The Science Notebook
  Lionel Chem-Lab - Chapter 3

  Home   Terms of Use   Safety  Contact Us   Experiment Pages   Downloads   Supplies   Useful Links!

NOTE:  This book was published in 1942 as a manual to accompany several Lionel Chemistry sets of the time.  While some of the experiments and activities here may be safely done as written, a number of them use chemicals and methods no longer considered safe.  In addition, much of the information contained in this book about chemistry and other subjects is outdated and some of it is inaccurate.  Therefore, this book is probably best appreciated for its historical value rather than as a source for current information and good experiments.  If you try anything here, please understand that you do so at your own risk.  See our Terms of Use.
Pages 40 - 57



Seldom do we realize that the entire earth is covered with an ocean of gases called the atmosphere, which is much deeper than the water that covers a large portion of the earth’s surface. It has been estimated that the ocean of gases comprising the atmosphere extends over one hundred miles in height compared to the greatest ocean depth of six or seven miles.

Man’s physical structure is best adapted to work at the bottom of this atmospheric ocean where the pressure is about fifteen pounds to the square inch. Mountain climbers and aviators who ascend to altitudes of more than four or five miles suffer because of the lack of oxygen in the thin air and from the lower pressure and frequently must carry oxygen tanks. Likewise deep sea divers require oxygen to breathe when they descend to a deep level under water.


The air of the atmosphere, is composed of numerous gases of which the principal ones are nitrogen and oxygen, plus certain amounts of water and dust. We have already seen that carbon dioxide is also present in the atmosphere but the proportion of this gas to the other gases is very small indeed. The remaining gases are known as rare gases and in recent years have become of considerable commercial importance. These gases include neon, krypton, argon and xenon.


Oxygen is the most abundant of all the elements and was discovered in England in the Eighteenth Century by the chemist, Joseph Priestley. One day he put a small amount of mercuric oxide in a tube filled with mercury, and then turned it upside down in a vessel of mercury. When the oxide rose to the top of the inverted tube, Priestley heated it with the rays of the sun through a burning glass. He soon found that a gas was given off in which a candle would burn with a remarkably brilliant flame. This "air” discovered by Priestley was the gas we now call oxygen and its discovery is considered one of the landmarks in the history of chemistry.




Free oxygen is found in the atmosphere. It is also found in many compounds comprising, for example, about 90% of the weight of water. It makes up nearly one half of the rocks that form the earth’s crust as well as over half of the plants and animals including ourselves. Animal life is entirely dependent on oxygen. When we inhale, we take oxygen into our lungs where it passes into the blood stream. As the blood goes through our circulatory system, oxygen is liberated. After it has been utilized by the tissues to provide us with energy and warmth, carbon dioxide is produced which is returned to the lungs by the blood and finally exhaled. The important part that carbon dioxide plays in this process is described more fully in our chapter on Carbon.

EXPERIMENT No. 54 Release Of Oxygen by Leaves

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Bottle, green leaves, pan.

PROCEDURE: Place some green leaves in a bottle. Fill the bottle to the brim with water. Pour two inches of water into a pan. Place a dish over the mouth of the bottle and invert it in the pan of water.  Remove the dish only after the mouth of the bottle is beneath the water level of the pan. Expose bottle and pan to the sunlight and allow to remain there for several hours.

SUMMARY: This experiment indicates how cooperative nature is. We human beings, together with the animals, breathe in oxygen to keep alive and breathe out carbon dioxide gas as waste material. Conversely, plant leaves absorb this gas, converting it to starch in the presence of sunlight and releasing oxygen as waste. Actually what has happened in this experiment is that the carbon dioxide in the air is absorbed by the leaves leaving only oxygen and nitrogen. This process tends to cause the water level to rise.

EXPERIMENT No. 55 Absorption Of Oxygen From The Air

(CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Powdered iron, wood chip, drinking glass, saucer.

PROCEDURE: Fill the saucer with water and then place the wood chip in it. Place four measures of powdered iron on the wood chip.  Sprinkle a little water over the filings. Invert the glass over the wood chip so that the air is enclosed within. Set aside for a few days.

