 Chapter 26 of the Romance of Modern Chemistry, this LibriVox recording, is in the public domain. The Romance of Modern Chemistry by James C. Philip. Chapter 26. Chemistry and Electricity. If the reader has had the patience to accompany us thus far, he will have learned that chemistry, so far from being an isolated system of facts, is intimately related to many other departments of scientific activity. The chemist has something to tell us about agriculture, about the composition of the stars, about the relation of animals and plants to the atmosphere, about the physiology of nutrition and other diverse matters. Especially when one considers the modern applications of science to industry and manufacturers, does the all-pervading influence of chemistry become apparent, for in the most unexpected quarters chemical changes are utilized and made to contribute to the requirements and comforts of life. It is not surprising to find that the bearing of chemistry on other branches of science has led to development of special study on the borderland of chemistry. Hence it comes that there is nowadays such specialization, as is indicated by the names agricultural chemistry, physical chemistry, and biochemistry. Another lateral branch of the science with a double-barreled name is electrochemistry, a subject which is of vast importance at the present time, not only from the point of view of the pure scientist, but also from that of the man who is mainly interested in applied science. The relation between chemistry and electricity is one of mutual indebtedness. It is a long time now, since Volta first showed how chemical forces might be utilized in the production of an electric current, how chemical energy might be converted into electrical energy. The chemical cell which Volta constructed consisted merely of a plate of copper, C, and a plate of zinc, Z, immersed in water to which a little sulfuric acid had been added. Volta found that if the two plates were joined by a wire outside the liquid, that an electric current pass through the wire. The electric current obtained from such a cell is not manufactured out of nothing. There is a quid pro quo. While the cell is running and producing current, chemical changes are going on, which means a lowering of the store of energy in the cell. As has been pointed out by the author of the Romance of Modern Electricity, this is exactly analogous to what happens in the case of a grandfather's clock. The store of energy in the clockweights at any time depends on the height to which they have been wound by muscular force, and the driving of the clockwork for any given time is possible only at the cost of so much of the energy residing in the weights. They will be lower down at the end of the period than they were at the beginning. Similarly, we can get a current out of a chemical cell only insofar as chemical changes go on which lower the amount of available energy in the cell. The nature of the changes which may thus be utilized in the production of an electric current are very well illustrated by reference to the Daniel cell, which is only slightly different from Volta's original one. The metals in the Daniel cell are the same as those in Volta's cell, zinc and copper. But instead of being immersed in acidulated water, the zinc plate dips in a solution of sulfate of zinc and the copper plate in a solution of sulfate of copper. The two solutions are prevented from mixing by a partition of porous earthenware, generally in the form of a cylindrical pot, inside which is the zinc sulfate and the zinc pole of the cell, and round which is the sulfate of copper solution with the copper pole. In the form of Daniel cell represented in the diagram, the copper pole is replaced by a copper pot, which holds the copper sulfate solution. If now, the zinc pole is connected with the copper pole by means of a wire, an electric current runs through this wire from the copper to the zinc. The passing of a current is evidence that work is being done by the cell, and the question therefore arises, what is the source of the energy? In the grandfather's clock the equivalent for the driving of the works is found in the gradual fall of the weights, a fairly obvious phenomenon. At a cursory inspection of the Daniel cell does not reveal any marked change which might be regarded as responsible for the electric current. Closer examination however shows that the current has been obtained only at the expense of certain alterations in the cell. If before allowing the cell to run safe for an hour, we were to weigh the two poles, then on weighing them afterwards, we should find that the zinc pole had become lighter and the copper pole heavier. Further, we should add that the solution around the zinc pole contained more sulfate of zinc than at the start, and that the solution in contact with the copper pole had lost some of its copper sulfate. The changes then which occur during the production of the current are one, the disappearance of some of the zinc to form zinc sulfate, and two, the deposition of copper on the other pole from the copper sulfate. All this might be represented very simply in the following way. Zinc plus copper sulfate going to copper plus zinc sulfate. The arrow indicating that the substances on the left are replaced by the substances named on the right. This may strike the reader as somewhat quite novel, but as a matter of fact a chemical change of exactly the same kind has already been considered in earlier chapters. One thing which, as was pointed out, served to support the alchemist belief in the transmutation of metals was the observation that when a clean steel knife blade has been dipped into a solution of copper sulfate, it looks as if it had been converted into copper. Things, however, are not always what they seem, and careful investigation has shown, one, that the formation of copper is only superficial, and two, that in exchange for the copper which has spontaneously settled on the blade, a certain quantity of iron has passed into solution as sulfate of iron. The change might in fact be represented as follows. Iron plus copper sulfate goes to copper plus iron sulfate. Now something exactly similar happens when a piece of zinc is employed instead of the knife blade. If we were to put a few bits of zinc foil in a solution of copper sulfate and leave them for some time, we should find that they had entirely disappeared, and that in their place a spongy mass of metallic copper lay at the bottom of the solution. This simple little experiment shows that the change zinc plus copper sulfate going to copper plus zinc sulfate is one which takes place spontaneously. A little reflection will convince the reader that the forces which bring about any spontaneous natural change can, if properly harnessed, be made to do work of various kinds. The force of gravitation under the influence of which an unsupported body falls to the ground is harnessed for the service of men in innumerable ways, as for instance in the grandfather's clock. The conversion of quick lime plus water into slicked lime is a change which takes place spontaneously, and as we have seen in an earlier chapter, is accompanied by a considerable increase in bulk. The force of this expansion has occasionally been utilized in blasting coal, by the simple device of packing quick lime into a hole in the coal and moistening it with water. The chemical forces set to work immediately, and the mechanical force of the expansion which accompanies the reaction suffices to split the coal apart. The Daniel cell is another illustration of this same general principle. It is simply a device whereby the spontaneous chemical change zinc plus copper sulfate going to copper plus zinc sulfate is harnessed and made to do work. The chemical energy of the cell is converted into electrical energy as evidenced by the production of an electric current. Besides the reaction which has just been discussed, there are many others which have been similarly harnessed. Among the better-known electrical cells, which, like the Daniel cell, are devices for transforming the energy of a chemical reaction into electrical energy, are the Grove cell, the Bicromate cell, and the LeClanche cell. Another very common form in which chemical energy is stored, ready for conversion into electrical energy, is the secondary cell or accumulator, sometimes called a storage cell. This is a sort of artificial chemical cell, and when complete, consists of two lead plates immersed in dilute sulfuric acid, one of the plates, however, being specially prepared and coated with peroxide of lead. In this condition, the cell is a store of chemical energy, and when the plates or poles are connected by a wire, a current passes through the ladder from the peroxide plate to the lead plate. If much current is taken out of the secondary cell, it gets run down, like the weights in the grandfather's clock. But like these, it can be wound up again. This is done by passing through the cell, say, from a dynamo, a current of electricity in the opposite direction to that of the current which the cell itself yields. The result of this is to put into the cell a fresh supply of electrical energy, which is there stored as chemical energy, ready for immediate use. From what has been said, it will be plain that chemistry has made some very important contributions to the development and application of electricity. This debt, however, has been amply repaid, and anyone who studies the modern development of chemistry will be struck with the part which electricity now plays in the chemical world. As was said at the beginning of the chapter, the relationship between chemistry and electricity is one of mutual indebtedness. We have seen how chemical changes have been utilized in the production of electrical energy. Suppose we glance now at one or two of the ways in which electricity has contributed to the advance of chemical knowledge and practice. It will be found that some of the most recent achievements of industrial chemistry have been rendered possible only by the cooperation of the chemist and the electrical engineer. It must be remembered that in some cases the electric current has been used only indirectly in order to bring about chemical changes. It is a familiar fact illustrated by the common electric glow lamp that the passage of a current through any