英文科學(xué)讀本 第六冊(cè)·Lesson 42 Phosphorus

Lesson 42 Phosphorus

In commencing the next lesson Mr. Wilson showed the class the bottle labelled phosphorus. "We have frequently," he said, "had occasion to employ the contents of this bottle during our experiments. We will now try and learn something about the substance itself."

He took one of the sticks of phosphorus from the bottle, and allowed the class to see that it is a pale yellow, almost transparent, solid substance. He cut it with a knife as easily as he could have cut a piece of wax.

As he cut it, and while they were examining it, he called attention to the fumes of greenish-white vapor surrounding it like a little cloud. As soon as it was taken out of the water small wreaths of this smoky vapor began to rise from it. This, he explained, was due to the rapid oxidizing of the substance, which was taking place even at the ordinary temperature of the room. A very slight increase of warmth—even the warmth of the hand—would be sufficient to cause it to burst into actual flame. He showed them that for this reason he wore a glove, as it is not safe to hold phosphorus in the naked hand. Even while he was speaking he let fall a little piece on a warm plate lying on the table, and in an instant it took fire.

This great inflammability of phosphorus, he said, "is its chief characteristic, and shows us clearly enough why it is necessary to always keep the substance in water.

Now watch this little experiment. I will put a small piece of dry phosphorus on the table together with a little powdered potassium chlorate, and strike them with a hammer. It burst into flame and exploded with a loud report.

It is this last property of phosphorus, he continued, as soon as the boys had recovered from the shock, "which makes it of such great use in the manufacture of lucifer matches.

In one kind of lucifer the paste containing the phosphorus is put on the end of the match itself. These matches will ignite when rubbed on any rough surface, and are, of course, dangerous. Many accidents have happened through carelessness in using them. In the 'safety' match the phosphorus paste is put on the box and not on the match. There is no danger of the match igniting until it is rubbed on this phosphorus rubber. Phosphorus is never met with in the free or uncombined state. It is chiefly found in combination with other elements, forming phosphates. Phosphate of lime is one of the most important of these; it is one of the materials of certain volcanic rocks. These are the rocks which by crumbling down produce our fertile soils. The plants which grow in such soils absorb the phosphorus compounds from them to help in building up their own material. These plants become the food of man and animals.

All the bony structures in the animal consist largely of phosphate of lime, which has been obtained in this way. The bones owe their hardness and rigidity to the phosphate of lime which they contain. You, no doubt, remember that we have several times, in our lessons in physiology, dissolved this phosphate of lime out from bone by soaking it in dilute hydrochloric acid. In those experiments our business was rather with the animal matter of the bone than with the mineral matter dissolved out from it. I have now, however, a bone in soak, and we will examine the liquid itself. It is quite clear; we can see nothing in it. Now, watch while I pour some ammonia into it, and you will see that the phosphate of lime will be precipitated to the bottom from the clear solution.

The chemist always prepares phosphorus for his purpose from bones. The bones are burned to a white ash in a clear, open fire. This burning, as you know gets rid of all the animal matter of the bone; the white ash left behind is the earthy or mineral matter and is known as bone-earth or bone-ash. It is this bone-ash which yields the phosphorus. Just one little experiment before we close.

I will cut off a small piece of phosphorus, dry it carefully in blotting-paper, place it on this metal plate, light it, and cover it with the bell-jar. Violent combustion at once begins, and the oxygen in the jar is rapidly consumed. In order that the whole of the phosphorus may be consumed, I will raise the jar a little from time to time to admit more air. All this time the jar is filled with dense white fumes, and at last the combustion ceases—the piece of phosphorus is all consumed.

Now let us watch those fumes. We shall see them gradually condense and fall on the plate, as a soft, white, snow-like powder. This powder is an oxide of phosphorus. Now that the fumes have all gone, we will remove the bell-jar, and pour a little water into the plate. The moment the water is put into the plate a violent hissing takes place.

This oxide of phosphorus has a powerful affinity for water. The energy with which the water sucked up the white powder caused that hissing. I will now add a little of the blue litmus solution, and you will see that the effect is the same as we have seen before in the oxides of carbon and sulphur. The blue litmus turns red, and if we dip the finger into it, we shall find again that this dissolved oxide has a sour taste."