Western Canada distribution

LFT Group is happy to announce that it has appointed Nationwide Natural Foods to distribute the Live for Tomorrow range of cleaning products in BC, Alberta, and Saskatchewan.

Nationwide Natural Foods

604-952-250-542-5540

sales@nationwidenaturalfoods.com

BC interior & Alberta distribution

LFT Group is happy to announce that it has appointed Summit Speciality Foods to distribute the Live for Tomorrow range of cleaning products in BC interior and Alberta.

Summit Speciality Foods
Dave Baker
250-542-5540

Vancouver Island & Gulf Islands distribution

LFT Group is happy to announce that it has appointed Kev's Wholesale as our distibutor for Vancouver Island and the Gulf Islands.

Kev’s Wholesale Market
Kev Carter
250-217-8565
kevswholesale@gmail.com

Live for Tomorrow exhibiting at CHFA

Live for Tomorrow will exhibiting at the CHFA show.

16 - 17 April 2011 - 10.00am - 5.00pm

Canadian Health Food Association (CHFA)

Booth 238

1055 Canada Place
Vancouver, BC, V6C 0C3

Potassium sorbate

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Jump to: navigation, search

Potassium sorbate[1][2]
IUPAC name[hide] Potassium (2E,4E)-hexa-2,4-dienoate
Other names[hide] E202 Sorbistat-K Sorbistat potassium
Identifiers
CAS number	24634-61-5 Y
PubChem	5282505
ChemSpider	4445644 Y
KEGG	D02411 N
SMILES [show] [K+].[O-]C(=O)\C=C\C=C\C
InChI [show] InChI=1S/C6H8O2.K/c1-2-3-4-5-6(7)8;/h2-5H,1H3,(H,7,8);/q;+1/p-1/b3-2+,5-4+; Y Key: CHHHXKFHOYLYRE-STWYSWDKSA-M Y InChI=1/C6H8O2.K/c1-2-3-4-5-6(7)8;/h2-5H,1H3,(H,7,8);/q;+1/p-1/b3-2+,5-4+; Key: CHHHXKFHOYLYRE-ZCSOUONQBI
Properties
Molecular formula	C6H7KO2
Molar mass	150.22 g/mol
Density	1.363 g/cm3
Melting point	270 °C (decomposition)
Solubility in water	58.2% at 20 °C
Solubility	soluble in ethanol, propylene glycol slightly soluble in acetone very slightly soluble in chloroform, corn oil, ether insoluble in benzene
Hazards
LD50	4920 mg/kg (rat, oral)
N(what is this?)  (verify) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Potassium sorbate is the potassium salt of sorbic acid. Its primary use is as a food preservative (E number 202).^[3] Potassium sorbate is effective in a variety of applications including food, wine, and personal care products.

Contents [hide] 1 Properties 2 Uses 3 Toxicology 4 See also 5 References

[edit] Properties

Potassium sorbate is produced by neutralizing potassium hydroxide with sorbic acid, an unsaturated carboxylic acid that occurs naturally in some berries. The colourless salt is very soluble in water (58.2% at 20 °C).

[edit] Uses

Potassium sorbate is used to inhibit molds and yeasts in many foods, such as cheese, wine, yogurt, dried meats, apple cider, soft drinks and fruit drinks, and baked goods.^[4] It can also be found in the ingredients list of many dried fruit products. In addition, herbal dietary supplement products generally contain potassium sorbate, which acts to prevent mold and microbes and to increase shelf life, and is used in quantities at which there are no known adverse health effects, over short periods of time^[5]. Labeling of this preservative on ingredient statements reads as "potassium sorbate". Also, it is used in many personal care products to inhibit the development of microorganisms for shelf stability. Some manufacturers are using this preservative as a replacement for parabens.

Also known as "wine stabilizer", potassium sorbate produces sorbic acid when added to wine. It serves two purposes. When active fermentation has ceased and the wine is racked for the final time after clearing, potassium sorbate will render any surviving yeast incapable of multiplying. Yeast living at that moment can continue fermenting any residual sugar into CO₂ and alcohol, but when they die no new yeast will be present to cause future fermentation. When a wine is sweetened before bottling, potassium sorbate is used to prevent refermentation when used in conjunction with potassium metabisulfite. It is primarily used with sweet wines, sparkling wines, and some hard ciders but may be added to table wines which exhibit difficulty in maintaining clarity after fining.

Some molds (notably some Trichoderma and Penicillium strains) and yeasts are able to detoxify sorbates by decarboxylation, producing 1,3-pentadiene. The pentadiene manifests as a typical odor of kerosene or petroleum.^[6]

[edit] Toxicology

Potassium sorbate is considered to be safe because of its long term safety record and non-toxic profile. Potassium sorbate is non-irritating and non-sensitizing. Allergic reactions are rare^{[citation needed]}, and it is well tolerated when administered internally.^[7]

Potassium sorbate exhibits low toxicity with LD₅₀ (rat) of 4.92 g/kg,^[8] similar to that of table salt.^[9] Typical usage rates of potassium sorbate are 0.025% to 0.1% (see sorbic acid), which in a 100 g serving yields intake of 25 mg to 100 mg. Acceptable daily intakes for human is 12.5 mg/kg, or 875 mg daily for a normal adult (70 kg), according to FAO/World Health Organization Expert Committee on Food Additives.^[5]

[edit] See also

• Sodium benzoate

[edit] References

1. ^ Merck Index, 11th Edition, 7661.
2. ^ Potassium sorbate at Sigma-Aldrich
3. ^ Nordic Food Additive Database Nordic Working Group on Food Toxicology and Risk Assessment
4. ^ Erich Lück, Martin Jager and Nico Raczek "Sorbic Acid" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2000.doi:10.1002/14356007.a24_507
5. ^ ^{a b} http://www.inchem.org/documents/jecfa/jecmono/40abcj15.htm
6. ^ The Soft Drinks Companion - A technical handbook for the beverage industry, Chapter 10
7. ^ Tulamait, Aiman; Laghi, F; Mikrut, K; Carey, RB; Budinger, GR (2005). "Potassium sorbate reduces gastric colonization in patients receiving mechanical ventilization". J Crit Care 20 (3): 281–287. doi:10.1016/j.jcrc.2005.03.002. PMID 16253799. 
8. ^ MSDS for potassium sorbate
9. ^ MSDS for sodium chloride

Retrieved from "http://en.wikipedia.org/wiki/Potassium_sorbate"

Categories: Potassium compounds | Preservatives | Sorbates



Now available at Gecko Green Living

The Live for Tomorrow range of cleaning products are now available at:

 

Gecko Green Living
103 McPhillips Avenue,
Salt Spring Island, BC V8K 2K3

Tel: 1 250-537-1151





Vegan certified

Live for Tomorrow cleaning products are now Vegan approved and have been certified by Vegan Action.

VEGAN CERTIFICATION

Vegan Action administers the Certified Vegan Logo, an easy-to-recognize symbol applied to foods, clothing, cosmetics and other items that contain no animal products and are not tested on animals.

What is the Certified Vegan Logo?
The Logo is a registered trademark, similar in nature to the "kosher" mark, for products that do not contain animal products and that have not been tested on animals. The Logo is easily visible to consumers interested in vegan products and helps vegans to shop without constantly consulting ingredient lists; it helps companies recognize a growing vegan market; and it helps bring the word "vegan"—and the lifestyle it represents—into the mainstream. (Please keep in mind, however, that the logo is not yet on every vegan product.)

The Logo is administered by The Vegan Awareness Foundation (official name of Vegan Action), a 501(c)3 nonprofit organization dedicated to educating the public about veganism and to assist vegan-friendly businesses.



Wholesale enquiry

Are you interested in selling Live for Tomorrow products?

We welcome enquiries. Please e-mail us  with your request and provide us with some basic information on your business.

Direct sales:

We do sell direct, where this does not conflict with distribution agreements.

Sales through Wholesalers and Distributors:

Vancouver Island: Distributor - please e-mail us and we will forward your contact information;

West Coast & Alberta: seeking a distributor. Please e-mail us for direct sales;

East Coast: seeking a distributor. Please e-mail us for direct sales;

USA: seeking a distributor. Please e-mail us for direct sales; and

UK & Europe: seeking a distributor. Please e-mail us for direct sales.

Bulk & custom sales:

We sell bulk product and are able to ofer some customization to products. Please e-mail us for more information.



