Chapter 14 of the Romance of Modern Chemistry—this LibriVox recording is in the public domain. The Romance of Modern Chemistry by James C. Phillip. Chapter 14. Flame What Is It? The reader will by this time have become fairly familiar with the conception of combustion, and he may be under the impression that, knowing what combustion is, he has nothing more to learn about flame. This would be a somewhat rash conclusion. For to begin with, the one thing does not always accompany the other. There are cases of undoubted combustion in which there is no real flame. A little piece of charcoal, for example, burning in air or oxygen, gives out no flame, it only glows. As a matter of fact, it is only when the burning substance is in the form of gas or vapor that we get flame produced. It is true that many liquid and solid substances give a flame when they burn, but the production of the flame is preceded by their conversion into vapor. By holding a match to the wick of a candle, we first melt and then vaporize the wax which is in the wick, the vapor catches fire and the candle is lit. Once this has been done, the heat of the flame keeps the wax around the base of the wick melted, the melted wax is sucked up the wick by capillary action, and at the top it is vaporized and ignited. Flame then is something different from combustion, and may be defined as gaseous matter which has been raised to such a high temperature that it is obvious to the eye. Solids begin to emit light when they are heated to about 900 degrees Fahrenheit, but vapors must be heated to a very much higher temperature before they become visible. When a combustible vapor reacts chemically with the oxygen in the air, which is the supporter of combustion, the heat produced is intense enough to raise the vapor to the point of incandescence, a flame is produced. We have spoken here of a combustible substance and a supporter of combustion as being necessary for the production of flame, but it is well to remember that these terms are purely relative. If we would picture this world and its inhabitants as quite different from what they are, and could imagine that the atmosphere around the globe was one of hydrogen instead of air, then the gas companies would have to supply us with oxygen for lighting and heating purposes. In such a world, oxygen would be regarded as the combustible substance and hydrogen as the supporter of combustion. It is in fact easy to show that air burns in coal gas quite as readily as coal gas burns in air. In the accompanying figure four, the apparatus necessary for this experiment is shown. A lamp glass is fitted at the bottom with a cork, through which pass two tubes. The one which ends just above the cork is connected with the gas supply. The longer one serves as an air passage. The top of the lamp glass is covered with an asbestos disk, in the middle of which there is a hole. When the gas is turned on, the air in the lamp glass is driven out, and the latter then contains an atmosphere of coal gas, the excess of which escapes through the disk at the top. This escaping gas may be lighted and gives the ordinary flame of coal gas burning in air. If now, the long tube which passes through the cork at the bottom is pushed up until it reaches the burning jet at the top, something at the end of this tube is seen to catch fire, and to remain alight, giving a visible flame even when the tube is drawn down again. The explanation of this interesting phenomenon is that air is being drawn up the long tube and is burning in the atmosphere of coal gas which surrounds the end of the tube. This apparatus then, in which we can see both coal gas burning in air and air burning in coal gas, shows that the terms combustible and supporter of combustion are interchangeable. There is no real distinction. The chemical process which goes on is the same in both the flames observed. When we come to look more closely at a flame, we find that it has a structure. It may seem odd to speak of a mobile elusive thing like a flame as having a structure, and certainly with a mixture such as coal gas being consumed at some ingenious modern burner, it is not easy to detect this structure. But if we take a simple gas such as hydrogen burning at the end of a plain round tube, we find the character of the flame to be exceedingly simple. The actual flame, as will be seen from the accompanying sketch, is confined to a certain zone or sheath in which the combustion is going on, and this cone-like sheath is hollow. That the cone of flame is hollow may be very easily and prettily shown by suspending the head of a match just above the end of the tube before lighting the gas. In spite of the burning gas, the match in the inside remains unaffected. See Figure 6A. But we can go a step further and show that this hollow part of the flame contains unburnt gas by carefully putting one end of a narrow tube in the center of the cone and applying a light to the other end some distance away. We get a flame there. See Figure 6B simply because with the tube we have succeeded in leading off some of the unburnt gas from the center of the cone. A candle flame and a coal gas flame differ from a hydrogen flame only in that their structure is a little more complicated. Their general characteristics are similar. In these two cases the dark hollow cone in which is the unburnt vapor is surrounded by a white luminous zone and this again by an outer envelope of flame which is non-luminous and very difficult to see. This outermost sheath is obviously the one for which there is an unlimited supply of oxygen and anything which has escaped combustion in the luminous zone is there completely burned to carbon dioxide and water. The luminosity of flame varies very remarkably with the nature of the combustible substance and with the conditions under which the combustion takes place. A hydrogen flame is quite non-luminous.  One cause of luminosity has already been referred to in a previous chapter namely the presence of solids which are made incandescent by the heat of the flame. A coal gas flame contains in its luminous zone a host of unburnt carbon particles which are raised to a very high temperature and so give out a strong light. By mixing the gas with air before it comes to the nozzle of the burner these carbon particles are completely oxidized and the flame becomes non-luminous. Such a non-luminous flame may however again be rendered useful for purposes of illumination by the artificial introduction of incombustible solids which are made incandescent by the heat of the flame. This is what is done in the ordinary incandescent gas burners and in the limelight. There are other causes which determine the luminosity of a flame besides the presence of solid particles. There are some flames known which are characterized by very high illuminating power and in which at the same time there cannot possibly be any solid particles present. For example, phosphorus burning in oxygen produces a dazzling light but the oxide of phosphorus which results from this combustion is converted into a vapor at a red heat and it is therefore impossible that it could exist in the solid state in the phosphorus flame, the temperature of which is far above the melting point of platinum. The well-known English chemist, Franklin, who made many experiments on the nature of flame and the cause of its luminosity, once took the trouble to carry candles up to the top of Mont Blanc and was much struck by the comparatively small amount of light which they emitted when burning there. He traced this decrease of luminosity to the small atmospheric pressure prevailing at such a high level and was able to show subsequently in his laboratory that the illuminating power of a candle is much reduced when it is burning in a partially exhausted vessel. Since diminution of pressure reduces the luminosity of a flame it might fairly be expected that increase of pressure would have the opposite effect and so it turns out. A spirit lamp, which as the reader knows gives practically no light when burning in air under ordinary conditions, gives a highly luminous flame when placed under a pressure of 4 atmospheres and Franklin estimated that under a pressure of 5 or 6 atmospheres its luminosity would be equal to that of sperm oil burning under ordinary atmospheric pressure. The influence of pressure on the luminosity of a flame is most strikingly illustrated by the effect of compression on burning hydrogen. This gas burns under ordinary conditions with a pale flame, absolutely useless for illuminating purposes and it might be supposed that the want of luminosity is due to the absence of any solid product of combustion. Water the compound which results from the union of hydrogen and oxygen is of course a vapor at the temperature of the flame. But if hydrogen is burned in oxygen at 10 atmospheres pressure the light emitted by the flame is sufficient to enable the observer to read a newspaper 2 feet away. Plainly therefore the presence of solid particles is not the only thing on which the luminosity of flame depends. The pressure under which the flame is burning has a decided influence. We must recognize also another factor which has a bearing on the luminosity of a flame and that is its temperature. The hotter a particular flame is the higher is its luminosity as a general rule. One way of raising the temperature of a flame is to feed it with oxygen instead of air and the result of doing this is sometimes surprising. The temperature for instance of a hydrogen flame in air is about 3600 degrees Fahrenheit while the same flame in an atmosphere of oxygen is some 1400 degrees hotter. It is true that in the case of hydrogen no increase in luminosity results from this very remarkable rise in temperature but the behavior of hydrogen is exceptional. If a candle burning in air is transferred to a jar of oxygen the flame shrinks in size but becomes distinctly more luminous owing to the higher temperature. We may get a hotter flame also by heating the air and the gas which are supplied to the burner and such a rise in the temperature of the flame leads to increased luminosity. This was the principle applied in the so called regenerative burners in which the usual glass chimney was surrounded by a wider one closed at the bottom. The air therefore which fed the flame had to pass down between the chimneys and was very considerably heated by contact with the inner one. Such devices however for securing increased luminosity have disappeared before the incandescent burner. Not only may a flame be made hotter in various ways it is possible also to lower its temperature. That mixture of an indifferent gas that is one which takes no part in the combustion produces a marked cooling effect and a similar result is obtained by introducing into the flame some body which is a good conductor of heat. Indeed the temperature may be so much lowered by this latter device that the flame is extinguished. If for example a coil of copper wire is carefully placed over the wick of a burning taper the flame goes out immediately. In order to understand the possibility of this phenomenon we must remember that every inflammable vapor has a certain ignition temperature. That is to say for each vapor there is a point to which it must be heated in presence of air before it will catch fire and give a flame. Once it has been ignited the heat given out by the flame as a result of the chemical action raises the incoming gas above the ignition temperature and so the combustion continues. Different substances have very different ignition temperatures. The vapor of carbon disulfide can be ignited by contact with a glass rod which has been heated only to 250 degrees Fahrenheit a little higher than the temperature of boiling water. A current of hydrogen issuing from a tube is ignited by sparks from a flint and steel whereas marsh gas is quite indifferent to such treatment. The possibility of cooling an inflamed vapor below its ignition temperature may be demonstrated in a very simple manner. If as shown in figure 7a a piece of copper wire gauze is pressed down on a flame of burning coal gas which as we have already seen contains a large proportion of marsh gas no combustion takes place above the gauze although it is very easy to show that there is inflammable vapor there by bringing up a lighted match. Again if we hold the gauze an inch or two above the nozzle from which the coal gas is issuing we may light the gas above the gauze without the flame passing through to the lower side. See figure 7b. The reason of this curious result is that copper is an excellent conductor of heat and the interposition of the gauze has such a cooling effect that the inflammable vapor on the other side from the flame is kept below its ignition temperature. As already stated the ignition temperature of hydrogen is lower than that of marsh gas and if we attempted to obtain with hydrogen the results just described we should not succeed. In all cases the hydrogen flame would strike through the gauze. The remarkable power of metal gauze to limit the extension of a marsh gas flame was utilized long ago in the well-known minor safety lamp devised by Sir Humphrey Davy. Coal measures are frequently highly charged with marsh gas and large quantities of this gas find their way into coal mines. Since this fire damp as it is called is inflammable and forms a very explosive mixture with air its presence in the mines is a source of great danger and has repeatedly led to serious disasters. The risk of using naked flames in such gassy mines had to be got over somehow and Davy was able to show that if the oil flame in the miner's lamp was surrounded by wire gauze the danger of explosions was very much reduced. An explosive mixture of fire, damp, and air will not as a rule be fired by such a lamp but will indicate its presence by burning inside and so warn the miner of danger. The action of the gauze in conducting away the heat prevents the explosive mixture outside reaching its ignition temperature. The old form of Davy lamp had been found effective in some respects and has been continuously improved thus the wire gauze cut off a great deal of the light so the lower part was replaced by a glass cylinder. Then it was found that a