تجزیه این ماده باعث ایجاد رادیکالهای OH میشود که بیش از چند ثانیه در دسترس نمیباشند و در این مدت با خاصیت شدید اکسیدکنندگی خود، مواد آلی و معدنی را اکسید میکند.
خصوصیات آب اکسیژنه[ویرایش]
به مرور آب اکسیژنه تجزیه و تبدیل به آب و اکسیژن میگردد. این عمل تجزیه در محیط بازی سریعتر و در محیط اسیدی کندتر از محیط خنثی صورت میگیرد. اگر مدت مدیدی آب اکسیژنه را انبار کنند، ممکن است کاملاً تجزیه و تبدیل به آب گردد. بر اثر گرد بعضی اجسام عمل تخریب آب اکسیژنه تسریع میگردد مانند گرد بیاکسید منگنز و گرد فلزات.
اگر بر روی محلول آب قدری از اجسام پایدار کننده مانند اسید فسفریک، اوره، اسید بنزوئیک و نظیر آنها بیفزایند، عمل تخریب بسیار کند میگردد. آب اکسیژنه اثر میکروب کشی ، بوبری و سفید کنندگی دارد چنانکه اگر یک تکه کالباس قرمز را درون ظرف محتوی آب اکسیژنه قرار دهیم پس از چند روز محتویات ظرف کاملاً بیبو است و بوی گندیده نمیدهد. آب اکسیژنه رنگها را نیز تخریب میکند بهمین دلیل تکه کالباس درون ظرف بعد از مدتی بیرنگ میشود.
موارد استعمال آب اکسیژنه[ویرایش]
لکه شراب قرمز، خون، قهوه و غیره را هم میتوان بوسیله آب اکسیژنه پاک نمود. محلول غلیظ H۲O۲ به عنوان یک اکسیدان برای سوخت موشکها نیز مورد استفاده قرار میگیرد. برخی از خمیر دندانها و سایر اجسامی که برای پاک کردن دندانها بکار میرود در موقع استعمال تولید آب اکسیژنه میکنند و اکسیژن این آب اکسیژنه دندان را سفید مینماید.
آب اکسیژنه در بیرنگ کردن شاخ، پشم گوسفند، پنبه، کتان، کنف، کاه، چوب، کاغذ، روغن، چربی، واکس، صابون، ابریشم، عاج، پر و غیره بکار میرود. رنگ بعضی لکههای صورت را هم آب اکسیژنه تخریب میکند. اگر موی سیاه را پس از شستن با کربنات سدیم (تا چربی آن برطرف شود) در محلول آب اکسیژنه بگذارند به رنگ روشن در میآید. اگر موی سیاهسر را با مخلوطی از ۱۰۰ گرم آب اکسیژنه ۳۰٪ و چهار قطره محلول ۲۵٪ آمونیاک تر نمایند و پس از ۱۰ تا ۲۰ دقیقه با آب خالص و سپس با محلول اسید استیکدار بشویند، بور مایل به قرمز میشود. وجود آمونیاک از این جهت لازم است که آب اکسیژنه در حضور قلیاییها سریعتر اکسیژن میدهد و در نتیجه موها تندتر بور میشوند. مصرف مکرر آب اکسیژنه برای مو مضر است زیرا که مو را شکننده مینماید. در جنگ جهانی دوم آب اکسیژنه ۸۵٪برای اکسیداسیون سریع الکل در زیر دریاییها و موشکها مصرف میکردند.
کاربرد در پزشکی[ویرایش]
آب اکسیژنه در گذشته به دلیل خاصیت ضدعفونی کننده آن در پانسمان زخمهای عفونی استفاده میشد ولی امروزه به دلیل آسیبی که به بافتهای مجاور وارد میکند دیگر در پانسمان استفاده نمیشود و فقط گاه برای ضدعفونی لوازم یا سطوح استفاده میشود. از آنجایی که آب اکسیژنه بوبر است گاه در درمان زخمهای بدبو مورد استعمال قرار میگیرد. در قرصهای اریتزون ۳۶٪ آب اکسیژنه به ۶۴٪ اوره متصل است و چون این قرصها را در دهان بگذارند، اکسیژن میدهد. پس هم میکروبهای دهان را میکشد و هم دندانها را سفید مینماید. آب اکسیژن رقیق را برای قرقره کردن هم بکار میبرند.
