1910
JOURNAL OF CHEMICAL EDUCATION
NOVEMBER, 1929
LECTURE DEMONSTRATIONS IN ANALYTICAL CHEMISTRY* I. M. KOLIHOFF, UNIVERSITYOF MINNESOTA. MINNEAPOLIS, MINNESOTA The elementary education in any subject requires in the first place the teaching of facts. If the student has an inquiring mind the mere knowledge of facts may not interest or satisfy him, but, as a rule, he has to acquire them. The success of a course to a great extent depends upon the way in which the material is presented. For example, if one teaches a child geography and geology, you can show him on the map the geographical conditions of a country, tell him about the population, economic conditions, imports and exports, but what a dead knowledge if the facts are offered without trying to appeal to his imagination! The student is likely to forget all the facts unless you try to tell him in a more lively way and show him by pictures how life and nature are in the different countries. He observes the facts and they will stick in his mind. Now let us consider the education in chemistry. The situation is somewhat different, because we are not satisfied by only observing facts; we want to interpret them. The highest achievement of a course is not that the student acquires a certain knowledge in a short time, but the success will depend upon his understanding of facts, their internal correlation and interpretation. From this point of view lecture experiments may have a great educational value, especially for the mediocre students, and I suppose we agree that the great majority belongs to this group. As a rule a student has a tendency to write down as much as possible of what the instructor says and he will study his notes in the same way as he does his work in English, history, etc. Simple, successful, practical demonstrations will impress him in different ways. What first was considered as a vague theory suddenly lives before his eyes; i t will stimulate his desire to understand the nature of things and implant facts in his mind. If we consider more particularly the significance of lecture experiments for a course in quantitative inorganic chemistry, we have to keep in mind that the teaching is more or less an introduction to the practical work. If the latter is carefully selected, i t may be questioned whether lecture experiments do not lose their significance,because the student has an opportunity to apply his knowledge in the laboratory himself. Still I believe that the demonstration of a few simple lecture experiments in connection with the description of facts or the understanding of some theoretical fundamentals is not only very useful, but necessary to the success of the course. Therefore, I have selected a few demonstration experiments in the field of gravi, metric and volumetric analyses.
* Presented before the Symposium on "Lecture Experimentation." Division of Chemical Education, at the 77th Meeting of the American Chemical Society, Columbus, Ohio, April 29 to May 3, 1929.
General Quantitative Analysis 1. Loss of Liquid by Entrainment in a Gas Evolution.-Moisten a filter paper with potassium ferrocyanide solution and lay i t on a glass plate. Place a small beaker on i t containing from one to two grams calcium carbonate and a few cubic centimeters of a solution of femc chloride. Add to this mixture about 20 cc. of 4 to 6 N hyd~ochloricadd. The filter paper is soon covered with blue spots. 2. A Flame from a Bunsen Gas Burner Contains Sulfur Trioxide.Place one gram of pure sodium chloride in a porcelain crucible and heat it for 45 to 60 minutes over a full Bunsen flame. After cooling dissolve the salt in water and show the presence of sulfate. Moreover, i t can be shown that by placing the crucible in a close fitting hole of an asbestos plate, the content is protected from the combustion gases. 3. Washing of a Filter.-Moisten a filter paper with potassium dicbromate and wash with water, until the wash water does not give a test for dichromate when treated with sulfuric acid and hydrogen peroxide in the presence of ether. Take the filter paper from the funnel, open it, and show some yellow spots. Then transfer it to a porcelain dish and add a little sulfuric acid, ether, and hydrogen peroxide. A blue color shows that dicbromate is still present. If the washing is continued, until the filtrate does not react with sulfuric acid and diphenyl carbazide, it can be shown that it takes a long time until all of the dichromate is washed out (quantitatively). If the test for dichromate in the filtrate is negative or very weak, the presence of dichromate on the filter can still be shown by the diphenyl carbazide test. 4. Solubility Product.-Shake some freshly prepared lead iodide with water, until the solution is nearly saturated and filter. Divide filtrate into two parts. To one part add some lead nitrate solution. To the other part add some potassium iodide solution. In both cases a precipitate of lead iodide appears. a small excess of 5. Colloidal Nature of Some hecipitates.-Add potassium chromate solution to about 0.01 N lead nitrate and divide the yellow suspension into two parts. (a) Add to one part a few drops of an aluminum chloride solution. The flocculation is immediately observed. (Compare with the original solution.) (b) Filter the original solution; the filtrate is cloudy. Filter the solution containing the aluminum salt; the filtrate is clear.
