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In 1808, Guy-Lussac and others in France produced boron by heating boric acid H3BO3 and potassium in a copper tube. Boron is an element of Group 3A in the second period, with the symbol B. Boron exists in nature as oxides, mainly boric acid and various borates. Boric acid is found in certain hot springs, and the most well-known borate is borax, Na2B4O7·10H2O. Furthermore, industrially important are ulexite 2MgO ·B2O3·H2O and borocalcite 2CaO ·3B2O3·5H2O. Since a series of borohydrides were synthesized in the 1960s of this century, the prospects for boron chemistry have become apparent. Subsequently, not only have significant advances been made in the study of the bonding properties of boron atoms as well as the structures of elemental boron and its compounds, but also a range of boron compounds with important applications in modern industry and defense have been developed, making boron the element that has been most extensively studied in the field of inorganic chemistry over the past 20 years. The valence electron configuration of a boron atom is 2s22p1; the number of valence electrons (3) is less than the number of valence orbitals (4), making it an electron-deficient atom that can form electron-deficient compounds. Boron exists in two forms: amorphous and crystalline. The amorphous form appears as a brownish-black to black powder, while the crystalline form is dark black to silver-gray in color, with a metallic luster; its hardness is similar to that of diamond. Amorphous boron has a density of 2.3 g·cm-3, while crystalline boron has a density of 2.31 g·cm-3. Amorphous boron oxidizes slowly in air at room temperature, and can catch fire spontaneously at around 800°C. Crystalline boron does not react with hydrochloric acid or hydrofluoric acid even when boiled for a long time, but it can be slowly eroded and oxidized by hot concentrated nitric acid and a mixture of sodium dichromate and sulfuric acid. Hydrogen peroxide and persulfuric acid can also slowly oxidize crystalline boron. The aforementioned reagent reacts violently with amorphous boron. When fused with mixtures of alkali metal carbonates and hydroxides, all forms of boron are completely oxidized. Fluorine, chlorine, bromine, and boron react to form the corresponding borohalides; at around 600°C, boron reacts violently with sulfur to produce a mixture of boron sulfides. When boron is heated to over 1,000°C in nitrogen or ammonia, boron nitride is formed. At temperatures of 1,800–2,000°C, boron and hydrogen do not react. Boron and silicon react at temperatures above 2,000°C to form silicoboride. At high temperatures, boron can react with many metals and metal oxides to form metal borides. Boron is particularly reactive at high temperatures; in metallurgy, it is used as an alloying agent for iron, manganese, and other metals, as a deoxidizing additive for steel and copper, and also in the heat treatment of wrought iron to enhance the high-temperature strength of alloy steels. Boron is also used in nuclear reactors and high-temperature technologies. Rod-shaped and strip-shaped boron steel are widely used as control rods in nuclear reactors. Due to its properties such as low density, high strength, and high melting point, boron can be used to manufacture certain structural materials in missiles and rockets. Boron compounds also have a wide range of applications in agriculture, medicine, the glass industry, and other fields. Boron is an essential micronutrient for plant growth. Elemental boron can be obtained by high-temperature reduction of boron oxides using active metals, reduction of boron halides with hydrogen, reduction of borax with carbon, thermal decomposition of borohydrides, and electrolysis of molten borates or other boron-containing compounds. The crude products obtained through these methods must undergo vacuum degassing or controlled halogenation in order to yield highly pure boron.