Dangerous ACIDS MADE SAFELY BY Home Chemist
By Raymond B. Wailes
BECAUSE they enter into a wide variety of reactions, acids form an interesting and important group of chemicals. By preparing them in small quantities, the home experimenter can learn a great deal about chemistry and its many mysterious reactions and valuable processes.
The fact that many acids are considered dangerous should in no way dampen the amateur chemist’s ardor. Handled cautiously, they are as safe and harmless as a sharp knife in the hands of one who is careful and dexterous. They should, of course, be stored in glass bottles and kept away from clothing and hands. If some acid is spilled accidentally, it should be neutralized immediately by applying a base such as ordinary baking soda.When diluting a strong acid, always pour the acid into the water, adding it slowly and stirring the mixture with a glass tube or rod. Never pour the acid in quickly. If you do, enough heat may be generated when the two liquids mix to form steam bubbles that will blow the acid and water out of the container.
Although the amateur chemist with his meager supply of equipment cannot prepare concentrated sulphuric acid in his home laboratory, he can manufacture it in a weak form that will illustrate the method and serve to introduce an important chemical phenomenon called catalysis.
To prepare sulphuric acid, you will need some sulphur, water, calcium chlo- ride, and iron (ferric) oxide. The experiment is a simple one and requires only homemade apparatus consisting of a bottle, a flask, glass tubing, a few corks, a glass funnel, a gas burner, and rubber tubing. The parts should be arranged as shown in the illustrations. Flowers of sulphur placed in the shallow lid from a tin can is burned under the funnel at the extreme right. The sulphur dioxide formed together with some air is collected by the funnel and then passes through a drying bottle, containing the calcium chloride, to the horizontal tube of hot iron oxide. The presence of the hot iron oxide causes the sulphur dioxide to steal oxygen from the air and become sulphur trioxide. Because in this reaction, it induces a chemical change in another substance and is unchanged itself, the iron oxide is said to be a catalyst.
Finally, the sulphur trioxide formed is bubbled through water in the absorbing flask at the left. Being soluble, it combines with the water and a weak solution of sulphuric acid results.
Unaided, the original sulphur dioxide formed by the burning sulphur would not follow the desired course through the various tubes and bottles. To pull it through the system, suction must be applied to the mouth of the absorbing flask. This can be done by allowing water to siphon from a gallon jug and applying the suction formed in the jug to the absorbing flask by means of a length of rubber tubing as shown in the drawing.
To prepare the iron oxide catalyst for this experiment, soak some asbestos fiber or pumice stone in iron chloride or some other iron chemical solution until the mass is well saturated. Then add ammonium hydroxide (ordinary household ammonia will serve). This will precipitate iron hydroxide in the pores of the asbestos or pumice. The liquid then can be poured off, fresh water added and shaken and also poured off.
Next heat the impregnated pumice or asbestos in a crucible or tin-can lid over a gas burner. This final operation will convert the iron hydroxide into the desired iron oxide. The finished catalyst then is placed in the horizontal tube and heated gently with a gas burner as the sulphur dioxide is pulled through.
After burning about a teaspoonful of the sulphur, remove the absorber from the system and test the liquid with a piece of blue litmus paper. If an acid is present, the paper will turn pink. To prove that it is sulphuric acid, place a small quantity of the liquid in a test tube and add two drops of hydrochloric acid followed by several drops of barium chloride solution. If sulphuric acid is present, a white precipitate will be formed.
Although sulphuric acid made by this simple process will be weak, it should dissolve bits of magnesium and attack pieces of zinc to produce tiny bubbles of hydrogen gas. Of course, the concentration of the liquid can be increased by boiling but even then the home chemist will find that the acid will be too weak -to be of any great value for experimental purposes. ‘ It is interesting to note, however, that this same type of contact process is used commercially to manufacture sulphuric acid. Of course, a more expensive substance, usually a form of platinum, is used as the catalyst.