body produces heat. The greater the opposition offered by the body to the passage of the electricity, the more intense is the heat generated by a given current. If therefore we employ very powerful currents and pass them through bodies which offer a stout resistance an enormous amount of heat is generated and a very high temperature is reached, much higher in fact than is attainable by any ordinary means. Many substances which are usually quite indifferent to each other react readily at such high temperatures so that the electric current merely by its heating action has been extremely useful in extending the chemist's field of knowledge. Some of the interesting facts which have thus been discovered at the high temperature of the electric furnace have already been described in Chapter 17. It is however not only by virtue of its heating effect that the electric current has been a service to the chemical discoverer and manufacturer. It has a remarkable power of splitting up compounds into simpler parts, provided it is supplied to these compounds while they are either in the dissolved or the molten condition. The value of the electric current for this purpose was demonstrated by the famous English chemist Sir Humphrey Davy, who succeeded in showing that potash and soda, which up to the time of his experiments had been regarded as elements, were really compounds. It was by passing an electric current through fused caustic potash that Davy first obtained potassium, a metal which is so ready to interact with air and moisture that it can be preserved only under naphtha. Potassium has the consistency of hard butter and it may easily be cut with a knife. The clean, fresh surface of the metal obtained by cutting is quite shiny, but it rapidly tarnishes, owing to the action of air and moisture. When a small piece of potassium is thrown into water, hydrogen gas and caustic potash are immediately generated and the heat of the reaction is so intense that the hydrogen catches fire. The pouring on of water, therefore, a process which is usually associated with the extinction of fire may in some cases actually lead to the production of flame. Sodium, the metal which Sir Humphrey Davy first isolated from caustic soda by the action of the electric current, is very similar to potassium, but rather less active. The decomposing action of the electric current is known as electrolysis and soon after Davy's time, another famous English investigator, Michael Faraday, discovered the laws which govern this phenomenon. He showed that when two wires connected with the poles of a battery were immersed in a solution of a salt, or the fused salt itself, decomposition took place, with the result that the metallic part of the salt was liberated at one wire, the cathode, and the acidic part of the salt at the other wire, the anode. Investigation indeed has shown that the passage of an electric current through a salt solution consists in a general movement of the metallic part towards the cathode and a general movement of the acidic part in the opposite direction, but most obvious to the onlooker is what happens at the wires or electrodes. What the observer sees taking place at the electrodes is sometimes only the secondary, not the direct result of electrolysis. For instance, if we were to pass a current through a solution of common salt or sodium chloride to give it the systematic chemical name, the metallic part of the salt, the sodium, would be liberated primarily at the cathode. Any particle of sodium, however, which was thus liberated would immediately be set upon by the surrounding water molecules, hydrogen gas would be evolved, and caustic soda would be formed in solution around the cathode. The action of water on sodium prevents our obtaining this metal by the electrolysis of an aqueous solution of any sodium salt. Sometimes the wire or plate which forms the anode is attacked and dissolved by the acidic part of the salt which is being electrolyzed. An interesting example of this is furnished by the electrolysis of a solution of copper sulfate between copper electrodes. During this process, the metallic part of the salt, the copper, is deposited on the cathode which therefore becomes gradually heavier. The sulfate or acidic part of the salt instead of being liberated at the copper anode attacks it, forming copper sulfate which dissolves in the water. So the net result of the electrolysis is that copper is transferred from the anode to the cathode, the latter increasing in weight exactly as fast as the former becomes lighter. This simple operation is really a very great technical importance. For the greater part of the world supply of copper is refined on the same principle. Plates of the impure copper which comes from this smelter are used as anodes in baths of acidified copper sulfate while sheets of pure copper act as the cathodes. When a current is passed through such a bath, the anode is gradually dissolved as already described and pure copper is deposited on the cathode. The impurities in the anode either pass into the solution and remain there or else settle down to the bottom of the bath as a sort of sludge. The small quantities of gold and silver which are present in crude copper are thus deposited in the sludge which is worked up for the sake of these valuable metals after the electrolysis is over. It is estimated that in the United States alone about 250,000 tons of copper are refined every year by this electrolytic process. 27 million ounces of silver and 346,000 ounces of gold being obtained as by-products from the sludge. Electrolysis, however, is applied not only in the purification of metals which have been produced by smelting but in obtaining the metals themselves from their compounds. Aluminum furnishes the best example of this operation for nowadays it is obtained exclusively by the electrolysis of alumina, the oxide of the metal. This material is found in various forms and in great abundance on the surface of the earth. But if it is to be employed in the electrolytic production of aluminum it must first be purified and separated from the dross which accompanies it in the natural state. The mineral which is used in this country as a source of alumina is bauxite obtainable in large quantities in the southeast of France. When pure alumina has been prepared from bauxite it is dissolved in a bath of molten cryolite, a Greenland mineral and subjected to the decomposing action of an electric current. Electrolysis takes place quietly at a temperature of about 1500 degrees Fahrenheit with the result that the aluminum from the alumina is separated at the cathode and the oxygen goes to the anode. The latter is made of carbon and at the comparatively high temperature which prevails it combines with the oxygen from the alumina and passes away as gaseous carbon monoxide. The metal on the other hand collects at the bottom of the bath in the molten condition and is runoff from time to time. There are some very interesting points about the production of aluminum in this country. As already stated the raw materials of the industry the bauxite and the cryolite are obtained from France and Greenland respectively. The bauxite is purified in Ireland where also by the way this mineral is to be found. While the actual production of the metal is carried on in some of the most outlying parts of Great Britain. We usually associate metallurgical processes such as tin and iron smelting with busy centres of population and it may seem strange that in order to find aluminum works we must go to the remote Scottish highlands. For this curious circumstance however there is a very sufficient explanation. Even the non-technical reader will perceive that in order to produce aluminum cheaply it is absolutely necessary to have inexpensive power for the dynamos which yield the electric current. Now the cheapest way of driving a dynamo is to utilize water power. This can be done on a large scale only where there is a big waterfall or where there is an adequate reservoir constantly replenished from natural sources. So far as this country is concerned these conditions are best realized in the highlands of Scotland and hence it comes that the aluminum industry is located at Foyers, in Invernessshire and it can look leaven on the borders of Argyllshire and Invernessshire. The water of the reservoir erected at the latter place is carried in a conduit to a point near the factory about 900 feet above it. From this point the water is run down to the turbines and pipes, 39 inches in diameter. Our scientific forefathers, could they see it would regard this new feature of the landscape with much curiosity. They would not understand what water pipes could possibly have to do with the aluminum industry. Perhaps also in the present stage of scientific development has long passed away our far off descendants will puzzle themselves over the ruins of these outlying industrial centers much as we today endeavour to read the riddle of druidical and Roman remains. End of chapter 26. Chapter 27 of the Romance of Modern Chemistry this LibriVox recording is in the public domain. The Romance of Modern Chemistry by James C. Phillip. Chapter 27, some interesting facts about solutions. If the reader were to glance at the titles of the papers published in any modern chemical journal he would probably be struck by a number of the most unpronounceable and incomprehensible names. It is perhaps a little difficult for him to realize how anyone can profitably spend time working at quote, the reduction of hydroxyl amino hydro umbilulone oxyne, unquote, or quote, the preparation of ethyl alfaciano gamma keto gamma phenyl butyrate unquote. But science is now so specialized that many of the advanced workers are necessarily engaged in fields which seem very remote from everyday concerns. Not only however is new and strange ground constantly being broken, the old problems which earlier workers thought they had settled are regularly coming up to review. New facts are daily discovered which bear on these problems and which if they do not clear up difficulties entirely at least contribute to their settlement. So it has been in recent years with the problem of solutions. The last two decades have witnessed an extraordinary activity on the part of chemists anxious to throw light on such questions as what happens to sugar and common salt