Standard Barium Solution

Barium Standard Solution (100 pig Ba in 1 mL) Dissolve 177.9 mg of barium chloride (BaCl2-2H20) in water in a 1000-mL volumetric flask, dilute to volume with water, and mix. 

Bacterial Alpha-Amylase Activity, 789 Bacterial (PC) Proteolytic Activity, 811 Baking Soda, 355 Balances and Weights, 729 Balsam Fir Oil, 156 Balsam Peru Oil, 38, 574 Barium Chloride TS, 850 Barium Diphenylamine Sulfonate TS, 850 Barium Hydroxide TS, (S 1)114 Barium Hydroxide, 0.2 N, 856 Barium Standard Solution, 849 Basil Oil, Comoros Type, 39, 574 Basil Oil, European Type, 39, 579 Basil Oil, Italian Type, 39 Basil Oil, Reunion Type, 39 Basil Oil Exotic, 39 Basil Oleoresin, 391, 392 Bay Leaf Oil, 217 Bay Oil, 40, 575 BCD, (S 1)15 Beeswax, White, 40 Beeswax, Yellow, 41 Beet Fiber, (S1 )45 Beet Sugar, 400, (S2)35 Benedict s Qualitative Reagent, 850, 851 Bentonite, 41 Benzaldehyde, 456, 607 Benzaldehyde Glyceryl Acetal, 456, 607, (S1)60... 

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. 

Hydrochloric acid and sulphuric acid are widely employed in the preparation of standard solutions of acids. Both of these are commercially available as concentrated solutions concentrated hydrochloric acid is about 10.5- 12M, and concentrated sulphuric acid is about 18M. By suitable dilution, solutions of any desired approximate concentration may be readily prepared. Hydrochloric acid is generally preferred, since most chlorides are soluble in water. Sulphuric acid forms insoluble salts with calcium and barium hydroxides for titration of hot liquids or for determinations which require boiling for some time with excess of acid, standard sulphuric acid is, however, preferable. Nitric acid is rarely employed, because it almost invariably contains a little nitrous acid, which has a destructive action upon many indicators. 

Discussion. The hydroxides of sodium, potassium, and barium are generally employed for the preparation of solutions of standard alkalis they are water-soluble strong bases. Solutions made from aqueous ammonia are undesirable, because they tend to lose ammonia, especially if the concentration exceeds 0.5M moreover, it is a weak base, and difficulties arise in titrations with weak acids (compare Section 10.15). Sodium hydroxide is most commonly used because of its cheapness. None of these solid hydroxides can be obtained pure, so that a standard solution cannot be prepared by dissolving a known weight in a definite volume of water. Both sodium hydroxide and potassium hydroxide are extremely hygroscopic a certain amount of alkali carbonate and water are always present. Exact results cannot be obtained in the presence of carbonate with some indicators, and it is therefore necessary to discuss methods for the preparation of carbonate-free alkali solutions. For many purposes sodium hydroxide (which contains 1-2 per cent of sodium carbonate) is sufficiently pure. 

This solution is widely employed, particularly for the titration of organic adds. Barium carbonate is insoluble, so that a clear solution is a carbonate-free strong alkali. The relative molecular mass of Ba(0H)2,8H20 is 315.50, but a standard solution cannot be prepared by direct weighing owing to the uncertainty of the... 

A slight excess of 10 per cent barium chloride solution is added to the hot solution to precipitate the carbonate as barium carbonate, and the excess of sodium hydroxide solution immediately determined, without filtering off the precipitate, by titration with the same standard acid phenolphthalein or thymol blue is used as indicator. If the volume of excess of sodium hydroxide solution added corresponds to timL of 1M sodium hydroxide and u mL 1M acid corresponds to the excess of the latter, then v — v = hydrogencarbonate, and V— v — v ) = carbonate. 

Pipette 25 mL barium ion solution (ca 0.01 M) into a 250 mL conical flask and dilute to about 100 mL with de-ionised water. Adjust the pH of the solution to 12 by the addition of 3-6 mL of 1M sodium hydroxide solution the pH must be checked with a pH meter as it must lie between 11.5 and 12.7. Add 50 mg of methyl thymol blue/potassium nitrate mixture [see Section 10.50(C)] and titrate with standard (0.01 M) EDTA solution until the colour changes from blue to grey. 