SUMMARY: Note that after several days the water has risen slightly inside the glass. This is due to the decrease in the amount of oxygen which has united with the iron to form iron oxide.

EXPERIMENT No. 56 Oxidation Of A Candle

(CL-11, CL-22, CL-33, CL-44. CL-55, CL-66, CL-77)


candle under a jar in a bowl of water


APPARATUS: Pan, candle, mason jar.

PROCEDURE: Place candle firmly on a bit of melted wax in the center of the pan and add two inches of water. Light the candle and place the inverted glass jar over it making certain that the flame is a few inches away from the top. Note that the flame goes out and the water levels change.

SUMMARY: The explanation of this interesting experiment is as follows: The candle burns so long as there is oxygen. When this vital gas is used up, the candle goes out. As the candle As the candle burns it forms carbon dioxide and steam. However, since the carbon dioxide gas dissolves in the water and the steam condenses into water, the water in the pan rises to fill the space previously occupied by oxygen.

EXPERIMENT No. 57 Preparation Of Oxygen From Potassium Nitrate

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Potassium nitrate, test tube, candle or alcohol lamp and a cotton string.

PROCEDURE: Put three measures of potassium nitrate in a dry test tube and heat until it begins to melt. Ignite the string and allow it to flame for a moment. Blow the flame out and immediately insert the glowing string into the test tube.

SUMMARY: Potassium nitrate contains oxygen, part of which is liberated upon heating. Since the air in the tube is richer in oxygen than the surrounding air, the glowing string bursts into flame.


When oxygen combines with another element the process of oxidation takes place. This process may be extremely slow or extremely rapid. For example, the rusting of iron is an illustration of slow oxidation, also the decay of organic matter. No light is given off and practically no heat. On the other hand, when a substance burns or undergoes combustion, oxidation is taking place very rapidly and the heat which is given off may cause the substance to glow like an ember or even to burst into flame.

EXPERIMENT No. 58 Slow Oxidation

(CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Powdered iron, tumbler.


PROCEDURE: Place three measures of powdered iron in a tumbler and add a few drops of water. Set aside for a few hours and note the rust formation.

SUMMARY: Sometimes oxidation takes place so slowly that no light is seen and unless careful measurements are made, no heat is noticed. This is an example of slow oxidation.

cutting metal

Linde Air Products Co.

Heavy pieces of metal are cut by the process of rapid oxidation. The metal is heated white hot by a blowpipe, then a jet of oxygen blown through a lance causes such rapid oxidation that the metal is cut.

EXPERIMENT No. 59 Oxidizing Hydrogen Peroxide

(CL-55, CL-66, CL-77)

APPARATUS: Fresh hydrogen peroxide (drug store), ferrous ammonium sulfate, test tube.

PROCEDURE: Dissolve two measures of ferrous ammonium sulfate in a test tube half full of water. Add three drops of hydrogen peroxide. Note the reaction and the color of the solution.

SUMMARY: Hydrogen peroxide is a stronger oxidizing agent than free oxygen, thus it oxidizes ferrous ammonium sulfate to ferric ammonium sulfate (rust-colored solution).

EXPERIMENT No. 60 Fire Writing

(CL-11, CL-22, CL-33. CL-44, CL-55, CL-66, CL-77)

APPARATUS: Potassium nitrate, candle or alcohol lamp, small soft brush and paper.


PROCEDURE: Dissolve half a spoonful of potassium nitrate in a test tube one-quarter filled with water. Use heat if necessary. Dip the brush in the solution and write on the paper heavily in one continuous line. Allow to dry. Light the end of the writing with a match to start a spark on the paper, blowing out the flame. Note how the spark moves along the line of the writing.

SUMMARY: The ease with which potassium nitrate burns shows that it is liberating oxygen rapidly as it burns along the written line.

EXPERIMENT No. 61 Making A Fuse

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Potassium nitrate, cotton string and test tube.

