30% Hydrogen peroxide

Peroxide solutions find use as oxidizing or bleaching agents in many sectors of industry. 

An on-site electrolytic process offers particular advantages for the pulp and paper industry where peroxide mixtures are used to brighten and bleacb pulp 1 

while chemical pulps contain 0-l5gdm NaOH 15-20gdm  

Na2Si03 and 5-10gdm H2O3, both liquors containing stabilizers  

3 Pure electrolytes, free from transition-metal ions which may decompose the peroxide and drastically reduce current efficiency. 

Peroxide or a combination of peroxide and peracetic acid is generally used to treat RO systems that are already contaminated with microbes. Due to its high ORP, however, a solution of only 0.2wt% peroxide is normally used (see Table 8.8). Temperature must be below 25°C and transition metals such as iron must be removed prior to treatment with peroxide to minimize oxidation of the membrane. Further, membrane should be cleaned free of deposits using an alkaline cleaner before peroxide is applied. Finally, a pH of 3 - 4 should be maintained and exposure limited to about 20 minutes for optimum result and maximum membrane life. Peroxide should not be used for storage of membrane modules. 

Sequestering agents (also known as scale inhibitors or antisealants) are used to minimize the potential for forming scale on the surface of an RO membrane. Antisealants work by one of three methods  

Early antisealants used sodium hexametaphosphate (SHMP) as a threshold agent to inhibit the growth of calcium carbonate and sulfate-based scales.6 Most antisealants on the market today contain sulfonate, phosphate, or carboxylic acid functional groups. Perhaps the most effective antisealants today contain and blend of polyacrylic acid (PAA) and phosphoric acid or polyacrylate and a hydroyethylidene diphosphonate (HEDP).12 The polyacrylate-HEDP blends also claim to have good dispersion qualities toward silts and clays.12 Some new inhibitors include a chelant and disper-ant to keep suspended solids such as iron and manganese oxides in solution. These newer antisealants are generally more effective than SHMP for a variety of potential scales.6 

Some antisealants, such as those containing HEDP are sensitive to chlorine and other oxidizers.12 They should be fed downstream of the dechlorination point. Most other antisealants are not affected by chlorine at typical 0.5 - + 1.0 ppm concentrations.13 

Antisealants are completely rejected by RO membranes and, therefore, are not a concern for product quality. 

Hydrogen peroxide was first made in 1818 by J. L. Thenard who acidified barium peroxide (p. 121) and then removed excess H2O by evaporation under reduced pressure. Later the compound was prepared by hydrolysis of peroxodisulfates obtained by electrolytic oxidation of acidified sulfate solutions at high current densities  

Such processes are now no longer used except in the laboratory preparation of D2O2, e.g.  

On an industrial scale H2O2 is now almost exclusively prepared by the autoxidation of 2-alkylanthraquinols (see Panel on next page). 

Hydrogen peroxide, when pure, is an almost colourless (very pale blue) liquid, less volatile than water and somewhat more dense and viscous. Its more important physical properties are in Table 14.11 (cf. H2O, p. 623). The compound is miscible with water in all proportions and forms a hydrate H2O2.H2O, mp —52°. Addition of water increases the already high dielectric constant of H2O2 (70.7) to a maximum value of 121 at 35% H2O2, i.e. substantially higher than the value of water itself (78.4 at 25°). 

In the gas phase the molecule adopts a skew configuration with a dihedral angle of 111.5° as 

Hydrogen peroxide is used as a topical antiseptic in dilute solutions, and as a water purifier in stronger solutions. 

Hydrogen peroxide breaks down into water and oxygen. A liter of 3 percent hydrogen peroxide will generate 10 liters of oxygen when a catalyst is used to facilitate the breakdown. Catalysts can be metals such as iron, copper, or silver, or organics such as the blood enzyme 

Catalase is an important enzyme in cells because hydrogen peroxide is a by-product of metabolism and can poison the cell unless it is decomposed quickly. Hydrogen peroxide is also produced by cells in the immune system, and catalase removes the excess. 

The 3 percent hydrogen peroxide you get at the drugstore is often protected from decomposing by the addition of sodium silicate, magnesium sulfate, or tin compounds. These stabilizers lock up the iron, copper, and other transition metals that can act as catalysts. 

Hydrogen peroxide is the most widely used peroxide compound. Originally, it was produced by the reaction of barium peroxide and sulfuric acid but this process and use have been superseded. 

The most important method of making hydrogen peroxide is by reduction of anthraquinone to the hydroquinone, followed by reoxidation to anthra-quinone by oxygen and formation of the peroxide. 

