I heard once that Oxygen (O2) is one of the most corrosive substances in the universe. Any truth to this?
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Oxygen, like a lot of substances our body needs, can be harmful in large quantities.  This is partly due to its highly electromagnetic and hydrophyllic nature.  As any patient that has been on oxygen therapy can tell you, your body is "set" to use a certain percentage of oxygen in the air we breathe.  If we get used to a higher amount, and then this amount is depreived, we feel light-headed and weak.  This is because at higher levels of oxygen, our lungs become accustomed to a smaller amount of effort to extract and use the oxygen from the air.  \n\nOxygen's corrosive nature is not found in the O2 form however, but the O and O3 (ozone) forms.  This is because oxygen is most stable in the O2 form.  This is where oxygen "wants" to be.  The O3 and O forms are unstable, and prone to reactions.    \n\n"Oxygen can be toxic at elevated partial pressures (i.e. high relative concentrations). This is important in some forms of scuba diving, such as with a rebreather.\n\nCertain derivatives of oxygen, such as ozone (O3), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. The body has developed mechanisms to protect against these toxic compounds. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin. To protect against the destructive nature of peroxides, nearly every organism on earth has developed some form of the enzyme catalase, which very quickly disproportionates peroxide into water and dioxygen.\n\nHighly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels. The fire that killed the Apollo 1 crew on a test launchpad spread so rapidly because the capsule was pressurized with pure oxygen as would be usual in an actual flight, but to maintain positive pressure in the capsule, this was at slightly more than atmospheric pressure instead of the 1/3 pressure that would be used in flight. (See partial pressure.) Similar hazards also apply to compounds of oxygen with a high oxidative potential, such as chlorates, perchlorates, and dichromates; they also can often cause chemical burns.\n\nOxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA, they form part of theories of carcinogenesis and aging." (Wikipedia 2006)\n\n"Oxygen is the most abundant element on this planet. The earth's crust is 46.6% oxygen by weight, the oceans are 86% oxygen by weight, and the atmosphere is 21% oxygen by volume. The name oxygen comes from the Greek stems oxys, "acid," and gennan, "to form or generate." Thus, oxygen literally means "acid former." This name was introduced by Lavoisier, who noticed that compounds rich in oxygen, such as SO2 and P4O10, dissolve in water to give acids.\n\n"Fluorine is the only element that is more electronegative than oxygen. As a result, oxygen gains electrons in virtually all its chemical reactions. Each O2 molecule must gain four electrons to satisfy the octets of the two oxygen atoms." ("Chemistry of Oxygen" - Bodner Research Group)
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strictly speaking, substances undergo combustion in the presence of oxygen and 'burn' would not be precise. \nhowever answering the question, oxygen is an oxidising substance and hence it is able to oxidise the organic subs. and usually converting them into carbon di oxide and water, and this is accompanied by the release of energy in the form of heat
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copper most easily interact with oxygen than the other metals\ncoins are made of
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Oxygen Facts\nAtomic Number  8\nSymbol  O\nAtomic Weight  15.9994 \nDiscovery  Joseph Priestley; Scheele 1774. \nElectron Configuration  [He]2s22p4 \nWord Origin  Greek: oxys: sharp or acid and Greek: genes: born, former... 'acid former'. \nIsotopes  Nine isotopes of oxygen are known. Natural oxygen is a mixture of three isotopes. \nProperties  Oxygen gas is colorless, odorless, and tasteless. The liquid and solid forms are a pale blue color and are strongly paramagnetic. Oxygen supports combustion, combines with most elements, and is a component of hundreds of thousands of organic compounds. Ozone (O3), a highly active compound with a name derived from the Greek word for 'I smell', is formed by the action of an electrical discharge or ultraviolet light on oxygen. \nUses  Oxygen was the atomic weight standard of comparison for the other elements until 1961 when the International Union of Pure and Applied Chemistry adopted carbon 12 as the new basis. It is the third most abundant element found in the sun and the earth, and it plays a part in the carbon-nitrogen cycle. Excited oxygen yields the bright red and yellow-green colors of the Aurora. Oxygen enrichment of steel blast furnaces accounts for the greatest use of the gas. Large quantities are used in making synthesis gas for ammonia, methanol, and ethylene oxide. It is also used as a bleach, for oxidizing oils, for oxy-acetylene welding, and for determining carbon content of steel and organic compounds. Plants and animals require oxygen for respiration. Hospitals frequently prescribe oxygen for patients. Approximately two thirds of the human body and nine tenths of the mass of water is oxygen. \n\nOxygen can exist in a number of physical forms. Most commonly it is found as a diatomic O-O gas and is called molecular oxygen (O2) or dioxygen. The gas exists as a triplet state but the singlet oxygen can also be formed and is more reactive. This gas can be liquefied at low temperatures to produce a blue liquid oxygen (LOX) and cooled further produces a blue solid. After electrical discharge or photolysis Oxygen may also form the triatomic ozone (O3) and under very extreme conditions atomic oxygen (O•) is produced.
