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Mountainninja

Joined 14 August 2010

Works edit

Mostly working on Chemistry articles.

Expands Trimer (chemistry) with 10 examples and edit related links and articles from sources in English Wikipedia, German Wikipedia and Polish Wikipedia -- Mountainninja (talk) 23:15, 4 May 2014 (UTC)

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1 edit

1.1 edit

Elegance of the peroxide processing of ammonia in making hydrazine, hydroxylamine, hyponitrous acid ?

Not co-producing salt(s) in this process aside, its role of acetone fixating and protecting the ammoniacal nitrogen atom so a hydrogen-nitrogen bond can participate in substitution and an carbon-nitrogen double bond addition reaction. The 3,3-disubstitutedoxaziridine can be used for: 1condensation reaction with more ammonia, as in the peroxide process for hydrazine. 2hydrolysis, giving hydroxylamine Ketazine is formed in the process, 3Yet more hydrogen peroxide can react with, to gives hyponitrous acid. As this ketazine-lysis takes two oxygen atoms, can even dioxygen molecule from purified air be a reactant, thus saving hydrogen peroxide ? I doubt. Thus, all these three compounds can be viewed essentially as oxidation products of ammonia by hydrogen peroxide.

2 edit

Salts of various sulfur oxoacids can be made from sulfur acid exhaust during the contact process. Bubbling excess sulfur dioxide through a solution of suitable base, such as sodium hydroxide, produces sodium bisulfite.

SO
2 + NaOH → NaHSO
3

Evaporating a solution of sodium bisulfite saturated with sulfur dioxide leaves the product sodium metabisulfite behind:

2 HSO
3− ⇌ H2O + S2O52−

Sodium dithionate is prepared from the reactions between sodium bisulfite with either manganese dioxide, silver oxide, chlorine or solution of sodium hypochlorite. All these reactions generate byproducts that are needed to be separated and not regenerated in the reaction. The following reactions would deviate from the traditional routes: Disproportionation of solid sodium metabisulfite generates sodium dithionite whose melting point is 52°C and sodium dithionate 190°C. By careful hitting of

2 S2O52− → S2O42− + S2O62−

Or, sodium thiosulfate is produced on an industrial scale chiefly from liquid waste products of sodium sulfide or sulfur dye manufacture. Oxidizing thiosulfate by adding one oxygen atom at a time to each thiosulfate can generate the desired salt of sulfur oxoacid from dithionite, metabisulfite and dithionate in individual pots of reactions.

Pot 1: S2O32− + [O] → S2O42−
Pot 2: S2O32− + 2[O] → S2O52−
Pot 3: S2O32− + 3[O] → S2O62−

Can heteropoly acids such as catalyze these oxidations of solid sodium thiosulfate? For reference, silicotungstic acid has been commercialized for the oxidation of ethylene to acetic acid:

C2H4 + O2 → CH3CO2H

Analogously,

Pot 1: S2O32− + H2O2 (Phosphomolybdic acid) → S2O42− + H2O
Pot 2: S2O32− + O2 (silicotungstic acid) → S2O52−
Pot 3: S2O32− + O3 (silicotungstic acid) → S2O62− or S2O32− + H2O2 + O2 (Phosphomolybdic acid) → S2O42− + H2O

3 edit

Hydroformylation of carbon dioxide CO
2 in supercritical or STP conditions yields glyoxylic acid: this synthesis creates a pathway to convert unused industrial gas or desulfurized and cleaned flue gas into biodegradable products or merchandise. Hydrogen, carbon monoxide and CO
2 can also come from plasma gasification(PG) of municipal solid waste(MSW). In essence, the proportion between burning syngas for electricity generation and manufacturing biodegradable products can be balanced to different needs of industry, commerce and communities.
Regarding PG and its commercialization, unused syngas and exhaust of energy generation can be converted to biodegradable products such as glyoxylic acid and the carbon dioxide and steam therein can be pressured into supercritical conditions to leach metal ions from the slag in PG before vitrification.
Instead of FIRST vitrifying the slag byproduct, could metal be first recovered before the vitrification ? For example, ADEQUATE AMOUNT of scrap aluminum and aluminium dross sacrificed by adding to the waste UPSTREAM so elements below aluminum on the reactivity series are reduced to metal when all inorganic materials are vitrified. The metal being denser than slag, can be recovered by centrifugal separation. Obviously, aluminum can reduce silica to produce silicon and the purpose of adding aluminum is to recover some other metals, say copper; so only adequate amount of scrap aluminum is needed. -- Mountainninja (talk) 23:32, 26 September 2014 (UTC)

