Trichloride, nitrogen

Determine nitrogen by transferring the solution to a distilling flask, adding excess sodium hydroxide, and distilling the ammonia into standard acid. 

For the determination of bromine, acidify with dilute nitric acid a sample treated as just described, oxidize the excess bisulfite with 1 per cent potassium permanganate solution, and determine the halogen by the Volhard or another suitable method. The bromine nitrogen ratio usually varies from 1.97 to 2.02. 

Ethereal solutions of dibromoamine are stable at — 70°C. for one hour or longer but decompose rather rapidly at 0°C. or at higher temperatures. The solution has a straw-yellow color and a sharp irritating odor. 

Nitrogen trichloride was fitst prepared by Dulong1 in 1811 by the action of chlorine gas upon a solution of an ammonium salt. It may also be obtained by the action of hypochlorous acid upon an ammonium salt.2 Hentschel3 prepared it in benzene solution by extracting the aqueous reaction mixture of calcium hypochlorite and ammonium chloride with benzene. 

The pure anhydrous compound may be prepared by passing chlorine into a serpentine tube containing a solution of ammonium sulfate, carrying the volatile compound 

Nitrogen trichloride forms a bright yellow oily liquid with a density of 1.653 g cm  

It is very volatile [5-9] and has a very unpleasant odor sometimes described as nauseous or pungent [3,6,7,10-12]. The vapors irritate the eyes. In its frozen state, it exists as a rhombohedral crystalline solid. The melting point is reported between —40 

Nitrogen trichloride is an extremely sensitive substance and it explodes violently with the slightest impact or friction [7,12], The explosion itself is characterized by 

Nitrogen trichloride decomposes to nitrogen and chlorine. Decomposition may be represented by the following equation [9]  

Warning—all equipment that may come into contact with nitrogen trichloride must be washed by alkali in order to clean it from grease. Not doing so will lead to explosion when contact between the two occurs. 

Griffiths and Norrish152 have studied the chlorine-photosensitized decomposition of NC13 by following the pressure increase from the overall reaction, viz. 

Although an aging effect was observed with some reaction cells, the rate of decomposition in the zero-order region was independent of the area of the il- 

While claiming no finality for their mechanism, the authors pointed out that it accounted satisfactorily for the zero-order dependence on NC13 concentration, the intensity exponent of unity and many other characteristics of the reaction. However, a serious defect lies in the apparent absence of any effect of pressure on the rate of diffusion of NC14 to the reaction vessel surface. Such an effect would be expected under their conditions and it would yield a dependence of the reaction rate on total pressure opposite to that observed. 

One further point deserves comment. Although reaction (65) is exothermic (see below) and probably has a low activation energy, it is inherently extremely complex. Indeed it would be surprising if the formation of the N N bond and the elimination of the chlorine atom and molecules occurred in a single step. Thus this reaction probably proceeds through a number of steps, e.g. 

There is unfortunately no information on N-Cl bond energies in the chlorine analogues of hydrazine. However, reaction (66) is likely to be highly exothermic, and some of the excess energy may reside in the N2C13 species. Collisionally 

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When chlorine gas is in excess a highly explosive substance, nitrogen trichloride, NCI3, is formed ... 

Nitrogen trichloride Ammonia, As, hydrogen sulflde, nitrogen dioxide, organic matter, ozone, phosphine, phosphorus, KCN, KOH, Se, dibutyl ether... 

The concentration of inorganic and organic chloramines in pool water is controlled by superchlorination or shock treatment. Because chloramines are decomposed by sunlight, their effects are more noticeable in indoor pools or spas. Nitrogen trichloride, the primary volatile chloramine, is a strong irritant similar to chlorine. Its effect is noticeable at >0.5 mg/m (>0.1 ppm) (73). The concentration of NCl depends on the extent of ventilation and typically varies from 0.2 to 0.5 mg /m (0.04 to 0.1 ppm) (74). 

The pH value is usually maintained above 9 to avoid formation of nitrogen trichloride. At lower pH values, aqueous solutions react with chlorine to form cyanogen chloride (52). 

Heating must be terminated when two-thirds of the nitrate has decomposed since explosive nitrogen trichloride may be formed from traces of ammonium chloride impurity. 

For the chlor-alkali industry, an emergency preparedness and response plan is mandatory for potential uncontrolled chlorine and other releases. Carbon tetrachloride is sometimes used to scrub nitrogen trichloride (formed in the process) and to maintain its levels below 4% to avoid fire and explosion. Substitutes for carbon tetrachloride may have to be used, as the use of carbon tetrachloride may be banned in the near future due to its carcinogenicity. 

A liquid chlorine tank was kept cool by a refrigeration system that used CFCs. In 1976 the local management decided to use ammonia instead. Management w as unaware that ammonia and chlorine react to form explosive nitrogen trichloride. Some of the armnonia leaked into the chlorine, and the nitrogen trichloride that was formed exploded in a pipeline... 

A potentially liazardous by-product of chlorine manufacture is nitrogen trichloride, which is unstable and liigldy explosive. It can be formed from a combination of chlorine with nitrogen compounds in the brine feed, ammonia in... 

Write the formulas for the following compounds and give the weight of one mole of each carbon disulfide, sulfur hexafluoride, nitrogen trichloride, osmium tetroxide. 

Chlorine Nitride. (Nitrogen Trichloride, Trichloramine, or Stickstofftrichlorid in Ger). 

Noyes, Nitrogen Trichloride in Inorganic Syntheses Vol I , McGraw-Hill (1939), 65 7) Korenschoof N.V., BritP631327 (1949) ... 

Nitrogen trichloride, which is highly endothermic, is also extremely unstable. All the accidents that are described below and involve this compound are linked to the catalysis of its decomposition. It is the same when it is in the pure state and exposed to a physical agent such as shock, light or ultrasound. 

Nitrogen trichloride detonates in the presence of nitrogen monoxide, dioxide or trioxide and also in contact with ammonia and potassium cyanide. 

Chlorine detonates in contact with hydrogen if the mixture is submitted to ultraviolet radiation or if a catalyst is present and if the concentration of chlorine is between 5 and 89% in volume. In the dark such an explosion occurs in the presence of 2% of nitrogen trichloride. This is to be expected, given the instability of nitrogen trichloride. 

Hydrochloric acid catalyses the explosive decomposition of nitrogen trichloride. Chlorine dioxide... 

With ammonium salts, sodium hypochlorite gives nitrogen trichloride, which detonates spontaneously. 

Arsenic catalyses the explosive decomposition of nitrogen trichloride. 

Finally, there was also spontaneous incandescence of a calcium car-bide/selenium mixture and the formation of an explosive compound in the presence of an alkali or alkaline earth amide. In addition, selenium catalyses the explosive decomposition of nitrogen trichloride. 

When ammonium chloride is used, there is a risk that nitrogen trichloride forms, which is explosive. However, this is not a dangerous reaction involving a nitrile group. 

During an attempt at destroying benzyl cyanide residues with sodium hypochlorite, a detonation was caused that was probabiy due to the formation of nitrogen trichloride. However, it might be asked if it was not due to the nitrile group oxidation by the hypochlorite present. 

The interaction of a sodium or calcium hypochlorite or phosphorus penta-chloride with urea causes a violent detonation, that has been put down to the decomposition of the nitrogen trichloride formed. 

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