PROCEDURE: Dissolve ten measures of potassium nitrate in a test tube one quarter filled with water. Dip the string into this solution for several minutes. Allow to dry thoroughly by suspending it from a hook.

SUMMARY: Since potassium nitrate is easily combustible, it makes an excellent fuse. Ignite and observe rapid travel of the combustion along the string.

EXPERIMENT No. 62 Oxidation Of Paper

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Potassium nitrate, test tube, alcohol lamp or candle, paper and test tube.

PROCEDURE: Place four measures of potassium nitrate in a test tube. Heat until the crystals melt. Drop into the tube a few small shreds of paper while continuing the heating.

SUMMARY: When potassium nitrate is heated, it liberates oxygen which causes the paper to burn.


Occasionally we read in the papers of fires which started by spontaneous combustion. What does this mean? Sometimes among heaps of rags, wool and cotton, usually soaked with oil; or wet hay or straw, a fire develops due to the internal development of heat. Sometimes coal in the bunkers of vessels catches fire in this same way. The name given to this phenomenon is spontaneous combustion which is the result of slow oxidation of combustible material in a confined space where the small amount of heat liberated at first cannot be dissipated. As the temperature rises, the oxidation proceeds more rapidly until glowing, or even flaming, is produced.



The name given to the compound formed when an element combines with oxygen is an oxide. The burning of an element is really its union with oxygen. When iron is burned in the presence of oxygen, iron oxide is formed.  Likewise the burning of sulfur forms an oxide of sulfur. Chemists, however, use the general term oxidation to refer to the combining of many of the elements with one another even though no oxygen happens to be involved in the reactions.

EXPERIMENT N0. 63 Zinc Oxide

(CL-44, CL-55, CL-66, CL-77)

APPARATUS: Zinc metal, test tube, alcohol lamp.

PROCEDURE: Place a small piece of zinc metal in a test tube and heat by means of the alcohol lamp. Continue heating until the formation of a white substance on the zinc appears.

SUMMARY: The white coating is zinc oxide formed when the zinc undergoes oxidation.

EXPERIMENT No. 64 Oxidation Of Sulfur By Potassium Nitrate

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Potassium nitrate, heating spoon, sulfur.

PROCEDURE: Heat one measure of potassium nitrate on the spoon until it begins to melt. Now add one measure of sulfur to the spoon.  Note the pungent odor.

SUMMARY: Sulfur is quickly oxidized by potassium nitrate to form sulfur dioxide recognizable by its penetrating odor.

EXPERIMENT No. 65 Oxidizing Hydrochloric Acid


APPARATUS: Sodium chromate, hydrochloric acid, and test tube.

PROCEDURE: Place one measure of sodium chromate in a test tube and add ten drops of hydrochloric acid.

SUMMARY: Note the color of the solution. Sodium chromate oxidizes the hydrochloric acid liberating chlorine gas which may be identified by its sharp odor.

CAUTION: While this gas is dangerous to inhale in large amounts, the amount liberated in this case is quite negligible. However, it would be well to avoid smelling the gas directly.

EXPERIMENT No. 66 Copper Oxide

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Clean copper wire, pliers, glass, gas flame.


PROCEDURE: Heat the wire over a gas flame holding it with a pair of pliers. Continue this for a few minutes, noting the color of the flame, then immerse the wire in a glass of water. Note the black formation.

SUMMARY: This is another example of oxidation, the copper being oxidized to form black copper oxide.

EXPERIMENT No. 67 Stannous Oxide

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Tin metal, test tube, alcohol lamp.

PROCEDURE: Put a small piece of tin into a test tube and apply heat by means of the alcohol lamp. Continue heating until a white formation on the tin may be noted.

SUMMARY: The element tin is oxidized as it combines with the oxygen to form the white coating of oxide which is a valuable abrasive used for fine polishing.

EXPERIMENT No. 68 Iron Oxide

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Bright iron nail, test tube holder, alcohol lamp, glass.

PROCEDURE: Heat the iron nail by holding it with the test tube holder in the flame of the alcohol lamp. Continue this for a few minutes then immerse the nail in a glass of water. Remove it when cool and note any changes.

