The objectives of this investigation are (1) to make a small quantity of solid barium sulfate, using exact quantities of barium chloride and ammonium sulfate so that there is no excess of either reagent in the remaining solution, (2) to demonstrate that there is, in fact no excess barium or sulfate ions in the residual solution, and (3) in the event that objective (2) fails, to explain reasons for the failure. These objectives should be achieved using only the provided samples of barium chloride and ammonium sulfate: no extra supplies will be made available.
Barium sulfate has such a low solubility in water that it is undetectable in solution without the use of elaborate equipment or methods. For practical purposes, it is quite insoluble in water.
If two soluble ionic compounds, such as barium chloride and ammonium sulfate, are mixed, the ions rearrange to form a precipitate of barium sulfate, leaving ammonium and chloride ions in solution. The reaction is:
BaCl2 (aq) + (NH4)2SO4 (s) -------------> BaSO4 (s) + 2NH4Cl (aq)
[Where "(aq)" is shown, the ions are actually separated in the solution, and not joined together as shown in the equation.]
If a mixture is made with an excess either of barium chloride or ammonium sulfate, then, of course, barium ions or sulfate ions will appear in the residual solution.
Formula masses: BaCl2 (H2O)2 = 244.3, and (NH4)2SO4 = 132.1
This suggests that the two materials should be mixed in a ratio 244.3/132.1 = 1.849/1.000 by mass.
Alternatively, if solutions of exactly equal molarity were provided, then exactly equal volumes could be mixed to achieve perfect results.
Before the attempt is made to make the barium sulfate precipitate using perfect quantities of reagents, and then to test whether the mixture was indeed perfect, some planning is necessary. What equipment will be needed? (Remember that the only chemicals available will be the initially-provided samples of barium chloride and ammonium sulfate, and purified water.) Is the procedure starting with solid reagents or with solutions? How much of each solid or solution should be measured and used? Should the whole supply provided by used? How can the residual solution be tested to demonstrate the absence of both sulfate and barium ions?
Carry out the planned preparation, test the residues as planned, and explain why the results achieved were perfect, or, if there was indeed excess sulfate or barium, why it was there. It may be useful to compare results across a number of replications, that is, to either repeat the procedure yourself several times, or to compare results with those of other students.
An underlying objective of this activity, not revealed to students at the outset, is to demonstrate the inherent limitations of measurement in the laboratory.
If students weigh solids to 0.01 G, there will be an unavoidable variability of sample mass at the milligram level. Over a number of replications, this will lead to a statistical distribution of relatively large excess of sulfate, slight excess of sulfate, "perfect mixture", slight excess of barium, and relatively large excess of barium. ( I have personally enjoyed the experience of trying to explain to a very bright and highly motivated student who was basking in the glow of being the only one in a group to achieve a "perfect mixture" that her achievement was in fact no better than those of students with a excess of barium: her "success" was merely a random statistical outcome.)
For this reason, it is better if students do not have access to balances accurate to 0.001 G or better. However, even an excess of 1 mG of one reagent may be enough to demonstrate clearly the "error" in the measurement.
Samples of no more than 2.1 G of each solid are ample. If 1.00 G of ammonium sulfate is used, then 1.85 G of barium chloride dihydrate will be needed. With no less than 1.00 G of ammonium sulfate, students lose the benefit of measuring to three significant figures.
To test the residual solution, which may be separated either by filtration or decantation (part of student planning), two samples each of a couple of milliliters should be taken. Dissolving a small amount of each leftover solid as testing reagents, ammonium sulfate solution added to a sample of the residual solution will show a white precipitate with excess barium, or barium chloride solution added to the residual will show a white precipitate with excess sulfate. This needs to have been planned: students should have saved some of each solid to do the testing.
The density of each precipitate should provide a rough estimate of the concentration of excess barium or sulfate in the residual solution. Unless special efforts are made to ensure that all test solutions are equal in molarity, however, comparisons, between replications will be only rough.
The quantities of solids suggested above are not compatible with good microscale practice. For laboratories fully committed to microscale chemistry, this activity could be varied by requiring that students carry out the precipitation using (for example) 2.00 mL of each of 0.0500 M barium chloride and ammonium sulfate solutions. The emphasis then would be on the limits of accuracy of the measurements of volume. Any bias in the statistical distribution of errors might then reflect upon the accuracy with which the stock solutions were prepared.