Chloride barium

Data reported for the protolysis constant of water in barium chloride media are listed in Table 5.23. The data are solely for a temperature of 25 °C, with the data being reported by Fischer and Bye (1964). 

Robinson and Stokes (1959) listed osmotic coefficient data for BaCl2 solutions. Water activity data were derived from these data using Eq. (5.18). The dependence of the water activity data on the ionic strength in BaCl2 solutions can be described using Eq. (5.19), and the values derived for and 2 in Eq. (5.19) are 0.04270 0.00013 x 10 and 6.67 0.09 x 10 kg mol , respectively. 

T(°C) mx (reported) Medium / (mol kg- ) rrix (molkg- ) log w (reported) log w (accepted) References 

Note the slightly differing equation for the ion interaction coefficients because of the 1 2 electrolyte. Also, for this electrolyte, the ionic strength / is three times the medium concentration. 

After the solution has been filtered it is made acid with hydrochloric acid. 

8-inch crystallizing dish and cover, two 5-inch watch glasses, iron ring and ring stand. 

Procedure The hydrochloric acid obtained in Preparation 7 may be used for this preparation or 166 cc. of 12 N HC1 (2 F.W.) may be used. Calculate the weight of barium carbonate required to convert your sample of hydrochloric acid into barium chloride. 

Pour the acid into an 8-inch porcelain evaporating dish and add enough water to make 600 cc. of solution. Remove 10 cc. of this dilute acid and save it for use later in the preparation. Add the 

How could you show that your preparation is a hydrate  

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Sulphuric acid is added to the electrolyte and the hydroxylamine is formed as hydroxylammonium sulphate, (NH30H)2S04 [cf, (NHJjSOj. Addition of barium chloride then precipitates barium sulphate and hydroxylammonium chloride, (NH30H)C1, is obtained. 

Addition of barium chloride precipitates white barium sulphite ... 

The sulphate ion is detected by addition of barium chloride in the presence of hydrochloric acid a white precipitate of barium sulphate is obtained. The same test can be used to estimate sulphate, the barium sulphate being filtered off, dried and weighed. 

Oxidation of a sulphur compound with concentrated nitric acid yields sulphuric acid or a sulphate, which can be tested for with barium chloride. This can be used to estimate the sulphur. 

Hydrated chromium(III) sulphate exhibits different colours and different forms from which varying amounts of sulphate ion can be precipitated by barium chloride, due to the formation of sulphato-complexes. Chromium(III) sulphate can form alums. 

Hydrolysis of Potassium Ethyl Sulphate. Dissolve about i g. of the crystals in about 4 ml. of cold distilled water, and divide the solution into two portions, a) To one portion, add barium chloride solution. If pure potassium ethyl sulphate were used, no precipitate should now form, as barium ethyl sulphate is soluble in water. Actually however, almost all samples of potassium ethyl sulphate contain traces of potassium hydrogen sulphate formed by slight hydrolysis of the ethyl compound during the evaporation of its solution, and barium chloride almost invariably gives a faint precipitate of barium sulphate. b) To the second portion, add 2-3 drops of concentrated hydrochloric acid, and boil the mixture gently for about one minute. Cool, add distilled water if necessary until the solution has its former volume, and then add barium chloride as before. A markedly heavier precipitate of barium sulphate separates. The hydrolysis of the potassium ethyl sulphate is hastened considerably by the presence of the free acid Caustic alkalis have a similar, but not quite so rapid an effect. 

The precipitation of the barium sulphate must be performed with care, otherwise high results are obtained owing to occlusion of barium chloride in the barium sulphate. This is avoided by the following method, which has the further advantage that the tedious initial removal of the excess of nitric acid by evaporation is unnecessary. 

Dilute the solution (if necessary) to about 150 ml., add i ml. of concentrated hydrochloric acid, and then heat in a covered beaker almost to boiling. Meanwhile dissolve a small excess (o 4-o 5 g.) of barium chloride in about 50 ml. of water, bring to the boil, and then transfer to a clean 50 ml. burette. Now run this solution slowly drop hy drop into the sulphuric acid solution, keeping the latter steadily stirred throughout the addition. Then boil the solution gently in the... 

Barium Chloride. 10% solution, i.e., 100 g. of BaCl2,2H20 dissolved in water and the solution made up to i litre. 

Extract the acidified solution with ether, remove the ether and identify the phenol in the usual manner (see Section IV,114).f Add a few drops of bromine water or nitric acid to the aqueous layer and test for sulphate with barium chloride solution. 

Materials. Beside inorganic materials (eg, barium chloride/fluoride crystals, doped with 0.05% samarium), transparent thermoplasts are preferred for the PHB technique, eg, poly (methyl methacrylate) (PMAIA), polycarbonate, and polybutyral doped with small amounts of suitable organic dyes, organic pigments like phthalocyanines, 9-arninoacridine, 1,4-dihydroxyanthraquinone [81-64-1] (quinizarin) (1), and 2,3-dihydroporphyrin (chlorin) (2). 