when they are dissolved in water? How is the behavior of water affected by their presence? These may at first appear to be questions of purely academic interest but they really have a direct bearing on many practical problems. To take one instance, a knowledge of the properties of solutions is essential to anyone who attempts to understand either plant or animal life for the vital processes are invariably associated with solutions. The ultimate unit in the plant is the cell and the cell sap is the seat of its life. Fresh food too is brought from outside always in dissolved form. In the animal again, solutions are everywhere in evidence. To wit, the blood, the digestive fluids, the urine, the lymph. From the biological point of view in fact, the study of solutions is to be regarded as of the utmost importance. One feature about solutions which is very characteristic and at the same time fairly easily detected is the property of diffusion. It must not be supposed that when we dissolve cane sugar in water and set the solution on one side, the sugar molecules remain absolutely at rest. On the contrary, we have every ground for believing that each sugar molecule surrounded it may be by a retinue of water molecules is constantly moving about through the solution ever in a non coming into collision with other molecules. We must picture a sugar solution therefore as a scene of bustling activity and the molecules as on the move in every direction limited only by the boundaries of the liquid for they cannot travel except where there is a waterway. In virtue of this molecular movement, it follows that when a strong sugar solution is put in contact with pure water, the sugar molecules will gradually distribute themselves throughout the water. In fact, the process of distribution, diffusion as it is called, continues until the strength of the sugar solution is everywhere the same. This may be shown by a very simple experiment in which some concentrated sugar solution is put at the bottom of a tall glass jar, the upper part being then carefully filled with water. If left to themselves, the sugar molecules gradually penetrate the water which occupies the upper part of the jar until there are as many of them at the top as at the bottom. This interesting phenomenon of diffusion is not peculiar to sugar in water. It is characteristic of all dissolved substances. The rate of diffusion, however, differs markedly from one case to another. For example, sodium chloride, common salt, diffuses three to four times as rapidly as cane sugar. The idea of diffusion is not new to the reader for at an earlier stage we have adopted the view that the molecules of a gas are in constant motion by virtue of which they are also ready to diffuse to expand and occupy fully any space which is put at their disposal. The magnitude of this diffusive and expansive force, the pressure in other words, can be ascertained by interposing some surface in the path of the expanding gas and thus stopping its further diffusion. Similarly, in the case of dissolved cane sugar, we may ascertain the magnitude of its diffusive force by interposing between the sugar solution and the water into which it naturally diffuses some diaphragm which shall allow only the water to pass and which, like a sieve, shall stop the diffusion of the sugar molecules. Such diaphragms have been discovered and are known as semipermeable membranes, the name having referenced to the fact that the membrane is permeable for water but not for the dissolved substance. The interposition of such a diaphragm between a strong sugar solution and water prevents the sugar molecules doing what they would naturally do, that is, diffusing into the water. As the water and the sugar solution are bound to come into equilibrium somehow or other and as the usual way of reaching equilibrium is barred, the principle of Mohammed and the mountain comes into play. Instead of sugar molecules diffusing into the water, the latter percolates through the membrane into the sugar solution and the membrane, if unsupported, would soon be ruptured. The diffusive force of the sugar thus assumes the guise of a water-attracting force. This force is known as the osmotic pressure of the sugar solution and although it is rather a difficult quantity to measure, several successful attempts have recently been made in this direction. The semi-permeable membrane used in these interesting experiments consisted of copper ferrocyanide deposited on and supported by the walls of a porous pot or tube. The necessity of giving the membrane some such support will be obvious when it is stated that the osmotic pressure of a 12% sugar solution is 142 pounds per square inch. A weaker solution has a smaller osmotic pressure and in fact, it has been found that this quantity is proportional to the concentration of the solution. Semi-permeable membranes are not only produced by the chemist in his laboratory. They occur frequently in the plant and animal worlds. A red blood corpuscle, for instance, consists of a delicate, flexible semi-permeable skin inside which is a solution of the coloring matter of the blood, the hemoglobin. While the latter is unable to pass out through the enclosing membrane, water can pass in and out freely. The corpuscle is, therefore, exactly comparable with a drop of a sugar solution surrounded by a semi-permeable membrane. Suppose now we put some blood corpuscles in pure water. What result may be expected? Obviously, if the contents of the corpuscle were able to penetrate their skin, they would diffuse out into the surrounding water. Owing, however, to the semi-permeable nature of the enclosing membrane, this is impossible. What really happens is that the water passes in through the skin, which, accordingly, expands as the contents increase in bulk. The membrane enclosing a blood corpuscle is, however, very delicate. A slight increase in the volume of the contents is sufficient to burst it so that the hemoglobin escapes and imparts its color to the water. The corpuscles are said to be leaked. If now, instead of using pure water, we put blood corpuscles in each of several solutions of common salt of gradually increasing strength, we should obtain a very interesting result. In all solutions of less than 0.5% strength, the corpuscles behave as in pure water, bursting and coloring the liquid. In all solutions containing more than 0.5% of salt, the corpuscles sink to the bottom and leave a colorless liquid above. The explanation of this latter behavior will be readily understood if we consider what would be the result of putting a drop of sugar solution surrounded by a semi-permeable membrane in a still stronger solution. There is a natural tendency constantly at work to equalize the osmotic pressures on the two sides of such a membrane so that water will always pass from the solution with the smaller osmotic pressure, that is from the weaker solution, to the one with the greater osmotic pressure, that is the stronger solution. In the case suggested, therefore, water will pass from the inside of the drop to the outside so that the sugar solution within becomes stronger. A blood corpuscle is affected in exactly the same way when it is put in a strong solution of salt. Water actually passes out through the skin. The corpuscle shrinks in size, becomes more dense, and sinks to the bottom of the salt solution. If we were to test a number of sugar solutions of gradually increasing strength in the same way as has been suggested for salt, we should find that the blood corpuscles were burst in all sugar solutions up to 5% strength but sank to the bottom in solutions of greater concentration. For both sugar and salt, therefore, we can, with the help of the corpuscles as indicators, pick out solutions which have the same osmotic strength. From this point of view, a 5% solution of cane sugar must be regarded as equivalent to a 0.5% solution of common salt, even although the amounts of dissolved matter are so different in the two cases. This may seem rather strange, but investigations which we cannot consider here have shown that the magnitude of the osmotic pressure is determined not by the weight of a substance which is dissolved in a given volume of solution but by the number of molecules present in that volume. Now, the sugar molecule is a very heavy one so that proportionately more of this substance must be taken in order to get a definite number of molecules. There are other interesting properties of solutions which are in reality closely connected with osmotic pressure. There is, for instance, the fact, perhaps already known to some readers, that a solution freezes at a lower temperature than water and boils at a higher temperature, more than that, the extent to which the freezing point of the solution is below 32 degrees Fahrenheit and its boiling point above 212 degrees is proportional to the number of molecules of the dissolved substance in a given volume so that the chemist can compare the number of molecules present in solutions of salt and sugar by finding out the temperatures at which these solutions freeze or boil. Working on the same principle, one could get a very fair idea of the amount of solid dissolved in seawater by comparing its freezing point with those of a number of common salt solutions of known strengths. The seawater must contain about as much dissolved solid as that particular salt solution which has the same freezing point. The topic of freezing and boiling of solutions is in reality closely related to the interesting question how the water can be separated from the dissolved substance, how, for example, pure fresh water can be obtained from seawater. The problem is not quite so simple as was thought by the examination candidate who suggested that in order to procure fresh from salt water, it was only necessary to set aside and skim off the salt after standing. This method certainly works in the case of milk but then milk is not simply a solution. The particles of fat which separate from milk on standing are not dissolved. They are only suspended and gradually come to the top because they are lighter than water. One way, however, of getting pure water from a solution of salt is to freeze it. When that takes place, the solution does not freeze as a whole. The solid crystals which separate consist of ice alone. That is, they are pure water. In