strong draft might blow the flame against and even through the gauze with the result that an explosive mixture outside would be ignited. The newest form of the safety lamp is therefore fitted not only with a glass cylinder at the bottom to let the light out but with an iron cylinder above to shield the lamp from drafts. When the Davy lamp is brought into an atmosphere in which fire damp is present a so-called cap of pale blue flame is seen surmounting the ordinary luminous flame in the lamp. The length of this cap increases as the percentage of fire damp in the surrounding atmosphere rises hence it will be seen that to the experienced eye the appearance of the Davy lamp flame serves as a means of estimating the amount of fire damp. It is only a rough estimate however which can be made in this way. In recent years a much more accurate method of estimating the amount of fire damp in mines or of petroleum in air has come into vogue. The apparatus used is really a safety lamp in which hydrogen is burned instead of oil. In an atmosphere containing fire damp caps appear on the hydrogen flame just as in the ordinary safety lamp but owing to the fact that the hydrogen flame is much less luminous than the oil flame the caps are much more easily seen and measured. The accompanying figure eight shows the nature of these flame caps and the way in which their length varies with the amount of fire damp in the atmosphere. The gradual development of the Davy lamp is an interesting example of the way in which scientific work has been directed to the detection of danger and the preservation of life. It would indeed be difficult to estimate the saving of human lives which has resulted from Davy's discovery of the valuable properties of metal wire gauze in relation to a marsh gas flame. CHAPTER XV The reader may at some time have seen or handled those curious little things known as Prince Rupert's drops. These are obtained by allowing drops of molten glass to fall into cold water where they solidify into a tadpole-like shape. If the tip of the tail of one of these drops is nipped off with the fingers the whole thing breaks up into dust with a loud explosion. The reason is that the glass which forms the solid drop is in a state of intense strain owing to the very sudden cooling which it has undergone. The outside and the inside of the drop have cooled at different rates. The particles of the glass are in a state of unstable equilibrium and the slightest jar upsets the whole structure. There are many chemical compounds which exhibit considerable analogy with Prince Rupert's drops. The molecules of these compounds have been formed by the combination of a number of atoms, but the equilibrium between the latter is an unstable one, libel to be disturbed by the most trivial, exciting cause. An example of this curious behavior is furnished by nitrogen iodide. This extraordinary substance is prepared by the action of iodine on ammonia, and although generally quite stable in the moist state, it has been known to explode even under water. As usually obtained, it is a chocolate brown powder which explodes violently on the slightest provocation. If it is dry, the falling of dust particles, the tread of a fly, or the nearest touch with a feather, will be sufficient to make it go off with a bang. The molecules fly to pieces, and a quantity of nitrogen gas and iodine vapor is generated, occupying much more space than the original solid substance. Such a sensitive material is obviously most dangerous to handle, but there are other compounds which exhibit the same character of unstable equilibrium, and which yet can be manipulated safely if due care is taken. As we shall see later, these readily exploded substances fulfill a useful function. One which is extremely employed, and which on that account deserves special notice, is mercury fulminate. This is prepared from mercury, nitric acid, and alcohol, and when pure is a shiny white crystalline substance containing the elements mercury, carbon, oxygen, and nitrogen. It cannot be kept in a glass-stoppered bottle, for the mere friction between the stopper and the neck would cause it to explode. When struck with a hammer, mercury fulminate goes off with a very sharp report, evolving a large quantity of gas, nitrogen, carbon monoxide, and mercury vapor. It is, of course, one of the essential characteristics of an explosive, that a small quantity of the substance should yield suddenly a very large volume of gas. In the case of fulminate, it is estimated that the gas produced by its explosion would occupy at the ordinary temperature thirteen hundred to fourteen hundred times the bulk of the substance itself. But the actual volume of the gas is produced is even much larger than that, for in the explosion of the fulminate a great amount of heat is liberated, in virtue of which the gases are raised to a high temperature and occupy a much larger space. The fact that mercury fulminate when it decomposes produces heat is worthy of notice, for it is a phenomenon rather different from what might be expected. We have seen that as a rule the chemical combination of elements is accompanied by the evolution of heat. The process is said to be exothermic. This being so, we may confidently anticipate that the reverse process, the decomposition of the compound into its elements, would use up heat, and therefore, if it took place spontaneously, it would be accompanied by an absorption of heat. This is quite a sound conclusion, but it is obvious that mercury fulminate, the decomposition of which leads to the production of much heat, must belong to a different category. The secret of the explanation is that, although the formation of most compounds is accompanied by the evolution of heat, there are some endothermic compounds, as they are called, the formation of which is accompanied by absorption of heat. In this case, the reverse process, in which the compound decomposes into its constituent elements, will be accompanied by the evolution of heat, so it is with mercury fulminate, which is an endothermic compound, and like others of this class, is peculiarly liable to sudden decomposition. As a matter of fact, the explosion of mercury fulminate is accompanied by the evolution of more heat, than is involved merely in the splitting of it into the constituents, for two of the elements liberated in this primary decomposition, namely the carbon and the oxygen, immediately unite to form carbon monoxide, and as this combination is an exothermic process, the heat produced by the explosion is much augmented. The explosive disruption of the molecules of nitrogen iodide and mercury fulminate is due to the want of cohesion between the constituent atoms. It is the old story of a house divided against itself. But most of the explosions which come about, intentionally or unintentionally, depend on an altogether different principle. They are simply combustions which take place with excessive rapidity, and which result in the production of quantities of gas. In such explosions the element oxygen plays an essential part. In the first place, any inflammable gas or vapor will form an explosive mixture with air. The reader must carefully distinguish between inflammable and explosive. It is not correct to speak of coal gas as explosive. It is certainly inflammable. And when ignited at a suitable nozzle, burns quietly as long as the supply lasts. Combustion takes place only where air and gas meet. A mixture of coal gas with air is, however, a very different thing. It is inflammable at every point, explosive in fact. Hydrogen once started is rapidly propagated through the bulk of the mixture. Hydrogen similarly forms an explosive mixture with air, and illustrations of this fact are not infrequent in a chemical laboratory. For it often happens that a beginner preparing hydrogen in a flask by the action of an acid on a metal applies a light to the issuing gas before all the air has been expelled. The result of this will probably be that part of the flask will adhere to the ceiling, and the rest will be converted into fine dust. That coal gas becomes explosive when mixed with air we are frequently reminded. As from time to time we read of someone who has gone to look for a leak of gas with a lighted match or candle, and has thereby brought disaster on himself and his surroundings. When we can smell gas through a house, the atmosphere there is a mixture, possibly explosive, of coal gas and air, and if we were to carry a naked flame in search of the leak we would be as foolish as the miner who goes into a gassy mine with a lighted candle. To produce an explosive mixture of air and coal gas about 6% of the latter is sufficient, so that one cannot be too careful. The only safe course is to begin by ventilating the house thoroughly, so that the proportion of gas may be reduced below the explosive limit. In explosions of this kind where both parties to the combustion are gashes, the amount of gas produced by the explosion is relatively less than in those cases where the original unexploded substance is a solid. The increase in volume is in fact due solely to the high temperature caused by the heat of combustion. If coal gas and air in the proportion of 1 to 5 by volume were exploded in a very strong closed vessel so that no expansion is possible, a pressure of 7 to 8 atmospheres is developed and the maximum temperature reached is nearly 3,500 degrees Fahrenheit. Explosions in which the oxygen necessary for the combustion is supplied in the form of air are actually employed as sources of energy in gas and motor engines. The pressure developed when a mixture of gas or petroleum with air is exploded is used to move a piston and the longitudinal motion of the piston is converted into circular motion as in a steam engine. It has been said that fire is a good servant but a bad master and the remark is true in reference to explosive as well as to ordinary combustion. If instead of using a mixture of two gases we take a solid combustible material and mix it intimately with some other substance which not only contains a large proportion of oxygen but is fairly ready to part with some of it, then on the supposition that the combustible material yields gaseous products when it is burned, the mixture of the two solids will be a compact explosive. It will be compact because its bulk will be small in comparison with the volume of the gases produced by its explosion. Common gunpowder is an explosive of this kind. It is an intimate mechanical mixture of the three substances, potassium nitrate, niter or saltpeter, charcoal and sulfur. The first and second of these are the essential constituents of gunpowder. The sulfur is present in a smaller proportion and is added for a special purpose which will be explained later. The charcoal and the sulfur as the reader will understand are the combustible constituents and the saltpeter which forms about three quarters of the gunpowder is a compound which contains a high proportion of oxygen and which moreover is easily induced to part with some of it. This being so, saltpeter may be regarded as a compact form of oxygen. Anyhow it is easy to show that charcoal and saltpeter while quite ready to lie down peacefully together at the ordinary temperature, act violently on each other when heated. Anyone can convince himself of this by throwing a pinch of saltpeter on a glowing coal fire. It is very easy to extract the potassium nitrate from gunpowder and it is worth the reader's while to try this, since the process illustrates very forcibly what was said in an earlier part of this volume about the separation of the constituents of a mechanical mixture and shows too the kind of simple operation of which the chemist makes daily use. If the gunpowder is boiled with water and the liquid is then filtered through a paper cone made of blotting paper, the sulfur and the charcoal being insoluble in water are held back and a colorless solution runs through the blotting paper into a vessel placed to receive it. This solution when allowed to cool or when evaporated a little will deposit white crystals of saltpeter. What takes place when gunpowder is fired is essentially a combustion of the charcoal, as a result of which large quantities of gas, carbon dioxide, carbon monoxide, and nitrogen are suddenly evolved. The presence of the sulfur makes it more easy to fire the gunpowder, a lower temperature being sufficient to set it off. The function of the sulfur is therefore similar to that which it used to fulfill when employed in coating matches. In addition, however, the presence of the sulfur contributes to the rapidity with which the explosion is propagated and its oxidation by the saltpeter adds materially to the heat evolved in the reaction. The advantage of having the explosive material in a compact solid form can be seen from the fact that when gunpowder is fired in a closed space, the pressure developed is about 2,000 atmospheres, quite a different magnitude from the pressure obtained in the explosion of coal gas and air. Gunpowder is the oldest explosive known, but it is largely displaced nowadays by so-called high explosives, which in addition to several other points of distinction are practically smokeless. Anyone can understand the long cherished desire of the military and naval specialist to find some substitute for gunpowder, which when fired envelops the operator of a gun in a dense cloud of smoke. In the chemical action which accompanies the explosion, the potassium from the saltpeter forms other salts, potassium carbonate and potassium sulfate. These substances are dissipated by the force of the explosion in a state of fine division and form the smoke. From such salts, the modern high explosives are free and they are consequently smokeless or very nearly so. The applications of gunpowder are therefore more restricted than they once were. For firearms large and small, it has been replaced by smokeless powders, but it is still largely employed for blasting purposes. It enters also into the composition of fireworks, in which however, potassium chlorate frequently acts as the oxygen-supplying constituent instead of saltpeter. In the manufacture of modern high explosives, a new and interesting principle has been