Title: Rust
Content:  When iron is in contact with water and oxygen, it rusts. If salt is present, for example in seawater or salt spray, the iron tends to rust more quickly, as a result of chemical reactions. Iron metal is relatively unaffected by pure water or by dry oxygen. As with other metals, like aluminium, a tightly adhering oxide coating, a passivation layer, protects the bulk iron from further oxidation. The conversion of the passivating ferrous oxide layer to rust results from the combined action of two agents, usually oxygen and water. Other degrading solutions are sulfur dioxide in water and carbon dioxide in water. Under these corrosive conditions, iron hydroxide species are formed. Unlike ferrous oxides, the hydroxides do not adhere to the bulk metal. As they form and flake off from the surface, fresh iron is exposed, and the corrosion process continues until either all of the iron is consumed or all of the oxygen, water, carbon dioxide or sulfur dioxide in the system are removed or consumed. When iron rusts, the oxides take up more volume than the original metal; this expansion can generate enormous forces, damaging structures made with iron. See economic effect for more details.
Title: Iron nanoparticle
Content: 2 Fe0(s) + 4 H+(aq) + O2(aq) → 2 Fe2+(aq) + 2 H2O(l) Fe0(s) + 2 H2O (aq) → Fe2+(aq) + H2(g) + 2 OH−(aq) When exposed to oxygen and water, iron oxidizes. This redox process can occur under either acidic or neutral/basic conditions:  
Title: Oxide
Content:  the process of electrolytic anodizing. Though solid magnesium and aluminum react slowly with oxygen at STP—they, like most metals, burn in air, generating very high temperatures. Finely grained powders of most metals can be dangerously explosive in air. Consequently, they are often used in solid-fuel rockets. In dry oxygen, iron readily forms iron(II) oxide, but the formation of the hydrated ferric oxides, Fe2O3−x(OH)2x, that mainly comprise rust, typically requires oxygen and water. Free oxygen production by photosynthetic bacteria some 3.5 billion years ago precipitated iron out of solution in the oceans as Fe2O3 in the economically important iron ore hematite.
Title: Rust
Content: O2 + 4 + 2  → 4 Fe → Fe2+ + 2 4 Fe2+ + O2 → 4 Fe3+ + 2 O2− Fe2+ + 2 H2O ⇌ Fe(OH)2 + 2 ; Fe3+ + 3 H2O ⇌ Fe(OH)3 + 3 Fe(OH)2 ⇌ FeO + ; Fe(OH)3 ⇌ FeO(OH) + ; 2 FeO(OH) ⇌ Fe2O3 + The rusting of iron is an electrochemical process that begins with the transfer of electrons from iron to oxygen. The iron is the reducing agent (gives up electrons) while the oxygen is the oxidizing agent (gains electrons). The rate of corrosion is affected by water and accelerated by electrolytes, as illustrated by the effects of road salt on the corrosion of automobiles. The key 
Title: Electrochemistry
Content:  iron rust to occur the metal has to be in contact with oxygen and water, although chemical reactions for this process are relatively complex and not all of them are completely understood. It is believed the causes are the following: Electron transfer (reduction-oxidation) Iron corrosion takes place in an acid medium; H+ ions come from reaction between carbon dioxide in the atmosphere and water, forming carbonic acid. Fe2+ ions oxidize, following this equation: Iron(III) oxide hydrate is known as rust. The concentration of water associated with iron oxide varies, thus the chemical formula is represented by Fe2O3·xH2O. An electric circuit is formed as passage of electrons and ions occurs, thus if an electrolyte is present it will facilitate oxidation, explaining why rusting is quicker in salt water.
Title: Rust
Content:  Rust is an iron oxide, a usually reddish-brown oxide formed by the reaction of iron and oxygen in the catalytic presence of water or air moisture. Rust consists of hydrous iron(III) oxides (Fe2O3·nH2O) and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3), and is typically associated with the corrosion of refined iron. Given sufficient time, any iron mass, in the presence of water and oxygen, could eventually convert entirely to rust. Surface rust is commonly flaky and friable, and provides no passivational protection to the underlying iron, unlike the formation of patina on copper surfaces. Rusting is the common term for corrosion of elemental iron and its alloys such as steel. Many other metals undergo similar corrosion, but the 
Title: Iron(III) oxide
Content: Fe2O3 + 3 CO → 2 Fe + 3 CO2 2 Al + Fe2O3 → 2 Fe + Al2O3 3 Fe2O3 + H2 → 2 Fe3O4 + H2O ZnO + Fe2O3 → Zn(FeO2)2 The most important reaction is its carbothermal reduction, which gives iron used in steel-making: Another redox reaction is the extremely exothermic thermite reaction with aluminium. This process is used to weld thick metals such as rails of train tracks by using a ceramic container to funnel the molten iron in between two sections of rail. Thermite is also used in weapons and making small-scale cast-iron sculptures and tools. Partial reduction with hydrogen at about 400 °C produces magnetite, a black magnetic material that contains both Fe(III) and Fe(II): Iron(III) oxide is insoluble in water but dissolves readily in strong acid, e.g. hydrochloric and sulfuric acids. It also dissolves well in solutions of chelating agents such as EDTA and oxalic acid. Heating iron(III) oxides with other metal oxides or carbonates yields materials known as ferrates (ferrate (III)):
Title: Properties of water
Content: 2 Na + 2 →  + 2  + 2 4 + 2  → 4 AgF + 4 HF + Water contains hydrogen in the oxidation state +1 and oxygen in the oxidation state −2. It oxidizes chemicals such as hydrides, alkali metals, and some alkaline earth metals. One example of an alkali metal reacting with water is:  Some other reactive metals, such as aluminum and beryllium, are oxidized by water as well, but their oxides adhere to the metal and form a passive protective layer. Note that the rusting of iron is a reaction between iron and oxygen that is dissolved in water, not between iron and water. Water can be oxidized to emit oxygen gas, but very few oxidants react with water even if their reduction potential is greater than the potential of. Almost all such reactions require a catalyst. An example of the oxidation of water is: 
Title: Rust
Content:  is the reduction of oxygen: Because it forms hydroxide ions, this process is strongly affected by the presence of acid. Likewise, the corrosion of most metals by oxygen is accelerated at low pH. Providing the electrons for the above reaction is the oxidation of iron that may be described as follows: The following redox reaction also occurs in the presence of water and is crucial to the formation of rust: In addition, the following multistep acid–base reactions affect the course of rust formation:  as do the following dehydration equilibria: From the above equations, it is also seen that the corrosion products are dictated by the availability of water and oxygen. With limited dissolved oxygen, iron(II)-containing materials are favoured, 
Title: Iron
Content:  a few dollars per kilogram or per pound (see Metal). Pristine and smooth pure iron surfaces are mirror-like silvery-gray. However, iron reacts readily with oxygen and water to give brown to black hydrated iron oxides, commonly known as rust. Unlike the oxides of some other metals, that form passivating layers, rust occupies more volume than the metal and thus flakes off, exposing fresh surfaces for corrosion. Although iron readily reacts, high purity iron, called electrolytic iron, has better corrosion resistance. The body of an adult human contains about 4 grams (0.005% body weight) of iron, mostly in hemoglobin and myoglobin. These two proteins play essential roles in vertebrate metabolism, respectively oxygen transport by 

Question: what happens when iron reacts with oxygen and water?
Answer: Rust.