The hydrogen peroxide is extracted with water and concentrated, and the quinone is recycled for reconversion to the hydroquinone. A second organic process uses isopropyl alcohol, which is oxidized at moderate temperatures and pressures to hydrogen peroxide and acetone. After distillation of the acetone and unreacted alcohol, the residual hydrogen peroxide is concentrated. 

Hydrogen peroxide applications include commercial bleaching dye oxidation, the manufacture of organic and peroxide chemicals. Hydrogen peroxide is also used in pulp and paper chemical synthesis, textiles, and environmental control, including municipal and industrial water treatment. 

Ibuprofen, which is sold under trade names such as Motrin and Advil , is an alternative to aspirin and acetaminophen because of its analgesic and antiinflammatory properties. 

Hydrogen peroxide (HY-druh-jin per-OK-side) is a clear, colorless, somewhat unstable liquid with a bitter taste. When absolutely pure, the compound is quite stable. Even small amounts of impurities (such as iron or copper), however, act as catalysts that increase its tendency to decompose, sometimes violently, into water and nascent oxygen (0). To prevent decomposition, small amounts of inhibitors, such as acetanilide or sodium stannate are added to pure hydrogen peroxide and hydrogen peroxide solutions. 

Hydrogen peroxide was discovered in 1818 by French chemist Louis Jacques Thenard (1777-1857). It was first used commercially in the 1800s, primarily to bleach hats. Today, industrial processes make about 500 million kilograms (1 billion pounds) of hydrogen peroxide annually for use in a wide variety of applications ranging from whitening of teeth to propelling rockets. 

Hydrogen peroxide. Red atoms are oxygen and white atoms are hydrogen, publishers 

Hydrogen peroxide occurs in very small amounts in nature. It is formed when atmospheric oxygen reacts with water to form H202. Hydrogen peroxide is also present in plant and animal cells as the byproduct of metabolic reactions that occur in those cells. 

Hydrogen peroxide is sold in concentrations ranging from 3 percent (for home use) to 90 percent (for industrial applications). 

Hydrogen Peroxide. Table 3.1 showed the reduction potential of oxygen in the following reactions  

There is an exception to the poor reactivity exhibited by hydrogen peroxide. When 158 was treated with 30% hydrogen peroxide, in the presence of benzonitrile and potassium hydrogen carbonate, epoxide 159 was formed in 86% yield, in what is known as the Payne epoxidatlon. l The reagents shown generate per-oxybenzimidic acid in situ. This transformation was taken from Smith s synthesis of (+)-calyculin A.250 Payne epoxidation is known to give a diastereofacial preference that is opposite to epoxidations with peroxy-acids. In this particular example, the epoxide 159 was a 3 1 (a P) mixture of isomers at the epoxy carbon. 

Hydrogen peroxide is one of the most powerful oxidants, together with the peroxides of alkali metals (sodium, potassium). Hydrogen peroxide is used for whitening textiles and as fuel for rockets and missiles. 

The most common solution contains 30 wt% (110 vol%). It is slightly acidic, with a pH of 4 at that concentration. 

Hydrogen peroxide is delivered with stabilisers, corrosion inhibitors, nitrates and a phosphate buffer. These additives do not modify the behaviour of aluminium. 

Aluminium, except copper-containing alloys of the 2000 and 7000 series, is one of the rare common metals that do not catalyse the decomposition of hydrogen peroxide [7]. This also applies to very high concentrations (80% and more), which are achieved in modem plants 

The resistance of aluminium depends on the concentration of hydrogen peroxide. Below a concentration of 10%, a slight superficial pitting attack with a depth of less than 0.1 mm is observed. The dissolution rate increases up to a concentration of 30% and reaches a value on the order of 0.20 mm per year. For concentrations of 30% or higher, the pitting attack ceases, and the decrease in thickness of the metal drops below 0.01 mm per year [8]. 

The use of hydrogen peroxide in conjunction with Fe(II) (Fenton s reagent) or ozone has already been noted. It has been nsed alone to examine the products from o - and m-phenylenediamines in the context of their mntagenicity (Watanabe et al. 1989). Successive reactions produced 3,4-diaminophenazine from o-phenylenediamine, and 3,7-diaminophenazine from m-phenylenediamine. 

CHEMICAL NAME = hydrogen peroxide CAS NUMBER = 7722-84-1 MOLECULAR FORMULA = H202 MOLAR MASS = 34.0 g/mol COMPOSITION = H (5.9%) 0(94.1%) 

The decomposition of hydrogen peroxide by metal ions of variable valency (Mez) was studied by Haber and Weiss 44 , whose mechanism is still the one accepted today  

The metal ion is a typical catalyst for a reaction of Class B, but because H+ and OH- ions are formed in the process, the reaction may be expected to be pH dependent as well. The catalyst does not necessarily have to be a metal ion. Every substrate that can change oxidation state at the right potential and is capable of reacting with the hydrogen peroxide molecule in its two valency states can be considered a catalyst for this reaction. This means that the best parameter to correlate this reaction with would be the redox potential of the catalyst, which unfortunately is very difficult to measure on a solid material. The best thing to do is to use a catalyst which can be dissolved in a liquid medium of some kind, and to study the redox properties in the dissolved state. Measurements of this kind will be discussed in the section on oxidation and dehydrogenation. 

