In the Classroom edited by
Tested Demonstrations
Ed Vitz
Using Large Glass Cylinders To Demonstrate Chemical Reactions
Kutztown University Kutztown, PA 19530
submitted by:
Wobbe de Vos* Centre for Science and Mathematics Education, University of Utrecht, Utrecht, The Netherlands
checked by:
Kim Kostka Department of Chemistry, University of Wisconsin–Rock County, Janesville, WI 53546
Precipitation reactions are among the most spectacular phenomena in elementary chemistry. There is, however, a tenacious misconception that precipitation reactions should be demonstrated exclusively in test-tubes. As always, the choice of a suitable reaction vessel depends on the aim of the experiment. This article describes a simple demonstration of chemical reactions, in particular precipitation reactions, in large glass cylinders. The experiment fulfills pedagogic as well as aesthetic purposes. It takes only a few minutes to set up. Once started, it goes on for an hour or more. The standard procedure for demonstrating the precipitation of an insoluble salt in elementary chemistry courses is to prepare two suitable salt solutions and mix them. Admittedly, the immediateness of the reaction contributes to the spectacle but this standard procedure has some serious drawbacks. Students see only the reaction product and are not given sufficient time to observe its actual formation. When the reaction is shown as a test-tube experiment, the only process that students can keep up with is the settling of the precipitate, and inexperienced observers unfortunately tend to interpret this sometimes as a decrease in the amount of the substance (1). Moreover, the size of a test-tube requires observation from close by and does not allow large groups to see details. The experiment described here gives students in introductory chemistry courses an opportunity to observe the actual formation of a precipitate. It is also suitable for a chemistry department in any school, college, university, or other institute that wishes to present some of the fascinating aspects of chemistry to visitors from outside. It was shown, for instance, at the Faculty of Chemistry of Utrecht University on the occasion of a “Science Week for Children”. The theme of the Science Week was “Movement”. In other cases, the visitors may be potential students or their parents, or members of the general public. Usually such visitors, curious but perhaps feeling a bit out of place in a laboratory, stroll around, cast a quick glance at everything in sight, and show special interest in things that are colorful and that move. Another possible target group consists of professional chemists, including chemistry teachers, who are to be made once more aware of the fascinating nature of their discipline. And last but not least, aside from pedagogical considerations, regular students in a course in introductory chemistry dealing with salts and salt formation will also appreciate the performance. This paper presents the principle of the experiment, followed by a detailed description of a concrete example. It *Email:
[email protected].
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discusses some factors that influence the process and presents some alternative versions. Principle The experiment makes use of controlled dissolution of salts in water. A precipitate is formed from two suitable soluble salts in a large glass cylinder that is almost completely filled with water. First, some crystals or grains of one of the salts are sprinkled into the water. They sink to the bottom, forming a dilute solution en route. The other salt is made to dissolve in the water inside a filter placed in the top part of the cylinder. Its solution, being denser than water, moves down through the stem of the funnel. Where the two solutions meet, a precipitate begins to form. With the passage of time a cloud of insoluble salt is formed in the cylinder. The formation process is slow and it may take hours before it is complete. Although the main course of events is fixed, the details of the process are unpredictable and irreproducible and the result therefore often surprises the demonstrator as much as the other spectators. In an educational setting, students have the opportunity to discuss the reaction process not, as in a test-tube, after it has finished, but while it takes place in front of their eyes. Depending on the level of the course, students can focus on the formation of the precipitate as such, or try to understand some of the details of the process. Procedure The following instructions apply for a typical example of the experiment. Fill a 500- or 1000-mL glass cylinder (e.g., a graduated cylinder) with tap water almost to the top. Then sprinkle a spoonful (say, 10 g) of crystalline washing soda (sodium carbonate decahydrate) into the water. Place a funnel provided with a filter paper in the top part of the cylinder. The lower part of the filter paper, about 1 cm, should be immersed in the water. Add a few milliliters of water into the filter to remove any soda solution from the funnel. This will prevent the formation of a precipitate inside the filter or in the stem of the funnel. Then add a small amount, say 1 or 2 g, of hydrated iron(III) chloride. After about a minute, the yellowish brown iron chloride solution, which is heavier than water, begins to pass through the stem of the funnel and on down into the liquid. On its way it reacts with the dissolved soda, forming a cloudy brown precipitate of iron hydroxide. The soda crystals that have reached the bottom of the cylinder slowly form a layer of concentrated solution that may be up to one centimeter thick.