Hydrogen peroxide

From Wikipedia, the free encyclopedia

Jump to: navigation, search

Hydrogen peroxide
IUPAC name[hide] dihydrogen dioxide
Other names[hide] Dioxidane
Identifiers
CAS number	7722-84-1 Y
PubChem	784
ChemSpider	763
EC number	231-765-0
UN number	2015 (>60% soln.) 2014 (20–60% soln.) 2984 (8–20% soln.)
IUPHAR ligand	2448
RTECS number	MX0900000 (>90% soln.) MX0887000 (>30% soln.)
Properties
Molecular formula	H2O2
Molar mass	34.0147 g/mol
Appearance	Very light blue color; colorless in solution
Density	1.463 g/cm3
Melting point	-0.43 °C, 273 K, 31 °F
Boiling point	150.2 °C, 423 K, 302 °F
Solubility in water	Miscible
Solubility	soluble in ether
Acidity (pKa)	11.62 [1]
Refractive index (nD)	1.34
Viscosity	1.245 cP (20 °C)
Dipole moment	2.26 D
Thermochemistry
Std enthalpy of formation ΔfHo298	-4.007 kJ/g
Specific heat capacity, C	1.267 J/g K (gas) 2.619 J/g K (liquid)
Hazards
MSDS	ICSC 0164 (>60% soln.)
EU Index	008-003-00-9
EU classification	Oxidant (O) Corrosive (C) Harmful (Xn)
R-phrases	R5, R8, R20/22, R35
S-phrases	(S1/2), S17, S26, S28, S36/37/39, S45
NFPA 704	0 3 2 OX
Flash point	Non-flammable
LD50	1518 mg/kg
Related compounds
Related compounds	Water Ozone Hydrazine Hydrogen disulfide
Y (what is this?)  (verify) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Hydrogen peroxide (H₂O₂) is a very pale blue liquid, slightly more viscous than water, that appears colorless in dilute solution. It has strong oxidizing properties, and is a powerful bleaching agent. It is used as a disinfectant, antiseptic, oxidizer, and in rocketry as a propellant.^[2] The oxidizing capacity of hydrogen peroxide is so strong that it is considered a highly reactive oxygen species.

Hydrogen peroxide is naturally produced in organisms as a by-product of oxidative metabolism. Nearly all living things (specifically, all obligate and facultative aerobes) possess enzymes known as peroxidases, which harmlessly and catalytically decompose low concentrations of hydrogen peroxide to water and oxygen.

Contents [hide] 1 Structure and properties 1.1 Comparison with analogues 1.2 Physical properties of hydrogen peroxide solutions 2 History 3 Manufacture 3.1 New developments 3.2 Availability 4 Reactions 4.1 Decomposition 4.2 Redox reactions 4.3 Formation of peroxide compounds 4.4 Alkalinity 5 Uses 5.1 Industrial applications 5.2 Chemical applications 5.3 Biological function 5.4 Domestic uses 5.5 Use as propellant 5.6 Therapeutic use 6 Safety 6.1 Historical incidents 7 References 7.1 Notes 7.2 Bibliography 8 External links

[edit] Structure and properties

H₂O₂ adopts a nonplanar structure of C2 symmetry. Although chiral, the molecule undergoes rapid racemization. The flat shape of the anti conformer would minimize steric repulsions, the 90° torsion angle of the syn conformer would optimize mixing between the filled p-type orbital of the oxygen (one of the lone pairs) and the LUMO of the vicinal O-H bond.^[3] The observed anticlinal "skewed" shape is a compromise between the two conformers.

Despite the fact that the O-O bond is a single bond, the molecule has a high barrier to complete rotation of 29.45 kJ/mol (compared with 12.5 kJ/mol for the rotational barrier of ethane). The increased barrier is attributed to repulsion between one lone pair and other lone pairs. The bond angles are affected by hydrogen bonding, which is relevant to the structural difference between gaseous and crystalline forms; indeed a wide range of values is seen in crystals containing molecular H₂O₂.

[edit] Comparison with analogues

Analogues of hydrogen peroxide include the chemically identical deuterium peroxide and hydrogen disulfide.^[4] Hydrogen disulfide has a boiling point of only 70.7°C despite having a higher molecular weight, indicating that hydrogen bonding increases the boiling point of hydrogen peroxide.

[edit] Physical properties of hydrogen peroxide solutions

The properties of aqueous solutions of hydrogen peroxide differ from those of the neat material, reflecting the effects of hydrogen bonding between water and hydrogen peroxide molecules. Hydrogen peroxide and water form a eutectic mixture, exhibiting freezing-point depression. Whereas pure water melts and freezes at approximately 273K, and pure hydrogen peroxide just 0.4K below that, a 50% (by volume) solution melts and freezes at 221 K.^[5]

[edit] History

Hydrogen peroxide was first isolated in 1818 by Louis Jacques Thénard by reacting barium peroxide with nitric acid.^[6] An improved version of this process used hydrochloric acid, followed by 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.^[7] Modern production methods are discussed below.

For a long time, pure hydrogen peroxide was believed to be unstable, because attempts to separate the hydrogen peroxide from the water, which is present during synthesis, failed. This instability was however due to traces of impurities (transition metals salts) that catalyze the decomposition of the hydrogen peroxide. One hundred percent pure hydrogen peroxide was first obtained through vacuum distillation by Richard Wolffenstein in 1894.^[8] At the end of 19th century, Petre Melikishvili and his pupil L. Pizarjevski showed that of the many proposed formulas of hydrogen peroxide, the correct one was H-O-O-H.

The use of H₂O₂ sterilization in biological safety cabinets and barrier isolators is a popular alternative to ethylene oxide (EtO) as a safer, more efficient decontamination method. H₂O₂ has long been widely used in the pharmaceutical industry. In aerospace research, H₂O₂ is used to sterilize satellites.

The U.S. FDA has recently granted 510(k) clearance to use H₂O₂ in individual medical device manufacturing applications. EtO criteria outlined in ANSI/AAMI/ISO 14937 may be used as a validation guideline. Sanyo was the first manufacturer to use the H₂O₂ process in situ in a cell culture incubator, which is a faster and more efficient cell culture sterilization process.

[edit] Manufacture

Formerly, hydrogen peroxide was prepared by the electrolysis of an aqueous solution of sulfuric acid or acidic ammonium bisulfate (NH₄HSO₄), followed by hydrolysis of the peroxodisulfate ((SO₄)₂)²⁻ that is formed. Today, hydrogen peroxide is manufactured almost exclusively by the autoxidation of a 2-alkyl anthrahydroquinone (or 2-alkyl-9,10-dihydroxyanthracene) to the corresponding 2-alkyl anthraquinone. Major producers commonly use either the 2-ethyl or the 2-amyl derivative. The cyclic reaction depicted below shows the 2-ethyl derivative, where 2-ethyl-9,10-dihydroxyanthracene (C₁₆H₁₄O₂) is oxidized to the corresponding 2-ethylanthraquinone (C₁₆H₁₂O₂) and hydrogen peroxide. Most commercial processes achieve this by bubbling compressed air through a solution of the 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.^[9][10]

This process is known as the Riedl-Pfleiderer process,^[10] having been first discovered by them in 1936. The overall equation for the process is deceptively simple:^[9]

H₂ + O₂ → H₂O₂

The economics of the process depend heavily on effective recycling of the quinone (which is expensive) and extraction solvents, and of the hydrogenation catalyst.

In 1994, world production of H₂O₂ was around 1.9 million tonnes and grew to 2.2 million in 2006,^[11] most of which was at a concentration of 70% or less^{[citation needed]}. In that year bulk 30% H₂O₂ sold for around US $0.54 per kg, equivalent to US $1.50 per kg (US $0.68 per lb) on a "100% basis^{[citation needed]}".