A neat illustration of how the requirements of the chemical reaction define the specificity of the catalyst was given by using polymeric chelates ofbis(hy-droxy-8-quinolyl)-4-methane. The order of reactivity for different complexed metal ions for H202-decomposition was 4S) 

This order of reactivity may be understood if we consider the ease with which Co and Mn change valency as compared with Cu, Ni and Zn. 

Because this is a reaction of Class C, the high activity of the copper complex compared with the other metal ions should probably be explained by its having the right paramagnetic properties for performing the reaction. 

Chlorosulfonic acid reacts with hydrogen peroxide in isoamyl alcohol to yield peroxymonosulfiiric acid 10 (Equation 17). The reaction involves nucleophilic attack by hydrogen peroxide on the electrophilic sulfur atom of chlorosulfonic acid. A peroxydisulfuric acid 11 can also be prepared by this route using excess chlorosulfonic acid (Equation 18),  

The reaction of chlorosulfonic acid with A -methylhydroxylamine to yield the corresponding O-sulfonic acid is described in Chapter 4, p 129. 

Hydroxylamine hydrochloride also reacts with chlorosulfonic acid (1.03 equivalents) in chloroform or 1,2-dichloroethane to give hydroxylamine O-sulfate.  

Siddiqi and N.A. Ansari, Proc. Symp. Chem. React. Non-Aqueous Media Molten Salts, Bombay, India, 1980 Chem. Abs., 96, 92689. 

Hardstaff and R.F. Langler, The Sulfur-Chlorine Bond in Sulfur in Organic and Inorganic Chemistry, A. Senning (ed), M. Dekker Inc., New York, 1982, Vol. 4, 263. 

LABORATORY CHEMICAL SAFETY SUMMARY HYDROGEN PEROXIDE  

Physical Properties Colorless liquid bp 150 °C, mp -0.4 °C Miscible in all proportions in water 

Major Hazards Contact with certain metals and organic compounds can lead to fires and explosions concentrated solutions can cause severe irritation or burns of the skin, eyes, and mucous membranes. 

Flammability and Explosibility Hydrogen peroxide is not flammable, bnt eoneentrated solutions may undergo violent decomposition in tbe presence of trace impurities or upon heating. 

Contact with many organic compounds can lead to immediate fires or violent explosions (consult Bretherick for references and examples). Hydrogen peroxide reacts with certain organic functional groups (ethers, acetals, etc.) to form peroxides, which may explode upon concentration. Reaction with acetone generates explosive cyclic dimeric and trimeric peroxides. Explosions may also occur on exposure of hydrogen peroxide to metals such as sodium, potassium, magnesium, copper, iron, and nickel. 

Homolytic scission of the 0-0 bond of hydrogen peroxide may be effected by heat or UV irradiation.245 The thermal reaction requires relatively high temperatures ( 90 Photolytic initiation generally employs 254 nm light. Reactions in organic media require a polar cosolvent (e.g. an alcohol). 

The Tf- CF system is preferred over Fenton s reagent because Ti4 is a less powerful oxidizing agent than Fc5+ and the above mentioned pathway and other side reactions are therefore of less consequence.252 Much of the discussion on redox initiation in Section 3.3.2.6.1 is also relevant to hydrogen peroxide. 

It has been suggested that the reactive species arc metal complexed hydroxy radicals rather than free hydroxyl radicals. The reactions observed show dependence on the nature of the metal ion and quite different product distributions can be obtained from reaction of organic substrates w ith Fc -H2O2 (Fenton s Reagent) and Ti -H202. However, it is not clear whether these findings reflect the involvement of a different active species or simply the different rates and/or pathways for destruction of the initially formed intermediates. Metal ions in 

The principal value of oxidations with hydrogen peroxide is for the degradation of aldonic acids to sugars with one less carbon atom ferric sulfate is used as a catalyst (see p. 118). It is noteworthy that although ferric salts catalyze this reaction, ferrous salts are used for sugars and ferric salts are not effective. 