Ideas for thought: Fractional crystallization (chemistry) Betterton-Kroll process Betts electrolytic process Parkes process -- Mountainninja (talk) 15:20, 19 February 2015 (UTC)

3.1 edit

Supercritical carbon dioxide and water treatment on lignosulfonates, tall oil and turpentines from sulfate and Kraft processes may yield syngas. As unused or exhaustive carbon dioxide, carbon monoxide, hydrogen gases are synthesized to glyoxylic acid, the acid can also be added into supercritical fluid to increase its hydrogen content. At any rate syngas is produced and thus synthesized into organic products thru common gas to liquids industrial procedures.

Pulp is a lignocellulosic fibrous material prepared by chemically or mechanically separating cellulose fibres from wood, fiber crops or waste paper. The wood fiber sources required for pulping are "45% sawmill residue, 21% logs and chips, and 34% recycled paper" (Canada, 2014).[1] Pulp is one of the most abundant raw materials worldwide.

[1]

4 edit

Sulfamic acid melts at 205 °C before decomposing at higher temperatures to H
2O, SO
3, SO
3, and N
2.[2] Sulfamic acid is produced industrially by treating urea with a mixture of sulfur trioxide and sulfuric acid (or oleum). The conversion is conducted in two stages:

OC(NH2)2 + SO
3 → OC(NH2)(NHSO3H)
OC(NH2)(NHSO3H) + H
2SO
4 → CO
2 + 2 H
3NSO
3

In this way, approximately 96,000 tons were produced in 1995.[3]

[4]

Can decomposing sulfamic acid with catalyst be a novel preparation of hydroxylamine ?

H
3NSO
3 → NH
2OH + SO
2

The would-be byproduct sulfur dioxide is recycled for preparing more sulfamic acid, sulfur trioxide, sulfuric acid or oleum. In essence this novel synthesis eliminates the use of chlorine and ammonia compounds and links more closely the industries for sulfuric acid and urea preparation.

5 edit

Chloralkali process, Solvay process, alkali metals, alkaline earth metals manufactures

Can carbon dioxide be pumped into sodium hydroxide product from the Chloralkali process to precipitate all formed sodium bicarbonate due to common ion effect in the brine ?

Besides manufacturing sodium bicarbonate and sodium carbonate, desired concentration of pure sodium hydroxide is re-manufactured by dissolving appropriate amount of sodium oxide into water, with heat recovery apparatus available for further heating of more sodium bicarbonate. Note that pure alkaline earth metal carbonates can be manufactured by reacting sodium carbonate with alkaline metal chlorides.

Above 50 °C, sodium bicarbonate gradually decomposes into sodium carbonate, water and carbon dioxide. The conversion is fast at 200 °C:[5]

2 NaHCO3 → Na2CO3 + H2O + CO2

Most bicarbonates undergo this dehydration reaction. Further heating converts the carbonate into the oxide (at over 850 °C):[5]

Na2CO3 → Na2O + CO2

[6]

Pure sodium chloride can then be manufactured by reacting the pure sodium hydroxide with the pure hydrogen chloride from reacting hydrogen and chlorine products from the chloralkali process.

In the chlor-alkali industry, brine (mixture of sodium chloride and water) solution is electrolyzed producing chlorine (Cl2), sodium hydroxide, and hydrogen (H2). The pure chlorine gas can be combined with hydrogen to produce hydrogen chloride in the presence of UV light.

Cl2(g) + H2(g) → 2 HCl(g)

As the reaction is exothermic, the installation is called an HCl oven or HCl burner. The resulting hydrogen chloride gas is absorbed in deionized water, resulting in chemically pure hydrochloric acid. This reaction can give a very pure product, e.g. for use in the food industry.

[7]

This pure sodium chloride can then be used for electrolysis for sodium metal. Thus the Chloralkali process is both synthesis and refining processes. Note that no ammonia is used, in contrast to usage in the Solvay proces, and no mercury is used in contrast to usage in the Castner–Kellner process. Pure alkaline earth metal chlorides can also be manufactured in similar processes for electrolysis for alkaline earth metals.

