How can Ba(OH)2 concentrations be determined using H2SO4, HCl, and Na2SO4 solutions?

A lake has been contaminated by 177,000 tons of barium hydroxide. After finding the volume of the lake, I have determined the molarity of the barium hydroxide to be 0.05 M. How can I assay the barium hydroxide by making up solutions using the following reagents: sulfuric acid, hydrochloric acid, and sodium sulfate at twice the molarity of the barium hydroxide in the lake water?
Brittney 12/31/98

If your lake had appreciable amounts of sulfate in it, much of the barium from the barium hydroxide will be locked up in bottom sediments as BaSO4. If the lake was brackish (with high concentrations of chloride), a bit of the BaSO4 will dissolve (and the methods that rely on BaSO4 formation won't work well).
  1. Titration of Ba(OH)2 with HCl. Barium hydroxide is a strong base*. The lake water can be titrated with an 0.1000 M HCl solution according to

    Ba(OH)2(aq) + 2 HCl(aq) rightarrow BaCl2(aq) + 2 H2O(ell)

    Phenolphthalein* can be used as an indicator* for this titration.

    Color change for phenolphthalein as a function of pH.
    color change for phenolphthalein
    pH scale

    It will change from pink to clear when the hydroxide in the lake water is completely neutralized. The endpoint can also be determined by monitoring pH over the course of the titration with a pH meter. A 20 mL sample of your lake water should require about 20 mL of 0.1000 M HCl solution.

    This method really determines the total ability of the lake water to neutralize acids- not the actual Ba(OH)2 concentration. Other basic substances like carbonate and bicarbonate in the water may not be distinguished from the barium hydroxide:

    2 HCl(aq) + CO32-(aq) rightarrow CO2(g) + H2O(ell)+ 2 Cl-(aq)
    HCl(aq) + HCO32-(aq) rightarrow CO2(g) + H2O(ell) + Cl-(aq)

    Try using phenolphthalein as an indicator in one titration, and using methyl orange as an indicator in another. If there is a significant difference between the endpoints of the two titrations then the carbonate or bicarbonate concentration in the lake water is high and should be accounted for. See Standard Methods for experimental details.

    Color change for methylorange as a function of pH.
    color change for methylorange
    pH scale

  2. Titration of Ba(OH)2 with H2SO4. The reaction between sulfuric acid and barium hydroxide is both a precipitation and a neutralization reaction:

    Ba(OH)2(aq) + H2SO4(aq) rightarrow BaSO4(s) + 2 H2O(ell)

    The endpoint of the titration can be detected using any of the methods outlined above. A 20 mL sample of lake water will require about 10 mL of 0.1000 M H2SO4 solution for complete neutralization.

  3. Gravimetric method using Na2SO4. Sulfate ion can be used to precipitate barium from the solution:

    Ba(OH)2(aq) + Na2SO4(aq) rightarrow BaSO4(s) + 2 NaOH(aq)

    The BaSO4 can be filtered, dried, and weighed. The mass of BaSO4 obtained can then be related to the molarity of Ba(OH)2 in the lake water by a simple calculation.

    The BaSO4 crystals are very fine and difficult to collect by filtration. The crystals can be made larger by heating the solution to a temperature just below boiling for an hour or two. The precipitate is filtered, washed, dried, and then fired to complete dryness in a muffle furnace. It's a very tedious analysis but done with care it is highly accurate. Given the very high concentration of Ba(OH)2 in your sample, gravimetric analysis should work well here.

    Carbonates may be a problem again. As the highly alkaline Ba(OH)2 solution is exposed to air, it will absorb carbon dioxide. The carbon dioxide absorbed ultimately precipitates as BaCO3. When you weigh the BaSO4, you may also be weighing some BaCO3.

    A better gravimetric method for determining barium in natural waters is given in Official Methods, protocol 33.100. The barium is precipitated as barium dichromate (BaCr2O7) and fired in a muffle furnace to constant weight BaCrO4.

  4. Titration of Na2SO4 with Ba(OH)2. Titration is much quicker than the gravimetric technique and almost as accurate. The endpoint can be detected using an adsorption indicator- a substance that forms a colored complex on the surface of a precipitate that changes color near the endpoint.

    If the lakewater sample is neutralized, adjusted to pH 3.5, and diluted by a known factor prior to this analysis, alizarin red S can be used as an adsorption indicator. Alizarin red S is yellow at this pH, but it forms a pink complex on the surface of BaSO4 crystals. It will only be absorbed when the surface of the particles have a positive charge (when excess barium is present). If a measured amount of Na2SO4 solution is titrated with the Ba(OH)2 unknown, the endpoint is signalled by a sharp color change from yellow to pink. To make the color change more intense, the surface area of the BaSO4 precipitate should be as large as possible. Adding methanol to the sulfate solution befor starting the titration keeps the BaSO4 from forming large crystals.

    Official Methods protocol 33.102 recommends precipitating the barium as barium dichromate. The precipitate is filtered off, washed, redissolved in HCl, and treated with 10% KI solution. The dichromate from the precipitate oxidizes the I- to I2 and I3-. Titrating with sodium thiosulfate (Na2S2O3) solution gives the amount of iodine, and so the amount of dichromate and the amount of barium.

References

Standard Methods for the Examination of Waste and Wastewater, APHA-AWWA-WPCF, 18th ed., 1994.
Official Methods of Analysis of the Association of Official Analytical Chemists, AOAC, 13th ed., 1980.

Author: Fred Senese senese@antoine.frostburg.edu


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