Barium chloride has the formula, BaCl2 and is an ionic chemical compound. It is one of the most important water-soluble salts of barium-containing compounds. Like other barium salts, it is toxic and imparts a yellow-green coloration to a flame. It is also hygroscopic. Barium chloride was the by-product of the discovery of radium by Madame Curie (1898). When refining radium, the final separation resulted in barium chloride and radium chloride.
Đang xem: So3 + bacl2 = baso3 + cl
BaCl2 crystallizes in both the cubic “fluorite” and “lead chloride” crystal structures, both of which accommodate the preference of the large Ba2+ ion for coordination numbers greater than six. There are four structure types known (Fig. 2.21).
Fe2P type: a = 4.059 Å, b = c = 7.083 Å, α = 120°, β = γ = 90°, cell volume = 176.31 Å, cell occupancy is shown in Table 2.68.
TABLE 2.68. Cell occupany is:
PbCl2 type, SG = mP24: a = 4.604 Å, b = 8.613 Å, c = 11.901 Å, α = 92.6°, β = γ = 90°, cell volume = 471.48 Å3, cell occupancy is shown in Table 2.69.
PbCl2 type, SG = 0P12: a = 4.732 Å, b = 7.872 Å, c = 9.425 Å, α = β = γ = 90°, cell volume = 651.08 Å3, cell occupancy is shown in Table 2.70.
In aqueous solution, BaCl2 behaves as a simple salt. In water, it is a 1:2 electrolyte and the solution exhibits a neutral pH. Electrolysis of molten barium chloride produces barium metal. This can be used for obtaining barium salts for commercial uses. Barium chloride reacts with sulfate ion to produce a thick white, insoluble precipitate of barium sulfate:
Ba2+(aq) + SO42−(aq) ⇒ BaSO4 (solid)
Oxalate effects a similar reaction:
BaCl2 (aq) + Na2C2O4 (aq) ⇒ BaC2O4 (solid) + 2NaCl (aq)
Barium chloride can be prepared from barium hydroxide or barium carbonate, the latter being found naturally as the mineral “Witherite”. These basic salts react to give hydrated barium chloride. On an industrial scale, it is prepared via a two-step process from the mineral “Baryte”:
BaSO4 + 4C ⇒ BaS + 4CO (gas)
This first step requires high temperatures. The second step requires fusion of the reactants:
BaS + CaCl2 ⇒ BaCl2 + CaS
The BaCl2 is then be leached out from the mixture with water. From water solutions of barium chloride, the dihydrate can then be crystallized as white crystals, BaCl2·2H2O, which are colorless, translucent rhomboidal tablets or lamellae. The dihydrate is stable in the air at room temperature, but loses one-half of its water above 55 °C (131 °F), and becomes anhydrous at 121 °C (250 F). As a cheap, soluble salt of barium, barium chloride finds wide application in the laboratory. It is commonly used as a test for sulfate ion. In industry, barium chloride is mainly used chiefly in the synthesis of pigments and in the manufacture of rodenticides and pharmaceuticals. It is also used:
In the purification of brine solution in caustic chlorine plants,
In the manufacture of heat treatment salts,
For case hardening of steel,
In the manufacture of pigments,
In the manufacture of other barium salts such as barium hydroxide,
In fireworks to give a bright green color. However, its toxicity limits its applicability,
As a flux in the manufacture of magnesium metal,
For making color kinescope glass ceramics,
In wastewater treatment,
For the production of PVC stabilizers, oil lubricants, barium chromate and barium fluoride,
For stimulating the heart and other muscles for medicinal purposes,
For softening water,
For manufacturing of other barium salts used as pesticides, pigments, boiler detergent,
As a mordant in dyeing and printing textiles,
In the manufacture of caustic soda, polymers, and stabilizers.
Its solubility in water in grams per 100 ml of water is shown in Table 2.71.
J.M. WagesJr, in Encyclopedia of Analytical Science (Second Edition), 2005
The direct amplification of silver can be carried out by precipitation with chromate. Ag2CrO4 is filtered and then treated with barium chloride for the precipitation of both BaCrO4 and AgCl. These precipitates are treated with Ag+ and thus 1 mol of Ag2CrO4 and 2 mol of AgCl are formed.
The process can be repeated, and on each treatment with BaCl2 an amount of silver chloride equivalent to twice the original silver chromate is produced; after n cycles 2(n+1) moles AgCl can be found.
This procedure can also be applied for the indirect determination of CrO42−, by measuring the AgCl obtained.
One other indirect amplification, similar to those mentioned, is the amplification of phosphate by precipitation of silver phosphate and its conversion to silver chromate.
A particular case of the alternative precipitation method, which Weisz called cyclic amplification by fixation of both ions of an amplifiable compound, requires that the substance to be amplified be subjected to a series of stoichiometric reactions in order to enhance its mass.