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One of these scientists, a man named Democritus, would sit here in the central market place of Athens, Greece, and explain to one and all how all things on earth and in the heavens are made of extremely tiny, simple particles. He called them "atoms.o/oo \n\n\nThus when the element hydrogen combines with the element oxygen, it must do so all or none. In the case of water, for instance, one atom of oxygen combines with two atoms of hydrogen. Now when oxygen combines with carbon to make carbon dioxide, two oxygen atoms combine with one carbon atom. And when they make carbon monoxide, it is one atom oxygen to one atom carbon. So, too, all other chemical compounds are made of a simple combination of a few basic atoms.
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Although the type of corrosion most of us are familiar with involves molecular oxygen, corrosion, in general, does *NOT* require the presence of free oxygen.  \n\nFor example, galvanic corrosion, which can occur when two dissimilar metals with different oxidation potentials are in electrical contact.  The less noble metal will form an anode, and be oxidized, while the more noble metal will form a cathode, where reduction reactions will occur.  In the absence of oxygen, reduction reactions like:\n\n  2H+ * 2e- -> H2(gas)\n\ncan occur.  This reaction is particularly important in acidic environments, (when the H+ concentration is high).  Such environments can develop in corrosion environments via hydrolysis reactions, e.g., \n\n  Fe3+ + 3H2O -> Fe(OH)3(solid) + 3H+\n\nwhere the Fe3+ ions are produced by the oxidation of iron metal, and Fe(OH)3 is the solid mineral "goethite", a major component of "rust".\n\nCorrosion can also be induced in the absence of free oxygen by the action of some species of bacteria. For example, there are bacteria that can reduce sulfate ions (SO4--) or nitrate ions (NO3-) to hydrogen sulfide (or metal sulfides) and molecular nitrogen.  To do so, they need to couple these reactions to a "source" of electrons, which can be a metal surface that is then oxidized (i.e., corroded). \n\n\nFinally, there are some metals that are themselves strong reductants -- so strong that they can reduce water.  Li, Mg, Na, and Al all have standard oxidation potentials that are lower than that of water, so they will react spontaneously (in the case of Li and Na, *violently*) with pure water (even water containing no free oxygen) to form metal ions or hydroxides and hydrogen gas.
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H2O2 or Hydrogen Peroxide is a corrosive and oxidizer and reacts with all types of metals (except stainless steel and aluminum), dirt, leather, ect. If you would like to see how 70% concentration reacts with a leather glove or laundry detergent. send me an E-Mail to m00nd0ct0r@yahoo.com and I'll send you the two videos later tonight when I get home form work.\n\nP.S. I'm working at a Hydrogen Peroxide plant right now.
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Defiantly, I think that's how the world will end. I think and can all but prove my theory. Ozone is what protects us from the sun. all of the nine planets in our solar sys have had ozone at one time. out of the the first 4 planets we are the only planet left with ozone the other 3 are all charred and burned. THAT'S THE  EXACT THING THAT WILL HAPPEN TO US IF THAT HOLE IN OUR OZONE CONTINUES TO GROW . WE WILL BURN AND BECOME THE FOURTH CHARRED PLANET. along with mercury Venus and mars. That's what happened to them. Their ozone depleted and they burned.
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"Iodine, it is sublime". (Because it sublimates-goes directly from solid to a gas).\n\nGet an "A".
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Peroxide is H2O2. If it were in a clear bottle, that second oxygen would evaporate and it would become H20...water.