virtue of this fact, an iceberg formed in seawater would, if melted, be found to yield approximately fresh water. Any salt which it still contained must have been merely imprisoned or entangled among the freezing particles of water. The separation of pure water from a salt solution, however, can be carried out not only by freezing but by boiling. The latter process indeed is more easily affected and gives a more perfect separation. It is therefore used on a large scale for the production of fresh water from seawater. The whole operation of heating the seawater and condensing the steam which comes off is known as distillation and the plant necessary for carrying this out is part of the regular equipment of an ocean liner. It will be obvious to the reader that one result of boiling a solution and thus getting rid of some of the water will be to increase the strength of the solution, provided, of course, that the dissolved substance itself has no tendency to volatize. Now, as already stated, the greater the concentration, the higher is the boiling point. If, therefore, a thermometer is put in a solution, say, of sugar, which is boiling in an open vessel, the readings of the thermometer will get higher and higher. Anyone who has made fondant for confectionary purposes will have observed this. In this operation, sugar and water are mixed in certain proportions in a pan and heated. The temperature rises rapidly to the point at which the mixture begins to boil and then slowly thereafter, as the water is boiled off and the sugar solution concentrates. The subsequent behavior of the sugar solution, when cold, varies with the extent to which this concentration has been allowed to proceed. It depends on the temperature at which the boiling has been stopped. In the endeavor to bring home to the reader the curious properties exhibited by solutions, the writer has taken sugar or salt as the dissolved substance. These compounds have the advantage of being perfectly well-known to everyone, but at the same time, they represent two quite distinct classes. So far, no reference has been made to this distinction between sugar and salt, but it is one which has led to much controversy and practical work on the part of the modern chemists and, as such, deserves our attention. There are various good grounds for believing that the molecule of cane sugar is nearly six times as heavy as the molecule of sodium chloride, common salt. If then we took equal weights of the two substances, we should have six times as many molecules of salt as of sugar. Since, as already stated, the extent to which the freezing point is lowered, the depression, as we may call it, is proportional to the number of molecules or dissolved units, it follows that salt ought to produce a depression six times as great as that caused by an equal weight of sugar, provided that each is dissolved in the same quantity of water. Expectations, however, are not realized in this case, and the salt gives a depression about 11 times as great as that due to the sugar. The salt in the process of solution seems to have yielded nearly twice as many dissolved units as we should expect. What interpretation can be given of this extraordinary behavior? Before any explanation is attempted, attention must be directed to another point of distinction between salt and sugar. If two wires connected with an electric battery are immersed in pure water, only an infinitesimally small current would pass. The water offers an enormous resistance to the passage of the current and may be described as practically a non-conductor. The moment, however, a pinch of salt is dissolved in the water, all this is changed, and the current experiences a comparatively small resistance. The addition of salt confers on water the power to conduct the electric current, and the salt solution may be electrolyzed as described in the previous chapter. The case of sugar is quite different. The resistance offered by water to the passage of a current is not diminished by the addition of sugar, nor can a solution of sugar be electrolyzed. Such differences in behavior, as those exhibited by salt and sugar, are shown by hosts of other substances. Of the chemical compounds which can be dissolved in water, many, comprising acids, bases, and salts behave like sodium chloride and are accordingly called electrolytes. The others, which like sugar, do not increase the conducting power of water, are known as non-electrolytes. Now the curious thing is that it is just those substances which make water a conductor, that is the electrolytes, which have an unexpectedly big effect on the freezing point of water. This coincidence was emphasized some 25 years ago by the well-known Swedish chemist, Hrenius, who also suggested an ingenious explanation. According to this, the molecule of an electrolyte when dissolved in water is liable to dissociate or split up into two parts or ions, one of which carries a positive electric charge and the other a negative charge. When, for instance, sodium chloride is added to water, the atom of sodium and the atom of chlorine, which have combined to form the molecule of that substance are instantly seized by a desire for divorce and they separate, so far at least, as the most of the molecules are concerned, giving rise to a positively charged sodium ion and a negatively charged chlorine ion. At first, this theory seems rather fantastic. There appears to be no sufficient reason why mere contact with water should induce sodium chloride in other electrolytes to commit molecular suicide as one critic has put it. The question of a motive for this suicide has been a difficulty, but recent work indicates that the ions have a greater affection for water than they have for each other and hence arises their apparent readiness to part company. Whether this be the correct explanation or not, it is certain that the erroneous hypothesis of ionic dissociation gives an excellent interpretation of many properties of solutions and has provided a basis for much valuable work. How, then, does it explain the fact that salt has an abnormally large influence on the freezing point of water? Simply in this way, that such an association of sodium chloride, as has been suggested, would mean an exceptionally large number of dissolved units. And since the depression of the freezing point is proportional to the number of dissolved units, the effect of the salt on the freezing point is unexpectedly great. Then again, the fact that sodium chloride makes water a conductor of the electric current becomes intelligible on the basis of Arhenius' hypothesis. For if the salt solution contains a large number of positively and negatively charged particles, the mere immersion of two battery wires will cause a streaming of the positive ions of the negative wire and of the negative ions in the opposite direction. Such a procession of ions carrying electric charges is nothing else than a transport of electricity and is therefore equivalent to the passage of a current through the solution. The presence of the salt, that is, has changed the water from a nonconductor to a conductor. It may sound rather fanciful to talk of a procession of ions, but ocular demonstration can be given of the fact that during electrolysis, the positive or metallic part of a salt actually moves towards one electrode while the negative or acidic part is at the same time traveling towards the other. For such a demonstration, some colored salt, permanganative potash, for instance, must be employed. Sometime perhaps the reader may have the opportunity of seeing this very interesting experiment. End of Chapter 27. Chapter 28 of the Romance of Modern Chemistry. This LibriVox recording is in the public domain. The Romance of Modern Chemistry by James C. Phillip. Chapter 28, From Solutions to Crystals. In the foregoing chapter, reference was made to the curious ways in which common substances affect the properties of water and to the methods of getting pure water from a solution. It was there suggested that by cooling a salt solution until it began to freeze, a separation of water from the dissolved substance could be affected, since it is pure ice which crystallizes first. Strictly speaking, this method would not work with a strong solution, for cooling in this case might result in the separation of the salt itself in the crystalline form before the freezing point was reached. This phenomenon of salt crystallization depends on the fact that substances as a rule are more soluble and hot than in cold water. Thus, for example, a saturated solution of salt-peter potassium nitrate, that is a solution which cannot dissolve any more nitrate, contains 24% of the salt at 68 degrees Fahrenheit, 48% at 130 degrees, and 71% at 212 degrees. Hence, if a saturated solution of salt-peter were prepared at 130 degrees Fahrenheit and was then cooled down, ultimately to 68 degrees, it would give up as crystals all the salt which is contained over and above 24%. In such a case, the salt-peter is said to have crystallized out from the solution. Crystallization is a common laboratory operation and is an efficient means of purifying salts and other substances. This depends on the fact that when an impure material is dissolved in water and crystallization is allowed to take place, the separated crystals are comparatively free from impurities. These are found to have accumulated in the liquid, which is alongside the crystals, the mother liquor as it is called. If the crystals are re-dissolved in the process of crystallization repeated, a still purer product is obtained. Sometimes it is necessary to carry out a re-crystallization repeatedly in order to get absolutely pure material. And cases are on record in which the operation has been performed 20 to 30 times. The reader will perceive therefore that patience is an essential part of a chemist's equipment. If a strong salt solution always behaved as it ought, and as we might reasonably expect it to behave, then when cooled to the temperature which it is saturated, it would begin to deposit crystals. But just as there are not a few individuals who have a great reluctance to get out of bed when they ought to be up, so there are some salts which exhibit a curious hesitancy to leave the dissolved condition. Their solutions deposit no crystals, even when cooled far below the saturation point. In