Volumetric Analysis-Acidmetry,
Alkalimetry, and Indicators
6. Color Change Interval of an Indicator: e. g., Phenol red with Buffers of pH from 6.8-8.0.-If there are two glass wedges of equal size
available, one may be filled with acid indicator solution, the other with an alkaline solution containing the same concentration of indicator. By placing the two wedges upon each other the transition colors can be observed, and the principle of the measurement of pH with di-color indicators without the use of buffer solutions can be demonstrated. 7. Classification of Indicators.-Depending upon the ionization constants of the acid or basic dyestuffs, the color change interval may be found in neutral, acid, or alkaline medium. Take some test tubes with the same volume of tapwater in each. (a) To two tubes add some phenolphthalein. (b) To three tubes add some phenol red or neutral red. (c) To two tubes add some methyl orange. The indicator in a is present in the acid form-addition of acid does not change the color; alkali is required to produce a red-violet color. Therefore, the indicator has its color-change interval in alkaline medium. The indicator in c is present in the alkaline form; addition of alkali does not change its color. Acid is required to change color to that of the acid form. The color change interval is on the acid side of neutral point. The indicator in b has an intermediate color. To one of the tubes add some alkali, to the other some acid. Phenol red has its color-change interval around the neutral point. The author prefers the use of tapwater in this experiment to that of distilled water, because the latter is always acid on account of carbon dioxide and has a pH of 5.7 or less. Tapwater is a buffer system and as a rule has a pH of about 7.5 + 0.5. 8. Necessary Excess of Reagent.-Demonstration of the titration error and the choice of the proper indicator for different cases. Use highgrade salts and distilled water for these demonstrations. Neutralization of a strong acid vith a strong base. Prepare about one liter of approximately 0.1 N sodium chloride. To two beakers each containing 100 cc. of 0.1 N sodium chloride add ten drops of a 0.1% solution of methyl orange. Add to one of the two beakers 0.1 to 0.2 cc. of 0.1 N hydrochloric acid and observe the ditference in color. Repeat, but use methyl red as an indicator. Addition of less than 0.1 cc. of 0.1 N acid changes the color completely to that of the acid form. Repeat, but use phenolphthalein as an indicator. Add to one of the two beakers 0.1 cc. of 0.1 N sodium hydroxide. A faint, violet color appears. If the solvent contains too much carbon dioxide, the color fades fairly rapidly. The demonstration shows the "jumpwise" change of pH a t both sides of the equivalence point in the neutralization of a strong acid with a strong base or vice versa. Neutralization of weuk acid with strong base. Prepare about one liter (approximately) 0.1 N sodium acetate. To two beakers each containing 10 cc. of 0.1 N sodium acetate, add a few drops of phenolphthalein. As a rule the solution does not change in color. Add to one of the two solutions
VOL.6, No. 11
SYMPOSIUM ON LECTURE EWERID~~NTATION
1913
0.1 cc. of 0.1 N sodium hydrcxide. The color turns red-violet. (Compare the test with NaCI.) Instead of using phenolphthalein as an indicator use phenol red. If the water is carbon-dioxide free, the solution will acquire the alkaline color; as a rule, however, an intermediate color of the indicator will be observed. In the latter case addition of 0.1 cc. of 0.1 N NaOH to 100 cc., changes the color completely to red. On the other hand addition of 0.1 cc. of 0.1 N HCI does not change the color completely to that of acid form; approximately 1 cc. is required. With methyl red as indicator: The solution is alkaline to this indicator. Approximately 4 cc. of 0.1 N HCl for 100 cc. of 0.1 N sodium acetate are required to produce a slight differencein color from the original solution. With more acid the color changes gradually to pink and red. (Color-change interval of methyl red.) A similar experiment can be made with bromocresol green, or methyl orange or bromophenol blue, but here it is better to start with 25 cc. of 0.1 N sodium acetate, as much acid is required to cause a change in color. Neutralization of weak base with strong acid. Prepare about one liter of (approximately) 0.1 N ammonium chloride. To two beakers containing 100 cc. of 0.1 N ammonium chloride add-afewdrops methyl red. An intermediate color is observed. Add to m e of the solutions 0.1 cc:.of 0;l N HCI; the color turns pure red. Add toithe other one. 0.1 cc. of 0.1 N NaOH; the color turns pure yellow :. Repeat the experiment, using methyl orange or bromophenol blue as an indicator. The ammonium chloride solution is alkaline to these indicators. The addition of 0.1 to 0.2 cc. of 0.1 N HCl causes a distinct color change. Use phenol red as an indicator. The solutions are acid to this indicator. About 1 cc. of 0.1 N NaOH is required to change the color of one of the two solutions from yellow to orange-red. Use phenolphthalein as an indicator. More than 5 cc. of 0.1 N sodium hydroxide are required to turn the solution a faint pink. With more NaOH the color changes gradually to a deeper red-violet. (Color change interval of phenolphthalein.) 9. Slow Neutralization of Carbon Dioxide to Bicarbonate.-Add a few drops of phenolphthalein to a dilute carbon dioxide solution. Then titrate with 0.1 N sodium hydroxide. The red color fades after standing a short time. With a little more sodium hydroxide the red color again appears, but fades on standing. A permanent end-point is obtained after all of the carbon dioxide has been transformed into bicarbonate. Precipitation Reactions 10. Titration to the Clearpoint.-Titrate 100 cc. of 0.025 N potassium iodide with 0.1 N silver nitrate carefully. About 1to 2% before the equivalence point the silver iodide starts to flocculate; the supernatant liquid is milky (cloudy). Continue the titration carefully drop by drop with shaking. At the equivalence point the supernatant liquid is no longer cloudy,