Discussion. The turbidity of a dilute barium sulphate suspension is difficult to reproduce it is therefore essential to adhere rigidly to the experimental procedure detailed below. The velocity of the precipitation, as well as the concentration of the reactants, must be controlled by adding (after all the other components are present) pure solid barium chloride of definite grain size. The rate of solution of the barium chloride controls the velocity of the reaction. Sodium chloride and hydrochloric acid are added before the precipitation in order to inhibit the growth of microcrystals of barium sulphate the optimum pH is maintained and minimises the effect of variable amounts of other electrolytes present in the sample upon the size of the suspended barium sulphate particles. A glycerol-ethanol solution helps to stabilise the turbidity. The reaction vessel is shaken gently in order to obtain a uniform particle size each vessel should be shaken at the same rate and the same number of times. The unknown must be treated exactly like the standard solution. The interval between the time of precipitation and measurement must be kept constant. 

Barium, D. of as chromate, (g) 474, (ti) 378 as sulphate, (g) 448 by EDTA, (ti) 324 Barium chloranilate 758 Barium hydroxide, standard solution.- 294 Barium sulphate sepn. from supersaturated soln., 421... 

C04-0034. While cleaning a laboratory, a technician discovers a large bottle containing a colorless solution. The bottle is labeled Ba (OH)2, but the molarity of the solution is not given. Concerned because of the toxicity of Ba ions, the technician titrates with a solution of hydrochloric acid standardized at 0.1374 M. A 25.00-mL sample of the barium hydroxide solution requires 36.72 mL of the HCl solution to reach the stoichiometric point. What is the concentration of Ba in the solution ... 

Carbon dioxide is not a common oxidation product in periodate work, but it does appear in the oxidation of ketoses,49 a-keto acids,14,39 and a-hydroxy acids,14 39 and it is often a product23 141 of overoxidation. Carbon dioxide analyses have been carried out using the Plantefol apparatus,49 the Warburg apparatus,14 23 and the Van Slyke-Neill mano-metric apparatus,39 and by absorption in standard sodium hydroxide141 followed by back-titration with acid. A most convenient method is the very old, barium hydroxide absorption scheme.16 The carbon dioxide is swept from the reaction mixture into a saturated, filtered barium hydroxide solution by means of a stream of pure nitrogen. The precipitated barium carbonate is filtered, dried, and weighed. This method is essentially a terminal assay. The manometric methods permit kinetic measurements, but involve use of much more complicated apparatus. 

Barium Sulphate To 1.0 g add a mixture of 3 ml 2M HNOs and 7 ml DW, heat on a water-bath for 5 minutes, filter, dilute the filtrate to 10 ml with DW, add 5 ml molybdovanadic reagent and allow to stand for 5 minutes. Any yellow colour produced is not more intense than that of a standard prepared simultaneously and in the same manner using 10 ml of phosphate standard solution (= 5 ppm P04) (50 ppm). 

The normality or molarity of the acid can be determined by titration with a standard solution of sodium hydroxide using a color indicator, or by potentio-metric titration using a pH meter or a millivoltmeter. The sulfate anion in dilute acid can be measured by precipitation with barium chloride or by ion chromatography. 

For solutions containing sulphuric acid only, direct titration with standard alkali, and measurement of the specific gravity, are possible as methods of estimation, provided that the process in either case is, if necessary, preceded by suitable dilution (see p. 165). Thermometric methods have also been suggested, depending on the rise in temperature when the acid is mixed with water, or when titrated with barium chloride solution.4 The water content of the concentrated acid may be determined by similar titration with oleum which has been standardised thermometrically by 80 per cent, sulphuric acid (see p. 147 ).5... 

Solutions of telluric acid give a quantitative precipitation of barium tellurate, BaTe04.3H20, on the addition of barium hydroxide solution, and the use of a standard barium hydroxide solution, followed by titration of the excess of alkali with a standard solution of oxalic acid, using phenolphthalein as indicator, forms a convenient process for... 

Step 1. Pipette exactly 5 mL of barium carrier into 100-mL beaker that contains 20 ml of deionized water. Add 5 drops of concentrated HC1. Pipette 1 ml of 226Ra standard solution and also pipette 1 mL of 228Ra standard solution into the beaker. Stir well. 

Estimation of Atmospheric Carbon Dioxide.—A convenient method is that of Pettenkofer,4 which consists in introducing a standard solution of barium hydroxide into a large bottle containing several litres of the air to be examined. The bottle is shaken from time to time to keep the sides moistened wit-h the solution, and after 5 or 6 hours the absorption of carbon dioxide may be regarded as complete. The baryta solution is decanted into a small stoppered bottle and allowed to stand until any suspended barium carbonate has settled. A portion of the clear liquid is then removed and titrated with dilute sulphuric acid, using phenol-phthalein as indicator. The diminution in alkalinity due to combination with carbonic acid is thus measured, and from the data obtained the percentage of carbon dioxide m the atmosphere may easily be calculated. 