these circumstances, we have what is known as a super saturated solution. We do not require to go very far afield to find a salt which exhibits this curious inertia. Sodium thiosulfate, better known perhaps as the hypo of the photographer, is a very good case in point. A strong solution of this substance may be made by nearly filling a flask with the crystals of the salt, adding a little water and then immersing the flask in hot water. This treatment renders the contents of the flask fluid and they remain in this condition even when cooled to the ordinary temperature. The solution is then super saturated and will deposit crystals only when it is irritated in some manner. This may be done by vigorous shaking or stirring, but most certainly by dropping in a crystal of sodium thiosulfate itself. This operation is described as inoculation or sowing and it is certainly a sowing which produces an immediate harvest. The presence of an already formed crystal acts as a stimulus to the molecules which have sluggishly lingered in the dissolved condition and they hasten to arrange themselves in the regular manner which is characteristic of the crystalline state. One result of this is that the contents of the flask, formerly fluid, appear to have become nearly solid and another obvious fact is a considerable rise in temperature. This evolution of heat which accompanies the crystallization of a super saturated solution is not to be wondered at. It is simply the repayment of a loan. For most salts absorb heat when they pass from the solid to the dissolved condition, a fact which anyone can realize by putting a quantity of salt peter in water and observing that the vessel containing the water becomes sensibly colder. This heat which the salt abstracts from the water and the containing vessel when it passes into solution is duly returned by it when it comes out of solution, hence the remarkable evolution of heat when a super saturated solution is suddenly stimulated into crystallization. Another substance which resembles sodium thiosulfate in readily forming super saturated solutions is acetate of soda. Indeed, its behavior in this respect has been turned to practical account in railway foot warmers. When these are filled with a hot strong solution of sodium acetate instead of hot water, the store of available heat is about four times as great. Reluctance to pass into the crystallized state is exhibited not only by dissolved substances but also by many fused compounds. It is possible in the latter case to cool the molten substance below its freezing point without crystallization setting in. The substance is said to be super cooled. As with super saturated solutions, mere contact with a crystal is sufficient to induce crystallization. There is however a remarkable difference in the rates at which different super cooled substances respond to this stimulus. This is very well shown by filling a long narrow tube with the super cooled liquid and touching one end of the column with a crystal of the solid material. Crystallization starts immediately at the point of contact and is propagated through the tube at a rate which is perfectly regular but differs from one case to another. For instance, yellow phosphorus and benzophenone, a substance well known to the student of organic chemistry, can both be obtained in the super cooled condition but the rates of solidification in narrow tubes are very different in the two cases. The crystallization of the phosphorus proceeds at a very great speed of 39 inches a second. In the case of benzophenone, the rate is only 1 25th of an inch per second. The most effective way of inducing crystallization in a super saturated solution or in a super cooled substance is, as already indicated, the addition of a crystal of the solid. But suppose no solid is available, what then? This is the position in which a chemist frequently finds himself when he is on the track of a new compound. He may have actually got it before him in a solution or in the form of an impure oil which will not crystallize. In such a case, seeing none of the already formed solid is available, the only thing to do is to try mechanical methods of inducing crystallization. If the reader has ever gone into a research laboratory for organic chemistry, he may have seen someone eagerly stirring and scratching at an oily looking substance on a watch glass. This is all with the object of persuading the substance to crystallize and it is wonderful how frequently this method is effective. An excellent illustration of the way in which scratching promotes crystallization is furnished by the behavior of potassium by tartrate. This substance is found in grape juice and is more familiar, especially to housewives, under the name of cream of tartar. It is only sparingly soluble in water. If a saturated solution is made at 90 degrees or 100 degrees Fahrenheit and is then cooled, it will be super saturated at the ordinary temperature. If the solution, as soon as it has cooled, is poured on a glass plate and the plate is scratched with a glass rod in such a way that the latter writes invisible letters on the plate, the writing soon becomes visible because especially rapid crystallization is induced along the lines with the glass rod and the plate were in contact. The letters are traced out by the deposited crystals. When a salt crystallizes out from its solution in water, it frequently happens that it carries water along with it. In the act of crystallization, each molecule of the salt hooks onto itself one or more molecules of water. This is not a mere mechanical adherence, for the crystals may be removed from the solution and pressed between blotting paper until they are absolutely dry without detaching any of these water molecules. They are, as a matter of fact, chemically combined with the salt to form a composite molecule and they will not be drawn by methods which suffice to dry up ordinary moisture. Water, which is held by a salt in this way, is described as water of crystallization. Although blotting paper fails to affect the separation of a salt and its water of crystallization, the bond of union is really not very strong and it may be said that the love between the two grows cold as the temperature rises so that by merely warming a salt which contains water of crystallization, the water is driven off as vapor and finally the salt alone, the anhydrous salt, as it is called, is left. Water molecules, however, are not all alike in the tenacity with which they cling to the salt molecule. Some can be detached only by application of a higher temperature than is required for others. Of this graded affection, blue vitriol or copper sulfate, to give it its chemical name, furnishes an interesting example. The ordinary crystals of this substance are blue in color and contain 36% of their weight of water. Each molecule of the salt carries with it five molecules of water of crystallization. If the crystals are exposed for some time to the temperature of 212 degrees Fahrenheit, say in a steam oven, four out of the five molecules go off and the residue is pale blue. The last molecule is more faithful but a rise of temperature to 400 degrees compels even this one to take its departure and a white powder is left as the anhydrous salt. In some cases the molecules of water in a crystallized salt begin to evaporate of their own accord even at the ordinary temperature. A domestic example of this curious behavior is available for washing soda which normally contains 10 molecules of water of crystallization to each molecule of sodium carbonate, loses some of them on mere exposure to air. This process is revealed to the observer by the fact that the crystals of washing soda, originally clear and transparent, become gradually opaque as if a white powder had been deposited on them. This is due simply to the partial removal of water from the surface layer of the crystals which therefore exhibit signs of disintegration. In blue vitriol an instance has already been cited of the way in which the color of a given salt varies according to the number of molecules of water of crystallization which it contains. An even more striking example of this phenomenon is furnished by a substance known to chemists as magnesium platinocyanide. This salt can be obtained with seven, six or two molecules of water as well as in the anhydrous state and these various products are respectively scarlet, lemon yellow, colorless, and orange yellow. The extraordinary influence which water thus has in altering the color of a salt explains the action of the so-called sympathetic or invisible inks. One of these is a solution of cobalt chloride, a salt which crystallizes with six molecules of water in the form of dark red crystals. The anhydrous salt on the other hand is deep blue in color. The water solution of cobalt chloride is merely pink and if this is used to write on paper instead of ordinary ink, the impression left is so slight as to be scarcely noticeable even when it has dried. The application of heat to the paper however makes the writing immediately visible for the cobalt chloride is thereby converted into the blue anhydrous salt. Curiously enough if left to itself the writing fades away for the blue salt gradually absorbs moisture from the air regenerating the pink salt which is almost invisible. On the same lines the reader will himself be able to explain the behavior of certain artificial flowers which are said to be made in Paris. Their petals are tinted with cobalt chloride with the result that while the flowers are usually of a rose color they turn blue in a very dry atmosphere. Nothing has yet been said about the strikingly regular and beautiful forms in which dissolved substances crystallize out from their solutions. These must be seen to be appreciated. As a rule each dissolved salt separates in a definite shape peculiar to itself and it is in fact this regularity of form which is the main distinguishing feature of the crystalline state. If the separate crystals are large it is easy not only to see distinctly the various shiny faces but also to count them and when the chemist has become familiar with the crystalline habits of a particular substance he can afterwards identify it even amongst many others merely by its appearance. The process of crystallization consists in an ordered fitting and packing together of