SUMMARY: The black coating is due to the formation of iron oxide. Thus, when iron is heated in the presence of oxygen, the iron is oxidized.


Nitrogen is one of the four essential constituents of air, there being about seventy-eight volumes of nitrogen in every one hundred volumes of dry air.

It is also found in a number of compounds, principally potassium nitrate (saltpeter), sodium nitrate (Chile saltpeter) and other mineral nitrates which develop from the life processes of animal and vegetable organisms. Large amounts of nitrogen exist in the proteins in meat, egg white, clover, peas, beans, etc.

Like most of the other gases, nitrogen is colorless, tasteless and odorless.  On a large scale it is manufactured from liquid air. The usual method of preparation in the laboratory is to make it from ordinary air by removing the oxygen and making an oxide with some substance that can then be easily separated from the nitrogen gas.

EXPERIMENT No. 69 Nitrogen Does Not Support Combustion

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL~77)


APPARATUS: Wooden splint, bottle, candle, pan.

PROCEDURE: Place a candle firmly upright in the center of a sauce pan. Fill with two inches of water. Light the candle and place the inverted bottle over it making sure that the flame is a few inches away from the top of the bottle. Note that the flame is extinguished and the water levels change. Remove the bottle and quickly insert a lighted splint into the mouth of the bottle.

SUMMARY: The lighted splint goes out proving that the nitrogen in the bottle will not support combustion.

EXPERIMENT No. 70 Extracting Nitrogen From The Air

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Candle, pan, bottle.

PROCEDURE: Repeat Experiment No. 69 except when the flame goes out and the water levels change, cover the mouth of the bottle so that the gas cannot escape once the bottle is placed in an upright position.

SUMMARY: The burning candle uses up oxygen, leaving essentially nitrogen in the bottle.


Nitrogen is vital to plant growth. A few plants such as beans, peas, clover and alfalfa make direct use of inactive nitrogen in the air with the aid of nitrogen-fixing bacteria which are attached to their roots. These minute organisms form many useful nitrogen compounds. Some of these are used by the plants directly and the others stay in the ground to enrich the soil.  Although nitrogen is taken from the air by plants, it is also returned to the air by the decay of plant and animal matter. This is known as the nitrogen cycle.

Nitrogen combines with hydrogen to form ammonia and the chief use of nitrogen in industry is in the preparation of this important substance - the first step in the manufacture of important  fertilizers and nitric acid. Nitrogen forms many compounds which are unstable and break up with a release of energy as, for example, tri-nitro toluene (TNT), ammonal, nitro-glycerine, dynamite, etc. Huge plants have been erected by the government for the purpose of fixing the nitrogen of the atmosphere by means of electricity.


Ammonia is composed of nitrogen and hydrogen and has the formula NH3.  It is a colorless gas the water solution of which is known as ammonium hydroxide, or ammonia water, commonly available at groceries and drug stores.


The usual laboratory method of preparing ammonia is by gently heating a mixture of ammonium chloride with either calcium hydroxide or sodium hydroxide. The interaction of these two compounds produces ammonium hydroxide, an unstable compound, which decomposes into ammonia and water.

EXPERIMENT No. 71 Preparing Ammonia Gas

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Ammonium chloride, calcium oxide, test tube and red litmus paper.

PROCEDURE: Place three measures of calcium oxide and an equal amount of ammonium chloride in a test tube. Note whether an odor is given off. Heat carefully. Remove test tube from flame. Carefully smell the gas present. Moisten some red litmus paper and hold it over the mouth of the test tube. Note the change in color.

SUMMARY: In the laboratory, ammonia can be made by heating calcium oxide with ammonium chloride to obtain ammonia, calcium chloride and water. Since ammonia in the presence of water forms a base, it turns red litmus paper blue.

EXPERIMENT No. 72 Extracting Ammonia From Proteins

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Red litmus paper, calcium oxide, test tube, protein matter (white of egg, cream, butter or gluten).