introduced. Gunpowder, as we have seen, is an intimate mixture of three solids, two of which are readily combustible, while the third supplies the oxygen necessary for combustion. In order that gunpowder may be a good explosive, it is manifestly essential that the mixing of the constituent should be very thorough. Provision must be made, as it were, for each combustible molecule finding near at hand another molecule out of which it can get the necessary oxygen so that when the powder is fired, no time may be lost and the explosion may be as rapid as possible. As a matter of fact, great pains are taken in the manufacture of gunpowder to secure the most thorough mixing of the constituents. Now in the modern high explosives, the oxygen is introduced not in the form of a compound which lies alongside the combustible constituent, but actually in the same molecule. In other words, chemical compounds are used as explosives instead of mechanical mixtures such as gunpowder is. This device secures an almost perfect mixing of the combustible elements with the oxygen. And the natural result is that these modern explosives go off with much greater rapidity than gunpowder. It may seem strange to the reader that we can prevail on atoms of carbon, hydrogen, nitrogen, and oxygen to form for a time a peaceable combination, which is ultimately to be rent asunder with such violence, and it must be confessed that the arrangement into which they are coaxed is somewhat of the nature of an unstable equilibrium, easily upset by irritating cause. In this respect, these high explosives have some resemblance to nitrogen iodide and mercury fulminate, but the process of explosion in the former cases is a real combustion, which the explosion of nitrogen iodide and fulminate is not. Gun-cotton, nitroglycerin, and picric acid, which either alone or in combination with other substances form the majority of the high explosives, are obtained by the action of nitric acid on cellulose, glycerin, and carbolyic acid, respectively. Nitric acid is a substance which contains a large proportion of oxygen, and its action on these materials is such that new substances are formed provided with a good deal of the oxygen, which was previously in the nitric acid. The various explosives just mentioned are alike in containing a particular grouping of nitrogen and oxygen atoms known to chemists as the nitro group, and gun-cotton and the rest of them are therefore frequently referred to as the nitro explosives. It is indeed very remarkable that a harmless thing like ordinary cotton, when treated with nitric acid, should undergo such a fundamental change and so be converted into a powerful explosive. Great care has to be taken in the process of manufacture, and only the best white cotton waste, perfectly free from grease and dirt, can be employed. During the time the cotton is in contact with the nitric acid, the temperature must be kept down, and subsequently every trace of acid must be washed away with water. It was owing to the want of attention to this last simple precaution that many of the explosions which attended the early manufacture of gun-cotton were due. Any trace of acid left in the finished article acts like an irritant and leads sooner or later to the decomposition of the explosive. Gun-cotton is a most curious substance. It takes fire much more easily than gunpowder, and the rate at which it burns altogether depends on the way in which it has been ignited and the conditions to which it is subject. A piece of loose gun-cotton may actually be burned on the hand without scorching the skin, merely by touching it with a hot glass rod. It can be fired on the top of a heap of gunpowder without igniting the latter. Under such conditions the combustion of the gun-cotton is rapid but not explosive. When, however, it is fired in a confined space and the flame from the portion first ignited is driven into the remaining mass, the temperature is forced up and the combustion becomes an explosion. It was therefore very naturally thought for a long time that in order to utilize the explosive force of gun-cotton it must be enclosed in some strong casing. Some 40 years ago, however, the very interesting discovery was made that this was unnecessary and that gun-cotton, which was merely compressed, not confined in an enclosed space could be exploded by a detonator, such as mercury fulminate. It is indeed a curious fact that if a little of this latter substance is exploded in the immediate neighborhood of a mass of unconfined, compressed gun-cotton, the gun-cotton explodes with extraordinary violence. It is well known that if a violin is made to emit a particular note, the string of a second instrument in the immediate neighborhood, if tuned to that note, will take it up and vibrate spontaneously. Some chemists have thought that something analogous to this takes place when mercury fulminate is exploded in contact with gun-cotton. The vibrations set up by the detonator are supposed to excite similar vibrations in the gun-cotton, so much so that the latter undergoes what might be called a sympathetic decomposition. Whether this is the correct explanation or not, there is no doubt that gun-cotton fired by a detonator gives a much greater effect than the same material fired in the ordinary way. This increased effectiveness is due to the greater rapidity of the explosion induced by the detonator. Thus, if a train of ordinary gun-cotton is touched with a hot rod, the resulting combustion advances only a few feet in several seconds. Whereas, if a train of compressed gun-cotton is detonated by mercury fulminate, it is estimated that the explosion is propagated along the train at the rate of 200 miles a minute. Perhaps still more curious and valuable was the discovery that wet gun-cotton, which is not explosive under ordinary conditions, could be detonated as easily as the dry material. A red-hot iron may be put into a mass of wet gun-cotton without setting it on fire, and a government committee in order to demonstrate incontestably the possibility of safely storing this explosive in the moist condition, once instituted experiments in which an iron case containing a ton of wet gun-cotton was put in a magazine and surrounded with shavings and other inflammable material. This was then ignited, and when the combustion was over, the case of wet gun-cotton was recovered, none the worse for its baptism of fire. Wet gun-cotton, however, can be at once exploded by detonation, provided only that a little of the dry material is in contact with the detonator. The old saying, therefore, keep your powder dry, is applicable only in a very limited sense to gun-cotton. It is, as a matter of fact, always stored in the wet state, containing about 20% of water, and it may be used in this condition in torpedoes and submarine mines. A more dangerous explosive than gun-cotton is nitroglycerin, a liquid obtained by the action of nitric acid on glycerin. The most extraordinary precautions have to be taken in the handling of this material, and it is only by a strict observance of these that a repetition of the disasters which marked the early years of nitroglycerin manufacture is avoided. So serious were the accidents which occurred with nitroglycerin some 40 years ago that several governments went to the length of altogether prohibiting the use of the explosive. Chemists soon discovered, however, the necessary precautions that have to be taken in the manufacture and handling of nitroglycerin, and at the present day, large quantities of this explosive are prepared. In a nitroglycerin factory, the sheds in which the various operations are carried on are well separated from each other and surrounded by banks of earth or sand. In order to avoid any risk of a spark being produced and setting off the nitroglycerin, all workers have to wear special clothing, boots with iron nails are absolutely prohibited, and in their place, shoes of rubber, felt, or sewn leather are employed. Girl operatives are forbidden hairpins, and no one is allowed to carry any article made of iron, such as knives or keys, for these by friction might give rise to a spark. Such precautions being necessary, the reader will understand that the handling and transport of nitroglycerin by the uninitiated person is fraught with great danger. Hence, before it leaves the factory, it is converted into various forms, which involve less risk. The commonest of the explosive materials, thus based on nitroglycerin, is dynamite. Certain substances have the power of soaking up or absorbing nitroglycerin, and one of these, which has been found very satisfactory, is an infusorial earth known as Kieselger, which takes in as much as three times its weight of nitroglycerin. The resulting product is dynamite, a material which is less violent than the parent substance and more easily and safely handled. Indeed, it was not until the little device of employing absorbent Kieselger was adopted that the manufacture of nitroglycerin assumed practical and commercial importance. This may be gauge from the fact that in 1870, the world's output of dynamite was only 11 tons, while 20 years later, it had risen to 12,000 tons. Dynamite, like gun-cotton, burns without danger when loose and in small quantity, but when fired by a detonating fuse of mercury-filminate, it explodes with extreme violence and rapidity. Indeed, it is estimated that the time occupied in the explosion of a dynamite cartridge is only 1.24 thousandths of a second. One consequence of this is that when dynamite is used for blasting rock, the usual boreholes may frequently be dispensed with, and the explosives may be laid on the top of the rock, covered merely with a little earth or clay. The greater sensitiveness of dynamite as compared with gun-cotton is well illustrated by their different behavior when fired at with a rifle. If a wooden box filled with dry compressed gun-cotton is exposed to this treatment, the contents of the box are generally inflamed but never exploded. Dynamite, on the contrary, and other explosives derived from nitroglycerin cannot contain themselves when treated in such a fashion and go off with great violence. End of Chapter 15. Chapter 16 of the Romance of Modern Chemistry. This LibriVox recording is in the public domain. The Romance of Modern Chemistry by James C. Phillip. Chapter 16, Below Zero. It is an old tale that Fahrenheit took as the zero of his thermometer the lowest temperature which was observed by him at Danzig during the winter of 1709 and one of his contemporaries remarks that nature never produced a cold beyond zero. This is quite a mistaken view for plenty of cases are on record in which considerably lower temperatures have been observed as the direct result of natural cold. By artificial methods, it is possible to realize a much greater degree of cold and within the last 10 years, temperatures of about minus 400 degrees Fahrenheit have been reached. As the mercury in our thermometers freezes at about minus 40 degrees Fahrenheit, the reader will see that a lowering of the temperature to minus 400 degrees brings us to altogether new conditions. The way in which chemists and physicists have gradually pushed forward in the region of low temperatures is very remarkable and their discoveries are not only a fascinating interest to the student of nature but have in some cases proved a practical and commercial value. The ambition to get farthest north has led to many thrilling adventures but the arctic exploration carried out recently in the laboratories of England and the continent is not a wit less romantic. The first step in the direction of low temperatures is taken when we can start with substances at the ordinary temperature and either by utilizing some inherent property of these substances or by treating them in some special way induce the temperature to fall, say below the freezing point of water. The mere bringing together of two substances may lead to either a rise or a fall of temperature. The reader may remember the reference made in a previous chapter to the fact that when sulfuric acid and water are mixed, so much heat is produced that the containing vessel becomes too hot to hold. The opposite effect is frequently observed when other substances are mixed with water. When salt, peter or salamoniac for example is stirred into water, the cooling effect is very noticeable and by this simple method quite a considerable fall of temperature is produced. A mixture of these two salts added to an equal weight of water at 50 degrees Fahrenheit brings the temperature down to 10 degrees Fahrenheit. More marked and more persistent cooling effects are obtained. If instead of adding salts to water we mix them with powdered ice or snow. Anyone who procures common salt and snow stirs them up well in the proper proportions and puts a thermometer in the mixture will see the mercury fall below zero Fahrenheit. Such a freezing mixture may be used not only for getting a low temperature in scientific experiments but also for the equally practical, if less exalted, object of making ices. The question may very naturally be put. Why should the mere bringing together of salt and snow result in such a marked fall of temperature? The answer to this question is very closely connected with what was said in a previous chapter about the melting point of alloys. Attention was then directed to the fact that the melting point of any metal is lowered by the presence of another metal. To the case of ice and salt, a similar rule applies. Every schoolboy knows that pure water freezes at 32 degrees Fahrenheit, zero degrees centigrade but it is a curious fact that water containing salt does not freeze until a lower temperature has been reached. That means that a mixture of snow and a little solid salt should, strictly speaking, be in the liquid condition at 32 degrees Fahrenheit. There cannot therefore be true equilibrium between snow and salt at this temperature. Now in nature, things are always trying to get into the most stable condition possible, in other words, to reach their true equilibrium. Water finds its own level. A hot and a cold object put side by side, gradually and of their own accord, assume the same temperature, while positive and negative electricity unite