the molecules of the solid. This regularity of arrangement is evident not only from a study of large well formed crystals but also from the appearance under the microscope of minute quantities of crystallized solutions. If for instance a drop of ammonium chloride solution is crystallized under a microscope slide the crystals are seen on close examination to have assumed a rather fern-like shape C figure 14. The reader must not suppose that it is only from solutions in water and other similar liquids that crystals are formed. Fused substances as already indicated solidify to crystalline masses but in addition we must be prepared to think of crystallization as having taken place from solvents which are solid at the ordinary temperature. A fused alloy for instance containing a little of one metal dissolved in another is quite analogous to a solution of a salt in water although the alloy must be kept at very much higher temperature if it is to remain in the liquid condition. Now just as the salt crystallizes out from its water solution so the one metal crystallizes out from the fused alloy. The difference is that in the alloy we cannot see the crystals which are formed first because the alloy as a whole has subsequently solidified. In fact an alloy which has crystallized and has been cooled down to the ordinary temperature is similar to what a salt solution would become if it were cooled a long way below the freezing point of water. In the latter circumstances we should not be able to see salt crystals because of the masses of ice with which they were surrounded. In spite of this difficulty there are ways and means of finding what sort and shape of crystals have primarily separated from an alloy. In a few cases the Ronkin rays are serviceable by virtue of the fact that metals differ in their permeability to these rays. Some metals are transparent to the rays others are opaque. Suppose for instance we had a small quantity of gold which is opaque dissolved in a large quantity of sodium a metal which is comparatively transparent to the Ronkin rays. Such an alloy if fused would be comparable with a dilute solution of a salt in water. In this latter case since the solution is dilute the primary crystallization on cooling would consist of ice. Similarly the fused alloy if cooled would deposit sodium. That is the first crystals to separate from the alloy would consist of the metal which is transparent to the Ronkin rays. While the spaces between these crystals would ultimately contain the gold which crystallizes last. The slower the cooling the better will be the opportunity for the primary crystals to grow large. If now a thin section of the cold alloy is placed on the top of a light tight envelope containing a sensitive plate and exposed to the Ronkin rays these are able to pass through the primary sodium crystals and act on the plate but are blocked to a large extent by the opaque gold in the spaces between the crystals. The impression therefore obtained on the plate differentiates between the primary crystals and the rest of the alloy. This extremely interesting method of studying the crystalline solution of alloys can obviously be employed only when there is a fairly well marked difference between the two metals in relation to the Ronkin rays. The scope of its application is therefore somewhat limited. Another way of revealing the crystalline condition of an alloy is to cut a section, polish one of the surfaces and treat it with an acid. This treatment brings out the details of the crystalline structure and with the help of microscope and camera a photomicrograph is obtained the surface of the alloy being illuminated by either oblique or reflected light. This method is not restricted in its scope like the previous one and it is largely applied at the present time notably in the investigation of the character of iron and steel which have been exposed to different conditions. Compare figure 15. Working on these lines the modern chemists can do really marvelous things in the way of deciphering the life history of an alloy. The tail is on the face of it if he has but the key to the language and the necessary patients. As aids he requires chiefly two instruments which in countless directions are invaluable to the scientific worker. It is indeed difficult to realize how much poorer natural science would be today had we no microscope and no camera. End of chapter 28. Chapter 29 of the Romance of Modern Chemistry this LibriVox recording is in the public domain. The Romance of Modern Chemistry by James C. Phillip. Great Effects from Small Causes It is a common place to say that incidents or persons may have an influence quite out of proportion to their apparent value. This is what everyone learns sooner or later but it is worthwhile noting here how very remarkably this principle is enforced by many facts with which the chemist is familiar. Nature herself in various striking cases reminds us that what is apparently insignificant is frequently of the utmost importance. Take the case of carbon dioxide. This gas is present in the atmosphere to the trifling extent of three parts in 10,000 and yet it is on this that the whole vegetable life of our globe depends. As was pointed out in a previous chapter it was not until the existence and significance of this three one-hundredths of one percent of carbon dioxide were appreciated that the miracle of vegetable growth could be rightly interpreted. Modern Chemistry furnishes many remarkable instances of the way in which the history of a chemical change or the behavior of a particular substance is profoundly modified by the presence of small quantities of foreign material. It is not necessary to go far afield in search of such cases. For water one of the commonest chemical compounds has recently been shown to have an extraordinary influence in promoting chemical action between other substances. The reader is probably familiar with the experiment in which a lighted taper is brought to the mouth of a soda water bottle containing a mixture of hydrogen and oxygen. A vigorous action marked by a violent explosion takes place between the gases and water is produced. The striking liberation of energy which accompanies the chemical action between hydrogen and oxygen is evidence of the extreme eagerness of the two elements to go for each other. Yet if care is taken to remove all traces of moisture from the original gases this lust of battle has apparently gone. A tube containing a mixture of perfectly dry hydrogen and perfectly dry oxygen may be strongly heated without the contents exploding. A really astounding result. It has been put on record that in 12 successive experiments on pairs of tubes, one of each pair containing perfectly dried hydrogen and oxygen, the other containing the imperfectly dried gases, the result of heating the tubes to redness in a Bunsen burner was invariably the same. There was an explosion in the tube containing the imperfectly dried gases but no explosion in the other. It is true that in order to secure this result the gases in the latter tube had been very carefully dried for 10 days but this does not detract from the striking character of the experiments. Dried for 10 days, the reader may exclaim. How is that done? He must of course dismiss from his mind the idea of using any ordinary means of drying wet objects. A gas can be dried only by letting it come in contact with some material which has an intense fondness for water and which will readily absorb it whenever it gets the chance. Such substances are quick lime, strong sulfuric acid, and phosphoric oxide. In the experiments just described the condition of perfect dryness was attained by putting some phosphoric oxide in the tube along with the hydrogen and oxygen. The oxide in these circumstances acts like so much bird lime and any water molecules that are flying about are gradually caught. In the absence of water molecules the mischief makers as we may term them or shall we say matchmakers the hydrogen and the oxygen are quite callous to each other so soon however as the merest trace of moisture is admitted into the tube the contents will explode when heated. A few water molecules in fact are responsible for all the difference between peace and war or between the single and wedded states according to the way in which the reader prefers to picture the interaction of hydrogen and oxygen. This extraordinary influence of water on chemical change is so remarkable that it is worthwhile to refer to another interesting experiment that has been made. As the reader is aware ammonia is a colorless alkaline gas whereas hydrogen chloride is a colorless acid gas. Like alkaline acids in general these two gases interact forming a salt, ammonium chloride or sal-ammoniac the characteristic and curious feature of the process being the production of this white solid substance from two colorless invisible gases. It turns out now that this combination between ammonia and hydrogen chloride which takes place so readily under ordinary circumstances is not observed when the gases have first been completely freed from all stray water molecules. The ordinary incandescent mantle is an excellent example of the value which may attach to small quantities of foreign material. The mantle consists to the extent of 99% of thoria which is the oxide of the metal thorium and is obtained chiefly from monazite sand found in Brazil and in the United States. The remaining 1% of the mantle is syria the oxide of the rare metal syrium and in spite of its small proportion it is on this admixture that the virtue of the mantle wholly depends. Mantles composed of pure thoria alone would be of no use for when put in a Bunsen flame they give out only a dull light. On the other hand if more than 1% of syria is added to the thoria a less brilliant effect is obtained. It is in fact possible to have too much of a good thing. Not only has this paltry 1% of syria made the incandescent mantle a brilliant success it has indirectly been the salvation of the gas industry. In competition with electricity gas would have been badly beaten as a source of light had it not been for the discovery of the incandescent mantle. By its agency the illuminating power of a cubic foot of coal gas is enormously increased. Another interesting fact in connection with the incandescent mantle deserving of passing notice