Manganese combined with other elements is widely distributed in the Earth's crust. Manganese is second only to iron among the transition elements in its abundance in the Earth's crust; it is roughly similar to iron in its physical and chemical properties but is harder and more brittle. It occurs in a number of substantial deposits, of which the most important ores (which are mainly oxides) consist primarily of manganese dioxide (MnO₂) in the form of pyrolusite, romanechite, and wad. Manganese is essential to plant growth and is involved in the reduction of nitrates in green plants and algae. It is an essential trace element in higher animals, in which it participates in the action of many enzymes. Lack of manganese causes testicular atrophy. An excess of this element in plants and animals is toxic.

More than 95 percent of the manganese produced is used in the form of ferromanganese and silicomanganese alloys for iron and steel manufacture. Manganese ores containing iron oxides are first reduced in blast furnaces or electric furnaces with carbon to yield ferromanganese, which in turn is used in steelmaking. Adding manganese, which has a greater affinity for sulfur than does iron, converts the low-melting iron sulfide in steel to high-melting manganese sulfide. Produced without manganese, steel breaks up when hot-rolled or forged. Steels generally contain less than 1 percent manganese. Manganese steel, also called Hadfield steel, is used for very rugged service; containing 12–14 percent manganese, it provides a hard, wear-resistant, and self-renewing surface over a tough unbreakable core. Pure manganese produced electrolytically is used mostly in the preparation of nonferrous alloys of copper, aluminum, magnesium, and nickel and in the production of high-purity chemicals. Practically all commercial alloys of aluminum and magnesium contain manganese to improve corrosion resistance and mechanical properties. (For detailed information on the extraction, refining, and applications of manganese, see manganese processing.)

All natural manganese is stable isotope manganese-55. It exists in four allotropic modifications; the complex cubic structure of the so-called alpha phase is the form stable at ordinary temperatures. Manganese somewhat resembles iron in general chemical activity. The metal oxidizes superficially in air and rusts in moist air. It burns in air or oxygen at elevated temperatures, as does iron; decomposes water slowly when cold and rapidly on heating; and dissolves readily in dilute mineral acids with hydrogen evolution and the formation of the corresponding salts in the +2 oxidation state.

Manganese is quite electropositive, dissolving very readily in dilute non-oxidizing acids. Although relatively unreactive toward nonmetals at room temperature, it reacts with many at elevated temperatures. Thus, manganese burns in chlorine to give manganese dichloride, MnCl₂; reacts with fluorine to give manganese di- and trifluorides, MnF₂ and MnF₃; and burns in nitrogen at about 1,200° C to give trimanganese dinitride, Mn₃N₂, and in oxygen to give trimanganese tetroxide, Mn₃O₄. Manganese also combines directly with boron, carbon, sulfur, silicon, or phosphorus, but not with hydrogen.

Of the wide variety of compounds formed by manganese, the most stable occur in oxidation states +2, +6, and +7. These are exemplified, respectively, by the manganous salts (with manganese as the Mn²⁺ ion), the manganates (MnO₄²⁻), and the permanganates (MnO₄⁻). As in the case of titanium, vanadium, and chromium, the highest oxidation state (+7) of manganese corresponds to the total number of 3d and 4s electrons. That state occurs only in the oxo species permanganate (MnO₄⁻), dimanganese heptoxide (Mn₂O₇), and manganese trioxide fluoride (MnO₃F), which show some similarity to corresponding compounds of the halogens—for example, in the instability of the oxide. Manganese in oxidation state +7 is powerfully oxidizing, usually being reduced to manganese in the +2 state. The intermediate oxidation states are known, but, except for some compounds in the +3 and +4 states, they are not particularly important.

The principal industrial compounds of manganese include several oxides. Manganous oxide, or manganese monoxide, MnO, is used as a starting material for the production of manganous salts, as an additive in fertilizers, and as a reagent in textile printing. It occurs in nature as the green mineral manganosite. It also can be prepared commercially by heating manganese carbonate in the absence of air or by passing hydrogen or carbon monoxide over manganese dioxide.

The most important manganese compound is manganese dioxide, in which manganese is in the +4 oxidation state, and the black mineral pyrolusite is the chief source of manganese and all of its compounds. It is also widely used as a chemical oxidant in organic synthesis. Manganese dioxide is used as the cathode material in electric dry cells. It is produced directly from the ore, although substantial amounts are also prepared synthetically. The synthetic oxide is prepared by decomposition of manganous nitrate; by reaction of manganous sulfate, oxygen, and sodium hydroxide; or by electrolysis of an aqueous solution of manganese sulfate.

Various manganese salts also have commercial importance. Manganous sulfate (MnSO₄) is added to soils to promote plant growth, especially of citrus crops. In addition, it is a good reducing agent, particularly useful in the manufacture of paint and varnish driers. Manganous chloride (MnCl₂) is widely employed as a catalyst in the chlorination of organic compounds and as a feed additive. The deep-purple compound potassium permanganate (KMnO₄) is used for disinfecting, deodorizing, and bleaching and as an analytical reagent.