Enjoying rather specialized applications, only about 100,000 tonnes of metallic sodium are produced annually.[8]

Metallic sodium was first produced commercially in 1855 by carbothermal reduction of sodium carbonate at 1100 °C,[citation needed] in what is known as the Deville process:[9][10][11]

Na2CO3 + 2 C → 2 Na + 3 CO

A related process based on the reduction of sodium hydroxide was developed in 1886.[9]

Sodium is now produced commercially through the electrolysis of molten sodium chloride, based on a process patented in 1924.[12][13] This is done in a Downs cell in which the NaCl is mixed with calcium chloride to lower the melting point below 700 °C. As calcium is less likely to be reduced based on its standard electrode potential than sodium, less calcium will be co-produced at the cathode. This method is less expensive than the previous Castner process of electrolyzing sodium hydroxide.

[14]

6 edit

The synthesis of adipic acid; one of the two reactants used in nylon manufacture, produces nitrogen oxides including nitric oxides[15][16] This might become a major commercial source, but will require the removal of higher oxides of nitrogen and organic impurities. Currently much of the gas is decomposed before release for environmental protection.

[17]

The production of adipic acid is linked to emissions of N
2O,[18] a potent greenhouse gas and cause of stratospheric ozone depletion. At adipic acid producers DuPont and Rhodia (now Invista and Solvay, respectively), processes have been implemented to catalytically convert the nitrous oxide to innocuous products:[19]

2 N2O → 2 N2 + O2

[20]

6.1 edit

Nitrogen oxides byproducts from nylon manufacturing can all be converted to nitrogen monoxide. It is then reacted with adipic acid to form N-oxoadipamide, that is later hydrogenated to hexamethylenediamine.

N
3O + NO
2 → 3NO
3NO + HO
2C(CH
2)
4CO
2H → ONOC(CH
2)
4CONO
ONOC(CH
2)
4CONO + 8H
2 → H
2N(CH
2)
6NH
2 + 4H
2O

In essence, nitrogen oxides byproducts can be reprocessed and then reacted with cyclohexene or cyclohexane to cyclohexanone and/or cyclohexanol; nitrogen gas and water are byproducts instead of nitrogen oxides.

6.2 edit

NO
2 is used to generate anhydrous metal nitrates from the oxides:[21]

MO + 3 NO
2 → M(NO
3)
2 + NO

How about reacting the organic impurities and oxides of nitrogen with alkali metal and/or alkaline metal carbonates into inorganic nitrates, carbonate esters and compounds of nitrogen that do not normally form salt, such as nitrous oxide. Because of the polar and nonpolar nature of these liquids and solids, they can be separated by distillation, fractional distillation and recrystallization. Both group of ionic and covalent compounds can be commercial sources of chemicals. Using alkali metal carbonates show two advantages. Explosive risks need to be taken by simply combining the organic impurities and oxides of nitrogen into nitrate esters; carbonate esters do not pose such risks. Alkali metal carbonates can be manufactured from chloralkali process, in which carbon dioxide is to react with the sodium hydroxide product to form sodium carbonate. Thus the recycling and reuse of nylon manufacture byproducts, chloralkali process, alkali metal carbonates manufacture, alkali metals manufacture can be combined into one greater manufacturing industry.

6.3 edit

Nitrous oxide can still be used for such as oxidation of benzene to phenol, or oxidation of ammonia to hydroxylamine, which is useful in the Beckmann rearrangement of cyclohexanone oxime to caprolactam in the manufacturing nylon 6.

C
6H
6 + N
2O → C
6H
5OH + N
2

NH
3 + N
2O → NH
2OH + N
2

Hydrogenation of nitric oxide dimer to hyponitrous acid, and more hydrogenation to hydroxylamine.

H
2 + N
2O
2 → H
2N
2O
2

2 H
2 + H
2N
2O
2 → 2 NH
2OH

Nitric oxide reacts in free-radical addition to hydrogen sulfide with elimination of water to tetrasulfur tetranitride.

NO + H
2S - H
2O → SN

NS + NO + H
2S - H
2O → S
2N
2

and so on...

Nitric oxide(NO) or NOx reacts on commercial metal sulfide ores to yield metal oxides and tetrasulfur tetranitride S
4N
4, which can be decomposed into nitrogen gas and sulfur. S4N4 itself can be an reagent introducing sulfur into chemical compounds. Also NO is one of the NOx. So in essence, the reaction converts three environmental problematic compounds to two other more valuable products.