is the extraordinary effect which the rapidly increasing use of thorium nitrate had on the price of that article. Early in 1894 an ounce of thorium nitrate sold for 55 shillings. By January 1895 on account of competition and improved methods of production the price had fallen to 25 shillings. By July 1895 to 14 shillings by November of the same year to eight shillings. In another six months the price was again halved while at the present time it has fallen to about one shilling. Seldom indeed has any chemical product undergone such a rapid change in price. Another interesting illustration of the extraordinary influence exerted by small quantities of foreign matter is furnished by the behavior of certain phosphorescent substances. Among them are the sulfides of the metals barium, strontium and calcium. And as the name phosphorescent implies these sulfides are luminous in a dark room after they have been brought out of the light. Luminous paints or luminous compositions generally are dependent for their characteristic behavior on the presence of such a phosphorescent substance. Curiously enough however the pure materials do not appear to be phosphorescent. It is only when minute traces of other matter are present that they are stimulated to luminous activity. Incandescent mantles and phosphorescent substances illustrate very well the striking modifications of properties which are attributable to small quantities of foreign material. But it is not only the properties of particular compounds which are affected by impurities. As we have seen in the case of imperfectly dried hydrogen and oxygen the speed at which a chemical change takes place may be remarkably modified by the presence of some alien substance which keeps so to speak in the background and does not itself suffer any apparent alteration. Such an acceleration of chemical change by an alien substance is known as catalysis and the alien substance itself is spoken of as a catalytic agent. Chemical changes which take place under the influence of a catalytic agent are not merely laboratory curiosities. They are of the utmost importance in the technical world. The most modern method of manufacturing sulfuric acid for instance depends on just such a change. And if the reader tries to realize the fact that about 3,000 tons of sulfuric acid are made in England every day he may appreciate the bearing of catalysis on chemical industry. The main thing to be done in making this important product is to persuade sulfur dioxide the suffocating gas which is produced when sulfur is burned to combine with more oxygen. This will not occur spontaneously when the gases are merely mixed even at a high temperature. They must be brought into contact with some third substance which plays the same part as water does in the combination of hydrogen and oxygen. In the case of sulfur dioxide and oxygen the third party which acts in some subtle way as mediator between the other two is platinum in a finely divided condition. This metal has quite a reputation for accelerating chemical actions in which it is not directly involved. It is a sort of chemical busy body. There is a well-known experiment which illustrates this characteristic of platinum very clearly indeed. A roll of platinum foil is suspended in the flame of a Bunsen burner until it is red hot. The gas is then turned off and immediately turned on again but not lighted. The observer sees that the foil which has begun to cool down whenever the gas was turned off begins to glow afresh although there is no visible flame. It remains in this condition so long as the mixture of air and gas from the burner is allowed to flow over it. What has happened is that the platinum induces the slow combustion of the gas and it is the heat given out in this process which keeps the metal visibly hot. The platinum itself is not affected so that we have here an excellent example of catalytic action. A similar part is played by the finely divided platinum in the modern method of making sulfuric acid, the contact process as it is appropriately called. Under the persuasive influence of the catalytic agent at a temperature of about 500 degrees Fahrenheit, sulfur dioxide and oxygen readily unite to form another compound named sulfur trioxide which need only be dissolved in water to produce sulfuric acid. This sounds all very simple and in fact this way of making sulfuric acid was discovered long ago. It could not however be employed on the manufacturing scale as the platinum turned out to be very sensitive to impurities in the sulfur dioxide and gradually became ineffective. Ways and means however have now been discovered for thoroughly removing the impurities and keeping the platinum in good condition so that what for long was merely a laboratory experiment has now become the basis of a very important manufacturing operation. The contact process is rapidly coming to the front and bids fair to oust the old and cumbersome method of manufacture, the prominent feature of which is the use of huge leaden chambers. In recent years platinum has been prepared in another condition in which it exhibits remarkable catalytic activity namely in solution in water. It may seem to the reader rather absurd to speak of dissolving a metal in water as if it were so much sugar or salt but it is indeed a fact that by the help of the electric current platinum has been got into water in such a state that it closely resembles a dissolved substance. If two pieces of stout platinum wire are immersed in water so that their points are very close together and an electric discharge is passed across the intervening space the water gradually assumes a deep brown color and is found then to contain platinum in solution. At least it seems to be in solution for the liquid may be filtered through a piece of blotting paper without leaving any particles behind and it may be kept for a long time without depositing any sediment. On grounds however into which we cannot go here the view is adopted that this platinum solution is really a suspension of exceedingly minute particles so tiny that they can find their way through the pores of filtering paper. However that may be there is no doubt that platinum in this condition is intensely active from the catalytic point of view as shown for instance by its effect in upsetting the equilibrium of hydrogen peroxide. This is a substance which in water solution is applied as a bleaching agent for hair ivory and old pictures. Chemically it is a very interesting substance being closely related to water. Its molecule in fact is a molecule of water to which an extra atom of oxygen has been tacked on. The attachment however is not very secure and the result is that hydrogen peroxide is readily decomposed into water and oxygen. This chemical action, this decomposition is accelerated in quite a remarkable manner by the addition of a little platinum solution to the hydrogen peroxide. Thus if we were to take dilute hydrogen peroxide and add to it so much platinum solution that a pint of the mixture contained one sixteenth thousandths of an ounce of platinum the decomposition of the hydrogen peroxide would be complete in about two hours. If no platinum solution were added the hydrogen peroxide would lose practically none of its oxygen in that time. Perhaps a still more convincing proof of the catalytic power of this platinum solution is obtained by shaking some of it in a flask with a mixture of hydrogen and oxygen. In ordinary circumstances these two gases require to be strongly heated before they will combine to form water but under the persuasive influence of the platinum solution they unite at the temperature of the room slowly but steadily and without any fuss. The extraordinary thing is that a minute quantity of platinum is able to bring about the combination of very large quantities of hydrogen and oxygen. One experiment has been recorded in which under the influence of the one sixteenth thousandths of an ounce of platinum over two gallons of a mixture of hydrogen and oxygen disappeared in 17 days and the platinum was as active at the end of this period as at the beginning. There are numerous catalytic agents in addition to those which are used in the laboratory or the factory. Our bodies are the scene of many chemical changes which are promoted and accelerated by the influence of certain agents called enzymes the exact nature of which is not yet known. These substances play an important part as catalytic agents notably in the processes of digestion but an excursion into this interesting field would take us too far. Perhaps enough has been said to convince the reader that in chemistry at least much that is apparently insignificant is of the greatest value and importance. End of chapter 29. Chapter 30 of the Romance of Modern Chemistry. This LibriVox recording is in the public domain. The Romance of Modern Chemistry by James C. Phillip. Chapter 30, how trifling observations lead to great discoveries. The scientist who is advancing into the unknown generally sets out with the object of searching for something which his theories lead him to believe is to be found in the unexplored region just ahead. It frequently happens, however, that as he steadily plods forward he discovers something by the way which is of much greater importance than the ultimate object of his search. The story of the ways in which some such unexpected discoveries have been made is interesting. If not romantic and the rehearsal of one or two of these will show the reader how much depends sometimes on a casual occurrence and on the observer's readiness to note what happens and to take advantage of it. The discovery of oxygen, the important element which forms one-fifth by volume of the air was made in a very casual sort of fashion about 140 years ago. Priestly, we are told, was very proud of a burning glass which had come into his possession and was going round his laboratory one day concentrating the sun's rays with this lens and focusing them on all sorts of substances. Among the materials which he had thus happened to expose to the heat of the concentrated solar rays was oxide of mercury, which as we now know is very readily split up by heating into its constituent elements, mercury and oxygen. Priestly observed that a gas was given off from the mercury oxide and when he had collected some of the gas he was able to show