Scheme 5 for the amplification of hexacyanoferrate(II), via precipitation of the silver salt and formation of silver chloride and Prussian Blue by reaction with iron chloride, shows that it is a particular case of the alternative precipitation method in which the amplification takes place by a cyclic procedure.
Its CAS number is 7787-39-5. It occurs as white, monoclinic crystals with a density of 4.44 g/cm3. Its solubility in water is low (0.001125 g/100 ml) and it is insoluble in ethanol. When heated, it decomposes at 480 °C to the oxide and SO2 gas.
It has been used as a weighting material on oil-drilling rigs to prevent “blow out” during drilling operations. This salt is used for such purposes, presumably because of its ability to absorb SO2 and H2S gases trapped in the rocks as drilling proceeds. The demand for barium sulfite is low.
Very little scientific studies of barium sulfite have ensued in the literature. The physical properties of this salt remain unspecified. One paper presents the formation of spherulites of SrSO3 and BaSO3, synthesized in agar–agar gels (0–40 °C). The reactants were sodium sulfite and the chloride of the respective metals. Each spherulite consists of fibrous crystals which are arranged minutely in a radial manner from the center. A linear relation was recognized between (mean diameter)2 and reaction time in the same manner as the CaSO3 0.5H2O spherulite reported. The slopes of the lines, namely the growth rates of the spherulites, were greater in the order of calcium sulfite > strontium sulfite > barium sulfite. The ratio of (mean diameter)2/time was dependent upon the concentration of the agar–agar gel (0.5%–2.0%) and the reaction temperature (0–40 °C); the ratios decreased linearly with an increase of the gel concentration and increased with an increase of temperature.
Barium sulfite is available in limited quantities, commercially.
|LD50||1||gm for 70 kg||p.o.(barium chloride)||Machata (1988)|
|LD50||630||mg/kg||p.o. (barium carbonate)||Lewis and Sweet (1984)|
|LD50||118||mg/kg||p.o. (barium chloride)||Lewis and Sweet (1984)|
|LD50||921||mg/kg||p.o. (barium acetate)||Lewis and Sweet (1984)|
|LD50||8.12||mg/kg||i.v. (barium chloride)||Syed and Hosain (1972)|
|LD50||8.49||mg/kg||i.v. (barium nitrate)||Syed and Hosain (1972)|
|LD50||11.32||mg/kg||i.v. (barium acetate)||Syed and Hosain (1972)|
|LD50||55||mg/kg||i.v. (barium chloride)||(http://rais.ornl.gov/tox/profiles/barium_f_V1.shtml%23t21)|
|LD50||96||mg/kg||i.v. (barium acetate)||(http://rais.ornl.gov/tox/profiles/barium_f_V1.shtml%23t21)|
Accurately weigh 1 g of the finely powdered rock material into a PTFE dish, moisten with water and add in succession, 15 ml of concentrated nitric acid, 2 ml of concentrated hydrochloric acid and 10 ml of concentrated hydrofluoric acid. Cover the dish and set aside overnight. Remove the cover and evaporate to dryness on a steam bath. If the rock powder is completely decomposed, add a few ml of water and 10 ml of concentrated nitric acid and again evaporate to dryness. Repeat this last operation once more. If the decomposition is not complete after the first evaporation, add 5 ml of concentrated perchloric acid (60% w/v) and, if organic matter is present, 100 mg of vanadium pentoxide. In addition if the material is low in calcium, add not less than 100 mg of calcium chloride. (There should be present at least two equivalents of calcium for each equivalent of sulphur). Evaporate to dryness, and, if any organic or black mineral particles remain, evaporate with further portions of perchloric acid until the decomposition is complete.
After the final evaporation, add 3 ml of concentrated hydrochloric acid and about 50 ml of water. Digest on the steam bath for 15 minutes with occasional stirring, then cool. Collect the residue on a small close-textured filter paper and wash well with warm water. Discard the residue. Transfer the filtrate to a 250-ml separating funnel. Dilute to a volume of about 100 ml, and extract iron and titanium with successive 50-ral portions of cupferron solution until the extract is no longer brown. Wash the solution twice with 50-ml portions of chloroform, and finally with 50 ml of light petroleum. Run off the aqueous layer into a 400-ml beaker and wash the light petroleum twice with 10-ml portions of water. Dilute the combined aqueous layer and washings to about 200 ml, and filter if necessary.
Heat the solution to boiling and add a slight excess of a hot barium chloride solution. Allow the solution to stand for an hour on a steam bath and then set the beaker aside overnight. Collect the precipitated barium sulphate on a small close-textured filter paper and wash with successive small quantities of cold water. Transfer the paper to a weighed platinum crucible, dry, ignite over the full flame of a Bunsen burner and weigh as BaSO4.
After ignition, the barium sulphate residue should be perfectly white. Blank values are usually insignificant.