whenever they get the opportunity. Similarly, snow and salt, when mixed together at 32 degrees Fahrenheit, do their best to get into that condition which nature has prescribed as the most stable one for them at that temperature. The result is that the snow melts and the salt dissolves in the melted ice. Now both these processes use up heat. As they take place spontaneously, this heat is taken from the surroundings and the temperature of the mixture and of the containing vessel falls. The reader will at once admit that heat is required to melt snow and he will see that the addition of salt is an ingenious way of persuading the snow to melt and so to abstract a definite amount of heat from its surroundings. For the same quantity of heat is always required to melt a pound of snow, whatever be the way in which we cause the melting to take place. So far as we have gone then, methods of producing cold depend either on dissolving a solid in a liquid or on making a solid melt by a little scientific stratagem. But just as we can utilize the change of solid into liquid as means of reaching lower temperatures, so we can employ another change of state for the same purpose, the change namely in which a liquid passes into the condition of a vapor. We usually convert a liquid into a gas or vapor by heating it. For the conversion of water at 212 degrees Fahrenheit into steam at 212 degrees, heat is as necessary as it is for the conversion of ice or snow at 32 degrees Fahrenheit into water at the same temperature. Evaporation then, that is, the process by which a liquid is changed into a vapor only takes place when heat is supplied. If by any means we can cause evaporation to take place without the external application of heat, then the necessary heat will be taken from the evaporating liquid itself and its surroundings. Under these circumstances, evaporation produces cold. A very simple way of causing a volatile liquid to evaporate rapidly without heating is to blow a strong current of air through it. That by this method a considerable reduction of temperature may take place can be shown by a very simple experiment. A small pool of water is made on the top of a flat wooden block and in this pool is set a flask containing strong ammonia solution. A strong current of air is then blown through the liquid with the aid of a bellows. The ammonia evaporates rapidly and before long the flask is frozen hard to the block. With these two general ways of producing cold at their disposal, Faraday and other chemists after him have been able to obtain in the liquid state many substances which exist ordinarily as invisible gases. The point to which the temperature of a gas must be lowered before it begins to liquefy will of course vary from one case to another. If we could imagine the temperature of our globe as being normally about 250 degrees Fahrenheit then water would exist only in the form of vapor or steam. And in order to liquefy it we should have to bring the temperature below 212 degrees the boiling point of water. At any temperature lower than 212 degrees steam will condense under the ordinary pressure of the atmosphere. Now we must remember that every other liquid has its own boiling point and substances which we know as gases are simply liquids whose boiling points are at a temperature lower than that prevailing on the surface of the globe. Sulphur dioxide for example, the colorless choking gas which is produced when sulfur is burned is very easily obtained as a liquid at temperatures not much below the freezing point of water. The boiling point of this liquid sulfur dioxide is 18 degrees Fahrenheit under the ordinary pressure so that when the gas is passed through a tube surrounded by a freezing mixture of ice and salt it condenses to the liquid form. Just as steam would do if it were passed through a tube surrounded by cold water. It is in fact quite easy to obtain liquid sulfur dioxide and it is now sold in siphons just as if it were so much soda water. When we come to gases like ammonia and carbon dioxide which are less easily condensed it is found advisable to use high pressure as an aid to liquefaction. The reader will understand the object of this if he remembers that the boiling point of the liquid gradually rises with the pressure to which it is exposed. For example, water boils at 212 degrees Fahrenheit under a pressure of one atmosphere but at 250 degrees when the pressure is two atmospheres. Conversely then, when a gas is kept under high pressure less cooling is necessary to bring it below its boiling point. By the combined application of cooling and compression both ammonia and carbon dioxide are readily obtained in the liquid form and they are now commercial articles sold in steel bottles or cylinders. With the aid of liquid ammonia and liquid carbon dioxide we are able to go a long step farther in realizing low temperatures. For the cold produced by their rapid evaporation is very intense. This is well shown by what happens when the tap of a liquid carbon dioxide bottle is opened. The liquid is forced out in a fine jet by the high pressure which prevails in the bottle and the cold produced by the evaporation of the outer portion of the jet is so great that the inner portions are solidified to a white, snow-like powder. If a coarse canvas bag is tied over the nozzle of the bottle while the liquid is escaping a quantity of this curious solid carbon dioxide may be collected. Carbonic acid snow, as we may call it can be placed on the hand without danger but if pressed into the skin a serious blister is produced the effect being pretty much the same as that caused by a red hot metal rod. A number of interesting experiments can be made with solid carbon dioxide. If, for example, some of it is placed on the top of a little mercury in a dish and some methylated spirit or ether is added the mercury is very quickly frozen to a hard mass. In fact, the temperature reached in this way is as low as minus 112 degrees Fahrenheit and if a mixture of ether and carbonic acid snow is made to evaporate very rapidly by connection with a suction pump the temperature reached is considerably lower still. The temperature of minus 112 degrees Fahrenheit just mentioned is the boiling point of liquid carbon dioxide under atmospheric pressure. This substance is a conspicuous example of what may be called cold boiling liquids and the reader will see that boiling does not necessarily mean a high temperature. That liquid carbon dioxide kept in an open vessel is very cold can be simply shown by thrusting a piece of metal into it. There is a hissing and a bubbling exactly similar to what is observed when a red hot poker is thrust into water. So that relatively to the piece of metal which is at the ordinary temperature liquid carbon dioxide is exceedingly cold. For purposes of refrigeration in ice making and cold storage liquid ammonia is very largely used nowadays. Rapid evaporation of this liquid under a suction pump gives a very low temperature and if Brian is circulated around the pipes in which the evaporation is taking place it is rendered so cold that water may be frozen by it in large quantities. The success of chemists in liquefying such gases as carbon dioxide and ammonia is now overshadowed by the greater achievements of the last 10 or 15 years during which period liquid air and liquid hydrogen have been produced in quantity. This has become possible by the introduction of an altogether new principle in gas liquefying machines, a principle which deserves a few words of explanation. We regard a gas as consisting of an enormous number of separate particles or molecules moving rapidly in all directions. Under ordinary conditions the total volume of the molecules is very much less than the space in which they move. In other words, the molecules are relatively and on the average not very close to each other. When however the gas is compressed the molecules are crowded together and they come within range of each other's attraction. Each molecule exerts an attractive force on its neighbors and is in turn attracted by them. So that when a highly compressed gas is allowed to expand there is a social force resisting the separation of the molecules which is involved in the expansion. In overcoming this social force work must be done and for the performance of this work heat is required. This heat is taken mostly from the gas itself which therefore exhibits the phenomenon of self-cooling. All this may be put more definitely and practically by saying that when highly compressed air is allowed to expand through a small nozzle or a porous plug it becomes slightly colder. In the actual machines for making liquid air the device is further adopted for allowing the expanded and slightly cooled air to circulate around the coil of tubing through which the next slot of compressed air is approaching the nozzle. In such a regenerative process the cooling effects are accumulated and the air which circulates through the machine alternately compressed and expanded becomes gradually cooler until at length it condenses and drops into a vessel placed to receive it. A vessel which is to contain liquid air or liquid hydrogen must be specially constructed if it is to be of any use at all. If we were to put liquid air which boils at minus 347 degrees Fahrenheit in an ordinary glass vessel we should very shortly see the last of it owing to the heat communicated through the walls of the vessel. That is in fact exactly what would happen if we put a glass of water in a hot air bath kept at 400 or 500 degrees. Such a communication of heat however may be very much diminished as Professor Dewar has shown by using double walled vessels and removing the air from the space between the walls. Sections of two such vessels, a tube and a flask are shown in figure nine. The importance of removing the air from the space between the walls will be realized when it is remembered that under ordinary circumstances that space is filled with molecules of oxygen and nitrogen rushing hither and thither. With the outer wall near the temperature of the atmosphere and the inner one in contact with liquid air these molecules act like an army of heat carriers. Each molecule as it strikes the outer wall will take up so much heat which sooner or later it delivers up to the inner wall only to return for a fresh supply. When the air is left in the intervening space the transfer of heat is therefore very rapidly affected and the liquid air in the vessel soon evaporates. When this space however is rendered free from air the heat carrying molecules are removed and the inner tube is more perfectly cut off from any heat exchange with the atmosphere. The insulation of the inner tube is made still more complete by silverying the inside of the outer one for at such a bright surface the heat rays are reflected. In a doer vacuum flask surrounded by a non-conducting material like cotton wool liquid air may be kept for over 24 hours and the examination of its properties is thus rendered possible. Not only is it possible to study the properties of liquid air itself but we can see how other substances behave when cooled to the temperature of liquid air. Their behavior then is frequently quite different from what it is under ordinary conditions. Grass, leaves of plants and India rubber for example become so brittle when kept for a short time at the temperature of liquid air that they can easily be powdered in a mortar. An egg after immersion in this wonderful medium becomes so hard-boiled that it may be severely knocked about without being damaged. Chemicals too which react vigorously at the ordinary temperature become mutually callous when cooled to the boiling point of liquid air. It has just been stated that liquid air boils at minus 347 degrees Fahrenheit but by boiling under reduced pressure the temperature is lower to a point at which the air of the atmosphere will condense straight away. This may be very simply and very beautifully shown in the following manner. A doer vacuum tube C figure 10 filled to the extent of two thirds with liquid air is provided with a cork. Through this cork there are past one an empty glass tube A closed at the bottom and dipping into the liquid air. Two a bent tube B open at both ends and leading to the vacuum pump. When the latter is turned on the liquid air in the vacuum vessel begins to boil vigorously under the reduced pressure and in consequence of the low temperature thus produced air gradually condenses and collects as a liquid in the previously empty tube A. That of course is the natural result of bringing the temperature of the air in A below its boiling point. The composition of liquid air is not quite the same as that of gaseous air for the simple reason that oxygen is rather more easily condensed than nitrogen so that liquid air contains a higher proportion of the former. Further if liquid air is allowed to evaporate slowly it becomes very much richer in oxygen for the nitrogen is the more volatile constituent and passes off more readily leaving behind a liquid with a higher proportion of oxygen. On this fact is based a method for the extraction of oxygen from the atmosphere. More wonderful even than liquid air is liquid hydrogen. It is more difficult to prepare for in applying the regenerative cooling process to hydrogen it is necessary first of all to cool the compressed gas to a low temperature by means of liquid air before it is allowed to issue from the nozzle of the apparatus. Dewar however has made considerable quantities of liquid hydrogen and on one occasion over a gallon of the substance made in his laboratory was carried through the streets of London to the rooms of the Royal Society. This quantity would weigh only about 11 ounces for liquid hydrogen is by far the lightest liquid known to the chemist. Bulk for bulk it is only one 14th as heavy as water. Some very interesting experiments have been made at these extremely low temperatures on the vitality of bacteria and seeds. Typical bacteria were exposed for a number of hours to the temperature of liquid air but their vitality was not destroyed by this treatment. Barley and peas have been kept for six hours in the liquid itself and yet when they were sown subsequently in the ordinary