PROCEDURE: Place a little protein matter in a test tube together with two measures of calcium oxide. Add two or three drops of water and heat gently. Smell the test tube cautiously to see if ammonia gas is being liberated. Moisten some red litmus paper and place it over the mouth of the tube.

SUMMARY: Proteins are nitrogen compounds, many of which liberate ammonia gas when heated with calcium oxide.

EXPERIMENT No. 73 Proving That Ammonia Is Soluble In Water

(CL-33, CL-44, CL-55, CL-66, LL-77)

APPARATUS: Calcium oxide, ammonium chloride, pan, test tube.

PROCEDURE: Place two measures of ammonium chloride and three measures of calcium oxide in a test tube. Heat tube and contents carefully keeping the tube partly closed with the index finger.  After heating, place your finger tightly over the mouth of the tube.  Invert the test tube so that its mouth is beneath the water level of the pan, taking your finger away only after the mouth of the tube is entirely immersed in the water. Now note the rise of water in the test tube. 

SUMMARY: Ammonia gas is soluble and is dissolved by water thus creating a partial vacuum which causes the water to be sucked into the test tube.


EXPERIMENT N0. 74 How To Make Aluminum Hydroxide

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

dissolving ammonia in water


APPARATUS: Calcium oxide, ammonium chloride, stopper and delivery tube, two test tubes, holder, candle or alcohol lamp.

PROCEDURE: Place three measures of ammonium chloride and an equal amount of calcium oxide in a test tube. Attach delivery tube and heat slowly holding the test tube with the holder. As the reaction develops and gas begins to escape, extend the long stem of the delivery tube into another test tube one fourth filled with water. Continue heating for a few minutes and then remove test tube containing the water and the gas. (Remove delivery tube from receiving test tube before cooling the delivery test tube to avoid sucking liquid back into a hot test tube and consequently breaking the glass.) Dip a piece of red litmus paper into this solution and observe the color change.

SUMMARY: Heating calcium oxide and ammonium chloride liberates ammonia gas. Since ammonia gas is soluble in water, it combines with water to form the base, ammonium hydroxide, which turns red litmus paper blue.

EXPERIMENT No. 75 How To Make Copper Ammonia Blue

(CL-55, CL-66, CL-77)

APPARATUS: Copper sulfate, ammonium hydroxide (or household ammonia), test tube.


refrigerator cutaway

Frigidaire Div. G. M. C.

An illustration in "phantom" of a modern electric refrigerator showing the motor compressor unit, refrigeration coils, and ice unit.  The compressor converts the liquid refrigerant to a volatile gas which is circulated through the coils, lowering the temperature in the refrigerator.


PROCEDURE: Dissolve three measures of copper sulfate in a test tube one third filled with water. Add a few drops of ammonium hydroxide and note a bluish-white precipitate. Add more ammonium hydroxide and note the deep blue color which appears.

SUMMARY: The bluish-white precipitate is copper hydroxide. When more ammonia is added, the copper hydroxide dissolves to form the beautifully colored solution of copper ammonia blue.

EXPERIMENT No. 76 Characteristics Of Ammonium Salts

(CL-66, CL-77)

APPARATUS: Magnesium sulfate, ammonium hydroxide (or household ammonia), ammonium chloride, two test tubes.

PROCEDURE: Dissolve (by shaking well) two measures of magnesium sulfate in a test tube one quarter filled with water. Add a few drops of ammonium hydroxide. Note the cloudy appearance when the tube is again shaken. Prepare an ammonium chloride solution by dissolving two measures of ammonium chloride in a test tube one quarter filled with water. Pour the contents of this tube into the first tube. Note that the resulting solution is clear.

SUMMARY: In the presence of excess ammonium salts, the precipitate of magnesium hydroxide dissolves because the ammonium salts have decreased the amount of hydroxide.

EXPERIMENT No. 77 Vaporizing Ammonium Chloride

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Ammonium chloride, test tube, candle or alcohol lamp.

PROCEDURE: Place two measures of ammonium chloride in a dry test tube and apply heat slowly. Note how the vapor condenses near the mouth of the tube.