کاربرد در بهداشت استخرها[ویرایش]
آب اکسیژنه برای حذف مواد آلی و معدنی که موجب فاسد شدن آب استخر میشوند بکار میرود. تزریق این عنصر قبل از دستگاه UV باعث ایجاد رادیکالهای OH میشود که بیش از چند ثانیه در دسترس نمیباشند و در این مدت با خاصیت شدید اکسید کنندگی خود، مواد باقیمانده آلی و معدنی را تجزیه میکند. بدین ترتیب نیاز به تعویض آب استخرها کاهش مییابد.
در صورتی که از آب اکسیژنه در آب استخرها استفاده شود دارای مزایا به شرح زیر است: ۱. آب اکسیژنه موجود در آب با دوز صحیح، برای شناگر غیرقابل تشخیص است. ۲. با آب بخوبی مخلوط میشود، غیر فرار است و تا زمان اکسید کردن مواد آلی در آب باقی میماند. ۳. خالص است و ایجاد املاح نمیکند. ۴. خورنده نیست و در نتیجه به تجهیزات و تاسیسات آسیب نمیرساند. ۵. ایجاد کف نمیکند، بیبو و بی طعم است. ۶. غیر سمی است ۷. ایجاد رسوب نکرده و در نتیجه آب کاملاً شفاف میماند.
شناسایی آب اکسیژنه[ویرایش]
در یک لوله آزمایشی که قبلاً چند سانتیمتر مکعب محلول بیکرمات پتاسیم و قدری اسید سولفوریک رقیق ریختهایم آب اکسیژنه میافزاییم در نتیجه رنگ آبی تند که بعداً تبدیل به سبز میشود، ظاهر میگردد. بهمین طریق میتوان وجود آب اکسیژنه را در اریتزن ثابت نمود.
تهیه آب اکسیژن در صنعت با روش خود اکسایش[ویرایش]
در این فرایند یکی از مشتقات آتراکینون بر اثر واکنش با هیدروژن در مجاورت کاتالیزور پالادیوم به آنتراهیدروکینون تبدیل میشود. با عبور هوا از ماده اخیر، محلول پراکسید هیدروژن ۲۰ درصد وزنی بدست میآید.
رقیه قدیم خانی و دکتر لطفعلی سقط فروش، آزمایشگاه شیمی معدنی ۱، انتشارات پیام نور
Hydrogen peroxide is a chemical compound with the formula H
Hydrogen peroxide is often described as being “water but with one more oxygen atom”, a description which can give the incorrect impression that there is a great deal of similarity between the two compounds. Pure hydrogen peroxide will explode if heated to boiling, will cause serious contact burns to the skin and can set materials alight on contact. For these reasons it is usually handled as a dilute solution (household grades are typically 3-6%). Its chemistry is dominated by the nature of its unstable peroxide bond.
Structure and properties
The boiling point of H
In aqueous solutions
In aqueous solutions hydrogen peroxide differs from the pure material due to the effects of hydrogen bonding between water and hydrogen peroxide molecules. Hydrogen peroxide and water form a eutectic mixture, exhibiting freezing-point depression; pure water has a melting point of 0 °C and pure hydrogen peroxide of −0.43 °C, but a 50% (by volume) solution of the two freezes at -51 °C. The boiling point of the same mixtures is also depressed in relation with the median of both boiling points (125.1 °C). It occurs at 114 °C. This boiling point is 14° greater than that of pure water and 36.2° less than that of pure hydrogen peroxide.
Hydrogen peroxide (H
The molecular structures of gaseous and crystalline H
Comparison with analogues
Hydrogen peroxide has several structural analogues with Hm-E-E-Hn bonding arrangements (Water also shown for comparison). It has the highest (theoretical) boiling point of this series (X = O, N, S). Its melting point is also fairly high, being comparable to that of hydrazine and water, with only hydroxylamine crystallising significantly more readily, indicative of particularly strong hydrogen bonding. Diphosphane and hydrogen disulfide exhibit only weak hydrogen bonding and have little chemical similarity to hydrogen peroxide. All of these analogues are thermodynamically unstable. Structurally, the analogues all adopt similar skewed structures, due to repulsion between adjacent lone pairs.
Hydrogen peroxide was first described in 1818 by Louis Jacques Thénard, who produced it by treating barium peroxide with nitric acid. An improved version of this process used hydrochloric acid, followed by addition of sulfuric acid to precipitate the barium sulfate byproduct. Thénard's process was used from the end of the 19th century until the middle of the 20th century.
