Heated Metal Low-poly 3D model
A metal (from Greek μέταλλον métallon, mine, quarry, metal) is a material that, when freshly prepared, polished, or fractured, shows a lustrous appearance, and conducts electricity and heat relatively well. Metals are typically malleable (they can be hammered into thin sheets) or ductile (can be drawn into wires). A metal may be a chemical element such as iron, or an alloy such as stainless steel.
In physics, a metal is generally regarded as any substance capable of conducting electricity at a temperature of absolute zero. Many elements and compounds that are not normally classified as metals become metallic under high pressures. For example, the nonmetal iodine gradually becomes a metal at a pressure of between 40 and 170 thousand times atmospheric pressure. Equally, some materials regarded as metals can become nonmetals. Sodium, for example, becomes a nonmetal at pressure of just under two million times atmospheric pressure.
In chemistry, two elements that would otherwise qualify (in physics) as brittle metals—arsenic and antimony—are commonly instead recognised as metalloids, on account of their predominately non-metallic chemistry. Around 95 of the 118 elements in the periodic table are metals (or are likely to be such). The number is inexact as the boundaries between metals, nonmetals, and metalloids fluctuate slightly due to a lack of universally accepted definitions of the categories involved.
In astrophysics the term metal is cast more widely to refer to all chemical elements in a star that are heavier than the lightest two, hydrogen and helium, and not just traditional metals. A star fuses lighter atoms, mostly hydrogen and helium, into heavier atoms over its lifetime. Used in that sense, the metallicity of an astronomical object is the proportion of its matter made up of the heavier chemical elements.
Metals comprise 25% of the Earth's crust and are present in many aspects of modern life. 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.
The history of metals is thought to begin with the use of copper about 11,000 years ago. Gold, silver, iron (as meteoric iron), lead, and brass were likewise in use before the first known appearance of bronze in the 5th millennium BCE. Subsequent developments include the production of early forms of steel; the discovery of sodium—the first light metal—in 1809; the rise of modern alloy steels; and, since the end of World War II, the development of more sophisticated alloys.
Metals are shiny and lustrous, at least when freshly prepared, polished, or fractured. Sheets of metal thicker than a few micrometres appear opaque, but gold leaf transmits green light.
The solid or liquid state of metals largely originates in the capacity of the metal atoms involved to readily lose their outer shell electrons. Broadly, the forces holding an individual atom’s outer shell electrons in place are weaker than the attractive forces on the same electrons arising from interactions between the atoms in the solid or liquid metal. The electrons involved become delocalised and the atomic structure of a metal can effectively be visualised as a collection of atoms embedded in a cloud of relatively mobile electrons. This type of interaction is called a metallic bond. The strength of metallic bonds for different elemental metals reaches a maximum around the center of the transition metal series, as these elements have large numbers of delocalized electrons.
Although most elemental metals have higher densities than most nonmetals, 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. 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. An iron ball would thus weigh about as much as three aluminium balls.
Metals are typically malleable and ductile, deforming under stress without cleaving. The nondirectional nature of metallic bonding is thought to contribute significantly to the ductility of most metallic solids. In contrast, in an ionic compound like table salt, when the planes of an ionic bond slide past one another, the resultant change in location shifts ions of the same charge into close proximity, resulting in the cleavage of the crystal. Such a shift is not observed in a covalently bonded crystal, such as a diamond, where fracture and crystal fragmentation occurs. Reversible elastic deformation in metals can be described by Hooke's Law for restoring forces, where the stress is linearly proportional to the strain.
Heat or forces larger than a metal's elastic limit may cause a permanent (irreversible) deformation, known as plastic deformation or plasticity. An applied force may be a tensile (pulling) force, a compressive (pushing) force, or a shear, bending or torsion (twisting) force. A temperature change may affect the movement or displacement 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 ensue.
The atoms of metallic substances are typically arranged 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.
The unit cell for each crystal structure is the smallest group of atoms which has the overall symmetry of the crystal, and from which the entire crystalline lattice can be built up by repetition in three dimensions. In the case of the body-centered cubic crystal structure shown above, the unit cell is made up of the central atom plus one-eight of each of the eight corner atoms.
The electronic structure of metals means they are relatively good conductors of electricity. Electrons in matter can only have fixed rather than variable energy levels, and in a metal the energy levels of the electrons in its electron cloud, at least to some degree, correspond to the energy levels at which electrical conduction can occur. In a semiconductor like silicon or a nonmetal like sulfur there is an energy gap between the electrons in the substance and the energy level at which electrical conduction can occur. Consequently semiconductors and nonmetals are relatively poor conductors. (The text accompanying the image in this subsection discusses this situation using more technical language.)
The elemental metals have electrical conductivity values of from 6.9 × 103 S/cm for manganese to 6.3 × 105 S/cm for silver. In contrast, a semiconducting metalloid such as boron has an electrical conductivity 1.5 × 10−6 S/cm. With one exception, metallic elements reduce their electrical conductivity when heated. Plutonium increases its electrical conductivity when heated in the temperature range of around −175 to +125 °C.
Metals are relatively good conductors of heat. The electrons in a metal's electron cloud are highly mobile and easily able to pass on heat-induced vibrational energy.
The contribution of a metal's electrons to its heat capacity and thermal conductivity, and the electrical conductivity of the metal itself can be calculated from the free electron model, albeit 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 mathematical models are applicable, the simplest being the nearly free electron model.
Metals are usually inclined to form cations through electron loss. Most will react with oxygen in the air to form oxides over various timescales (potassium burns in seconds while iron rusts over years). Some others, like palladium, platinum and gold, do not react with the atmosphere at all. The oxides of metals are generally basic, as opposed to those of nonmetals, which are acidic or neutral. Exceptions are largely oxides with very high oxidation states such as CrO3, Mn2O7, and OsO4, which have strictly acidic reactions.
Painting, anodizing or plating metals are good ways to prevent their corrosion. However, a more reactive metal in the electrochemical series must be chosen for coating, especially when chipping of the coating is expected. Water and the two metals form an electrochemical cell, and if the coating is less reactive than the underlying metal, the coating actually promotes corrosion.