[edit] New developments

A new, so-called "high-productivity/high-yield" process, based on an optimized distribution of isomers of 2-amyl anthraquinone, has been developed by Solvay. In July 2008, this process allowed the construction of a "mega-scale" single-train plant in Zandvliet (Belgium). The plant has an annual production capacity more than twice that of the world's next-largest single-train plant. An even-larger plant is scheduled to come onstream at Map Ta Phut (Thailand) in 2011. It is likely that this will lead to a reduction in the cost of production due to economies of scale.^[12]

A process to produce hydrogen peroxide directly from the elements has been of interest for many years. The problem with the direct synthesis process is that, in terms of thermodynamics, the reaction of hydrogen with oxygen favors production of water. It had been recognized for some time that a finely dispersed catalyst is beneficial in promoting selectivity to hydrogen peroxide, but, while selectivity was improved, it was still not sufficiently high to permit commercial development of the process. However, an apparent breakthrough was made in the early 2000s by researchers at Headwaters Technology. The breakthrough revolves around development of a minute (nanometer-size) phase-controlled noble metal crystal particles on carbon support. This advance led, in a joint venture with Evonik Industries, to the construction of a pilot plant in Germany in late 2005. It is claimed that there are reductions in investment cost because the process is simpler and involves less equipment; however, the process is also more corrosive and unproven. This process results in low concentrations of hydrogen peroxide (about 5–10 wt% versus about 40 wt% through the anthraquione process).^[12]

In 2009, another catalyst development was announced by workers at Cardiff University.^[13] This development also relates to the direct synthesis, but, in this case, using gold–palladium nanoparticles. Under normal circumstances, the direct synthesis must be carried out in an acid medium to prevent immediate decomposition of the hydrogen peroxide once it is formed. Whereas hydrogen peroxide tends to decompose on its own (which is why, even after production, it is often necessary to add stabilisers to the commercial product when it is to be transported or stored for long periods), the nature of the catalyst can cause this decomposition to accelerate rapidly. It is claimed that the use of this gold-palladium catalyst reduces this decomposition and, as a consequence, little to no acid is required. The process is in a very early stage of development and currently results in very low concentrations of hydrogen peroxide being formed (less than about 1–2 wt%). Nonetheless, it is envisaged by the inventors that the process will lead to an inexpensive, efficient, and environmentally friendly process.^{[12][13][14][15]}

A novel electrochemical process for the production of alkaline hydrogen peroxide has been developed by Dow. The process employs a monopolar cell to achieve an electrolytic reduction of oxygen in a dilute sodium hydroxide solution.^[12]

[edit] Availability

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 (see decomposition); 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 >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 submit to inspection by the small number of commercial manufacturers.

[edit] Reactions

[edit] Decomposition

Hydrogen peroxide decomposes (disproportionates) exothermically into water and oxygen gas spontaneously:

2 H₂O₂ → 2 H₂O + O₂

This process is thermodynamically favorable. It has a ΔH^o of −98.2 kJ·mol⁻¹ and a ΔG^o of −119.2 kJ·mol⁻¹ and a ΔS of 70.5 J·mol⁻¹·K⁻¹. The rate of decomposition is dependent on the temperature and concentration of the peroxide, as well as the pH and the presence of impurities and stabilizers. Hydrogen peroxide is incompatible with many substances that catalyse its decomposition, including most of the transition metals and their compounds. Common catalysts include manganese dioxide and silver. The same reaction is catalysed by the enzyme catalase, found in the liver, whose main function in the body is the removal of toxic byproducts of metabolism and the reduction of oxidative stress. The decomposition occurs more rapidly in alkali, so acid is often added as a stabilizer.

The liberation of oxygen and energy in the decomposition has dangerous side-effects. Spilling high concentrations of hydrogen peroxide on a flammable substance can cause an immediate fire, which is further fueled by the oxygen released by the decomposing hydrogen peroxide. High test peroxide, or HTP (also called high-strength peroxide) must be stored in a suitable,^{[citation needed]} vented container to prevent the buildup of oxygen gas, which would otherwise lead to the eventual rupture of the container.

In the presence of certain catalysts, such as Fe²⁺ or Ti³⁺, the decomposition may take a different path, with free radicals such as HO· (hydroxyl) and HOO· being formed. A combination of H₂O₂ and Fe²⁺ is known as Fenton's reagent.

A common concentration for hydrogen peroxide is 20-volume, which means that, when 1 volume of hydrogen peroxide is decomposed, it produces 20 volumes of oxygen. A 20-volume concentration of hydrogen peroxide is equivalent to 1.667 mol/dm³ (Molar solution) or about 6%.

[edit] Redox reactions

This section does not cite any references or sources. Please help improve this article by adding citations to reliable sources. Unsourced material may be challenged and removed. (July 2010)

In acidic solution, H₂O₂ is one of the most powerful oxidizers known—stronger than chlorine, chlorine dioxide, and potassium permanganate. Also, through catalysis, H₂O₂ can be converted into hydroxyl radicals (.OH), which are highly reactive.

Oxidant/Reduced product	Oxidation potential, V
Fluorine/Hydrogen fluoride	3.0
Ozone/Oxygen	2.1
Hydrogen peroxide/Water	1.8
Potassium permanganate/Manganese dioxide	1.7
Chlorine dioxide/HClO	1.5
Chlorine/Chloride	1.4

In aqueous solution, hydrogen peroxide can oxidize or reduce a variety of inorganic ions. When it acts as a reducing agent, oxygen gas is also produced.

In acidic solutions Fe²⁺ is oxidized to Fe³⁺ (hydrogen peroxide acting as an oxidizing agent),

2 Fe²⁺(aq) + H₂O₂ + 2 H⁺(aq) → 2 Fe³⁺(aq) + 2H₂O(l)

and sulfite (SO₃²⁻) is oxidized to sulfate (SO₄²⁻). However, potassium permanganate is reduced to Mn²⁺ by acidic H₂O₂. Under alkaline conditions, however, some of these reactions reverse; for example, Mn²⁺ is oxidized to Mn⁴⁺ (as MnO₂).

Another example of hydrogen peroxide's acting as a reducing agent is the reaction with sodium hypochlorite, which is a convenient method for preparing oxygen in the laboratory.

NaOCl + H₂O₂ → O₂ + NaCl + H₂O

Hydrogen peroxide is frequently used as an oxidizing agent in organic chemistry. One application is for the oxidation of thioethers to sulfoxides.^{[citation needed]} For example, methyl phenyl sulfide was oxidized to methyl phenyl sulfoxide in 99% yield in methanol in 18 hours (or 20 minutes using a TiCl₃ catalyst):^{[citation needed]}

Ph-S-CH₃ + H₂O₂ → Ph-S(O)-CH₃ + H₂O

Alkaline hydrogen peroxide is used for epoxidation of electron-deficient alkenes such as acrylic acids, and also for oxidation of alkylboranes to alcohols, the second step of hydroboration-oxidation.

[edit] Formation of peroxide compounds

Hydrogen peroxide is a weak acid, and it can form hydroperoxide or peroxide salts or derivatives of many metals.

For example, on addition to an aqueous solution of chromic acid (CrO₃) or acidic solutions of dichromate salts, it will form an unstable blue peroxide CrO(O₂)₂. In aqueous solution it rapidly decomposes to form oxygen gas and chromium salts.

It can also produce peroxoanions by reaction with anions; for example, reaction with borax leads to sodium perborate, a bleach used in laundry detergents:

Na₂B₄O₇ + 4 H₂O₂ + 2 NaOH → 2 Na₂B₂O₄(OH)₄ + H₂O

H₂O₂ converts carboxylic acids (RCOOH) into peroxy acids (RCOOOH), which are themselves used as oxidizing agents. Hydrogen peroxide reacts with acetone to form acetone peroxide, and it interacts with ozone to form hydrogen trioxide, also known as trioxidane. Reaction with urea produces carbamide peroxide, used for whitening teeth. An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for H₂O₂ in some reactions.

[edit] Alkalinity

Hydrogen peroxide is a much weaker base than water, but it can still form adducts with very strong acids. The superacid HF/SbF₅ forms unstable compounds containing the [H₃O₂]⁺ ion.

[edit] Uses

[edit] Industrial applications

ISO tank container for hydrogen peroxide transportation.

About 50% of the world's production of hydrogen peroxide in 1994 was used for pulp- and paper-bleaching.^[11] Other bleaching applications are becoming more important as hydrogen peroxide is seen as an environmentally benign alternative to chlorine-based bleaches.

Other major industrial applications for hydrogen peroxide include the manufacture of sodium percarbonate and sodium perborate, used as mild bleaches in laundry detergents. It is used in the production of certain organic peroxides such as dibenzoyl peroxide, used in polymerisations and other chemical processes. Hydrogen peroxide is also used in the production of epoxides such as propylene oxide. Reaction with carboxylic acids produces a corresponding peroxy acid. Peracetic acid and meta-chloroperoxybenzoic acid (commonly abbreviated mCPBA) are prepared from acetic acid and meta-chlorobenzoic acid, respectively. The latter is commonly reacted with alkenes to give the corresponding epoxide.

In the PCB manufacturing process, hydrogen peroxide mixed with sulfuric acid was used as the microetch chemical for copper surface roughening preparation.