Pure hydrogen peroxide was long believed to be unstable as early attempts to separate it from the water, which is present during synthesis, all failed. This instability was due to traces of impurities (transition metals salts) which catalyze the decomposition of the hydrogen peroxide. Pure hydrogen peroxide was first obtained in 1894 — almost 80 years after its discovery — by Richard Wolffenstein, who produced it via vacuum distillation.
Determination of the molecular structure of hydrogen peroxide proved to be very difficult. In 1892 the Italian physical chemist Giacomo Carrara (1864–1925) determined its molecular weight by freezing point depression, which confirmed that its molecular formula is H2O2. At least half a dozen hypothetical molecular structures seemed to be consistent with the available evidence. In 1934, the English mathematical physicist William Penney and the Scottish physicist Gordon Sutherland proposed a molecular structure for hydrogen peroxide which was very similar to the presently accepted one.
Previously, hydrogen peroxide was prepared industrially by hydrolysis of the ammonium peroxydisulfate, which was itself obtained via the electrolysis of a solution of ammonium bisulfate (NH
Today, hydrogen peroxide is manufactured almost exclusively by the anthraquinone process, which was formalized in 1936 and patented in 1939. It begins with the reduction of an anthraquinone (such as 2-ethylanthraquinone or the 2-amyl derivative) to the corresponding anthrahydroquinone, typically via hydrogenation on a palladium catalyst; the anthrahydroquinone then undergoes to autoxidation to regenerate the starting anthraquinone, with hydrogen peroxide being produced as a by-product. Most commercial processes achieve oxidation by bubbling compressed air through a solution of the derivatized anthracene, whereby the oxygen present in the air reacts with the labile hydrogen atoms (of the hydroxy group), giving hydrogen peroxide and regenerating the anthraquinone. Hydrogen peroxide is then extracted and the anthraquinone derivative is reduced back to the dihydroxy (anthracene) compound using hydrogen gas in the presence of a metal catalyst. The cycle then repeats itself.
The simplified overall equation for the process is deceptively simple:
A process to produce hydrogen peroxide directly from the elements has been of interest for many years. Direct synthesis is difficult to achieve as, in terms of thermodynamics, the reaction of hydrogen with oxygen favours production of water. Systems for direct synthesis have been developed, most of which are based around finely dispersed metal catalysts. None of these has yet reached a point where they can be used for industrial-scale synthesis.
Hydrogen peroxide is most commonly available as a solution in water. For consumers, it is usually available from pharmacies at 3 and 6 wt% concentrations. The concentrations are sometimes described in terms of the volume of oxygen gas generated; one milliliter of a 20-volume solution generates twenty milliliters of oxygen gas when completely decomposed. For laboratory use, 30 wt% solutions are most common. Commercial grades from 70% to 98% are also available, but due to the potential of solutions of more than 68% hydrogen peroxide to be converted entirely to steam and oxygen (with the temperature of the steam increasing as the concentration increases above 68%) these grades are potentially far more hazardous, and require special care in dedicated storage areas. Buyers must typically allow inspection by commercial manufacturers.
In 1994, world production of H
The rate of decomposition increases with rising temperature, concentration and pH, with cool, dilute, acidic solutions showing the best stability. Decomposition is catalysed by various compounds, including most transition metals and their compounds (e.g. manganese dioxide, silver, and platinum). Certain metal ions, such as Fe2+
Hydrogen peroxide exhibits oxidizing and reducing properties, depending on pH.
In acidic solutions, H
In acidic solutions Fe2+
and sulfite (SO2−
In basic solution, hydrogen peroxide can reduce a variety of inorganic ions. When it acts as a reducing agent, oxygen gas is also produced. For example, hydrogen peroxide will reduce sodium hypochlorite and potassium permanganate, which is a convenient method for preparing oxygen in the laboratory.
Alkaline hydrogen peroxide 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 peroxide compounds
It also converts metal oxides into the corresponding peroxides. For example, upon treatment with hydrogen peroxide, chromic acid (CrO
A study published in Nature found that hydrogen peroxide plays a role in the immune system. Scientists found that hydrogen peroxide presence inside cells increased after tissues are damaged in zebrafish, which is thought to act as a signal to white blood cells to converge on the site and initiate the healing process. When the genes required to produce hydrogen peroxide were disabled, white blood cells did not accumulate at the site of damage. The experiments were conducted on fish; however, because fish are genetically similar to humans, the same process is speculated to occur in humans. The study in Nature suggested asthma sufferers have higher levels of hydrogen peroxide in their lungs than healthy people, which could explain why asthma sufferers have inappropriate levels of white blood cells in their lungs.