way no falling off in the power of growth could be detected. It is possible to get a very high vacuum in a closed glass tube by simply immersing one end of it in liquid hydrogen. The boiling point of the ladder is about 100 degrees Fahrenheit lower than the boiling point of liquid air and the mere contact of one end of the tube with the liquid hydrogen is sufficient to condense the air which it contains so completely that none is left in the upper part. The attainment of such low temperatures has raised the very interesting question as to what prospect there is of ever reaching the absolute zero. On various grounds, chemists and physicists believe that at a certain temperature minus 460 degrees Fahrenheit the existence of a gas as such would cease to be possible. The movements of the molecules which we have learned to regard as characteristic of a gas would be so paralyzed by the intense cold as to stop altogether. The chill of death would settle on their activity. This temperature is called the absolute zero and is in the eyes of low temperature investigators what the North Pole is to the Arctic explorers. In this connection, the year 1908 will be remembered as the year in which helium, the most obstinately gaseous substance known, was reduced to the liquid state. The labor expended in procuring even so little as two ounces of liquid helium can hardly be appreciated by the lay reader but it may be mentioned that the preliminaries consisted in the preparation of 16 gallons of liquid air and four and a half gallons of liquid hydrogen. By boiling liquefied helium under reduced pressure the temperature of minus 454 degrees Fahrenheit was reached only six degrees from the absolute zero. It might therefore be thought that this interesting point was practically within reach for an interval of six degrees does not seem a very serious obstacle. At these low temperatures however an advance of even one degree is a very great matter and it must be confessed that there is no immediate prospect of reaching the chilly goal. End of chapter 16. Chapter 17 of the Romance of Modern Chemistry this LibriVox recording is in the public domain. The Romance of Modern Chemistry by James C. Phillip. Chapter 17, Chemistry at High Temperatures. It is not only into the region of low temperatures that such a surprising advance has recently been made much has been achieved also in the other direction and it has lately become possible to realize within a limited space a degree of heat far beyond what can be produced with the aid of ordinary fuel alone. This attainment of extremes of heat and cold has immensely widened the range of temperature over which the chemist can study the properties of matter and as a result many new substances as well as new methods of making old substances have been discovered. One result of low temperature research as we have seen is that all the known gases have been reduced to the liquid state and in many cases even solidified. Similarly, by the recent application of very high temperature the most refractory solids have been melted and even vaporized. Apart however from the extremely high temperatures reached by recent methods there are easily attainable temperatures at which many substances existing ordinarily as stable solids are first melted and then converted into vapor. Now the possibility of changing any substance into vapor without decomposing it involves a great deal. For it means that a distillation can be carried out and as a method of separating and purifying chemical compounds this is one of the most ancient and valuable laboratory operations. When for instance a salt solution in other words a mixture of a salt and water is boiled and the steam condensed it is found to be pure water perfectly free from salt. This operation of boiling and then condensing the vapor distillation as it is called obviously makes it possible to separate salt and water simply because the water is easily vaporized in contrast to the salt. The same principle may be applied in numberless other cases. Metals for instance which are comparatively volatile such as mercury and zinc may be separated by distillation from others such as copper and iron which are mixed with them and which are much less easily vaporized. A rise of temperature however not only makes it possible to melt and then vaporize many solid substances but it also has the general effect of weakening the bonds which hold together the atoms in a molecule. On heating a chemical compound the chances are that when a certain temperature is reached it begins to break up into simpler compounds or even into the constituent atoms. This change is known as decomposition or dissociation. The former term is applied to the case in which the atoms are simpler molecules having been once separated by heat show no signs of coming together again on cooling. They have done with each other for good and all. But in many cases the interesting observation has been made that the separation caused by heating the compound molecule is spontaneously reversed on cooling and the compound is reformed provided of course that the atoms or simpler molecules have been allowed to remain side by side. The effect of heating such a compound is described as dissociation and this is followed on cooling by an association of the separated atoms or molecules. These cases of dissociation are of great interest and there are many common substances which undergo this change on heating. Carbonate of lime in its various forms, limestone, chalk and marble is one of them. When heated it breaks up into quick lime, calcium oxide and carbon dioxide. If the latter were left in contact with the quick lime then on cooling recombination would take place and the carbonate of lime would be regenerated. This being so, the reader may ask how it is possible to convert limestone into lime by heating or burning in kilns. The explanation is quite simple for in the lime kilns the carbon dioxide is constantly being removed by the draft so that when the lime begins to cool the carbon dioxide with which it would gladly have combined is not there. This interesting phenomenon of a chemical change taking place in one direction at a particular temperature and in the opposite direction at another temperature is very well illustrated by one of the common methods for obtaining oxygen from the atmosphere on the large scale. There is a solid compound somewhat similar to quick lime known as barium oxide which at a temperature of 1,100 degrees Fahrenheit or there about readily takes in more oxygen forming a new substance, barium dioxide. The grip of the latter however on the extra atom of oxygen is not very secure and by raising the temperature to 1,560 degrees it can be so weakened that the gas is released and may be collected. In the actual manufacturing process a current of air is pumped into retorts heated to 1,300 degrees Fahrenheit and containing barium oxide which takes up the oxygen and allows the nitrogen to pass on. When the charging is completed the current of air is shut off and the retorts which now contain a certain proportion of barium dioxide are connected with a suction pump. The effect of this diminution of pressure is the same as that of a rise of temperature and from the engineering point of view it is better to alter the pressure than to alter the temperature. The barium dioxide accordingly gives up the extra oxygen which it extracted from the air and the oxygen so obtained is compressed in steel bottles under a pressure of 120 atmospheres and sent into the market. The barium oxide may be used over and over again in the same fashion and theoretically at least a given quantity of this substance should suffice for the winning of unlimited quantities of oxygen from the air by alternate association and dissociation. The highest temperatures reached in furnaces fed with ordinary fuel the furnaces employed for technical purposes lie about 3,200 degrees Fahrenheit but it is possible to get a few hundred degrees beyond that with the oxy hydrogen blowpipe. When we feed a coal gas flame with a blast of air as in an ordinary blowpipe we get a very high temperature but the effect is wonderfully increased by substituting oxygen for air. The reason of this is not far to seek. Roughly speaking air consists of one part oxygen to four parts of nitrogen. The latter gas although it takes no part in the combustion yet passes through the flame and has to be warmed up thereby absorbing a considerable proportion of the heat produced by the combustion. In the oxy hydrogen or oxy coal gas flame the nitrogen is not there to dilute the active oxygen so that the temperature reached is very much higher. The increased heating effects secured in this way makes it possible to melt platinum and the operation is actually carried out in the winning of this metal from its ores. The furnace in which the platinum is melted must obviously be of some material which has a higher melting point still and quick lime is found to fulfill this requirement. Pipe clay when put in the oxy hydrogen flame is immediately fused to a sort of glass while gold and silver not only melt but vaporize into a dense smoke. Even the temperatures of the oxy hydrogen or oxy coal gas flame however are comparatively chilly in comparison with those which are now attainable in the electric furnace. Within the last 15 or 20 years the efficiency of this furnace has been so improved that temperatures of 6,000 degrees Fahrenheit can be reached. And under these conditions many common substances are found to behave in a most extraordinary manner. Such a furnace consists essentially of a hollow box made of some non-conducting material into the cavity of which project two carbon rods. An electric arc is established between these rods with the result that an extraordinary degree of heat is attained in the cavity of the furnace. As was said in a previous chapter electricity is generated as a rule in a dynamo driven by an engine, which in its turn depends for its power on the chemical process of combustion. In the electric furnace we merely get back a certain fraction of the heat which was produced in the combustion and the reader might be inclined to consider the whole affair a wasteful cycle of operations. Economical it certainly is not but the advantage lies in this that in the electric furnace the heat which originally was distributed over a considerable space is concentrated in a fraction of a cubic foot. The effect is locally intensified, the temperature is higher and this is of the utmost importance as it turns out for certain processes. It must be remembered too that these objections on the score of economy lose their force in cases where water power is available for driving the dynamos as it is for example at Niagara and in Norway. One of the chief difficulties in working at such high temperatures as are reached in the electric furnace is to find a suitably refractory substance out of which the enclosing box may be constructed. Up to a certain point quick lime is an excellent material. As its employment in the oxy-hydrogen limelight shows it is not easily fused and it has further the recommendation of being a very poor conductor of heat. This latter property is well demonstrated in an experiment carried out by the French chemist Moussin whose name will always be associated with the utilization of the electric furnace. In one of the lime furnaces which he employed the top consisted of a slab of quick lime rather less than one and a quarter inch thick. The electric arc was allowed to play for 10 minutes in the cavity below the slab the temperature rising probably to over 5,000 degrees Fahrenheit. In spite of this the slab could be handled on the outside without discomfort while examination of the lower surface which had been in contact with the arc showed that the quick lime had actually been melted over an area of several square inches. The tremendous heat therefore which had been generated in the cavity of the furnace had been completely kept in by a layer of lime one and one quarter inch thick. With bigger currents and more powerful arcs even lime furnaces become useless. The lime fuses and runs like water and ultimately it boils producing clouds of smoke. The difficulty may be partly surmounted by enlarging the cavity of the lime furnace and making a little platform of one half inch plates of magnesium and carbon arranged alternately. Using a device of this sort Monson was able to study the behavior of a large number of substances at temperatures up to 6,000 degrees Fahrenheit. In the earlier part of this chapter it was said that compared with zinc at least copper was not volatile. Things are quite different however at the temperature of the electric furnace as appeared from Monson's experiments.  