Journal of Chemical Education • Vol. 76 No. 4 April 1999 • JChemEd.chem.wisc.edu
In the Classroom
Due to its high density, this layer will not mix well with the cloudy precipitate and will remain colorless throughout the experiment. In this specific example the situation is sometimes complicated by the formation of small gas bubbles. Apparently, the acidic iron chloride solution releases some carbon dioxide gas from the dissolved soda. The gas bubbles rise to the surface, dragging along parts of the iron hydroxide cloud. At the same time other parts move downward through the solution. The result is a fascinating two-way traffic that may continue for a long time. Although the process is normally easy to observe, its visibility can be improved by a suitable background and lighting. A sheet of white paper is usually sufficient. The performance is more impressive when the cylinder is illuminated from below (e.g., by placing it on top of an overhead projector). Variables There are a few variables that strongly influence the result of this experiment. The initial concentration of soda and the supply rate of iron chloride are the most influential. Like most other hydrated salts, soda dissolves fairly rapidly in water. The total amount of dissolved soda in the upper part of the cylinder is largely determined by the amount of soda added, by the size of the soda crystals, and to some extent, by the temperature of the water. If small crystals or powder is used, 5 g will be sufficient. Grinding the soda and sprinkling it into the water will help to produce a more concentrated solution. The effect is that a denser cloud of iron hydroxide is formed. Obviously, more soda dissolves on its way to the bottom when the water in the cylinder is hot. The overall effect of using hot water, however, is rather complex; it does not merely accelerate the process but also affects it in other ways. Like soda, iron chloride normally dissolves rapidly but the filter paper slows down the dissolution rate. As passage through the filter is the ratedetermining step of the entire process, it provides a means to manipulate the duration of the performance. Slower precipitate formation results when a double filter is used. On the other hand, the reaction can be considerably accelerated during the experiment by adding a few milliliters of water to the filter or, using a pipet, carefully removing some of the water outside the funnel. In general, changing any of the variables hardly ever ruins the effect of the experiment. Instead, it leads to a different and, in its details, unpredictable outcome. Variations The advantages of the soda and iron chloride example are that it is inexpensive and, from both an environmental and a personal point of view, very safe. Normally, the whole content of the cylinder
can be poured down the washbasin and washed away with plenty of water. The reader is advised, however, to check the local regulations on this point. In principle, every precipitation from two soluble salts can be demonstrated in a glass cylinder. Among the salts that are relatively harmless from an ecological viewpoint are calcium chloride hexahydrate, several iron(II) and iron(III) salts, and, for example, sodium oxalate. If chemical waste containers are available many more reactions are possible. In the following examples, the first salt is to be added to the cylinder and the second salt is placed in the filter. Small quantities, say 5 g of the first salt and 1 or 2 g of the second salt are sufficient in a 1000-mL cylinder in all cases. The name and appearance of the resulting precipitate are given in parentheses. Potassium iodide and lead nitrate (lead iodide, glittering yellow crystals; shown in the sequence of images to the left and in color on the cover of this issue) Soda and copper sulfate (copper hydroxide, light blue) Ir on(III) chloride and potassium hexacyanoferrate(II) (Prussian blue, dark blue; shown in color on page 450) Sodium sulfide and copper sulfate (copper sulfide, black) Calcium chloride hexahydrate and silver nitrate (silver chloride, white, slowly turning purple or grey) Potassium iodide and silver nitrate (silver iodide, light yellow) Copper sulfate pentahydrate and potassium chromate (copper chromate, brown) Sodium hydroxide and aluminium nitrate (aluminium hydroxide, white)
The beautiful yellow lead iodide precipitate also features in other instructive experiments described in this Journal (1, 2). Silver and lead salts may require demineralized water instead of tap water in the cylinder. When the demonstration is finished, most heavy metal salts are best precipitated completely by adding excess soda. The upper layer can be decanted after the precipitate has settled, leaving only a small volume to be transferred to the waste container. An interesting variation consists of replacing the filter by a new one and then adding a third salt to the system. In the first two examples, for instance, sodium sulfide produces a dramatic effect. The experiment is, of course, not restricted to precipitation reactions. It includes, in principle, all reactions in aqueous solutions that produce a visible effect. A few examples of other reaction types are mentioned here as suggestions. Again, small quantities of chemicals are sufficient in all cases. The possibilities are almost endless.
JChemEd.chem.wisc.edu • Vol. 76 No. 4 April 1999 • Journal of Chemical Education
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In the Classroom
Complex formation. The red iron(III) thiocyanate complex is soluble in water. A beautiful red cloud can be obtained using potassium thiocyanide and iron(III) chloride. Some other transition metal salts also produce beautifully colored complexes. Gas evolution reactions. With a slightly acidic solution, say 0.1 M hydrochloric acid or diluted vinegar instead of water in the cylinder and soda in the filter a steady stream of carbon dioxide bubbles is produced. A few drops of phenolphthalein solution added to the filter embellish the effect. Redox reactions. In principle, all spontaneous redox reactions that produce a visible effect in aqueous solutions are suitable. Some examples are (first substance in cylinder, second substance in filter): Potassium iodide and copper sulfate or iron(III) chloride. The addition, in advance, of starch solution to the cylinder enhances the effect. Hydrogen peroxide and potassium iodide or (a few crystals only) potassium permanganate. The cylinder is filled with a very dilute (say 0.1%) solution of hydrogen peroxide. If this solution has been acidified with dilute sulfuric acid, an almost colorless reaction product can be obtained from
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the permanganate. It is surprising to see the trail of potassium permanganate disappear somewhere part way down the cylinder. Sodium sulfite and potassium permanganate. The permanganate reacts to form brown clouds of manganese(IV) oxide, MnO2.
Miscellaneous. In a very simple setup, the cylinder contains only water. A few crystals of potassium permanganate are added to the filter. The result is a seemingly ever-moving purple trail and a slowly growing layer of purple solution at the bottom of the cylinder. Another very simple variation starts with a cylinder full of water. No funnel is used. First a few drops of one salt solution are added, then a few drops of a solution of the second salt. The reader is invited to go to the laboratory and explore other possibilities. Literature Cited 1. De Vos, W.; Verdonk, A. H.. J. Chem. Educ. 1985, 62, 648– 649. 2. Cortel, A. J. Chem. Educ. 1997, 74, 297.
Journal of Chemical Education • Vol. 76 No. 4 April 1999 • JChemEd.chem.wisc.edu