A combination of a powdered precious metal-based catalyst, hydrogen peroxide, methanol and water can produce superheated steam in one to two seconds, releasing only CO₂ and high-temperature steam for a variety of purposes.^[16]

Recently, there has been increased use of vaporized hydrogen peroxide in the validation and bio-decontamination of half-suit and glove-port isolators in pharmaceutical production.

Nuclear pressurized water reactors (PWRs) use hydrogen peroxide during the plant shutdown to force the oxidation and dissolution of activated corrosion products deposited on the fuel. The corrosion products are then removed with the cleanup systems before the reactor is disassembled.

Hydrogen peroxide is also used in the oil and gas exploration industry to oxidize rock matrix in preparation for micro-fossil analysis.

[edit] Chemical applications

A method of producing propylene oxide from hydrogen peroxide has been developed. The process is claimed to be environmentally friendly, since the only significant byproduct is water. It is also claimed the process has significantly lower investment and operating costs. Two of these "HPPO" (hydrogen peroxide to propylene oxide) plants came onstream in 2008: One of them located in Belgium is a Solvay, Dow-BASF joint venture, and the other in Korea is a EvonikHeadwaters, SK Chemicals joint venture. A caprolactam application for hydrogen peroxide has been commercialized. Potential routes to phenol and epichlorohydrin utilizing hydrogen peroxide have been postulated.^[12]

[edit] Biological function

Hydrogen peroxide is also one of the two chief chemicals in the defense system of the bombardier beetle, reacting with hydroquinone to discourage predators.

A study published in Nature found that hydrogen peroxide plays a role in the immune system. Scientists found that hydrogen peroxide is released after tissues are damaged in zebra fish, 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 be explained by asthma sufferers having inappropriate levels of white blood cells in their lungs.^[17][18]

[edit] Domestic uses

Skin immediately after exposure to 30% H₂O₂

• Diluted H₂O₂ (between 3% and 12%) is used to bleach human hair when mixed with ammonium hydroxide, hence the phrase "peroxide blonde".
• It is absorbed by skin upon contact and creates a local skin capillary embolism that appears as a temporary whitening of the skin.
• It is used to whiten bones that are to be put on display.
• 3% H₂O₂ is used medically for cleaning wounds, removing dead tissue, and as an oral debriding agent. Peroxide stops slow (small vessel) wound bleeding/oozing, as well. However, recent studies have suggested that hydrogen peroxide impedes scarless healing as it destroys newly formed skin cells.^[19] Most over-the-counter peroxide solutions are not suitable for ingestion.
• If a dog has swallowed a harmful substance (e.g., rat poison), small amounts of hydrogen peroxide can be given to induce vomiting.^[20]
• 3% H₂O₂ is effective at treating fresh (red) blood-stains in clothing and on other items. It must be applied to clothing before blood stains can be accidentally "set" with heated water. Cold water and soap are then used to remove the peroxide treated blood.
• The United States Food and Drug Administration (FDA) has classified hydrogen peroxide as a Low Regulatory Priority (LRP) drug for use in controlling fungus on fish and fish eggs. (See ectoparasite.)
• 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.^[21][22] There is some peer-reviewed academic research to back up some of the claims.^[23]
• 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.^[24][25] Hydrogen peroxide releases oxygen by decomposition when it is exposed to catalysts such as manganese dioxide.
• Hydrogen peroxide is a strong oxidizer effective in controlling sulfide and organic-related odors in wastewater collection and treatment systems. It is typically applied to a wastewater system where there is a retention time of 30 minutes to 5 hours before hydrogen sulfide is released. Hydrogen peroxide oxidizes the hydrogen sulfide and promotes bio-oxidation of organic odors. Hydrogen peroxide decomposes to oxygen and water, adding dissolved oxygen to the system, thereby negating some Biochemical Oxygen Demand (BOD).
• Mixed with baking soda and a small amount of hand soap, hydrogen peroxide is effective at removing skunk odor.^[26]
• Hydrogen peroxide is used with phenyl oxalate ester and an appropriate dye in glow sticks as an oxidizing agent. It reacts with the ester to form an unstable CO₂ dimer, which excites the dye to an excited state; the dye emits a photon (light) when it spontaneously relaxes back to the ground state.

[edit] Use as propellant

For more details on this topic, see High test peroxide.

Rocket Belt hydrogen peroxide propulsion system (see Jet packs).

High concentration H₂O₂ is referred to as HTP or High test peroxide. It can be used either as a monopropellant (not mixed with fuel) or as the oxidizer component of a bipropellant rocket. Use as a monopropellant takes advantage of the decomposition of 70–98+% concentration hydrogen peroxide into steam and oxygen. The propellant is pumped into a reaction chamber where a catalyst, usually a silver or platinum screen, triggers decomposition, producing steam at over 600 °C, which is expelled through a nozzle, generating thrust. H₂O₂ monopropellant produces a maximum specific impulse (I_sp) of 161 s (1.6 kN·s/kg), which makes it a low-performance monopropellant. Peroxide generates much less thrust than hydrazine, but is not toxic. The Bell Rocket Belt used hydrogen peroxide monopropellant.

As a bipropellant H₂O₂ is decomposed to burn a fuel as an oxidizer. Specific impulses as high as 350 s (3.5 kN·s/kg) can be achieved, depending on the fuel. Peroxide used as an oxidizer gives a somewhat lower I_sp than liquid oxygen, but is dense, storable, noncryogenic and can be more easily used to drive gas turbines to give high pressures using an efficient closed cycle. It can also be used for regenerative cooling of rocket engines. Peroxide was used very successfully as an oxidizer in World-War-II German rockets (e.g. T-Stoff, containing oxyquinoline stabilizer, for the Me-163), and for the low-cost British Black Knight and Black Arrow launchers.

In the 1940s and 1950s, the Walter 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, but this was dangerous and has been discontinued by most navies. Hydrogen peroxide leaks were blamed for the sinkings of HMS Sidon and the Russian submarine Kursk. It was discovered, for example, by the Japanese Navy in torpedo trials, that the concentration of H₂O₂ in right-angle bends in HTP pipework can often lead to explosions in submarines and torpedoes. 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.^[27]

While rarely used now as a monopropellant for large engines, small hydrogen peroxide attitude control thrusters^{[disambiguation needed]} are still in use on some satellites. They are easy to throttle, and safer to fuel and handle before launch than hydrazine thrusters. However, hydrazine is more often used in spacecraft because of its higher specific impulse and lower rate of decomposition.

[edit] Therapeutic use

Hydrogen peroxide is generally recognized as safe (GRAS) as an antimicrobial agent, an oxidizing agent and for other purposes by the FDA.^[28]

Hydrogen peroxide has been used as an antiseptic and anti-bacterial agent for many years due to its oxidizing effect. While its use has decreased in recent years with the popularity of readily available over the counter products, it is still used by many hospitals, doctors and dentists.

• Like many oxidative antiseptics, hydrogen peroxide causes mild damage to tissue in open wounds, but it also is effective at rapidly stopping capillary bleeding (slow blood oozing from small vessels in abrasions), and is sometimes used sparingly for this purpose, as well as cleaning.
• Hydrogen peroxide can be used as a toothpaste when mixed with correct quantities of baking soda and salt.^[29]
• Hydrogen peroxide and benzoyl peroxide are sometimes used to treat acne.^[30]
• Hydrogen peroxide is used as an emetic in veterinary practice.^[31]

Alternative uses

• The American Cancer Society states that "there is no scientific evidence that hydrogen peroxide is a safe, effective or useful cancer treatment", and advises cancer patients to "remain in the care of qualified doctors who use proven methods of treatment and approved clinical trials of promising new treatments." ^[32]
• Another controversial alternative medical procedure is inhalation of hydrogen peroxide at a concentration of about 1%. Intravenous usage of hydrogen peroxide has been linked to several deaths.^[33][34]
• Advocates of internal use say that they never use "High Concentration" but only diluted. The confusion arises because of the use of High Purity "Food Grade" Hydrogen peroxide which is commonly only sold in bulk concentrations of 35% to industry and is used because it is approved for internal use in food preparation by the FDA. But it is NEVER used at the concentration of 35% or even 1% to 3% internally. Hydrogen Peroxide when used internally is diluted down to strengths of less of than 0.1% (i.e. adding 2-20 drops per 6-8 ounce (250ml) of distilled or spring water).^[35]
• See also Liquid Oxygen (supplement)

[edit] Safety

Regulations vary, but low concentrations, such as 3%, are widely available and legal to buy for medical use. 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 H₂O₂ will react violently.