With sugars, the products depend upon the conditions and the presence and nature of the catalyst. In any case, the products are usually mixtures. At low temperatures and in the presence of ferrous sulfate, glucose and fructose are converted to glucosone and on further oxidation to glycolic acid, glyoxylic acid, and trihydroxy butyric acid 250). At low temperatures in 

The nature of the products formed under various conditions and the mechanism of the reaction have been investigated by Kiichlin 252). At low temperatures and for dilute solutions in the presence of ferrous sulfate, the following products were formed from glucose and identified as derivatives glucosone, 2-ketogluconic acid, and 2,3-diketogluconic acid in concentrated solutions, formaldehyde also was found. The formation of these products at low temperatures was ascribed to the following series of reactions  

The catalytic effect of ferrous salts is ascribed by Kiichlin to the formation of a complex between ferrous ions and the carbonyl group and its 

According to Haber and Weiss ferrous salts bring about the de- 

The specific gravity of an 80% solution of hydrogen peroxide is 1.34. The solution is fairly stable at room temperature and decomposes only at an elevated temperature (the stability of hydrogen peroxide will be descussed later). To increase the stability of this solution, stabilizing substances such as phosphoric acid and its salts or 8-hydroxyquinoline were added. 

The manufacture of 90% hydrogen peroxide was started after World War II in the U.S.A. and Great Britain. 

Davis and Keefe quote the following physico-chemical constants characteristic of 90% hydrogen peroxide [29]  

Specific gravity at 20°C Viscosity at 18 C Vapour pressure at 30°C Freezing point Boiling point Heat of formation of liquid 

Chemically pure hydrogen peroxide can be stored for a long time without noticeable decomposition. Loss on storage of hydrogen peroxide may amount to 1% per annum. 

The rate law of the oxidation by Fe(III) is dependent on the ratio of the concentrations of the reactants. When peroxide is in excess and when the acidity is sufficient to suppress hydrolysis of Fe(IlI) the rate expression 

Haber and Weiss ° proposed a one-equivalent oxidation of HO2 to H02 later modified by Weiss to 

However, Anderseti noted a new rate law at lower values of [H202]/[Fe(III)] with a higher order dependence on peroxide. 

The reaction scheme as presented is of a non-chain variety provided OH is removed by Fe(II) rather than by peroxide. Barb et have included further steps in the scheme, viz. 

Subsequent analysis of Andersen s data confirms a [HjOj] dependence but not a [Fe(III)] dependence, and this has been confirmed for these conditions by Barb et who obtained a first-order dependence on oxidant concentration. 

Use of catalysts, generally ferric or ferrous salts, promotes radical reactions with hydrogen peroxide the oxidizing action produced by the ferrous ions is more vigorous. 

Ferric ion catalyzes the formation of the hydroperoxyl radical, according to Eq. (35) such a radical appears to constitute the oxidant in the Ruff method of degrading aldonic acids to the next lower aldoses. A number of examples of the use of this reagent in the laboratory are given in a review article by Moody.108 The hydroperoxyl radical, which is not so effective an oxidant as the hydroxyl radical, does not attack aliphatic alcohols accordingly, a substantial yield (about 50%) of the aldose is obtained from the higher aldonic acid. In the presence of an excess of hydrogen peroxide, however, the accumulation of ferrous ions in solution catalyzes the production of hydroxyl radicals and lowers the yield of aldose [see Eq. (36)]. 

A possible mechanism for this reaction suggests the formation of a carboxyl radical, which undergoes degradation to liberate C02. The resulting radical produces the corresponding next lower aldose plus H . 

Hydrogen peroxide in acetonitrile has been shown to convert a pentofuranosidulose oxime into the corresponding nitro derivative.109 

The Fenton reagent is a mixture of hydrogen peroxide and a ferrous salt, which leads to the formation of hydroxyl radicals, according to Eqs. (36)-(40). Hydroxyl radicals are very effective in abstracting hydrogen atoms, in contrast to the weaker action ascribed (Section VI. 1) to hydroperoxyl radicals. 

Preparation. Pour 20 ml of a 5% sulphuric acid solution into a 50-ml flask and cool it to 0 °C. While shaking the contents of the flask during 5-10 min, add about 1 g of barium peroxide in small portions. Filter the obtained hydrogen peroxide solution from the precipitate. Study the properties of hydrogen peroxide. 

Properties. 1. Introduce manganese(IV) oxide into a test tube containing a hydrogen peroxide solution. What gas is evolved What role is played by the manganese(IV) oxide Write the equation of the reaction. 

Pour several drops of a potassium iodide solution into a hydrogen peroxide solution. What is observed Dilute the solution with water and add one or two drops of a starch solution to it. What happens Write the molecular and electron equations of the reaction. 