Barium Hydroxide, 0.2 N [17.14 g Ba(OH)2 per 1000 mL] Dissolve about 36 g of barium hydroxide [Ba(0H)2-8H20] in 1 L of recently boiled and cooled water, and quickly filter the solution. Keep this solution in bottles with well-fitted rubber stoppers with a soda-lime tube attached to each bottle to protect the solution from carbon dioxide in the air. Standardize as follows Transfer quantitatively about 60 mL of 0.1 A hydrochloric acid, accurately measured, to a flask add 2 drops of Phenolphthalein TS and slowly titrate with the barium hydroxide solution, with constant stirring, until a permanent pink color is produced. Calculate the normality of the barium hydroxide solution and, if desired, adjust to exactly 0.2 A with freshly boiled and cooled water. 

Note Solutions of alkali hydroxides absorb carbon dioxide when exposed to air. Connect the buret used for titrations with barium hydroxide solution directly to the storage bottle, and provide the bottle with a soda-lime tube so that air entering must pass through this tube, which will absorb carbon dioxide. Frequently restandardize standard solutions of barium hydroxide. 

Indirect determinations of several types can be carried out. Sulfate has been determined by adding an excess of standard barium(II) solution and back-titrating the excess. By titrating the cations in moderately soluble precipitates, other ions can be determined indirectly. Thus sodium has been determined by titration of zinc in sodium zinc uranyl acetate, and phosphate by determination of magnesium in magnesium ammonium phosphate. Quantitative formation of tetracyano nickel-ate(II) has been used for the indirect determination of cyanide. ... 

Crystalline /3-D-glucose (4.5 g.) was added to a mixture of 500 cc. of barium bromide solution (containing 30 g. of BaBr2-2HjO and saturated with carbon dioxide), 10 cc. of bromine and 30 g. of barium carbonate. The average bromine concentration, as determined by titration with standard sodium thiosulfate, was 0.3S1 moles per liter, the average bromide content was 0.412 moles per liter and the concentration of free bromine was 0.0692 moles per liter. 

Standard solutions of bases are ordinarily prepared from solid sodium, potassium, and occasionally barium hydroxides. Again, eye protection should always be used when handling dilute solutions of these reagents. 

Sodium hydroxide is the most common base for preparing standard solutions, although potassium hydroxide and barium hydroxide are also encountered. Because none of these is obtainable in primary-standard purity, standardization of these solutions is required after preparation. 

The product is boiled with water, and the resulting precipitate of silver sulfide is collected by filtration, washed, dried, and weighed. Another sample of product is dissolved completely in concentrated nitric acid, any excess acid is removed by boiling, and the total silver ion is determined by titration with standard ammonium thiocyanate solution. Another sample of product is oxidized with concentrated nitric acid in a bomb tube, and after removal of silver and nitrate the sulfur is determined by precipitation as barium sulfate by addition of barium chloride solution. 

For most analyses, standard solutions are aqueous solutions containing several percent of a mineral acid such as HCl or nitric acid. When mixing different elements in acids, it is important to remember basic chemistry and solubility rules for inorganic compounds. The elements must be compatible with each other and soluble in the acid used so that no precipitation reactions occur. Such reactions would change the solution concentration of the elements involved in the reaction and make the standard useless. Combinations to be avoided are silver and HCl, barium and sulfuric acid, and similar... 

Silver chloride is an example of a soft crystal and is therefore not susceptible to the problems seen when precipitating hard crystals such as barium sulfate and calcium oxalate. This, however, does not mean that the opalescence obtained in a test or standard solution is independent of the operational parameters of the precipitation procedure. But it can be considered a more reproducible and rugged determination compared to hard crystal precipitations, and the steps in the procedure contributing to loss of reproducibility are more easily standardized. The most obvious difference is that a fairly reproducible test procedure can be obtained without the use of a seeded standard, as is the case in 6.3. Calcium and 6.12. Sulfates. 

Step one in the test is to prepare a crystal seed by adding 1 ml of a 250 g/1 solution of barium chloride R to 1.5 ml of sulfate standard solution (10 ppm SO4) Rl. Shake and allow to stand for 1 min. Then, 15 ml of the solution to be examined and 0.5 ml of acetic acid R are combined. Prepare a standard in the same manner using 15 ml of sulfate standard solution (10 ppm SO4) Rl instead of the solution to be examined. After 5 min, any opalescence in the test solution should not be more intense than that in the standard. 

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