] which is a consequence of delocalized states at the Fermi energy.[1][2] Many elements and compounds become metallic under high pressures, for example, iodine gradually becomes a metal at a pressure of between 40 and 170 thousand times atmospheric pressure. Sodium becomes a nonmetal at pressure of just under two million times atmospheric pressure, and at even higher pressures it is expected to become a metal again.

When discussing the periodic table and some chemical properties the term metal is often used to denote those elements which in pure form and at standard conditions are metals in the sense of electrical conduction mentioned above. The related term metallic may also be used for types of dopant atoms or alloying elements.

In astronomy metal refers to all chemical elements in a star that are heavier than helium. In this sense the first four "metals" collecting in stellar cores through nucleosynthesis are carbon, nitrogen, oxygen, and neon. A star fuses lighter atoms, mostly hydrogen and helium, into heavier atoms over its lifetime. The metallicity of an astronomical object is the proportion of its matter made up of the heavier chemical elements.[6][7]

The strength and resilience of some metals has led to their frequent use in, for example, high-rise building and bridge construction, as well as most vehicles, many home appliances, tools, pipes, and railroad tracks. Precious metals were historically used as coinage, but in the modern era, coinage metals have extended to at least 23 of the chemical elements.[8] There is also extensive use of multi-element metals such as titanium nitride[9] or degenerate semiconductors in the semiconductor industry.

Most metals are shiny and lustrous, at least when polished, or fractured. Sheets of metal thicker than a few micrometres appear opaque, but gold leaf transmits green light. This is due to the freely moving electrons which reflect light.[1][2]

Although most elemental metals have higher densities than nonmetals,[10] there is a wide variation in their densities, lithium being the least dense (0.534 g/cm3) and osmium (22.59 g/cm3) the most dense. Some of the 6d transition metals are expected to be denser than osmium, but their known isotopes are too unstable for bulk production to be possible[11] Magnesium, aluminium and titanium are light metals of significant commercial importance. Their respective densities of 1.7, 2.7, and 4.5 g/cm3 can be compared to those of the older structural metals, like iron at 7.9 and copper at 8.9 g/cm3. The most common lightweight metals are aluminium[12][13] and magnesium[14][15] alloys.

Metals are typically malleable and ductile, deforming under stress without cleaving.[10] The nondirectional nature of metallic bonding contributes to the ductility of most metallic solids, where the Peierls stress is relatively low allowing for dislocation motion, and there are also many combinations of planes and directions for plastic deformation.[16] Due to their having close packed arrangements of atoms the Burgers vector of the dislocations are fairly small, which also means that the energy needed to produce one is small.[3][16] In contrast, in an ionic compound like table salt the Burgers vectors are much larger and the energy to move a dislocation is far higher.[3] Reversible elastic deformation in metals can be described well by Hooke's Law for the restoring forces, where the stress is linearly proportional to the strain.[17]

A temperature change may lead to the movement of structural defects in the metal such as grain boundaries, point vacancies, line and screw dislocations, stacking faults and twins in both crystalline and non-crystalline metals. Internal slip, creep, and metal fatigue may also ensue.[3][16]

The atoms of simple metallic substances are often in one of three common crystal structures, namely body-centered cubic (bcc), face-centered cubic (fcc), and hexagonal close-packed (hcp). In bcc, each atom is positioned at the center of a cube of eight others. In fcc and hcp, each atom is surrounded by twelve others, but the stacking of the layers differs. Some metals adopt different structures depending on the temperature.[18]

The electronic structure of metals means they are relatively good conductors of electricity. The electrons all have different momenta, which average to zero when there is no external voltage. When a voltage is applied some move a little faster in a given direction, some a little slower so there is a net drift velocity which leads to an electric current.[1][2] This involves small changes in which wavefunctions the electrons are in, changing to those with the higher momenta. Quantum mechanics dictates that one can only have one electron in a given state, the Pauli exclusion principle.[20] Therefore there have to be empty delocalized electron states (with the higher momenta) available at the highest occupied energies as sketched in the Figure. In a semiconductor like silicon or a nonmetal like strontium titanate there is an energy gap between the highest filled states of the electrons and the lowest unfilled, so no accessible states with slightly higher momenta. Consequently, semiconductors and nonmetals are poor conductors, although they can carry some current when doped with elements that introduce additional partially occupied energy states at higher temperatures.[21]

All of the metallic alloys as well as conducting ceramics and polymers are metals by the same definition; for instance titanium nitride has delocalized states at the Fermi level. They have electrical conductivities similar to those of elemental metals. Liquid forms are also metallic conductors or electricity, for instance mercury. In normal conditions no gases are metallic conductors. However, a plasma (physics) is a metallic conductor and the charged particles in a plasma have many properties in common with those of electrons in elemental metals, particularly for white dwarf stars.[23]

The contribution of a metal's electrons to its heat capacity and thermal conductivity, and the electrical conductivity of the metal itself can be approximately calculated from the free electron model.[2] However, this does not take into account the detailed structure of the metal's ion lattice. Taking into account the positive potential caused by the arrangement of the ion cores enables consideration of the electronic band structure and binding energy of a metal. Various models are applicable, the simplest being the nearly free electron model.[2] Modern methods such as density functional theory are typically used.[26][27]

The elements which form metals usually form cations through electron loss.[10] Most will react with oxygen in the air to form oxides over various timescales (potassium burns in seconds while iron rusts over years) which depend upon whether the native oxide forms a passivation layer that acts as a diffusion barrier.[28][29] Some others, like palladium, platinum, and gold, do not react with the atmosphere at all; gold can form compounds where it gains an electron (aurides, e.g. caesium auride). The oxides of elemental metals are often basic. However, oxides with very high oxidation states such as CrO3, Mn2O7, and OsO4 often have strictly acidic reactions; and oxides of the less electropositive metals such as BeO, Al2O3, and PbO, can display both basic and acidic properties. The latter are termed amphoteric oxides.

The elements that form exclusively metallic structures under ordinary conditions are shown in yellow on the periodic table below. The remaining elements either form covalent network structures (light blue), molecular covalent structures (dark blue), or remain as single atoms (violet).[30] Astatine (At), francium (Fr), and the elements from fermium (Fm) onwards are shown in gray because they are extremely radioactive and have never been produced in bulk. Theoretical and experimental evidence suggests that almost all these uninvestigated elements should be metals,[31] though there is some doubt for oganesson (Og).[32]

MetallicNetwork covalentMolecular covalentSingle atomsUnknownBackground color shows bonding of simple substances in the periodic table. If there are several, the most stable allotrope is considered.

The situation changes with pressure: at extremely high pressures, all elements (and indeed all substances) are expected to metallize.[31] Arsenic (As) has both a stable metallic allotrope and a metastable semiconducting allotrope at standard conditions. A similar situation affects carbon (C): graphite is metallic, but diamond is not.

Most pure metals are either too soft, brittle, or chemically reactive for practical use. Combining different ratios of metals and other elements in alloys modifies the properties to produce desirable characteristics, for instance more ductile, harder, resistant to corrosion, or have a more desirable color and luster. Of all the metallic alloys in use today, the alloys of iron (steel, stainless steel, cast iron, tool steel, alloy steel) make up the largest proportion both by quantity and commercial value. Iron alloyed with various proportions of carbon gives low-, mid-, and high-carbon steels, with increasing carbon levels reducing ductility and toughness. The addition of silicon will produce cast irons, while the addition of chromium, nickel, and molybdenum to carbon steels (more than 10%) results in stainless steels with enhanced corrosion resistance.

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