Prepare a lead sulphide precipitate by reacting solutions of a lead salt and ammonium sulphide. Treat a small amount of the precipitate, washed with water by decantation, with a hydrogen peroxide solution. Why does the colour of the precipitate change Write the equations of the reactions. What properties does hydrogen peroxide exhibit in this and the preceding reactions  

H2O2 is one of the most powerful and widely used oxidizers in present-day industrial CMP slurries [5-9]. Hirabayashi and coworkers [10] were among the first to describe a Cu CMP slurry that contains hydrogen peroxide, glycine, and 

FIGURE 7.2 Material removal and static etch rate of copper as a function of H2O2 concentration (from Ref. 5). 

FIGURE 7.3 A schematic illustration of hydrogen peroxide concentration effect on material removal rate of copper with the addition of complexing agent such as glycine and passivating agent such as BTA (from Ref. 11). 

FIGURE 7.4 Material removal rate on 8 tungsten blanket test wafers as a function of hydrogen peroxide concentration (from Ref. 12). 

FIGURE 7.5 A correlation between hydrogen peroxide concentration and NiP removal rate for a slurry containing 1 wt % glycine, 5 wt % alumina, and various amount of hydrogen peroxide at pH 4 (from Ref. 12). 

H2O2 by an acid-dependent pathway/ The reaction is first order in both species and a mechanism including the formation of various protonated activated complexes is postulated. The dinuclear complex [(CN)5Ru(II)—CN—Ru(III)(CN)5] may be formed by chemical oxidation of [Ru(CN)6] by hydrogen peroxide, among other oxidants.  

Interstitial carbide compounds of transition metals (such as WC and TiC) do not readily hydrolyze. However, these compounds when dissolved in H2O2 solutions give oxalato complexes of W(VI), etc. This implies formation of a 

Kinetic measurements on the oxidation of dichlorotetracarbonyldir-hodium(I) by H2O2 in ethanol/hydrochloric acid indicate that the reaction proceeds through an intermediate, Rh(CO)Cl3. The reaction is first order with respect to the rhodium complex, hydrogen peroxide, and hydrochloric acid but the kinetics are complicated by complex formation between [Rh(CO)2Cl2] and chloride ions, and between the intermediate and hydrogen peroxide. 

Data were last reviewed in lARC (1985) and the eompound was classified in LARC Monographs Supplement 7 (1987). 

7722-84-1 Chem. Abstr. Name Hydrogen peroxide lUPAC Systematic Name Hydrogen peroxide 

Synonyms Dihydrogen dioxide hydrogen dioxide hydrogen oxide hydroperoxide peroxide 

2 Structural and molecular formulae and relative molecular mass 

Description Colourless liquid with a bitter taste (Budavari, 1996 Lide, 997) Boiling-point 150.2°C (Lide, 1997) 

The 51V NMR spectra of bisW-mcthylhydroxylaminc) complexes show signals for methyl groups in orientations assignable to the possible positions available on the nitrogen [7,8]. Unfortunately, detailed kinetics experiments have not been carried out. It would be of interest to know whether the end-for-end rotation is correlated with methyl group reorientation. 

Minor products that are not shown in the figure include anionic V2L, WL2, and V2L3 (Section 8.1). 

Vanadium Chemistry, Biochemistry, Pharmacology and Practical Applications 

FIGURE 5.1 Distribution diagram showing the formation of vanadate and peroxovanadate species as a function of the concentration of hydrogen peroxide and of pH. Conditions for the simulation 2mmol/L total vanadate 0.1 tmol/L to 10 mmol/L total hydrogen peroxide 0.15 mol/L ionic strength with NaCl pH values, as indicated. The formation constants are from reference 11. 

Kinetics studies carried out in acetonitrile solutions provided a similar picture for the mechanism of peroxidation. Studies were carried out for three related tertiary amine-derived ligands /V,/V-bis(2-pyridylmcthyl)glycinc, /V-(2-pyridylmcthyl)imin-odiacetic acid, and AK2-amidomethyl [iminodiacetic acid [14], The kinetics studies 

RATE COEFFICIENTS FOR H202 + S20g AT 30 °C (Units l.mole and sec) 

Considerable sensitivity to contaminants was evident this leads to induction periods. The proposed mechanism involves the chain cycle (7)-(10) with some or all of the termination steps (11)-(14), viz. 

Both Gray and Symons drew attention to the more favourable parameters to be expected for (15) relative to (7). 

Ozone is reputedly formed in the reaction, and its origins have been investigated by isotopic labelling. The measurements reveal that two oxygen atoms come from the H2O2 and one from S20 . It is argued that peroxomono-sulphate, SO4 and 80 are all insignificant as reaction intermediates. 

Catalysis of the reaction by silver - and copper ions has been discussed. 

H2O2 commercially mainly available in concentrations of 35. 50 and 70% by weight 

Commercially hydrogen peroxide is available in a variety of concentrations. The most important are aqueous solutions with 35, 50 and 70% by weight H2O2. The following statistics refer to 100% hydrogen peroxide . 