Hydrogen peroxide has important roles as a signalling molecule in the regulation of a wide variety of biological processes. The compound is a major factor implicated in the free-radical theory of aging, based on how readily hydrogen peroxide can decompose into a hydroxyl radical and how superoxide radical byproducts of cellular metabolism can react with ambient water to form hydrogen peroxide. These hydroxyl radicals in turn readily react with and damage vital cellular components, especially those of the mitochondria. At least one study has also tried to link hydrogen peroxide production to cancer. These studies have frequently been quoted in fraudulent treatment claims.
About 60% of the world's production of hydrogen peroxide is used for pulp- and paper-bleaching. The second major industrial application is the manufacture of sodium percarbonate and sodium perborate which are used as mild bleaches in laundry detergents.
It is used in the production of various organic peroxides with dibenzoyl peroxide being a high volume example. This is used in polymerisations, as a flour bleaching agent and as a treatment for acne. Peroxy acids, such as peracetic acid and meta-chloroperoxybenzoic acid are also typically produced using hydrogen peroxide.
Hydrogen peroxide is used in certain waste-water treatment processes to remove organic impurities. This is achieved by advanced oxidation processes, such as the Fenton reaction, which use it to generate highly reactive hydroxyl radicals (·OH). These are able to destroy organic contaminates which are ordinarily difficult to remove, such as aromatic or halogenated compounds. It can also oxidize sulfur based compounds present in the waste; which is beneficial as it generally reduces their odour.
Hydrogen peroxide can be used for the sterilization of various surfaces, including surgical tools and may be deployed as a vapour (VHP) for room sterilization. H2O2 demonstrates broad-spectrum efficacy against viruses, bacteria, yeasts, and bacterial spores. In general, greater activity is seen against gram-positive than gram-negative bacteria; however, the presence of catalase or other peroxidases in these organisms can increase tolerance in the presence of lower concentrations. Higher concentrations of H2O2 (10 to 30%) and longer contact times are required for sporicidal activity.
Hydrogen peroxide is seen as an environmentally safe alternative to chlorine-based bleaches, as it degrades to form oxygen and water and it is generally recognized as safe as an antimicrobial agent by the U.S. Food and Drug Administration (FDA).
Historically hydrogen peroxide was used for disinfecting wounds, partly because of its low cost and prompt availability compared to other antiseptics. It is now thought to slow healing and lead to scarring because it destroys newly formed skin cells. Only a very low concentration of H2O2 can induce healing, and only if not repeatedly applied. Surgical use can lead to gas embolism formation. Despite this it is still used for wound treatment in many developing countries.
Use in alternative medicine
Practitioners of alternative medicine have advocated the use of hydrogen peroxide for the treatment of various conditions, including emphysema, influenza, AIDS and in particular cancer. The practice calls for the daily consumption of hydrogen peroxide, either orally or by injection and is, in general, based around two precepts. First, that hydrogen peroxide is naturally produced by the body to combat infection; and second, that human pathogens (including cancer: See Warburg hypothesis) are anaerobic and cannot survive in oxygen-rich environments. The ingestion or injection of hydrogen peroxide is therefore believed to kill disease by mimicking the immune response in addition to increasing levels of oxygen within the body. This makes it similar to other oxygen-based therapies, such as ozone therapy and hyperbaric oxygen therapy.
Both the effectiveness and safety of hydrogen peroxide therapy is disputed by mainstream scientists. Hydrogen peroxide is produced by the immune system but in a carefully controlled manner. Cells called by phagocytes engulf pathogens and then use hydrogen peroxide to destroy them. The peroxide is toxic to both the cell and the pathogen and so is kept within a special compartment, called a phagosome. Free hydrogen peroxide will damage any tissue it encounters via oxidative stress; a process which also has been proposed as a cause of cancer. Claims that hydrogen peroxide therapy increase cellular levels of oxygen have not been supported. The quantities administered would be expected to provide very little additional oxygen compared to that available from normal respiration. It should also be noted that it is difficult to raise the level of oxygen around cancer cells within a tumour, as the blood supply tends to be poor, a situation known as tumor hypoxia.
Large oral doses of hydrogen peroxide at a 3% concentration may cause irritation and blistering to the mouth, throat, and abdomen as well as abdominal pain, vomiting, and diarrhea. Intravenous injection of hydrogen peroxide has been linked to several deaths.