Soon after the current was turned on dazzling flames 18 inches long burst out violently through the openings at the ends of the furnace. These flames were due to copper vapor burning in the air and it was found after the experiment was over that the copper left in the furnace now weighed only three ounces. One ounce of the metal having been converted into vapor. Similar and equally surprising results were obtained with such refractory metals as silver, gold and platinum. One of the most surprising things accomplished in Monson's electric furnace was the vaporization of silica. This substance as the reader may already be aware is the oxide of the element silicon and is the main constituent of sand and quartz. Indeed quartz is nearly pure silica. It is melted with the greatest ease in the electric furnace and after seven to eight minutes issues through the openings as a bluish smoke or vapor. Another proof of the really fervent heat which is generated in this way. While the electric furnace has astonished us by revealing the volatility of even the most staid and refractory materials known to the chemist it has at the same time brought to light a number of substances which are quite at home at these high temperatures. Indeed it is the electric furnace alone which has enabled us to prepare them. Among these substances are the carbides, compounds of the metals with carbon. And it is in the preparation of one of these namely calcium carbide that the electric furnace is most extensively employed at the present time. Monson showed that by heating a mixture of pure lime and carbon in the electric furnace calcium carbide could be readily obtained and this is the method now employed on the manufacturing scale except that limestone and coke are used as crude materials instead of lime and carbon. The use of limestone instead of lime does not really involve any difference for at the high temperature employed the limestone loses its carbon dioxide and is converted into lime. The other material the coke is at best a very impure form of carbon so that the calcium carbide obtained in the manufacturing process is not a pure product. The essential chemical change which goes on in the electric furnace during the formation of carbide is an extremely simple one. Lime is a compound of two elements calcium and oxygen but this union is broken by the interposition of carbon at the high temperature of the furnace. The latter element combines with both the calcium and the oxygen so that these two are separated. The new compounds formed calcium carbide and carbon monoxide are quite distinct in their properties for the former remains in the furnace in a fused condition while the latter is a gas and escapes at once. As the reader probably knows the characteristic feature of calcium carbide is that it gives off an inflammable gas acetylene on contact with water. One usually regards flame and water as essentially antagonistic but here is a case where water is a sine qua non in the production of an inflammable gas. The curious action between water and calcium carbide molecules consists simply in a change of partners. The hydrogen of the water unites with the carbon of the carbide forming acetylene while the oxygen of the water combines with the calcium of the carbide forming quick lime which promptly slakes in excess of water. Acetylene when burned at specially constructed nozzles gives a very brilliant flame more like sunlight in its character than any other artificial illuminant. On this ground there is much to be said for the use of acetylene for lighting purposes. The portable nature of calcium carbide and the ease with which the gas can be obtained from this material are circumstances also which have favored the introduction of acetylene as an illuminant especially in places where electricity and coal gas are not available. The eagerness of carbon to unite with both calcium and oxygen at the temperature of the electric furnace as illustrated by the formation of calcium carbide has found a recent interesting application in the manufacture of phosphorus. The chief source of this element is bone ash which consists to a large extent of calcium phosphate a compound of calcium, phosphorus, and oxygen. In the older process for obtaining phosphorus from bone ash it was put through quite a number of distinct operations but nowadays with the aid of the electric furnace a much more straightforward plan is feasible. By simply mixing the bone ash with carbon and heating in the furnace the carbon annexes both the calcium and the oxygen forming calcium carbide and carbon monoxide. The phosphorus on the other hand escapes as a vapor and is condensed under water in the usual manner. It is not only in the electric furnace that the high temperature of the electric arc has been utilized but also in connection with the interesting problem of the utilization of nitrogen from the atmosphere for agricultural purposes. For the fertilization of the soil large quantities of nitrogenous material are required which are at present derived to a great extent from Chile where extensive deposits of sodium nitrate Chile saltpeter as it is called are found. Those who should know best are of opinion that these nitrate beds will be exhausted in 30 years or there about and hence it was that Sir William Crookes in his presidential address to the British Association in 1898 insisted on the necessity of discovering some way by which the great store of nitrogen in the atmosphere could be made available. The problem is by no means easily solved for nitrogen is very slow to enter into combination with other elements. With the aid of the electric arc however it is possible to induce some of the oxygen and nitrogen in the air to unite forming nitric oxide which in its turn can easily be converted into nitric acid or nitrates. This has been known to chemists for a long time but it is only recently that the difficulties in the way of making the process a commercial success have been overcome. Within the last few years the necessary plan for carrying out this process on the large scale has been set up in Norway where power is cheap. The factories there are now turning out large quantities of nitrate of lime suitable for fertilizing purposes and capable of replacing the natural nitrate brought from Chile. End of chapter 17. Chapter 18 of the romance of modern chemistry. This LibriVox recording is in the public domain. The romance of modern chemistry by James C. Phillip. Chapter 18, chemistry of the stars. From a study of the electric furnace and of the curious effects which very high temperatures have on the farthest substances known to the chemist it is but a short step to a consideration of the conditions which prevail on the sun and other heavenly bodies where nature herself has concentrated so much heat. What is the constitution of the sun and stars? Do the elements of which they are composed differ from those with which we are familiar? How is their condition affected by the high temperatures which prevail there? Such are some of the questions which occur to us in this connection. To these and similar inquiries the older astronomy had no reply. It displayed a marvelous power of calculating times and seasons of accurately predicting the movements of the celestial army but as to the materials of which these other worlds were built up it had nothing to say. At one time indeed it looked as if astronomical science had come to the end of its tether. It had attained such a thorough mastery of the problems connected with the movements, the size and the distances of the heavenly bodies that no very startling advance was to be expected in that direction and there was no hope that the constitution of these bodies would ever be discovered by working on the old lines. How different is the outlook nowadays? Much information is available as to the actual elements of which the sun and stars are composed and it may with truth be said that we know more about the chemical composition of the heavens above than about that of the earth beneath. For man with all his wonderful achievements have scratched only the surface of the globe and we can but speculate about the materials of which the interior is composed. It is indeed exceedingly probable that large quantities of iron exist in the interior of the earth. Firstly, on account of the fact of terrestrial magnetism and secondly, because the average density of the earth as a whole is considerably greater than the average density of the crust pointing to the presence of some heavy metallic material at lower depths, but no direct evidence is forthcoming as to the actual composition of the interior of our globe. If a scientist were asked, however, to name some of the materials of which the sun is composed, he would be ready with an unhesitating answer and this would be the case also in regard to many of the stars. How has this come about? How is it that we can speak now so confidently about the constitution of heavenly bodies whose distance is measured in millions of miles and whose very presence in the sky speaks so eloquently of the unattainable and the mysterious? We certainly cannot travel to the heavenly bodies in order to study their chemical composition, but we do have occasional visitors to our planet from celestial spaces. These are the meteorites, the falling of which from the sky has excited both fear and wonder in the breast of man and the life history of which scientists so much desire to know. Some consider that meteorites are terrestrial in their origin and have been projected into space from active volcanoes in long past ages of the Earth's history. But the opposite opinion is most widely held, namely that they are genuine samples of celestial matter. An inspection of the fine collection of meteorites in the Natural History Museum at South Kensington will show that many of them consist to a large extent of iron, or rather of an alloy of iron with a small percentage of nickel. Other meteorites contain but little iron and are more like stones in their composition. Altogether there is considerable similarity in the composition of meteorites which have fallen at different times and in different places, and this uniformity has suggested to some scientists that most meteorites, if not all, have come from a common source and are possibly chips of one heavenly body. No single element has been found in a meteorite which is not obtainable also from terrestrial sources. The ones which occur most commonly in these other world chips, in addition to iron and nickel, are aluminum, calcium, carbon, magnesium, oxygen, phosphorus, silicon and sulfur, all in a state of combination. Some indeed of the compound minerals occurring in the meteorites are new and it is curious that quartz, the most common of terrestrial minerals, should not be found in meteoric stones. The study of the composition of these celestial visitors is of the greatest interest, but it is not from them that our trustworthy information about the Constitution of the Sun and Stars is derived. This information is obtained in a much more wonderful fashion. It is based not on any laboratory examination of celestial specimens, but on a study of the light which comes to us from the heavenly bodies, in other words, on the use of spectrum analysis. When white light, such as is obtained from the upper part of a candle flame, is passed through a slit at the end of a telescope and then through a glass prism, it is seen as a strip which is red at one end, violet at the other, and between these two extremes passes continuously through the various shades of orange, yellow, green and blue. This strip of graded color is known as a continuous spectrum and it results from the splitting up of the white light into its various components which is affected by the prism. The apparatus consisting of all the necessary parts for the production and observation of a spectrum is known as a spectroscope and this is the instrument which has yielded such marvelous results in the study of the sun and stars. If we were to examine with a spectroscope the light given out by a red hot poker, we should see only the red end of the spectrum. If the poker were put in the fire again and its temperature were raised, the spectrum observed would show some orange and yellow as well as red. While if we brought the poker to a white heat and examined it in this condition with a spectroscope, we should see a spectrum perfectly continuous from the red to the violet end. Molten iron also would exhibit a continuous spectrum and one can say generally that the spectrum of light emitted by any incandescent solid or liquid is continuous. It is quite easy however to get an incomplete spectrum, one which consists only of isolated lines or bands of different colors. In order to do that, we have merely to examine with the spectroscope the light which is emitted by an incandescent vapor. One of the simplest spectra of this kind is obtained by introducing a little common salt, sodium chloride say, on the previously charred and moistened end of a match into the non-luminous flame of a spirit lamp or a Bunsen burner. To the naked eye, the flame will assume an intense yellow color and if the spectroscope is directed towards it, the spectrum is seen to consist of a single yellow line as shown in figure 11. The flame in fact is giving out only one particular kind of light. With a first class instrument, this line turns out to be two distinct lines very close together, but this division is not apparent with an ordinary spectroscope and need not concern us here. Suppose now we were to introduce into the Bunsen burner flame some other salt of sodium, washing soda for example, we should get exactly the same spectrum. This is a fact of the greatest significance indicating that whatever be the form in which sodium is introduced into the non-luminous flame, its presence is invariably marked by the yellow line at a definite position in the spectrum. From this simple case, the reader will easily appreciate the power of detection with which the spectroscope equips the chemist. For if the question arises whether a given substance contains sodium or not, he has but to introduce some of it into a Bunsen burner flame and see whether that incriminating yellow line appears in the spectrum. It has actually been found that as little as one 10 millionth of a grain of a sodium salt can easily be detected in this way. Other incomplete spectra, generally more complex than that of sodium, are observed by introducing salts of various metals into the non-luminous flame of a Bunsen burner, say figure 11. Barium salts for example, impart a green color to the flame, and their spectrum is characterized by a number of green lines. Strontium salts on the other hand, tinge the Bunsen burner flame a brilliant crimson, and their spectrum contains a series of lines and bands, mostly at the red end. It is possible that every reader, perhaps without knowing it, has seen the colors which barium and strontium salts impart to a flame. For the green and red lights which figure so largely in firework displays, are produced by adding these salts to combustible mixtures containing sulfur. For the detection of sodium, barium, or strontium, the reader might think it's sufficient to observe the color which a substance under examination imparts to the Bunsen flame. So it would be provided only one of the metals were present. This condition however will not always hold good, and when two or more are present, the color of the flame will give no certain indication. But it is just here that the full value of the spectroscope becomes apparent, for each constituent in a mixture contributes to the spectrum its own quota of lines, uninfluenced by the others