World Europe USA/ Canada Latin Aus-America tralia Asia excl. Japan Japan Africa 

Hydrogen peroxide demand is currently increasing comparatively strongly, the production capacity having increased by 70% between 1979 and 1991 from 882 10- to 1492 - 10- t/a. Further growth is expected. 

Producers of hydrogen peroxide are e.g. Degussa, DuPont, EKA Nobel, EMC, Kemira, Mitsubishi Gas Chemical, Oxysynthese and Solvay-Interox. 

Carsky, M. Machacek, and R. Zahradnik, Coll. Czech. Chem. Comm., 1973, 38, 3067. 

The pyrolysis of H202 has been followed mass spectrometrically under conditions such that a steady-state concentration of the HOz radical, equal to or exceeding 10-2 times the concentration of the H202, would have been detected.34 The actual concentration of the H02 in the reacting mixture was found to be below the detection limit from which it was concluded that the simplest mechanism that provides an entirely quantitative description of the pyrolysis is  

The decomposition of hydrogen peroxide on the surface of cobalt sulphide has been studied.35 An investigation of the kinetics of the catalytic decomposition of hydrogen peroxide in the presence of the hydroxides of the Group III p, d, and f elements, and of the synthesis and properties of the peroxide products, has been carried out.36 The principal conclusions drawn from a combination of experimental and literature data were  

The hydrolysis of peroxodisulfate (produced by electrolytic oxidation of [HS04] at high current densities 

After treatment, deionized water is produced. The resins are reactivated by treatment with dilute H2SO4 and Na2C03 solutions respectively. Reverse osmosis at high pressures is also an important process in water purification, with cellulose acetate as the usual membrane the latter prevents the passage of dissolved solutes or insoluble impurities. The removal of nitrates is highlighted in Box 14.10. 

The purification of drinking water is a complicated industrial process. Water may be abundant on the Earth, 

Jiang and N.J.D. Graham (1997) Chemistry Industry p. 388 - Pre-polymerized inorganic coagulants for treating water and waste water . 

Delyannis and E.A. Delyannis (1979) Gmelin Elandbook of Inorganic Chemistry, O Water Desalting, Supplement Volume 1, System-Number 3, Springer-Verlag, Berlin. 

Ligand oxidation may be affected by molecular oxygen in the nickel complex of JV,iV-bis(2-mercaptoethyl)-2-methylthioethylamine) forming a sulfinato complex. O2 has been detected by ESR in dry Nafion perfluorinated membranes neutralized by Ti with the most likely source of the radical being electron transfer from The interaction of Pt(II) with creatinine generates a monomeric 

Reactions of [Fe(II)(NTA)] and [Fe(II)(EDDA)] (NTA = nitrilotriacetate EDDA = ethylenediamine-N,7V-diacetate) with H2O2 at neutral pH and in the 

The reaction is postulated to proceed through a hydroperoxide intermediate, [Cu2(L-H)(OOH)]. Experimental studies into Fenton and Fenton-like reactions point out that the detailed mechanism involves the formation of a low-valent cation.However, while [(H20)5Cr 02H ] decomposes directly to [Cr(H20)6] and OH in the presence of 0.1 M EtOH, 2-PrOH, and 2-BuOH, [(H20), Cu 02H ] reacts directly with the alcohols. In the absence of organic substrate, the latter complex does yield [Cu(H20) ] and OH. 

The oldest method for the preparation of H2O2 is reaction 16.28. The hydrolysis of peroxodisulfate (produced by electrolytic oxidation of [HS04] at high current densities using Pt electrodes) has also been an important route to H2O2 (equation 16.29). 

Ba02 b H2SO4 —— BaS04 -b H2O2 2[NH4][HS04] [NH4]2[S20g] 

Aspects of the chemistry of water have already been covered as follows  

Flocculator-clarifier units in a water treatment plant in Florida, US. Rick PoleyI Visuals Unlimited, Inc 

At many forms of carbon, the reduction of oxygen occurs by a 2 X 2e mechanism and the potentials for the reactions 

Electrolysis cells with a planar GDE have also been developed to produce H202 + Na0H solutions for the same market [34]. In such cells, the cathodes 

Good performance is also reported by a Spanish group for the combination of hydrogen and oxygen to give hydrogen peroxide in an alkaline fuel cell with an uncatalyzed carbon cathode. With a 1M KOH electrolyte, the current efficiency was close to 100% at a current density of 130 mA cm [37]. 