High-concentration hydrogen peroxide streams, typically above 40%, should be considered a D001 hazardous waste, due to concentrated hydrogen peroxide's meeting the definition of a DOT oxidizer, if released into the environment. The EPA Reportable Quantity (RQ) for D001 hazardous wastes is 100 pounds, or approximately ten gallons, of concentrated hydrogen peroxide.

Hydrogen peroxide should be stored in a cool, dry, well-ventilated area and away from any flammable or combustible substances.^[36] 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).^[37] 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.^[38]

Hydrogen peroxide, either in pure or diluted form, can pose several risks:

• Explosive vapors. Above roughly 70% concentrations, hydrogen peroxide can give off vapor that can detonate above 70 °C (158 °F) at normal atmospheric pressure.^{[citation needed]} This can then cause a boiling liquid expanding vapor explosion (BLEVE) of the remaining liquid. Distillation of hydrogen peroxide at normal pressures is thus highly dangerous.
• Hazardous reactions. Hydrogen peroxide vapors can form sensitive contact explosives with hydrocarbons such as greases. Hazardous reactions ranging from ignition to explosion have been reported with alcohols, ketones, carboxylic acids (particularly acetic acid), amines and phosphorus.^{[citation needed]}
• Spontaneous ignition. Concentrated hydrogen peroxide, if spilled on clothing (or other flammable materials), will preferentially evaporate water until the concentration reaches sufficient strength, at which point the material may spontaneously ignite.^[39][40]
• Corrosive. Concentrated hydrogen peroxide (>50%) is corrosive, and even domestic-strength solutions can cause irritation to the eyes, mucous membranes and skin.^[41] 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.^{[citation needed]}
• Bleach agent. Low concentrations of hydrogen peroxide, on the order of 3% or less, will chemically bleach many types of clothing to a pinkish hue. Caution should be exercised when using common products that may contain hydrogen peroxide, such as facial cleaner or contact lens solution, which easily splatter upon other surfaces.
• Internal ailments. Large oral doses of hydrogen peroxide at a 3% concentration may cause "irritation and blistering to the mouth, (which is known as Black hairy tongue) throat, and abdomen", as well as "abdominal pain, vomiting, and diarrhea".^[42]
• Vapor pressure. Hydrogen peroxide has a significant vapor pressure (1.2 kPa at 50 ^oC[CRC Handbook of Chemistry and Physics, 76th Ed, 1995-1996]) and exposure to the vapor is potentially hazardous. Hydrogen peroxide vapor is a primary irritant, primarily affecting the eyes and respiratory system and the NIOSH Immediately dangerous to life and health limit (IDLH) is only 75 ppm. Documentation for Immediately Dangerous to Life or Health Concentrations (IDLH): NIOSH [http://www.cdc.gov/NIOSH/National Institute for Occupational Safety and Health] Chemical Listing and Documentation of Revised IDLH Values (as of 3/1/95). Long term exposure to low ppm concentrations is also hazardous and can result in permanent lung damage and OSHA Occupational Safety and Health Administration 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 ACGIH American Conference of Industrial Hygienists (ACGIH) as a "known animal carcinogen, with unknown relevance on humans.[2008 Threshold Limit Values for Chemical Substances and Physical Agents & Biological Exposure Indices, ACGIH] In applications where high concentrations of hydrogen peroxide are used, suitable personal protective equipment should be worn and it is prudent in situations where the vapor is likely to be generated, such as hydrogen peroxide gas or vapor sterilization, to ensure that there is adequate ventilation and the vapor concentration monitored with a continuous gas monitor for hydrogen peroxide. Continuous gas monitors for hydrogen peroxide are available from several suppliers. Further information on the hazards of hydrogen peroxide is available from OSHA Occupational Safety and Health Guideline for Hydrogen Peroxide and from the ATSDR. Agency for Toxic Substances and Disease Registry
• Skin disorders. Vitiligo is an acquired skin disorder with the loss of native skin pigment, which affects about 0.5-1% of the world population. Recent studies have discovered increased H₂O₂ levels in the epidermis and in blood are one of many hallmarks of this disease.^[43]

[edit] Historical incidents

• On July 16, 1934 in Kummersdorf, Germany a rocket engine using hydrogen peroxide exploded, killing three people. As a result of this incident, Werner von Braun decided not to use hydrogen peroxide as an oxidizer in the rockets he developed afterward.
• Several people received minor injuries after a hydrogen peroxide spill on board Northwest Airlines flight 957 from Orlando to Memphis on October 28, 1998 and subsequent fire on Northwest Airlines flight 7.^[44]
• During the Second World War, doctors in Nazi concentration camps experimented with the use of hydrogen peroxide injections in the killing of human subjects.^[45]
• Hydrogen peroxide was said to be one of the ingredients in the bombs that failed to explode in the July 21, 2005 London bombings.^[46]
• The Russian submarine K-141 Kursk sailed out to sea to perform an exercise of firing dummy torpedoes at the Pyotr Velikiy, a Kirov class battlecruiser. On August 12, 2000 at 11:28 local time (07:28 UTC), there was an explosion while preparing to fire the torpedoes. The only credible report to date is that this was due to the failure and explosion of one of the Kursk's hydrogen peroxide-fueled torpedoes. It is believed that HTP, a form of highly concentrated hydrogen peroxide used as propellant for the torpedo, seeped through rust in the torpedo casing. A similar incident was responsible for the loss of HMS Sidon in 1955
• On August 16, 2010 a spill of about 10 gallons of cleaning fluid spilled on the 53rd floor of 1515 Broadway, in Times Square, New York City. The spill, which a spokesperson for the New York City fire department said was of Hydrogen Peroxide, shut down Broadway between West 42nd and West 48th streets as a number of fire engines responded to the hazmat situation. There were no reported injuries.^[47]