which are present. This marvelously sensitive spectroscopic method of analysis can be applied not only to metallic salts which are volatile in the Bunsen flame, but also to substances like hydrogen, which are gases at the ordinary temperature, and to refractory metals such as iron. Ingenious devices have been adopted for bringing these into the state of incandescent vapor, from which alone we may expect to obtain a characteristic discontinuous spectrum. Hydrogen for example is filled into a glass tube at low pressure, and an electric discharge is passed through the rarefied gas. The spectrum of the glowing hydrogen is then found to be characterized by three main lines, red, green, and blue respectively, C figure 11. To obtain the spectrum of iron on the other hand, the metal or one of its compounds is placed between the poles of an electric arc. At the high temperature of this discharge, the iron is partly converted into incandescent vapor, and its spectrum containing an enormous number of lines is visible. When once the characteristic spectra of the elements, obtained by one or other of the methods just described, have been properly mapped out, then each line which we observe in any new spectrum may be referred to the element which is responsible for it. All this is more or less by way of introduction, and we come now to the celestial problem. If a telescope is directed towards a star, a nebula, or a comet, and the light proceeding from this heavenly body is examined spectroscopically, we find in a certain number of cases at least a spectrum consisting of definite lines or bands, and on comparing these with the spectrum already mapped, we can with confidence affirm that such and such elements are present in the far-off heavenly body. A bold step this, right out to the confines of space, and yet one which is fully warranted by the scientific evidence. In the spectrum of a nebula, there are bright lines which are identical with the characteristic hydrogen lines so that the latter element must be one of the constituents of a nebula. The spectrum of a comet is closely similar to that of the element carbon, as obtained by examination of the blue base of a candle flame. Cometary matter, therefore, contains carbon. Curiously enough, as a comet approaches the sun, its spectrum alters in character, and evidence is obtained that sodium and iron also enter into its composition. Surprise is in store for us when we come to examine spectroscopically the light which comes from the sun and the great majority of the stars. Instead of getting isolated colored lines or bands on a dark background, we observe a complete reversal of this, namely dark lines on a colored background. See figure 11. The explanation of this puzzling phenomenon is best understood perhaps by reference to an actual experiment. If we were to direct a spectroscope towards an electric arc light in which there is incandescent solid carbon, we should observe a continuous spectrum. Suppose now that between the spectroscope and the arc light, we interpose a Bunsen flame colored yellow by incandescent sodium vapor. The effect on the spectrum is rather surprising. The continuity of the spectrum is seen to be broken by a dark line, which occupies the exact position of the bright line in the ordinary sodium spectrum, as might easily be shown by momentarily screening off the arc light behind. The arc light is at a much higher temperature than the Bunsen flame, and what has happened is that the sodium vapor in the ladder has absorbed or picked out of the light from the hotter source exactly those rays which it itself usually emits. The light which passes on is therefore bereft of those particular rays and the spectrum shows the deficiency. The reader must remember that light is a species of vibration and that just as the string of a musical instrument will respond alone to that particular note out of many which has its own pitch, so an incandescent vapor will absorb exactly those rays which it emits. Provided therefore that the source of white light behind is hot enough, the passage of the light through various incandescent vapors at a lower temperature will be revealed by a number of dark lines on the spectrum exactly at those positions which bright lines from the vapors themselves would occupy. The solar spectrum then, consisting as it does of a very large number of dark lines on a colored background, tells us that the center of the sun is at a white heat and that this incandescent core is surrounded by an atmosphere of incandescent vapor at a somewhat lower temperature. By comparing the positions of the dark lines in the solar spectrum with the bright lines in spectra which have already been mapped, we learn what is the composition of the sun's atmosphere, the chromosphere as it is called. Among the elements which are thus proved to be present in the sun are hydrogen, sodium, calcium, barium, magnesium, iron, zinc, and copper. The reader will see that so far as the mere elements go, there is nothing very strikingly novel about the composition of the sun, but there is probably a considerable difference between the earth and the sun in the extent to which the elements are combined. On the earth, the elements just named are almost without exception found in the form of compounds, but at the high temperature of the sun, all the ordinary compounds will have undergone dissociation into their constituent elements. It appears pretty certain that the temperature of the sun is not below 10,000 degrees Fahrenheit. In comparison with which the electric furnace, our best attempt at producing a high temperature is miserably cold. Similar conclusions as to the composition of the stars have been drawn from their spectra, and it appears that the elements entering into their composition are pretty much those with which we are familiar. Astronomers have actually ventured a step farther and endeavored to estimate the approximate temperature of each star from the character of its spectrum. This attempt is based on the observation that the spectrum of an element varies somewhat according to the way in which it is vaporized. According as the substance is exposed to the action of a flame, of the electric arc or of the electric spark, a different spectrum is produced. The conclusions which astronomers have drawn from these observations are deeply interesting, but that is another story. A very curious fact in connection with the application of spectrum analysis has been the discovery of an element in the sun before it was known on the earth. In 1868, attention was drawn to a conspicuous bright line in the spectrum of the sun's atmosphere, which did not correspond to a line of any element which was then known. Lockyer and Franklin did not hesitate to assert that there must be a new element in the sun and immediately proceeded to its christening, they called it helium, from the Greek for the sun. This is an excellent illustration of the confidence which scientists have in the trustworthiness of the spectroscopic method, a confidence which in this particular case was justified after the lapse of nearly 30 years. In 1895, Sir William Ramsey, working with the rare mineral cleavite, discovered a gas the spectrum of which contains a line coincident with the mysterious bright line already mentioned. This gas is in fact helium, and although it is an element of comparative rarity on our globe, it appears to play an important part in the constitution of the sun and stars. Examples of the wonderful detective power of the spectroscope might be multiplied. One might quote, for instance, the discovery of two new alkali metals, rubidium and cesium, by Bunsen and Kirchhoff, some 50 years ago. These workers, whose names are so closely associated with the marvelous development of spectrum analysis, detected some new lines in the spectrum of a liquid obtained by concentrating a certain German mineral water. They boldly concluded that there was in this water some previously undiscovered element, and they forthwith proceeded to search for it. And this element, cesium, took some finding. 40 tons of the mineral water had to be evaporated and operated on before as much as one quarter of an ounce of cesium chloride could be collected. A splendid tribute this, not only to the sensitiveness of the spectroscopic method, but also to the confidence and patience of the searchers. Among the many problems raised by the spectroscope is that concerning the peculiar light of the aurora borealis, its spectrum is characterized, specially, by a bright yellowish-green line which has given rise to much discussion, and which, for long, could not be referred to any known element. Again, the discoveries of chemists have supplied the clue, and it appears that Krypton, one of the recently detected gases of the atmosphere, is the responsible party. End of Chapter 18. Chapter 19 of the Romance of Modern Chemistry. This LibriVox recording is in the public domain. The Romance of Modern Chemistry by James C. Phillip. Chapter 19. Chemistry and Agriculture. It is becoming more and more obvious, as time goes on, that there is scarcely any department of nature's activity, scarcely any useful art practiced by man, in which the laws and principles of chemistry are not involved. In the delicate processes which go on in our bodies, in the roaring of the blast furnace, in the silent growth of the tiniest blade of grass, chemical forces are at work, merely on atoms and molecules, and yet producing changes which in their sum total can be described only as mighty and marvelous. The activity of these forces has often remained unsuspected for long ages, and man's skill in many useful arts has been acquired, not from any scientific knowledge of the underlying principles, but by long experience and practice. Agriculture is a case in point. Since Adam delved, the art of tilling the soil has been a common occupation, and a vast store of practical knowledge of agriculture has been gradually accumulated. In these days of competition, however, rule-of-thumb methods, handed down from father to son, are not sufficient to command success, and the aid of the chemist has to be invoked. We can well imagine how ancient and hoary agriculture might resent the intrusion into its domain of the modern upstart chemistry, as if it had any right to teach agriculture why and how it ought to do this and that. This struggle between the practice of the past and the knowledge of the new age is always occurring, and we are slow to learn the lesson that the science of the laboratory cannot in the long run be kept out of the field, the factory, the workshop, or even the kitchen. Suppose then we consider for a little what chemistry has to teach us about the growth and culture of the vegetable world, about the yearly marvel of wood and field and garden, for it is a marvel. Look at the fields in the time of sowing. They are brown and bare and dead. Look at them five months later. They are clothed with an abundant garment of living green or gold. In the interval, no influence, but that of soil and sun, of wind and rain, has played upon the seed in the growing plant. Whence, then, all this wealth of fresh material? Is it a new creation, or is it an equally marvelous transformation? If the latter, what are the substances which have been changed as biomegicians wand into stem and leaf and flower? Surprising as it may seem, it is only 300 years ago since a chemist of repute endeavored to show that vegetable substances were produced from water alone. The experiment by which he sought to prove this was a very simple one and is worth rehearsal. The story shows how easy it is for any traveller into the unknown to miss the right path. This chemist took a willow weighing five pounds and planted it in a quantity of dried earth, which weighed 200 pounds. For five years he did nothing to the willow except water it occasionally. At the end of that time, it was pulled up and found to weigh 169 pounds, two ounces. The earth in which the willow had grown was dried as before and was found to be only two ounces lighter than at the start. Our chemist therefore drew the conclusion that 164 pounds of wood, bark, roots and leaves had been produced from water alone. The experiment was straightforward enough and at that time before the composition and influence of the air were discovered, it was quite convincing. Even now when we know for a fact that the solid part of a plant is largely derived from the carbon dioxide in the atmosphere, the average person may find it difficult to realize that a gas which is present in the air only to the extent of three parts in 10,000 should really be responsible for all this. But so it is. Under the influence of light, the green coloring matter in the leaves, the chlorophyll as it is called, has the power of dealing with the carbon dioxide which is taken in from the atmosphere, liberating the oxygen and converting the carbon into various compounds which form the substance of the plant. It is difficult to appreciate the prodigious quantity of carbon dioxide which is consumed by the vegetable world in this way. But some idea of it may be gained from the fact that an acre of a good wheat crop obtains from the atmosphere in the course of four months as much as one ton of carbon. Plants, then, are the great purifiers of the air. If it were not for their activity in removing the carbon dioxide, there would very shortly be an unhealthy excess of this gas in the atmosphere. From the lungs of men and animals and from our house and factory chimneys, huge quantities of carbon dioxide are being constantly poured out. And with all this, the vegetable world must deal. It has been estimated that an acre of forest will just about balance 15 men. And although the 15 men are very seldom in the same place as the acre of forest, the winds of heaven secure a wonderfully rapid and even distribution of the carbon dioxide. To be strictly accurate, we must also bear in mind the fact that plants resemble animals in contributing to the contamination of the air. They, too, use up oxygen and breathe out carbon dioxide. In daylight, however, this process of plant breathing is quite outbalanced by the reverse operation, a characteristic of plants alone, whereby they give out oxygen