Ideas for thought: Herz reaction Sulfilimine Sulfamide

6.4 edit

Dinitrogen monoxide replaces silver oxide as a mild oxidizing agent. For example, as N
2O is a by-product of manufacturing nylon which uses hydroxlamine (NH
2OH), N
2O can be used to oxidize ammonia to yield hydroxylamine and dinitrogen which is recycled. NH
3 + N
2O → NH
2OH + N
2

7 edit

Condensation and dehydration of acrolein and acetonitrile afford pyridine. C
3H
4O + C
2H
3N - H
2O → C
5H
5N

Acetonitrile is a byproduct from the manufacture of acrylonitrile. Most is combusted to support the intended process but an estimated several thousand tons are retained for the above-mentioned applications.[22]

Acrolein is prepared industrially by oxidation of propene. The process uses air as the source of oxygen and requires metal oxides as heterogeneous catalysts:[23] CH
2CHCH
3 + O
2 → CH
2CHCHO + H
2O

Useful template edit

Template:Clear

Reference edit

  1. ^ "Pulp (paper)". en.wikipedia. Retrieved 14 February 2016.
  2. ^ Yoshikubo, K.; Suzuki, M. (2000). "Sulfamic Acid and Sulfamates". Kirk-Othmer Encyclopedia of Chemical Technology. doi:10.1002/0471238961.1921120625151908.a01. ISBN 0471238961.
  3. ^ A. Metzger "Sulfamic Acid" in Ullmann's Encyclopedia of Industrial Chemistry, 2012, Wily-VCH, Weinheim. doi:10.1002/14356007.a25_439
  4. ^ "Sulfamic acid". en.wikipedia. Retrieved 15 February 2014.
  5. ^ a b "Decomposition of Carbonates". General Chemistry Online.
  6. ^ "Sodium bicarbonate". en.wikipedia. Retrieved 17 May 2015.
  7. ^ "Hydrogen chloride". en.wikipedia. Retrieved 17 May 2015.
  8. ^ "Sodium". en.wikipedia. Retrieved 29 July 2018.
  9. ^ a b Eggeman, Tim; Updated By Staff (2007). "Kirk-Othmer Encyclopedia of Chemical Technology". John Wiley & Sons. doi:10.1002/0471238961.1915040912051311.a01.pub3. ISBN 0-471-23896-1. {{cite journal}}: |chapter= ignored (help); Cite journal requires |journal= (help)
  10. ^ Oesper, R. E.; Lemay, P. (1950). "Henri Sainte-Claire Deville, 1818–1881". Chymia. 3: 205–221. doi:10.2307/27757153. JSTOR 27757153.
  11. ^ Banks, Alton (1990). "Sodium". Journal of Chemical Education. 67 (12): 1046. Bibcode:1990JChEd..67.1046B. doi:10.1021/ed067p1046.
  12. ^ Pauling, Linus, General Chemistry, 1970 ed., Dover Publications
  13. ^ "Los Alamos National Laboratory – Sodium". Retrieved 2007-06-08.
  14. ^ "Sodium". en.wikipedia. Retrieved 17 May 2015.
  15. ^ Reimer R. A.; Slaten C. S.; Seapan M.; Lower M. W.; Tomlinson P. E.; (1994). "Abatement of N2O emissions produced in the adipic acid industry". Environmental progress. 13 (2): 134–137. doi:10.1002/ep.670130217.{{cite journal}}: CS1 maint: extra punctuation (link) CS1 maint: multiple names: authors list (link)
  16. ^ Shimizu, A.; Tanaka, K. and Fujimori, M. (2000). "Abatement of N2O emissions produced in the adipic acid industry". Chemosphere – Global Change Science. 2 (3–4): 425–434. doi:10.1016/S1465-9972(00)00024-6.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  17. ^ "Nitrous oxide". en.wikipedia. Retrieved 17 May 2015.
  18. ^ US EPA. "U.S. Greenhouse Gas Inventory Report, Chapter 4. Industrial Processes" (PDF). Retrieved 2013-11-29.
  19. ^ Reimer, R. A.; Slaten, C. S.; Seapan, M.; Koch, T. A. and Triner, V. G. (2000). "Adipic Acid Industry — N2O Abatement". Non-CO2 Greenhouse Gases: Scientific Understanding, Control and Implementation. Netherlands: Springer. pp. 347–358. doi:10.1007/978-94-015-9343-4_56. ISBN 978-94-015-9343-4.{{cite book}}: CS1 maint: multiple names: authors list (link)
  20. ^ "Adipic acid". en.wikipedia. Retrieved 14 March 2016.
  21. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  22. ^ "Acetonitrile". en.wikipedia. Retrieved July 29, 2018.
  23. ^ "Acrolein". en.wikipedia. Retrieved July 29, 2018.
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