hut perfectly clear. This experiment shows, moreover, the colloidal nature of silver iodide and its peptization by iodide ions. In connection with this experiment discuss the precision titration of silver according to Gay-Lussac. Mulder. 11. Silver Thiocyanate Adsorbs Silver Ions.-Titrate 25 cc. of 0.1 N silver nitrate according to Volhard with 0.1 N thiocyanate, using ferric alum as an indicator. The first color change to pink-brown occurs before the equivalence point is reached (about 0.8%). On shaking, the color disappears again. A permanent end-point is reached a t the equivalence point, where all of the adsorbed silver has reacted with the thiocyanate in the solution. 12. Demonstration of Behavior of Adsorption Indicators.-To 25 cc. of approximately 0.025 N silver nitrate add a few drops of a 0.1% fluorescein solution. The indicator assumes the same color as in pure water. Add a little 0.1 N sodium chloride to the silver solution; the color turns red immediately. With an excess of chloride the color changes again to greenyellow. This experiment shows that the indicator is not sensitive to silver ions in the solution, hut reacts in a very delicate way with silver ions adsorbed on the silver chloride lattice. Oxidation-Reduction Reactions
13. Diphenylamine as Oxidation-Reduction Indicator.-Prepare approximately '/40 M potassium ferrocyanide, 1% ferricyanide, '/zo molar zinc sulfate in water and 1%diphenylamine in strong sulfuric acid. To 25 cc. ferrocyanide solution add two drops ferricyanide, 5 cc. of 4 N sulfuric acid, and two drops diphenylamine. The indicator does not change the color of the solution. In another beaker add to about 50 cc. of the zinc sulfate solution, 5 cc. of 4 N sulfuric acid and one to two drops indicator; the liquid remains colorless. Add the zinc sulfate solution slowly to the ferrocyanide solution. After the complete formation of K2Zn3(Fe(CN)6)2, the color suddenly turns blue-violet. With a little excess ferrocyanide the color disappears again. The color change is reversible. The experiment demonstrates that the indicator is not specific for the reacting ions, hut behaves as an oxidation-reduction indicator. The ferrocyanide solution, containing a little ferricyanide does not have an oxidation potential high enough to oxidize the indicator via diphenylhenzidine to the blue quinoid form. By removing the ferrocyanide ions with zinc the oxidation potential increases and a t the equivalence point where practically no ferrocyanide ions are left, the oxidation potential is so large that the indicator is transformed into the blue oxidation product. Addition of ferrocyanide again lowers the potential and the color disappears. 14. Influence of Complex Ion Formation on Oxidation Potential.(a) Prepare an approximately 0.1 M ferric chloride solution, about 20%
phosphoric acid, 4 N hydrochloric acid, and 1 N potassium iodide. Take three beakers and add 50 cc. oi ferric chloride to each and 2 cc. hydrochloric acid. Add to two of the solutions 20 cc. phosphoric acid. The yellow color disappears due to a complex formation between the femc ions and phosphoric acid. Add to the original solution and one of the solutions containing phosphoric acid 10 cc. of potassium iodide. The former is colored dark brown from the iodide which is formed in the reaction: Fe+++ I- 5=s Fe++ '1% In; the iron solution containing phosphoric acid does not change. As a result of the complex formation, the oxidation potential of the femc iron solution has decreased so much that the iodide ions are no longer discharged. (b) With regard to the iron titration according to Knop the following experiment is recommended. Take 25 cc. of the ferric chloride solution and add 15 cc. of 4 N hydrochloric acid. The solution is yellow. Add a few drops 1% diphenylamine; the color changes to dark, brownish red violet. Add 20 cc. of phosphoric acid; the dark color turns greenish yellow. The oxidation potential of the ferric ions in hydrochloric acid is large enough to oxidize the indicator. Phosphoric acid cuts the potential down. 15. Abnormal Reaction between Iodine and Thiosulfate in Weakly Alkaline Solution.-(a) Titrate 25 cc. of 0.1 N iodine in a neutral or weakly acid solution with a 0.1 N solution of sodium thiosulfate. (b) Add to 25 cc. of 0.1 N iodine, one gram of sodium bicarbonate and titrate with 0.1 N sodium thiosulfate. The titration number in the latter case is much smaller than in the former. Acidify both solutions after the titration with acetic acid and add a few cubic centimeters of barium chloride; solution a remains clear, whereas b gives a distinct precipitate of barium sulfate. The fact is often overlooked in the literature that in a weakly alkaline medium iodine oxidizes a part of the thiosulfate to sulfate, which 41% 100H- + 2S0481may involve a large error: SzOs= 5HsO.
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