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Well to start off any reaction that produces oxygen is potentially combustible, especially one that requires energy (endothermic) so be careful.\n\nTwo suggestions: \nFirst, electrolysis, this is where you break the covalent bond between oxygen and hydrogen with electricity (energy).  The by products are oxygen and hydrogen gas.\n\nMy second idea is simple rust.  Put iron (Fe) into water (H2O) and you will produce Iron-oxide or rust.  The Oxygen combines with the iron.\n\nGood luck.
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AlthoughO2 is a strong oxidizing agent, its oxidizing reactions are usually slow kinetically. Many flammable materials will exist in the presence of air until a flame or spark starts the reaction. The very electronegative nonmetals are poor reducing agents. \n\nF2 will oxidize all metals and most other substances. It must be stored in either nickel or copper containers because the fluorides of these metals form a film passivating the metal surface. \n\nFor more please visit:\nhttp://www.wou.edu/las/physci/ch412/nonmetal.htm
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You have a misconception as to what an ozone layer is.  Ozone is a gas and is found throughout the atmosphere, even at ground level.  The ozone layer is the range in the atmosphere where ozone levels are the highest.  Oxygen O2 is present in the ozone layer as well (as are all other atmospheric gasses).  You should also note that in the ozone layer there isn't actually a whole lot of ozone, just more than there is anywhere else.  I think it's around 5 parts per million; meaning there is 1 molecule of ozone for every 200,000 other molecules.  There isn't a blanket of ozone, it's just an area where ozone is 5-10ppm instead of 1ppm or less.  \n\nWhy does most of it stay where it is?  Ozone is a somewhat unstable molecule and the same mechanism that typically forms ozone also breaks it down.  Simply, O2 gets hit by some photons of the right energy (some kind of UV light) and is split into O and O, one (or both) of those O's bonds to another O2 molecule to form O3, ozone.  Sooner or later, that O3 gets hit by the right energy of UV light and one of the O's break off to form O2 and a spare O which will find another O2 to bond with and repeat the cycle.  A small percent of the ozone does drop into the statosphere (where we live).
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carbon- obviously carbon\nmonoxide- mono-one oxide-oxygen
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Yes.  The methane and hydrogen, the main ingredients, are flamable.  The perfect circumstances?  That's a tough one.  Perhaps if you were a bit stoned so that it would be humorous?
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Ozone (O3) is a triatomic molecule, consisting of three oxygen atoms. It is an allotrope of oxygen that is much less stable than the diatomic species O2. It is present in low concentrations throughout the Earth's atmosphere. It has many industrial and consumer applications as well as being used in ozone therapy.\n\nOzone, the first allotrope of a chemical element to be described by science, was discovered by Christian Friedrich Schönbein in 1840, who named it after the Greek word for smell (ozein), from the peculiar odour in lightning storms.[1] The odour from a lightning strike is from electrons freed during the rapid chemical changes, not the ozone itself.[2]\n\nPhysical properties\nUndiluted ozone is a pale blue gas at standard temperature and pressure; it forms a dark blue liquid below −112 °C and a violet-black solid below −193 °C [3]. At concentrations found in the atmosphere it is colorless [4]. The concentration above which it can be smelt (odor threshold) is between 0.0076 and 0.036 ppm[5].