[edit] References

[edit] Notes

1. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
2. ^ Hill, C. N. (2001). A Vertical Empire: The History of the UK Rocket and Space Programme, 1950-1971. Imperial College Press. ISBN 9781860942686. http://books.google.com/?id=AzoCJfTmRDsC. 
3. ^ Dougherty, Dennis A.; Eric V. Anslyn (2005). Modern Physical Organic Chemistry. University Science. p. 122. ISBN 1-891389-31-9. 
4. ^ Landolt-Börnstein Substance - Property Index
5. ^ 60% hydrogen peroxide msds 50% H2O2 MSDS
6. ^ L. J. Thenard (1818). Annales de chimie et de physique 8: 308. 
7. ^ C. W. Jones, J. H. Clark. Applications of Hydrogen Peroxide and Derivatives. Royal Society of Chemistry, 1999.
8. ^ Richard Wolffenstein (1894). "Concentration und Destillation von Wasserstoffsuperoxyd". Berichte der deutschen chemischen Gesellschaft 27 (3): 3307–3312. doi:10.1002/cber.189402703127. 
9. ^ ^{a b} Jose M. Campos-Martin, Gema Blanco-Brieva, Jose L. G. Fierro (2006). "Hydrogen Peroxide Synthesis: An Outlook beyond the Anthraquinone Process". Angewandte Chemie International Edition 45 (42): 6962–6984. doi:10.1002/anie.200503779. PMID 17039551. 
10. ^ ^{a b} H. Riedl and G. Pfleiderer, U.S. Patent 2,158,525 (October 2, 1936 in USA, and October 10, 1935 in Germany) to I. G. Farbenindustrie, Germany
11. ^ ^{a b} Ronald Hage, Achim Lienke (2005). "Applications of Transition-Metal Catalysts to Textile and Wood-Pulp Bleaching". Angewandte Chemie International Edition 45 (2): 206–222. doi:10.1002/anie.200500525. PMID 16342123. 
12. ^ ^{a b c d e} Hydrogen Peroxide 07/08-03 Report, ChemSystems, May 2009.
13. ^ ^{a b} G.J. Hutchings et al, Science, 2009, 323, 1037
14. ^ "Gold-palladium Nanoparticles Achieve Greener, Smarter Production Of Hydrogen Peroxide". Sciencedaily.com. 2009-03-03. http://www.sciencedaily.com/releases/2009/02/090219141507.htm. Retrieved 2010-09-05. 
15. ^ Jennifer K. Edwards, Benjamin Solsona, Edwin Ntainjua N, Albert F. Carley (Feb 2009). "Switching off hydrogen peroxide hydrogenation in the direct synthesis process.". Science 323 (5917): 1037–41. doi:10.1126/science.1168980. PMID 19229032. 
16. ^ Instant steam puts heat on MRSA, Society Of Chemical Industry
17. ^ "Natural bleach 'key to healing'". BBC News. 6 June 2009. http://news.bbc.co.uk/1/hi/health/8078525.stm. Retrieved 2009-07-02. 
18. ^ Niethammer, Philipp; Clemens Grabher, A. Thomas Look & Timothy J. Mitchison (3 June 2009). "A tissue-scale gradient of hydrogen peroxide mediates rapid wound detection in zebrafish". Nature 459 (7249): 996–999. doi:10.1038/nature08119. ISSN doi=10.1038/nature08119. PMID 19494811. PMC 2803098. http://www.nature.com/nature/journal/v459/n7249/full/nature08119.html. Retrieved 2009-07-02. 
19. ^ "Hydrogen peroxide disrupts scarless fetal wound repair". Cat.inist.fr. http://cat.inist.fr/?aModele=afficheN&cpsidt=17151171. Retrieved 2010-09-05. 
20. ^ How to Induce Vomiting (Emesis) in Dogs
21. ^ Fredrickson, Bryce. "Hydrogen Peroxide and Horticulture". http://www.socalplumeriacare.com/Faqs/F-7.pdf. Retrieved 2009-01-25. 
22. ^ Ways to use hydrogen peroxide in the garden
23. ^ Oxygation Unlocks Yield Potentials of Crops in Oxygen-Limited Soil Environments Advances in Agronomy, Volume 88, 2005, Pages 313-377 Surya P. Bhattarai, Ninghu Su, David J. Midmore
24. ^ Great-lakes.org
25. ^ fws.gov
26. ^ Chemist Paul Krebaum claims to have originated the formula for use on skunked pets at Skunk Remedy
27. ^ Scott, Richard (November, 1997). "Homing Instincts". Jane's Navy Steam generated by catalytic decomposition of 80-90 % hydrogen peroxide was used for driving the turbopump turbines of the V-2 rockets, the X-15 rocketplanes, the early Centaur RL-10 engines and is still used on Soyuz for that purpose to-day. International. http://babriet.tripod.com/articles/art_hominginstinct.htm. 
28. ^ "Sec. 184.1366 Hydrogen peroxide". U.S. Government Printing Office via GPO Access. 2001-04-01. http://a257.g.akamaitech.net/7/257/2422/04nov20031500/edocket.access.gpo.gov/cfr_2001/aprqtr/21cfr184.1366.htm. Retrieved 2007-07-07. 
29. ^ Shepherd, Steven. "Brushing Up on Gum Disease". FDA Consumer. Archived from the original on May 14, 2007. http://web.archive.org/web/20070514102017/http://www.fda.gov/bbs/topics/CONSUMER/CON00065.html. Retrieved 2007-07-07. 
30. ^ Milani, Massimo; Bigardi, Andrea; Zavattarelli, Marco (2003). "Efficacy and safety of stabilised hydrogen peroxide cream (Crystacide) in mild-to-moderate acne vulgaris: a randomised, controlled trial versus benzoyl peroxide gel" (^{[dead link]}). Current Medical Research and Opinion 19 (2): 135–138(4). doi:10.1185/030079902125001523. PMID 12740158. http://www.medscape.com/viewarticle/452990. 
31. ^ "Drugs to Control or Stimulate Vomiting". Merck Veterinary manual. Merck & Co., Inc. 2006. http://www.merckvetmanual.com/mvm/index.jsp?cfile=htm/bc/190303.htm. 
32. ^ "Questionable methods of cancer management: hydrogen peroxide and other 'hyperoxygenation' therapies". CA: a cancer journal for clinicians 43 (1): 47–56. 1993. doi:10.3322/canjclin.43.1.47. PMID 8422605. 
33. ^ Cooper, Anderson (2005-01-12). "A Prescription for Death?". CBS News. http://www.cbsnews.com/stories/2005/01/12/60II/main666489.shtml. Retrieved 2007-07-07. 
34. ^ Mikkelson, Barbara (2006-04-30). "Hydrogen Peroxide". Snopes.com. http://www.snopes.com/medical/healthyself/peroxide.asp. Retrieved 2007-07-07. 
35. ^ Template:Cite title=Hydrogen Peroxide Therapy - Drinking Hydrogen Peroxide
36. ^ Hydrogen Peroxide MSDS
37. ^ Ozonelab Peroxide compatibility
38. ^ "The Many Uses of Hydrogen Peroxide-Truth! Fiction! Unproven!". http://www.truthorfiction.com/rumors/h/hydrogen-peroxide.htm. Retrieved 2008-06-30. 
39. ^ NTSB - Hazardous Materials Incident Brief
40. ^ Armadilloaerospace material tests with HTP
41. ^ For example, see an MSDS for a 3% peroxide solution.
42. ^ Hydrogen Peroxide, 3%. 3. Hazards Identification Southeast Fisheries Science Center, daughter agency of NOAA.
43. ^ "forschung". Vitiligo.eu.com. http://www.vitiligo.eu.com/turmeric.htm. Retrieved 2010-09-05. 
44. ^ Hazardous Materials Incident Brief DCA-99-MZ-001, "Spill of undeclared shipment of hazardous materials in cargo compartment of aircraft". pub: National Transportation Safety Board. October 28, 1998; adopted May 17, 2000.
45. ^ "The Nazi Doctors: Medical Killing and the Psychology of Genocide". Robert Jay Lifton. http://www.holocaust-history.org/lifton/LiftonT257.shtml. Retrieved 1 November 2007. 
46. ^ Four Men Found Guilty in Plot to Blow Up London's Transit System, "FOXNews.com". (July 9, 2007)
47. ^ Times Sq. cleaning fluid spill brings fire trucks, [1]. (August 17, 2010)

[edit] Bibliography

• J. Drabowicz et al., in The Syntheses of Sulphones, Sulphoxides and Cyclic Sulphides, p112-116, G. Capozzi et al., eds., John Wiley & Sons, Chichester, UK, 1994. ISBN 0-471-93970-6.
• N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997. A great description of properties & chemistry of H₂O₂.
• J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
• W. T. Hess, Hydrogen Peroxide, in Kirk-Othmer Encyclopedia of Chemical Technology, 4th edition, Wiley, New York, Vol.13, 961-995 (1995).

[edit] External links

This article's use of external links may not follow Wikipedia's policies or guidelines. Please improve this article by removing excessive and inappropriate external links or by converting links into footnote references. (June 2010)

• Hydrogen Peroxide Distillation for rocket fuel
• Material Safety Data Sheet
• ATSDR Agency for Toxic Substances and Disease Registry FAQ
• Negative effects of Hydrogen Peroxide as an oral rinse
• Food Grade Hydrogen Peroxide Information
• Experimental Rocket Propulsion Society
• International Chemical Safety Card 0164
• NIOSH Pocket Guide to Chemical Hazards
• IARC Monograph "Hydrogen Peroxide"
• General Kinetics Inc. Hydrogen Peroxide Rocket Engines and Gas Generators
• Oxygenation Therapy:Unproven Treatments for Cancer and AIDS
• Hydrogen Peroxide in the Human Body
• Information on many common uses for hydrogen peroxide, especially household uses.
• Hydrogen peroxide in tooth whiteners summary by GreenFacts of the European Commission SCCP assessment

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Sodium sulfate

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Sodium sulfate
Other names[hide] Thenardite (mineral) Glauber's salt (decahydrate) Sal mirabilis (decahydrate) Mirabilite (decahydrate)
Identifiers
CAS number	7757-82-6 Y,  7727-73-3 (decahydrate)
PubChem	24436
RTECS number	WE1650000
Properties
Molecular formula	Na2SO4
Molar mass	142.04 g/mol (anhydrous) 322.20 g/mol (decahydrate)
Appearance	white crystalline solid hygroscopic
Density	2.664 g/cm3 (anhydrous) 1.464 g/cm3 (decahydrate)
Melting point	884 °C (anhydrous) 32.4 °C (decahydrate)
Boiling point	1429 °C (anhydrous)
Solubility in water	47.6 g/L (0 °C) 427 g/L (100 °C)
Solubility	insoluble in ethanol
Refractive index (nD)	1.468 (anhydrous) 1.394 (decahydrate)
Structure
Crystal structure	orthorhombic or hexagonal (anhydrous) monoclinic (decahydrate)
Hazards
MSDS	External MSDS
EU Index	Not listed
Main hazards	Irritant
NFPA 704	0 1 0
Flash point	Non-flammable
Related compounds
Other anions	Sodium selenate Sodium tellurate
Other cations	Lithium sulfate Potassium sulfate Rubidium sulfate Caesium sulfate
Related compounds	Sodium bisulfate Sodium sulfite Sodium persulfate
Supplementary data page
Structure and properties	n, εr, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Y (what is this?)  (verify) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Sodium sulfate is the sodium salt of sulfuric acid. When anhydrous, it is a white crystalline solid of formula Na₂SO₄ known as the mineral thenardite; the decahydrate Na₂SO₄·10H₂O has been known as Glauber's salt or, historically, sal mirabilis since the 17th century. Another solid is the heptahydrate, which transforms to mirabilite when cooled. With an annual production of 6 million tonnes, it is a major commodity chemical and one of the most damaging salts in structure conservation: when it grows in the pores of stones it can achieve high levels of pressure, causing structures to crack.