that a candle burned in it with a remarkably vigorous flame. To Priestly, this was something quite new and fascinating as he says himself, This surprised me more than I can well express. I was utterly at a loss how to account for it." Further experiment showed him that the gas, quote, possessed all the properties of common air only in much greater perfection, unquote. He had in fact discovered oxygen and all as the result of curiosity about the powers of his newly acquired lens. He was, it is true, on the lookout for new gases at that time, but after all the concentration of the sun's rays by a lens is a most unusual way of producing heat and would not naturally be chosen for that purpose. If however the investigator's mind is occupied with a definite subject, it is wonderful how the most trifling occurrences are seen by him to have a bearing on the problem and are made to contribute to its solution. So it was with Priestly and so it has been in many other cases which might be quoted. One of those which has been put on record occurred in connection with the discovery of blasting gelatin by Nobel. As has been stated in a previous chapter, the dangerously explosive substance, nitroglycerin, cannot by itself be safely handled and transported. The difficulty maybe got over by soaking up the liquid nitroglycerin into Kieselger and so converting it into the product known as dynamite. It was obvious to Nobel that this operation involved a reduction of the explosive force of nitroglycerin, for the absorbent Kieselger is a neutral, harmless, non-explosive material. So although it can take up as much as three times its quantity of nitroglycerin, the explosive power of the latter is lowered by one fourth. Nobel was therefore anxious to find as a substitute for Kieselger, some substance which would convert nitroglycerin into a form suitable for safe handling and transport and which at the same time, being itself explosive would not diminish the effectiveness of the nitroglycerin. The discovery of a material with the desired properties came quite by accident. Nobel cut his finger one day in the laboratory and procured some collodion to paint over the cut and so form an artificial protective skin. Collodion, it should be stated, is a solution of a substance resembling gun-cotton in a mixture of alcohol and ether. As these two liquids are very volatile, a film of collodion exposed to the air soon dries up and forms a skin. After Nobel had used a little of the collodion to paint over the wound, it occurred to him to pour what was left into a vessel containing nitroglycerin. He did this and observed that the collodion mixed with the nitroglycerin and formed a jelly-like mass. This little observation was enough to show him the way in which the problem of the replacement of Kieselger by a more active substance could be solved. Experiments were carried out on the large scale and these led to the manufacture of the explosive known as blasting gelatin which is a mixture of nine parts of nitroglycerin and one part of soluble gun-cotton. Pure blasting gelatin is so violent in its action that it cannot be used except for the hardest rocks. It was employed, for instance, in parts of the St. Goddard tunnel. For ordinary practical purposes, however, the explosive power of blasting gelatin is modified by introducing a certain amount of non-explosive absorbent material. Some discoveries have actually been made through an accident happening to the apparatus with which the experiments were being carried out. This was the case with one important series of investigations into the behavior of gases and the famous chemist Graham has explained what it was that led him to make his wonderful experiments on gaseous diffusion. It appears that an earlier worker, Dobariner, had occasion to prepare large quantities of hydrogen and one day accidentally used as gas holder a glass jar which had a tiny crack in it. Now it is a well-known fact that if an undamaged glass jar or tumbler containing hydrogen or air is inverted in a dish of water so that the level of the water outside and inside the jar or tumbler is the same, then no appreciable change will take place in the position of the water level even after a considerable time. But Dobariner, to his great surprise, found that with his cracked jar inverted in water and containing hydrogen, the water gradually rose inside, one and a half inches in 12 hours, two and two-thirds inches in 24 hours. It was left to Graham to give the correct interpretation of this very striking observation. He showed that hydrogen, as the lightest known gas, can get through minute apertures more rapidly than any other gas so that what occurred in Dobariner's cracked glass jar was an escape of hydrogen from the inside to the outside accompanied by a slower entrance of air through the crack. As the hydrogen escape more rapidly than the air got in, the pressure of the gas inside the jar was lowered and the level of the water rose. Thus it was that the use of a cracked vessel instead of a sound one led on to Graham's famous investigations on the diffusion of gases. A more recent and equally striking instance of a breakage leading directly to a valuable discovery has been recorded in connection with the manufacture of artificial indigo, a manufacturer which, as we have already seen, furnishes a conspicuous case of the chemist's successful attempt to build up natural products and to compete with nature herself. One of the most important steps in the manufacturing process is the production of phthalic acid from naphthalene, the chief raw material of synthetic indigo. This change can be affected by the action of hot sulfuric acid upon naphthalene, but only slowly. In the course, however, of experiments carried out with the object of improving the method of converting naphthalene into phthalic acid, the bulb of a thermometer was accidentally broken and the mercury ran out into the heated mixture. It was at once noticed that in presence of mercury, the conversion of naphthalene into phthalic acid was much accelerated, and this chance observation led at once to the desired improvement of the process. The use of mercury at this stage of indigo manufacture is now an established custom. The reader must, of course, remember that without adequate knowledge on the part of the investigator and without keenness of observation, these chance occurrences would have been of no account. The observer, even supposing he has the necessary equipment, must always be on the lookout for what is strange and unexpected, always eager to see nature and unfamiliar garb. The difficulty is that people sometimes make a valuable observation without attaching importance to it. It may be difficult to bring their new discovery into harmony with what they already know, and so they come to the conclusion that their observation must have been wrong and that their senses must have deceived them or else, by some forced explanation, they seek to fit the newly observed facts into some of the mental pigeonholes which are already available. When such difficulties crop up, the remedy is to have recourse to fresh observation and to collect more facts. In this connection there is an interesting story of Liebik whose fame as a chemist rests on many other things than extractive meat. On one occasion he prepared a liquid which in many of its properties resembled chloride of iodine, although in other respects quite different. He was struck by the differences, but without making any further experiments devised an explanation which satisfied him at the time. He was at least sufficiently satisfied to label the bottle of liquid chloride of iodine. The reader can imagine Liebik's disappointment and chagrin a few months later when he heard of the discovery by a Frenchman of the new element bromine and realized that it was this element which he had had before his eyes all the time and had labeled chloride of iodine. Liebik tells the story himself and quotes it as showing the result of adopting explanations not founded on experiment. As an example of the persistent and successful following up of puzzling observations by further experiments, nothing better can be quoted than the work which led to the discovery that there was in atmospheric air a gas, the presence of which had not previously been suspected. That argon as this gas is now called should have so long remained undiscovered is due to the fact that it is extremely similar to nitrogen. It is therefore difficult to find any way of distinguishing and separating the two gases when they are mixed together as an ordinary air. As a matter of fact, argon is rather heavier, bulk for bulk, than nitrogen. And it was this slight difference which Lord Raleigh observed and followed up. Suppose the reader tries to realize how very small was the difference in weight actually observed. The globe which Lord Raleigh used in weighing gases was filled firstly with nitrogen, atmospheric nitrogen, obtained from air by removal of oxygen, moisture and carbon dioxide. Secondly, with nitrogen prepared from various chemical compounds. Although these two samples would naturally be expected to exhibit the same behavior, the weight of the atmospheric nitrogen filling the globe was one seventh of a grain heavier than the weight, 35 and a half grains of the chemical nitrogen filling the same globe. This is obviously quite a small difference. And probably many investigators would have attributed the discrepancy to some error in their experiments and thought no more about it. Not so, Lord Raleigh. After showing that numerous possible sources of error were excluded, he succeeded in cooperation with Professor Ramsey in separating and examining the argon which is responsible for the greater weight of atmospheric nitrogen as compared with chemical nitrogen. From all this, the reader will see what a high value attaches to close and trustworthy observation, even of trifling occurrences. Elaborate apparatus and costly materials are all very well, but what is primarily essential for the true investigator is the learning and observing attitude towards nature. Anyone indeed who cultivates the habit of careful and patient observation rediscovers many things for himself and may hope to add his contribution to the romance of science. End of chapter 30. End of the Romance of Modern Chemistry by James C. Phillip.