Since the publication of the Upjohn procedure in 1976, the use of N-methylmorpho-line N-oxide (NMO) based oxidants has become one of the standard methods for osmium-catalyzed dihydroxylations. However, NMO has not been fully appreciated in the asymmetric dihydroxylation for a long time since it was difficult to obtain high enantiomeric excess (ee). This drawback was significantly improved by slow addition of the alkene to the aqueous tert-BuOH reaction mixture, in which 97% ee was achieved with styrene [15]. 

Although hydrogen peroxide was one of the first stoichiometric oxidants used in osmium-catalyzed dihydroxylation [10a], it was not employed efficiently until recently. When hydrogen peroxide is used as a reoxidant for transition metal catalysts, a very common big disadvantage is that a large excess of H2O2 is required to compensate for the major unproductive peroxide decomposition to O2. 

By using a chiral Sharpless ligand, high enantioselectivities were obtained. In the flavin system, an increase in the addition time for alkene and H2O2 has a positive 

Nowadays, H2O2 is manufactured by the oxidation of 2-ethylanthraquinol (or a related alkyl derivative). The H2O2 formed is extracted into water and the organic product is reduced back to starting material. The process is summarized in the catalytic cycle in Fig. 16.9.  

Melting point/K Boiling point/K Af// (298K)/kJmol AfG (298K)/kJmol  

In aqueous solution, H2O2 is partially ionized (eq. 16.31), and in alkaline solution, is present as the [H02] ion. 

Tracer studies using show that in these redox reactions H2( 0)2 is converted to ( 0)2, confirming that no oxygen from the solvent (which is not labelled) is incorporated and the 0—0 bond is not broken. 

Marilyn Monroe in a scene from the movie Gentlemen Prefer Blondes 

Debbie Harry of Blondie ( Press Association Images. With permission.) 

The hydrogen peroxide oxidizes the melanin pigment inside the hairs to a colorless substance. Using very slightly alkaline hydrogen peroxide solution softens the cuticle and makes the hair more permeable to the H2O2, so the peroxide can reach the melanin. 

It depends on what you mean by unstable . It is stable with respect to the elements, hydrogen and oxygen, of which it s composed  

However, its decomposition products, water and (di) oxygen gas are even more stable, so its decomposition is exothermic. 

Invariably, the electrolysis medium contains sulphuric acid but the concentrations of both the acid and the chromium(iii) vary widely depending on the source of the solution from the oxidation of organic compounds, a typical solution may contain 1 mol dm Cr(iii) and 3 mol dm H2SO4, while from a plating bath to be conditioned, the chromium(iii) may be as dilute as 0.02 mol dm with 0.005 mol dm H2SO4 and a large excess of chromium(vi). 

The old design of cell is based on a lead-lined tank filled with aqueous sulphuric acid, and the spent chromic acid solution in porous ceramic pots was placed in the tank. The anodes were lead rods dipped into the ceramic pots and the cathode was the lining to the tank. The plant was obviously operated totally as a batch process. More modern and economic cells are now available they are often based on a bipolar filterpress concept with lead alloy anodes, steel cathodes and a Nafion proton-conducting membrane. The energy consumption and space-time yields with such cells are clearly likely to be superior. Moreover, such cells can be operated as continuous or semi-continuous processes. 

The cathodic reduction of dissolved oxygen to peroxide was first demonstrated by Traube in 1882. Commercialization of an electrolytic process for hydrogen peroxide has been retarded by several factors related to the complex electrochemistry of oxygen reduction, together with a poor understanding of the influence of electrode materials and cell design on the process efficiency. There has been a gradual awareness of the desirable factors for a successful process over the last 10-15 years, which include  

This chapter wiU summarize the recent developments in the area of osmium-catalyzed dihydroxylations, which bring this transformation closer to a green reaction . Hence, special emphasis is given to the use of new reoxidants and recycling of the osmium catalyst. 

Ever since the Upjohn procedure was pubhshed in 1976 the N-methyhnorpholine N-oxide-based procedure has become one of the standard methods for osmium-catalyzed dihydroxylations. However, in the asymmetric dihydroxylation NMO has not 

The flavin is an efficient electron-transfer mediator, but rather unstable. Several transition metal complexes, for instance vanadyl acetylacetonate, can also activate hydrogen peroxide and are capable of replacing the flavin in the dihydroxylation reaction [18]. 

---

Trimethylamine, CjH N, (CH3J3N. Colourless liquid with a strong fishy odour, miscible with water, m.p. — I24 C, b.p. 3-5°C. It occurs naturally in plants, herring brine, bone oil and urine. It reacts with hydrogen peroxide to give trimethylamine oxide and with ethylene oxide to give choline its commercial importance stems chiefly from this latter reaction. 