\n\n\n[edit] Chemistry\nOzone is a powerful oxidizing agent. It is also unstable at high concentrations, decaying to ordinary diatomic oxygen:\n\n2 O3 → 3 O2. \nThis reaction proceeds more rapidly with increasing temperature and decreasing pressure. Ozone will oxidize metals (except gold, platinum, and iridium) to oxides of the metals in their highest oxidation state:\n\n2 Cu2+ + 2 H+ + O3 → 2 Cu3+ + H2O + O2 \nOzone converts oxides to peroxides:\n\nSO2 + O3 → SO3 + O2 \nIt also increases the oxidation number of oxides:\n\nNO + O3 → NO2 + O2 \nThe above reaction is accompanied by chemiluminescence. The NO2 can be further oxidized:\n\nNO2 + O3 → NO3 + O2 \nThe NO3 formed can react with NO2 to form N2O5:\n\nNO2 + NO3 → N2O5 \nOzone reacts with carbon to form carbon dioxide, even at room temperature:\n\nC + 2 O3 → CO2 + 2 O2 \nOzone does not react with ammonium salts but it reacts with ammonia to form ammonium nitrate:\n\n2 NH3 + 4 O3 → NH4NO3 + 4 O2 + H2O \nOzone reacts with sulfides to make sulfates:\n\nPbS + 4 O3 → PbSO4 + 4 O2 \nSulfuric acid can be produced from ozone, either starting from elemental sulfur or from sulfur dioxide:\n\nS + H2O + O3 → H2SO4 \n3 SO2 + 3 H2O + O3 → 3 H2SO4 \nAll three atoms of ozone may also react, as in the reaction with tin(II) chloride and hydrochloric acid:\n\n3 SnCl2 + 6 HCl + O3 → 3 SnCl4 + 3 H2O \nIn the gas phase, ozone reacts with hydrogen sulfide to form sulfur dioxide:\n\nH2S + O3 → SO2 + H2O \nIn an aqueous solution, however, two competing simultaneous reactions occur, one to produce elemental sulfur, and one to produce sulfuric acid:\n\nH2S + O3 → S + O2 + H2O \n3 H2S + 4 O3 → 3 H2SO4 \nIodine perchlorate can be made by treating iodine dissolved in cold anhydrous perchloric acid with ozone:\n\nI2 + 6 HClO4 + O3 → 2 I(ClO4)3 + 3 H2O \nSolid nitryl perchlorate can be made from NO2, ClO2, and O3 gases:\n\n2 NO2 + 2 ClO2 2 O3 → 2 NO2ClO4 + O2 \nOzone can be used for combustion reactions and combusting gases in ozone provides higher temperatures than combusting in dioxygen (O2). Following is a reaction for the combustion of carbon subnitride:\n\n3 C4N2 + 4 O3 → 12 CO + 3 N2 \nOzone can react at cryogenic temperatures. At 77 K (-196 °C), atomic hydrogen reacts with liquid ozone to form a hydrogen superoxide radical, which dimerizes[6]:\n\nH + O3 → HO2 + O \n2 HO2 → H2O4 \nOzonides can be formed, which contain the ozonide anion, O3-. These compounds are explosive and must be stored at cryogenic temperatures. Ozonides for all the alkali metals are known. KO3, RbO3, and CsO3 can be prepared from their respective superoxides:\n\nKO2 + O3 → KO3 + O2 \nAlthough KO3 can be formed as above, it can also be formed from potassium hydroxide and ozone:[7]\n\n2 KOH + 5 O3 → 2 KO3 + 5 O2 + H2O \nNaO3 and LiO3 must be prepared by action of CsO3 in liquid NH3 on an ion exchange resin containing Na+ or Li+ ions:[8]\n\nCsO3 + Na+ → Cs+ + NaO3 \nTreatment wit
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*Advertisement*\n*voice-over*\n"Are you tired of factually incorrect answers?"\n"Are you tired of answers that make a mockery of your question?"\n"Are you disappointed that your good question didn't get the response it deserved?"\n\nNow, there is no need to worry, as we introduce you to the most reliable answering machine- \nBATMAN!! *Cheers*\n\nTa daaaa..\n\n*End of commercial*\n\nHi Jade,\n\nI've got the answer! Here it comes- \nPeroxide bubbles because of constant production of oxygen gas as a result of spontaneous decomposition of peroxide.\n\nThe oxidation state of oxygen in peroxide(H2O2) is (-1) i.e. one electron for each oxygen atom.\n\nIn water (H20), the oxidation state is (-2) i.e two electrons for one oxygen atom. \n\nIn an oxygen molecule (O2), the oxidation state is zero. \n\nNow, the aqueous solutions of hydrogen peroxide spontaneously decomposes like this-\n\nOne molecule of Peroxide (liquid) ----> One molecule of Water (Liquid) + Half molecule of Oxygen (gaseous)\n\nSymbolically-\nH202 (l)   -------> H2O (l) + (1/2)O2 (g)\n(Dammit, how do I put numbers in sub-script and super-script?)\n\nNow, the Oxygen produced here causes production of bubbles.\n\nThis decomposition occurs slowly but the presence of metal surfaces, traces of alkali (present in glass) strongly catalyze the reaction. \n\nSo, if you don't want all your peroxide to convert into water, don't keep it in metal or glass bottles! Keep it in wax-lined glass or plastic vessels in the presence of stabilizers like Urea.\n\nIt is also decomposed by exposure to light. \n\nAlso, keep it away from dust, dust (even in small quantities) induces explosive decomposition.