Decomposition<unk>  decomposes to form water and oxygen with a ΔHo of –2884.5 kJ/kg and a ΔS of 70.5 J/(mol·K): :2 → 2 + The rate of decomposition increases with rise in temperature, concentration, and pH ( being unstable under alkaline conditions), with cool, dilute, and acidic solutions showing the best stability. Decomposition is catalysed by various redox-active ions or compounds, including most transition metals and their compounds (e.g. manganese dioxide (MnO2), silver, and platinum). Certain metal ions, such as or , can cause the decomposition to take a different path, with free radicals such as the hydroxyl radical (HO•) and hydroperoxyl (HOO•) being formed. Non-metallic catalysts include potassium iodide (KI), which reacts particularly rapidly and forms the basis of the elephant toothpaste demonstration.<unk>  can also be decomposed biologically by the enzyme catalase. The decomposition of<unk>  liberates oxygen and heat; this can be dangerous, as spilling high-concentration<unk>  on a flammable substance can cause an immediate fire. Redox reactions The redox properties of<unk>  depend on pH as acidic conditions exacerbate the power of oxidizing agents and basic conditions exacerbate the power of reducing agents. As<unk>  exhibits ambivalent redox properties, being simultaneously an oxidizer or a reductant, its redox behavior immediately depends on pH. In acidic solutions, is a powerful oxidizer, stronger than chlorine, chlorine dioxide, and potassium permanganate. When used for cleaning laboratory glassware, a solution of<unk>  and sulfuric acid is referred to as Piranha solution. is a source of hydroxyl radicals (•OH), which are highly reactive. is used in the Briggs–Rauscher and Bray–Liebhafsky oscillating reactions. Oxidizingreagent Reducedproduct Oxidationpotential(V) F2 HF 3.0 O3 O2 2.1<unk>  H2O 1.8 KMnO4 MnO2 1.7 ClO2 HClO 1.5 Cl2 Cl− 1.4 In acidic solutions, is oxidized to<unk>  acting as an oxidizing agent): :2 (aq) + + 2 (aq) → 2 (aq) + 2 (l) and sulfite () is oxidized to sulfate (). However, potassium permanganate is reduced to by acidic . Under alkaline conditions, however, some of these reactions reverse; for example, is oxidized to (as ). In basic solutions,<unk>  is a strong reductant and can reduce a variety of inorganic ions. When acts as a reducing agent, oxygen gas is also produced. For example,<unk>  will reduce sodium hypochlorite and potassium permanganate, which is a convenient method for preparing oxygen in the laboratory: :NaOCl + → + NaCl + :2 + 3 → 2 + 2 KOH + 2 + 3 Organic reactions<unk>  is frequently used as an oxidizing agent. Illustrative is oxidation of thioethers to sulfoxides: :Ph + → Ph + Alkaline<unk>  is used for epoxidation of electron-deficient alkenes such as acrylic acid derivatives, and for the oxidation of alkylboranes to alcohols, the second step of hydroboration-oxidation. It is also the principal reagent in the Dakin oxidation process. Precursor to other<unk>  compounds<unk>  is a weak acid, forming hydroperoxide or<unk>  salts with many metals. It also converts metal oxides into the corresponding<unk> s. For example, upon treatment with<unk> , chromic acid ( and ) forms a blue<unk>  CrO(. This kind of reaction is used industrially to produce peroxoanions. For example, reaction with borax leads to sodium perborate, a bleach used in laundry detergents: : + 4 + 2 NaOH → 2 + converts carboxylic acids (RCO2H) into peroxy acids (RC(O)O2H), which are themselves used as oxidizing agents.<unk>  reacts with acetone to form acetone<unk>  and with ozone to form trioxidane.<unk>  forms stable adducts with urea<unk>  - urea), sodium carbonate (sodium percarbonate) and other compounds. An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for in some reactions.<unk>  is both an oxidizing agent and reducing agent. The oxidation of<unk>  by sodium hypochlorite yields singlet oxygen. The net reaction of a ferric ion with<unk>  is a ferrous ion and oxygen. This proceeds via single electron oxidation and hydroxyl radicals. This is used in some organic chemistry oxidations, e.g. in the Fenton's reagent. Only catalytic quantities of iron ion is needed since<unk>  also oxidizes ferrous to ferric ion. The net reaction of<unk>  and permanganate or manganese dioxide is manganous ion; however, until the<unk>  is spent some manganese ions are reoxidized to make the reaction catalytic. This forms the basis for common monopropellant rockets.
Hydrogen peroxide
Carbon monoxide
Hydroxyl radical
Cyanide
Nitrosamine
Isothiocyanate
Nicosan
Hydrogen fluoride