CH3CH2OHCH3. B.p. 82 C. Manufactured by hydrolysis of propene. Used in the production of acetone (propanone) by oxidation, for the preparation of esters (e.g. the ethanoate used as a solvent), amines (diisopropylamines, etc.), glycerol, hydrogen peroxide. The alcohol is used as an important solvent for many resins, aerosols, anti-freezes. U.S. production 1978 775 000 tonnes. 

In these methods, the sulfur oxides produced during combustion are, before detection, either converted into sulfuric acid by bubbling in a hydrogen peroxide-water solution or converted into sulfates. 

Quartz Tube Method the sample is burned in a quartz tube and a stream of purified air carries the combustion gases into a hydrogen peroxide solution. 

Fehrensen B, Luckhaus D and Quack M 1999 Mode selective stereomutation tunnelling in hydrogen peroxide isotopomers Chem. Phys. Lett. 300 312-20... 

Figure B2.5.22. Potential V along the minimum energy path for the stereonuitation of hydrogen peroxide. Adapted from [103].

Figure B2.5.23. Mode-specific stereomiitation tunnelling in hydrogen peroxide time-dependent probability...

Kuhn B, Rizzo T R, Luckhaus D, Quack M and Suhm M A 1999 A new six-dimensional analytical potential up to chemically significant energies for the electronic ground state of hydrogen peroxide J. Chem. Phys. Ill 2565-87... 

The product is a solid yellow hydrated oxide. If prepared by a method in the absence of water, a black anhydrous product is obtained. Germanium(II) oxide is stable in air at room temperature but is readily oxidised when heated in air or when treated at room temperature with, for example, nitric acid, hydrogen peroxide, or potassium manganate(VII). When heated in the absence of air it disproportionates at 800 K ... 

On the industrial scale oxygen is obtained by the fractional distillation of air. A common laboratory method for the preparation of oxygen is by the decomposition of hydrogen peroxide. H Oj, a reaction catalysed by manganese(IV) oxide ... 

Oxygen can be produced by certain reactions in solution, for example the oxidation of hydrogen peroxide by potassium manganate(VII) acidified with sulphuric acid ... 

In addition to the hydrides of formula HjX, oxygen forms the hydride H2O2, hydrogen peroxide, and sulphur forms a whole series of hydrides called sulphanes. These are yellow liquids which are thermodynamically unstable with respect to hydrogen sulphide and sulphur. 

Hydrogen peroxide is probably unique in the very large number of reactions by which it is formed. Some of these may be mentioned ... 

In many of the processes, it is believed that hydroxyl radicals, OH-, are formed and that some of these unite to form hydrogen peroxide ... 

In the laboratory, hydrogen peroxide can be prepared in dilute aqueous solution by adding barium peroxide to ice-cold dilute sulphuric acid ... 

Hydrogen peroxide is obtained in aqueous solution at the bottom of the column. This is a good method of preparation. 

On a large scale, hydrogen peroxide is produced by the electrolysis of ammonium hydrogensulphate, using a platinum anode and a lead cathode separated by a diaphragm. The essential process occurring is ... 

The ammonium hydrogensulphate is returned to the electrolytic cell. A process such as this yields an aqueous solution containing about 30% hydrogen peroxide. The solution can be further concentrated, yielding ultimately pure hydrogen peroxide, by fractional distillation but the heating of concentrated hydrogen peroxide solutions requires care (see below). 

The above method has now been largely replaced by a newer process, in which the substance 2-ethylanthraquinone is reduced by hydrogen in presence of a catalyst to 2-ethylanthraquinol when this substance is oxidised by air, hydrogen peroxide is formed and the original anthraquinone is recovered ... 

Pure hydrogen peroxide is a colourless, viscous liquid, m.p. 272.5 K, density l,4gcm . On heating at atmospheric pressure it decomposes before the boiling point is reached and a sudden increase of temperature may produce explosive decomposition, since the decomposition reaction is strongly exothermic ... 

This is a disproportionation reaction, and is strongly catalysed by light and by a wide variety of materials, including many metals (for example copper and iron) especially if these materials have a large surface area. Some of these can induce explosive decomposition. Pure hydrogen peroxide can be kept in glass vessels in the dark, or in stone jars or in vessels made of pure aluminium with a smooth surface. 

Rotation about the O—O bond is relatively easy. Hydrogen bonding causes even more association of liquid hydrogen peroxide than occurs in water. 

Because of the instability of pure and concentrated aqueous solutions of hydrogen peroxide, it is usually used in dilute solution. The concentration of such solutions is often expressed in terms of the volume of oxygen evolved when the solution decomposes ... 

Aqueous solutions of hydrogen peroxide decompose slowly the decomposition is catalysed by alkalis, by light and by heterogeneous catalysts, for example dust, platinum black and manganese... 

IV) oxide, the latter being used in the eommon laboratory preparation of oxygen from hydrogen peroxide (p. 260. ... 

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