The American Cancer Society states that "there is no scientific evidence that hydrogen peroxide is a safe, effective or useful cancer treatment." Furthermore, the therapy is not approved by the U.S. FDA.
For more details on this topic, see High-test peroxide.
High concentration H
As a bipropellant H
In the 1940s and 1950s, the Hellmuth Walter KG-conceived turbine used hydrogen peroxide for use in submarines while submerged; it was found to be too noisy and require too much maintenance compared to diesel-electric power systems. Some torpedoes used hydrogen peroxide as oxidizer or propellant. Operator error in the use of hydrogen peroxide torpedoes were named as possible causes for the sinkings of HMS Sidon and the Russian submarine Kursk. SAAB Underwater Systems is manufacturing the Torpedo 2000. This torpedo, used by the Swedish navy, is powered by a piston engine propelled by HTP as an oxidizer and kerosene as a fuel in a bipropellant system.
Hydrogen peroxide has been used for creating organic peroxide based explosives, such as acetone peroxide, for improvised explosive devices, including the 7 July 2005 London bombings. These explosives tend to degrade quickly and hence are not used as commercial or military explosives.
Hydrogen peroxide has various domestic uses, primarily as a cleaning and disinfecting agent.
Some horticulturalists and users of hydroponics advocate the use of weak hydrogen peroxide solution in watering solutions. Its spontaneous decomposition releases oxygen that enhances a plant's root development and helps to treat root rot (cellular root death due to lack of oxygen) and a variety of other pests.
Laboratory tests conducted by fish culturists in recent years have demonstrated that common household hydrogen peroxide can be used safely to provide oxygen for small fish. The hydrogen peroxide releases oxygen by decomposition when it is exposed to catalysts such as manganese dioxide.
Regulations vary, but low concentrations, such as 6%, are widely available and legal to buy for medical use. Most over-the-counter peroxide solutions are not suitable for ingestion. Higher concentrations may be considered hazardous and are typically accompanied by a Material Safety Data Sheet (MSDS). In high concentrations, hydrogen peroxide is an aggressive oxidizer and will corrode many materials, including human skin. In the presence of a reducing agent, high concentrations of H
High-concentration hydrogen peroxide streams, typically above 40%, should be considered hazardous due to concentrated hydrogen peroxide's meeting the definition of a DOT oxidizer according to U.S. regulations, if released into the environment. The EPA Reportable Quantity (RQ) for D001 hazardous wastes is 100 pounds (45 kg), or approximately 10 US gallons (38 L), of concentrated hydrogen peroxide.
Hydrogen peroxide should be stored in a cool, dry, well-ventilated area and away from any flammable or combustible substances. It should be stored in a container composed of non-reactive materials such as stainless steel or glass (other materials including some plastics and aluminium alloys may also be suitable). Because it breaks down quickly when exposed to light, it should be stored in an opaque container, and pharmaceutical formulations typically come in brown bottles that filter out light.
Hydrogen peroxide, either in pure or diluted form, can pose several risks, the main one being that it forms explosive mixtures upon contact with organic compounds. Highly concentrated hydrogen peroxide itself is unstable, and can then cause a boiling liquid expanding vapour explosion (BLEVE) of the remaining liquid. Distillation of hydrogen peroxide at normal pressures is thus highly dangerous. It is also corrosive, especially when concentrated, but even domestic-strength solutions can cause irritation to the eyes, mucous membranes and skin. Swallowing hydrogen peroxide solutions is particularly dangerous, as decomposition in the stomach releases large quantities of gas (10 times the volume of a 3% solution), leading to internal bleeding. Inhaling over 10% can cause severe pulmonary irritation.
With a significant vapour pressure (1.2 kPa at 50 °C[CRC Handbook of Chemistry and Physics, 76th Ed, 1995–1996]), hydrogen peroxide vapour is potentially hazardous. According to U.S. NIOSH, the Immediately Dangerous to Life and Health (IDLH) limit is only 75 ppm. The U.S. Occupational Safety and Health Administration (OSHA) has established a permissible exposure limit of 1.0 ppm calculated as an eight-hour time weighted average (29 CFR 1910.1000, Table Z-1) and hydrogen peroxide has also been classified by the American Conference of Governmental Industrial Hygienists (ACGIH) as a "known animal carcinogen, with unknown relevance on humans." For workplaces where there is a risk of exposure to the hazardous concentrations of the vapours, continuous monitors for hydrogen peroxide should be used. Information on the hazards of hydrogen peroxide is available from OSHA and from the ATSDR.