and purify the air. It is only when they are kept in the dark that the action of plants in giving out carbon dioxide becomes noticeable. Taken all together, their services in purifying the atmosphere quite outbalance what they contribute to its contamination. The fact that a plant is really able, under the stimulating influence of light, to liberate oxygen from carbon dioxide may be demonstrated by a very simple experiment. A bit of a growing plant, a sprig of mint, for instance, is put in a glass tube, which is then filled with tap water and inverted in a dish, also containing tap water. The latter is employed in preference to distilled water in this experiment because it is charged to some extent with carbon dioxide. This simple piece of apparatus is then exposed to sunlight for several hours. It will be noticed that gas bubbles are formed on the surface of the leaves and that these frequently ascend and collect at the top of the tube. After a few hours have passed, the gas which has collected in the tube may be examined. To do this, the thumb is put on the end of the tube while it is still under water. The tube may then be taken out and inverted, the gas in this way being brought to the mouth of the tube. If a glowing slip of wood is thrust into the gas while the thumb is removed for a moment, it will be relit, showing that the gas which was collected was oxygen. So much then is fairly established that the carbon dioxide of the atmosphere is taken in by the plant and that the carbon is retained while the oxygen is given off. But chemists have not been able to discover the actual chemical process to which the carbon dioxide is subjected in the mysterious laboratories of the plant leaf. It is indeed certain that water also is involved, so that the leaves may be set to feed mainly on carbon dioxide and water, a simple life diet which produces the most extraordinary results. When we speak of the carbon of the carbon dioxide being retained in the plant, we must not suppose that it is actually found in that form. It is no sooner extracted from the carbon dioxide than it passes into some form of combination with hydrogen and oxygen, probably formaldehyde in the first instance. As to the methods by which the living plant subsequently builds up more complicated products, such as starch, sugar and cellulose, we know very little. The experiment was convinced the scientists of 300 years ago that vegetable matter could be produced from water alone has been shown to be incomplete and inconclusive. But we must admit at the same time that water does enter very largely indeed into the composition of living plants. Some succulent vegetables contain over 90% of their weight of water, and even trees felled in the driest period of the year will have as much as 40%. If we suppose the plant's supply of water completely removed, the remainder which we may call the dry material of the plant consists partly of combustible and partly of incombustible matter. In the combustible part, there are to be found five chemical elements, namely carbon, hydrogen, oxygen, nitrogen and sulfur, the first three of which are present in by far the largest proportion. These three combined in a variety of ways constitute the woody matter or cellulose, the sugar, the starch and the fats of the plant. Other ingredients of the combustible part of the plant are one, the nitrogenous bodies, which contain nitrogen as well as carbon, hydrogen and oxygen, and two, the albuminoids, in which sulfur is found as well as the other four elements. In the incombustible ash of the plant, there are also five elements found, namely potassium, magnesium, calcium, iron and phosphorus. These are essential to the life of the plant and exist in its tissues largely as carbonates, sulfates and phosphates. These constituents as well as some others which often occur but are not essential are derived from the soil in which the plant grows so that the nature and composition of the soil are all important factors in the vitality of the plant. In order that the reader may get an idea of the relative proportions of the water, the ash and the combustible matter in such a common vegetable product as metalgrass, the following figures are quoted. The crop to which the figures refer weighed five tons when freshly cut and produced one and a half tons of hay. Out of the five tons, that is 11,200 pounds of metalgrass, 8,378 pounds were water while the combustible matter weighed 2,613 pounds and the ash 209 pounds. We have discussed the marvelous way in which the living plant procures its carbon but the origin of some of the other constituents is also full of interest. Let us consider first the nitrogen which although it is present in vegetable tissue only to a small extent is an important and indeed essential constituent. The reader might suppose that the natural source of nitrogenous food for the plant would be the atmosphere with its vast stock of nitrogen. It is conceivable that the leaves might take in and assimilate the nitrogen of the air just as they deal with the carbon dioxide which is so much more scarce. There are some who have supposed that this really takes place but the bulk of the evidence shows that the leaves of plants generally are unable to digest nitrogen when it is presented to them in the form of the element itself. Atmospheric nitrogen however, does ultimately reach the tissues of some plants but by a very indirect road via the soil and the roots. Laguminous plants such as peas and veggies are provided with exceptional apparatus for assimilating nitrogen in the shape of swellings or nodules on their roots. These nodules contain microorganisms which have the power of taking in atmospheric nitrogen and so manipulating it as to render it suitable for use as food by the plant. The majority of plants however are destitute of these parasitic attendants and are unable to utilize atmospheric nitrogen directly. They appear to find this element most digestible when it is presented to them in the form of a salt such as a nitrate. Nitrates are readily taken up from the soil by plants and the nitrogen is subsequently transformed into the complex nitrogenous constituents of the plant tissues by various chemical processes which at present are not within our knowledge. Far less our power of imitation. In comparison with the practical chemistry which goes on in the cells of plants, the methods of the chemist are elementary and crude and he may well feel humble in view of the complex and delicate processes which are carried out in the wonderful little laboratories of the plant. All very well the reader will say the plant may take in the bulk of its nitrogen in the form of nitrates from the soil but how did the nitrates come to be there at all? To understand this it is necessary to remember that the atmosphere contains small quantities of nitrogen in the combined form, namely as ammonia, a compound of nitrogen and hydrogen and as nitric acid which as already stated is a compound of nitrogen, hydrogen and oxygen. The ammonia in the atmosphere has been given off from decaying organic matter and the nitric acid is due to the power of an electric discharge such as lightning is to induce the nitrogen and oxygen of the air to combine to some small extent. Now these two nitrogenous substances, ammonia and nitric acid, the one in alkali and the other in acid, dissolve readily in water and are either absorbed by the soil direct or are washed down into it by the rain. Quite a large amount of combined nitrogen gets into the soil in this fashion in addition to what is already there as the remains of earlier vegetation. Experiments carried out a rathamstead have shown that the total quantity of nitrogen carried to the soil by rain in one year is between four and five pounds per acre. When ammonia compounds get into the soil, their latter end is near for there they are tackled by microorganisms whose object in life is to convert all other nitrogenous bodies into nitrates. Since from the point of view of the plant, a nitrate is a much more digestible form of nitrogen than any ammonia compound, these nitrifying bacteria are valuable agents in the nourishment of the plant. Apart from the carbon, nitrogen, hydrogen and oxygen, the origin of which we have discussed, the elements which are essential to the building up of the plant are derived from the soil itself. Compounds containing potassium, magnesium, calcium, iron, sulfur and phosphorus are found in the rocks of the Earth's crust and it is through the breaking down of these rocks that the various ingredients of soils have been produced, except indeed the humus, which has quite a different origin. The humus is that part of the soil which represents the decayed vegetation of an earlier age. It is organic in origin and contains the ruins and remains of the nitrogenous compounds which were built up in that vegetation. Hence arises the fertility of virgin soil from which no crops have ever been taken. It is rich in nitrogenous humus and is practically a storehouse of food for the first crop which the new settler grows upon it. When the crops which grow on a given piece of ground are removed year after year, the soil must obviously become impoverished in the chemical materials on which the crops have fed. There need be no anxiety about the supply of carbon. The source of this element is in the atmosphere and fresh quantities of carbon dioxide are always being produced. Nor is there likely to be any shortage of hydrogen and oxygen. They come from water and we are not often seriously troubled in this country at least with a deficiency of that commodity. It is really in regard to nitrogen, phosphorus, lime and potassium that the soil becomes most rapidly impoverished and if the crops are to be kept up in quality and quantity we must replenish the store of these elements that is, manuring becomes essential. The necessity for this was of course recognized long before the agricultural chemist came on the scene. But since his appearance the materials which must be added to the soil have been definitely ascertained and their effects on various crops have been studied. The waste products of the animal body contain much of the material which is required for the enrichment of the soil and hence farm yard dung is an excellent general manure. Guano, the dried excrement of seabirds also contains nitrogen, phosphate and potash and so has been largely employed for the same purpose. Occasionally for special crops and in special circumstances it becomes necessary to supply to the soil a particular plant food, nitrogen for instance. In this case one may use as manure either sulfate of ammonia from the gasworks or nitrate of soda from chili. The nitrogen from ammonium sulfate is not so rapidly available for the use of the plant as the nitrogen from the chili salt peter in as much as the ammonia in the former has first to be interviewed by the nitrifying bacteria and converted into nitrate. The approaching exhaustion of the chili salt peter beds has stimulated chemists to discover ways and means of utilizing the nitrogen in the atmosphere for plant feeding purposes and the reader may remember the reference made in a previous chapter to the work already done in this direction. At the high temperature of the electric arc the nitrogen and oxygen of the atmosphere combine to a small extent and the compound so formed is easily converted into nitric acid. As already indicated the small amount of nitric acid occurring in the atmosphere is to be traced to the influence of electric discharges so that the method now in vogue for the manufacture of nitric acid from the atmosphere depends really on the production of artificial lightning. Mention should also be made of another modern electrical method of capturing the nitrogen of the atmosphere for agricultural purposes. This method results in the production of a compound known as calcium cyanamide which readily yields up its nitrogen for the use of crops. For certain soils and particular crops it is not necessary to manure with nitrate so much as phosphate. One of the simplest ways of supplying this constituent is to grind bones and scatter the bone dust in the soil. Phosphate of lime of which there is a considerable proportion in bones is an insoluble substance and as the plant prefers to have its food in dissolved form the effect of bone dust is not obvious at once. Such a phosphate manure however may be made more readily available by treating the bones or other substances containing phosphate of lime with sulfuric acid. This brings some at least of the phosphate into a soluble condition and the product super phosphate as it is called is extensively employed as an ingredient of artificial manures. It does not sound very probable that any product connected with the blast furnace could assist the growth of plants but here again it is the unexpected that happens. The slag produced in presence of lime when molten pig iron containing phosphorus is subjected to a blast of air and so purified is relatively rich in that element. It is accordingly used to a large extent as a phosphate manure for which purpose it must be very finely ground. By such artificial additions to the soil as the foregoing we are able to stimulate the growth of the plant but we must not run away with the idea that we are masters of the situation. Although the processes which go on during the growth of a plant seem to be purely chemical changes the fact stares us in the face that we cannot turn out a plant in our laboratories. The thing is absolutely beyond us. The attempt so far may to imitate the processes of plant growth can scarcely be called successful and our failure seems to be most complete in connection with the most important and wonderful process of all namely the assimilation of carbon dioxide. The chemistry of plant life and growth is, in fact, one of those mysterious chambers which have as yet been only partially explored. In spite of failure, however, to imitate the actual processes of growth chemists have been wonderfully successful in producing by artificial methods the substances which are found in plants. The story of the interesting advance that has been made in this direction will be told in another chapter. End of chapter 19.