While the home chemist will be interested particularly in the chemical uses of sulphuric acid, he can perform a novel experiment to illustrate one of its important physical properties. In a concentrated form, sulphuric acid is capable of absorbing large quantities of moisture from the air. For this reason, it is often referred to as being hygroscopic. To understand this action more clearly, place some strong sulphur- ic acid in a small vessel and expose it to the air. The acid will absorb so much water from the surrounding air that it soon will overflow the container.
Besides many of its other valuable uses, concentrated sulphuric acid can be used to produce another useful chemical —nitric acid. This is done by placing some sodium nitrate or potassium nitrate in a glass retort containing a quantity of sulphuric acid made by mixing equal parts -of the acid and water. When the chemicals are heated, nitric acid vapors will be given off and can be condensed to a liquid by cooling.
To condense these vapors, the best procedure is shown in the photograph. Insert the end of the retort outlet tube in the mouth of a flask and rest the flask in a glass funnel. A stream of water directed on the upper face of the flask then will serve to cool it and condense the vapors leaving the retort. The funnel will serve to catch the cooling water which can be led through a rubber tube to a drain or a large pan or bottle placed on the floor.
Nitric acid manufactured by this method will be found to be quite energetic in its action with metals, carbonates, and other chemicals. Because of its activity, it should be stored in glass-stoppered bottles. It will attack both cork and rubber.
By using sulphuric acid and a small amount of iron sulphate solution, the home experimenter can test for the presence of nitric acid or nitrates. Simply place about a quarter of an inch of the sulphuric acid in a test tube, add an equal amount of iron sulphate solution, being careful not to shake the tube, and then slowly add the liquid to be tested by allowing it to run down the walls of the tube. If a brown ring is formed when the solution reaches the area between the acid and the iron sulphate and gentle heating causes the ring to disappear, it is proof that either nitric acid or a nitrate is present.
Hydrochloric acid, a third member of the important acid family, can be produced by adding ordinary table salt to sulphuric acid and heating the mixture.
Like nitric acid, hydrochloric acid also should be made in an all-glass retort. The end of the exit tube dipped into a water-cooled flask of water then will lead the gas through the water where it will be dissolved to form liquid hydrochloric acid. Although the home chemist can manufacture hydrochloric acid by this method, it will be less expensive and troublesome to use commercial muriatic acid (slightly impure hydrochloric acid).
It is a simple matter to test the distillate formed for hydrochloric acid. If a drop of silver nitrate solution is added to any solution of a chloride, a white curdy precipitate will be formed. Exposed to the sunlight, this precipitate of silver will change to a dark brown owing to decomposition.
An interesting experiment showing how heating may decompose a substance can be performed with some sal ammoniac (ammonium chloride). Being produced when hydrochloric acid gas comes in contact with ammonia gas, it can be made to break apart again by applying heat.
To separate the two gases when they are set free, the home chemist must employ a niterlike wad of asbestos fibers or other nonflammable substance rammed into a glass tube to form a plug. Ammonium chloride then is inserted into the tube at one side of the plug and the tube is mounted horizontally above the small flame of a gas burner.
IN A few seconds, the ammonium chloride will begin to decompose to form hydrochloric acid gas and ammonia gas. Being lighter than the hydrochloric acid gas, the ammonia will diffuse, spread, or travel faster and will issue from the open end of the tube nearest the porous plug. The presence of the gas can be shown by holding a moist strip of red litmus paper near the mouth of the tube until it turns blue. Similarly, the hydrochloric acid gas will issue from the other end of the tube and will give evidence of its presence by coloring damp blue litmus red. In these experiments with acids, and in fact in any experiment where a chemical in a long tube must be heated evenly, the flame-spreading attachment shown in the photograph will form a valuable addition to your gas burner. If you made the burner previously described (P.S.M., May ’33, p. 53) you will recall that the stack was made from a six-inch piece of three-eighths-inch iron pipe. To make a flame spreader, simply select a three-eighths-inch pipe cap, saw three slots across the top of the cap sixty degrees apart, drill holes at the ends of each slot, and finally screw the cap into place on threads cut in the upper end of the burner.