Sodium sulfate is mainly used for the manufacture of detergents and in the Kraft process of paper pulping. About two-thirds of the world's production is from mirabilite, the natural mineral form of the decahydrate, and the remainder from by-products of chemical processes such as hydrochloric acid production.

Contents [hide] 1 History 2 Physical and chemical properties 3 Production 3.1 Natural sources 3.2 Chemical industry 4 Applications 4.1 Commodity industries 4.2 Thermal storage 4.3 Small-scale applications 5 Safety 6 References 7 External links

[edit] History

The hydrate of sodium sulfate is known as Glauber's Salt after the Dutch/German chemist and apothecary Johann Rudolf Glauber (1604–1670), who discovered it 1625 in Austrian spring water. He named it sal mirabilis (miraculous salt), because of its medicinal properties: the crystals were used as a general purpose laxative, until more sophisticated alternatives came about in the 1900s.^[1][2]

In the 18th century, Glauber's salt began to be used as a raw material for the industrial production of soda ash (sodium carbonate), by reaction with potash (potassium carbonate). Demand for soda ash increased and supply of sodium sulfate had to increase in line. Therefore, in the nineteenth century, the Leblanc process, producing synthetic sodium sulfate as a key intermediate, became the principal method of soda ash production.^[3]

[edit] Physical and chemical properties

Sodium sulfate is chemically very stable, being unreactive toward most oxidising or reducing agents at normal temperatures. At high temperatures, it can be reduced to sodium sulfide.^[4] It is a neutral salt, which forms aqueous solutions with pH of 7. The neutrality of such solutions reflects the fact that Na₂SO₄ is derived, formally speaking, from the strong acid sulfuric acid and a strong base sodium hydroxide. Sodium sulfate reacts with an equivalent amount of sulfuric acid to give an equilibrium concentration of the acid salt sodium bisulfate^[5][6]:

Na₂SO₄(aq) + H₂SO₄(aq) ⇌ 2 NaHSO₄(aq)

In fact, the equilibrium is very complex, depending on concentration and temperature, with other acid salts being present.

Sodium sulfate is a typical ionic sulfate, containing Na⁺ ions and SO₄²⁻ ions. Aqueous solutions can produce precipitates when combined with salts of Ba²⁺ or Pb²⁺, which form insoluble sulfates

Na₂SO₄(aq) + BaCl₂(aq) → 2 NaCl(aq) + BaSO₄(s)

Sodium sulfate has unusual solubility characteristics in water.^[7] Its solubility in water rises more than tenfold between 0 °C to 32.384 °C, where it reaches a maximum of 497 g/L. At this point the solubility curve changes slope, and the solubility becomes almost independent of temperature. This temperature at 32.384 °C, corresponding to the release of crystal water and melting of the hydrated salt, serves as an accurate temperature reference for thermometer calibration.

Sodium sulfate decahydrate is also unusual among hydrated salts in having a measureable residual entropy (entropy at absolute zero) of 6.32 J·K⁻¹·mol⁻¹. This is ascribed to its ability to distribute water much more rapidly compared to most hydrates.^[8]

Sodium sulfate displays a moderate tendency to form double salts. The only alums formed with common trivalent metals are NaAl(SO₄)₂ (unstable above 39 °C) and NaCr(SO₄)₂, in contrast to potassium sulfate and ammonium sulfate which form many stable alums.^[9] Double salts with some other alkali metal sulfates are known, including Na₂SO₄·3K₂SO₄ which occurs naturally as the mineral glaserite. Formation of glaserite by reaction of sodium sulfate with potassium chloride has been used as the basis of a method for producing potassium sulfate, a fertiliser.^[10] Other double salts include 3Na₂SO₄·CaSO₄, 3Na₂SO₄·MgSO₄ (vanthoffite) and NaF·Na₂SO₄.^[11]

[edit] Production

The world production of sodium sulfate, mostly in the form of the decahydrate amounts to approximately 5.5 to 6 million tonnes annually (Mt/a). In 1985, production was 4.5 Mt/a, half from natural sources, and half from chemical production. After 2000, at a stable level until 2006, natural production had increased to 4 Mt/a, and chemical production decreased to 1.5 to 2 Mt/a, with a total of 5.5 to 6 Mt/a.^{[12][13][14][15]} For all applications, naturally produced and chemically produced sodium sulfate are practically interchangeable.

[edit] Natural sources

Two thirds of the world's production of the decahydrate (Glauber's salt) is from the natural mineral form mirabilite, for example as found in lake beds in southern Saskatchewan. In 1990, Mexico and Spain were the world's main producers of natural sodium sulfate (each around 500,000 tonnes), with Russia, US and Canada around 350,000 tonnes each.^[13] Natural resources are estimated as over 1 billion tonnes.^[12][13]

Major producers of 200,000–1,500,000 tonnes/a in 2006 include Searles Valley Minerals (California, US), Airborne Industrial Minerals (Saskatchewan, Canada), Química del Rey (Coahuila, Mexico), Criaderos Minerales Y Derivados and Minera de Santa Marta, also known as Grupo Crimidesa (Burgos, Spain), FMC Foret (Toledo, Spain), Sulquisa (Madrid, Spain), and in China Chengdu Sanlian Tianquan Chemical (Sichuan), Hongze Yinzhu Chemical Group (Jiangsu), Nafine Chemical Industry Group (Shanxi), and Sichuan Province Chuanmei Mirabilite (Sichuan), and Kuchuksulphat JSC (Altai Krai, Siberia, Russia).^[12][14]

Anhydrous sodium sulfate occurs in arid environments as the mineral thenardite. It slowly turns to mirabilite in damp air. Sodium sulfate is also found as glauberite, a calcium sodium sulfate mineral. Both minerals are less common than mirabilite.

[edit] Chemical industry

About one third of the world's sodium sulfate is produced as by-product of other processes in chemical industry. Most of this production is chemically inherent to the primary process, and only marginally economical. By effort of the industry, therefore, sodium sulfate production as by-product is declining.

The most important chemical sodium sulfate production is during hydrochloric acid production, either from sodium chloride (salt) and sulfuric acid, in the Mannheim process, or from sulfur dioxide in the Hargreaves process.^[16][17] The resulting sodium sulfate from these processes are known as salt cake.

Mannheim: 2 NaCl + H₂SO₄ → 2 HCl + Na₂SO₄

Hargreaves: 4 NaCl + 2 SO₂ + O₂ + 2 H₂O → 4 HCl + 2 Na₂SO₄

The second major production of sodium sulfate are the processes where surplus sulfuric acid is neutralised by sodium hydroxide, as applied on a large scale in the production of rayon. This method is also a regularly applied and convenient laboratory preparation.

2 NaOH(aq) + H₂SO₄(aq) → Na₂SO₄(aq) + 2 H₂O(l)

Formerly, sodium sulfate was also a by-product of the manufacture of sodium dichromate, where sulfuric acid is added to sodium chromate solution forming sodium dichromate, or subsequently chromic acid. Alternatively, sodium sulfate is or was formed in the production of lithium carbonate, chelating agents, resorcinol, ascorbic acid, silica pigments, nitric acid, and phenol.^[12]

Bulk sodium sulfate is usually purified via the decahydrate form, since the anhydrous form tends to attract iron compounds and organic compounds. The anhydrous form is easily produced from the hydrated form by gentle warming.

Major sodium sulfate by-product producers of 50–80 Mt/a in 2006 include Elementis Chromium (chromium industry, Castle Hayne, NC, US), Lenzing AG (200 Mt/a, rayon industry, Lenzing, Austria), Addiseo (formerly Rhodia, methionine industry, Les Roches-Roussillon, France), Elementis (chromium industry, Stockton-on-Tees, UK), Shikoku Chemicals (Tokushima, Japan) and Visko-R (rayon industry, Russia).^[12]

[edit] Applications

Sodium sulfate used to dry an organic liquid. Here clumps form, indicating the presence of water in the organic liquid.