These rosin-based sizes, whether paste, Hquid, or emulsions, can be used to size all grades of paper that are produced at acid pH. The latter include bleached or unbleached kraft Hnerboard and bag paper, bleached printing and writing grades, and cylinder board. In addition, polyaluminum compounds have been used in place of alum, most notably, polyaluminum chloride (48), which can reduce barium deposits where these have been a problem. The barium chloride by-product is more water-soluble than barium sulfate. Other polyaluminum compounds such as polyhydroxylated forms of alum and polyaluminum siHcosulfate have been evaluated as alum replacements. 

Sodium thiosulfate is determined by titration with standard iodine solution (37). Sulfate and sulfite are determined together by comparison of the turbidity produced when barium chloride is added after the iodine oxidation with the turbidity produced by a known quantity of sulfate iu the same volume of solution. The absence of sulfide is iadicated when the addition of alkaline lead acetate produces no color within one minute. 

Anhydrous BaCl2 exists as monoclinic or cubic crystals. The transition to cubic occurs at 925 °C. Barium chloride melts at 962°C the dihydrate, which has monoclinic crystals, loses water at 113 °C. Barium chloride, which is very hygroscopic, is sold in moisture-proof bags and steel or fiber dmms. 

Barium chloride finds use in the production of barium colors, such as the diazo dyes barium hthol ted [50867-36-2] and barium salt of Red Lake C [5160-02-1], a mordant for acid dyes and dying of textiles. Other uses include aluminum refining and boiler water treatment. 

Barium nitrate is prepared by reaction of BaCO and nitric acid, filtration and evaporative crystallization, or by dissolving sodium nitrate in a saturated solution of barium chloride, with subsequent precipitation of barium nitrate. The precipitate is centrifuged, washed, and dried. Barium nitrate is used in pyrotechnic green flares, tracer buUets, primers, and in detonators. These make use of its property of easy decomposition as well as its characteristic green flame. A small amount is used as a source of barium oxide in enamels. 

Barium nitrite [13465-94-6] Ba(N02)2, crystallines from aqueous solution as barium nitrite monohydrate [7787-38-4], Ba(N02)2 H2O, which has yellowish hexagonal crystals, sp gr 3.173, solubihty 54.8 g Ba(NO2)2/100 g H2O at 0°C, 319 g at 100°C. The monohydrate loses its water of crystallization at 116°C. Anhydrous barium nitrite, sp gr 3.234, melts at 267°C and decomposes at 270 °C into BaO, NO, and N2. Barium nitrite may be prepared by crystallization from a solution of equivalent quantities of barium chloride and sodium nitrite, by thermal decomposition of barium nitrate in an atmosphere of NO, or by treating barium hydroxide or barium carbonate with the gaseous oxidiation products of ammonia. It has been used in diazotization reactions. 

The more soluble forms of barium such as the carbonate, chloride, acetate, sulfide, oxide, and nitrate, tend to be more acutely toxic (50). Mean lethal doses for ingested barium chloride were 300—500 mg/kg in rats and 7—29 mg/kg in mice (47). 

The threshold of a toxic dose in adult humans is about 0.2—0.5 g Ba the lethal dose in untreated cases is 3—4 g Ba, LD q about 66 mg/kg (47). The fatal dose of barium chloride for humans is reported to be between 0.8 and 0.9 g (0.55—0.60 g of Ba) (50). However, for most of the acid-soluble salts of barium, doses greater than 1 g have been tolerated (51). Lethal doses are summarized in Table 5. Dusts of barium oxide are considered potential dermal and nasal irritants (52). 

Ethyl chloride can be dehydrochlorinated to ethylene using alcohoHc potash. Condensation of alcohol with ethyl chloride in this reaction also produces some diethyl ether. Heating to 625°C and subsequent contact with calcium oxide and water at 400—450°C gives ethyl alcohol as the chief product of decomposition. Ethyl chloride yields butane, ethylene, water, and a soHd of unknown composition when heated with metallic magnesium for about six hours in a sealed tube. Ethyl chloride forms regular crystals of a hydrate with water at 0°C (5). Dry ethyl chloride can be used in contact with most common metals in the absence of air up to 200°C. Its oxidation and hydrolysis are slow at ordinary temperatures. Ethyl chloride yields ethyl alcohol, acetaldehyde, and some ethylene in the presence of steam with various catalysts, eg, titanium dioxide and barium chloride. 

The sodium carbonate content may be deterrnined on the same sample after a slight excess of silver nitrate has been added. An excess of barium chloride solution is added and, after the barium carbonate has setded, it is filtered, washed, and decomposed by boiling with an excess of standard hydrochloric acid. The excess of acid is then titrated with standard sodium hydroxide solution, using methyl red as indicator, and the sodium carbonate content is calculated. 

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