By further application of sodium sulfate the liquid may be brought to dryness, indicated here by the absence of clumping.

[edit] Commodity industries

With US pricing at $30 per tonne in 1970, in 2006 up to $90 per tonne for salt cake quality and $130 for better grades, sodium sulfate is a very cheap material. The largest use is as filler in powdered home laundry detergents, consuming approx. 50% of world production. This use is waning as domestic consumers are increasingly switching to compact or liquid detergents that do not include sodium sulfate.^[12]

Another formerly major use for sodium sulfate, notably in the US and Canada, is in the Kraft process for the manufacture of wood pulp. Organics present in the "black liquor" from this process are burnt to produce heat, needed to drive the reduction of sodium sulfate to sodium sulfide. However, this process is being replaced by newer processes; use of sodium sulfate in the US and Canadian pulp industry declined from 1.4 Mt/a in 1970 to only approx. 150,000 tonnes in 2006.^[12]

The glass industry provides another significant application for sodium sulfate, as second largest application in Europe. Sodium sulfate is used as a fining agent, to help remove small air bubbles from molten glass. It fluxes the glass, and prevents scum formation of the glass melt during refining. The glass industry in Europe has been consuming from 1970 to 2006 a stable 110,000 tonnes annually.^[12]

Sodium sulfate is important in the manufacture of textiles, particularly in Japan, where it is the largest application. Sodium sulfate helps in "levelling", reducing negative charges on fibres so that dyes can penetrate evenly. Unlike the alternative sodium chloride, it does not corrode the stainless steel vessels used in dyeing. This application in Japan and US consumed in 2006 approximately 100,000 tonnes.^[12]

[edit] Thermal storage

The high heat storage capacity in the phase change from solid to liquid, and the advantageous phase change temperature of 32 °C (90 °F) makes this material especially appropriate for storing low grade solar heat for later release in space heating applications. In some applications the material is incorporated into thermal tiles that are placed in an attic space while in other applications the salt is incorporated into cells surrounded by solar–heated water. The phase change allows a substantial reduction in the mass of the material required for effective heat storage (83 calories per gram stored across the phase change, versus one calorie per gram per degree Celsius using only water), with the further advantage of a consistency of temperature as long as sufficient material in the appropriate phase is available.

[edit] Small-scale applications

In the laboratory, anhydrous sodium sulfate is widely used as an inert drying agent, for removing traces of water from organic solutions.^[18] It is more efficient, but slower-acting, than the similar agent magnesium sulfate. It is only effective below about 30 °C, but it can be used with a variety of materials since it is chemically fairly inert. Sodium sulfate is added to the solution until the crystals no longer clump together; the two video clips (see above) demonstrate how the crystals clump when still wet, but some crystals flow freely once a sample is dry.

Glauber's salt, the decahydrate, was historically used as a laxative. It is effective for the removal of certain drugs such as acetaminophen from the body, for example, after an overdose.^[19][20]

In 1953, sodium sulfate was proposed for heat storage in passive solar heating systems. This takes advantage of its unusual solubility properties, and the high heat of crystallisation (78.2 kJ/mol).^[21]

Other uses for sodium sulfate include de-frosting windows, in carpet fresheners, starch manufacture, and as an additive to cattle feed.

Lately, sodium sulfate has been found effective in dissolving very finely electroplated micrometre gold that is found in gold electroplated hardware on electronic products such as pins, and other connectors and switches. It is safer and cheaper than other reagents used for gold recovery, with little concern for adverse reactions or health effects.^{[citation needed]}

At least one company, ThermalTake, makes a laptop computer chill mat (iXoft Notebook Cooler) using sodium sulfate decahydrate inside a quilted plastic pad. The material slowly turns to liquid and recirculates, equalizing laptop temperature and acting as an insulation.

[edit] Safety

Although sodium sulfate is generally regarded as non-toxic,^[22] it should be handled with care. The dust can cause temporary asthma or eye irritation; this risk can be prevented by using eye protection and a paper mask. Transport is not limited, and no Risk Phrase or Safety Phrase apply.^[23]

[edit] References

1. ^ Szydlo, Zbigniew (1994). Water which does not wet hands: The Alchemy of Michael Sendivogius. London-Warsaw: Polish Academy of Sciences. 
2. ^ Westfall, Richard S. (1995). "Glauber, Johann Rudolf". The Galileo Project. http://galileo.rice.edu/Catalog/NewFiles/glauber.html. 
3. ^ Aftalion, Fred (1991). A History of the International Chemical Industry. Philadelphia: University of Pennsylvania Press. pp. 11–16. ISBN 0-8122-1297-5. 
4. ^ Handbook of Chemistry and Physics (71st ed.). Ann Arbor, Michigan: CRC Press. 1990. 
5. ^ The Merck Index (7th ed.). Rahway, New Jersey, US: Merck & Co.. 1960. 
6. ^ Nechamkin, Howard (1968). The Chemistry of the Elements. New York: McGraw-Hill. 
7. ^ Linke, W.F.; A. Seidell (1965). Solubilities of Inorganic and Metal Organic Compounds (4th ed.). Van Nostrand. 
8. ^ Brodale, G.; W.F. Giauque (1958). "The Heat of Hydration of Sodium Sulfate. Low Temperature Heat Capacity and Entropy of Sodium Sulfate Decahydrate". Journal of the American Chemical Society 80: 2042–2044. doi:10.1021/ja01542a003. 
9. ^ Lipson, Henry; C.A. Beevers (1935). "The Crystal Structure of the Alums". Proceedings of the Royal Society of London Series A 148 (865): 664–80. doi:10.1098/rspa.1935.0040. 
10. ^ Garrett, Donald E. (2001). Sodium sulfate : handbook of deposits, processing, properties, and use. San Diego: Academic Press. ISBN 9780122761515. 
11. ^ Mellor, Joseph William (1961). Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry. Volume II (new impression ed.). London: Longmans. pp. 656–673. 
12. ^ ^{a b c d e f g h i} Suresh, Bala; Kazuteru Yokose (May 2006). Sodium sulfate. Zurich: Chemical Economic Handbook SRI Consulting. pp. 771.1000A–771.1002J. http://www.sriconsulting.com/CEH/Public/Reports/771.1000/?Abstract.html. 
13. ^ ^{a b c} "Statistical compendium Sodium sulfate". Reston, Virginia: US Geological Survey, Minerals Information. 1997. http://minerals.usgs.gov/minerals/pubs/commodity/sodium_sulfate/stat. Retrieved 2007-04-22. 
14. ^ ^{a b} The economics of sodium sulphate (Eighth ed.). London: Roskill Information Services. 1999. 
15. ^ The sodium sulphate business. London: Chem Systems International. November 1984. 
16. ^ Butts, D. (1997). Kirk-Othmer Encyclopedia of Chemical Technology. v22 (4th ed.). pp. 403–411. 
17. ^ Hargreaves, J. (1873). Chem. News 27: 183. 
18. ^ Vogel, Arthur I.; B.V. Smith, N.M. Waldron (1980). Vogel's Elementary Practical Organic Chemistry 1 Preparations (3rd ed.). London: Longman Scientific & Technical. 
19. ^ Cocchetto, D.M.; G. Levy (1981). "Absorption of orally administered sodium sulfate in humans". J Pharm Sci 70 (3): 331–3. doi:10.1002/jps.2600700330. PMID 7264905. 
20. ^ Prescott, L.F.; J.A.J.H. Critchley (1979). "The Treatment of Acetaminophen Poisoning". Annual Review of Pharmacology and Toxicology 23: 87–101. doi:10.1146/annurev.pa.23.040183.000511. PMID 6347057. 
21. ^ Telkes, Maria (1953). Improvements in or relating to a device and a composition of matter for the storage of heat. http://v3.espacenet.com/textdes?DB=EPODOC&IDX=GB694553&F=0&QPN=GB694553. 
22. ^ "Sodium sulfate (WHO Food Additives Series 44)". World Health Organization. 2000. http://www.inchem.org/documents/jecfa/jecmono/v44jec07.htm. Retrieved 2007-06-06. 
23. ^ "MSDS Sodium Sulfate Anhydrous". James T Baker. 2006. http://www.jtbaker.com/msds/englishhtml/S5022.htm. Retrieved 2007-04-21. 

[edit] External links

• Sodium sulfate information of Airborne Industrial Minerals
• Sodium sulfate website of Elementis Chromium