SUMMARY: Many ammonium salts, including ammonium chloride, are vaporized directly from a solid to a gas.


Because of its great industrial importance, large quantities of ammonia are required. It is manufactured largely by either the Haber process which depends upon the formation of ammonia by the direct union of nitrogen and hydrogen, or from soft coal by the process known as destructive distillation.

One of the most important uses of ammonia is as a refrigerant. When a gas such as ammonia is liquefied, heat is liberated, and likewise when the liquid passes again into the gaseous state, a great deal of heat is absorbed. This fact is utilized in the manufacture of ice and in maintaining a low temperature in cold storage plants.


Other uses for ammonia are: to make household ammonia, to make ammonium salts (fertilizers and explosives), to make nitric acid and to make sodium carbonate (washing soda) by the Solvay Process.


Ammonia was quite satisfactory for the commercial manufacture of ice and certain other industrial applications, but with the advent of air-conditioning and domestic refrigerators, a new refrigerant was required which was not poisonous, explosive, or inflammable. It must also have no odor when mixed with air, even in fairly high concentrations, so that panic would not result in crowded theaters or department stores should a leak occur in the air-conditioning system. Applied science, consequently, has synthesized a new family of refrigerants fulfilling every requirement. These new materials which are fluorinated-chlorinated hydrocarbons, are safe, and because of their safety are now widely used not only in domestic refrigerators but in the air-conditioning of theaters, office buildings, and a rapidly increasing number of homes.


Nitric acid is the most common acid of nitrogen and its chief use is in the making of explosives, such as, nitro-glycerine, dynamite, gun cotton and TNT. It also is used in the manufacture of fertilizers, artificial silk, drugs, dyes and celluloid and to make such important nitrates as silver nitrate employed in the manufacture of photographic film.

Nitric acid is very strong and reacts with bases to form salts called nitrates.

EXPERIMENT No. 78 Preparing Nitric Acid

(CL-11, CL-22, CL-33, CL-44, CL-55, CL-66, CL-77)

APPARATUS: Potassium nitrate, sodium bisulfate, test tube, blue litmus paper, alcohol lamp or candle.

PROCEDURE: Place four measures of potassium nitrate in a test tube.  Add four measures of sodium bisulfate and four or five drops of water. Put a strip of moistened blue litmus paper over the mouth of the test tube and apply heat. Remove from flame when fumes are beginning to be given off and smell cautiously. Note that the litmus paper turns red.

SUMMARY: Nitric acid is manufactured commercially by treating saltpeter with sulfuric acid. In the laboratory, however, it is made by the reaction of potassium nitrate with sodium bisulfate.


The name given to the process of utilizing free nitrogen in the air to make useful compounds such as nitric acid, nitrates and ammonia, is the fixation of


nitrogen. Since the supply of natural nitrates such as Chile saltpeter (sodium nitrate) is limited, and the demand for nitrogen and its compounds is large, it becomes necessary to draw on the vast resources of the air as a source.

Nearly half of the world’s requirements of nitrogen compounds are fulfilled by artificial fixation processes: (1) the electric arc process, (2) the synthetic ammonia process and (3) the cyanamide process.


Hydrogen gas, the lightest of all elements, was used for a good many years to inflate balloons and dirigibles but because of its inflammability is no longer used for this purpose. The most important use of hydrogen, today, is in the treatment of certain oils to produce fats for household purposes. This is known as the hydrogenation process. Huge quantities of hydrogen are also used in the manufacture of acids, ammonia and in metal refining.


Hydrogen is a gas without color, taste or odor and does not occur in such quantities as oxygen and nitrogen. It differs from oxygen in that only very small amounts of it occur in the free state. The compounds of hydrogen, on the other hand, are everywhere about us: in water, in foods and in our bodies. It combines with carbon in all proportions to make hundreds of compounds including petroleum and natural gases. These are known as hydrocarbons.

We have already described the process of oxidation. Hydrogen is a very important reducing agent because of its tendency to combine with oxygen and thus remove oxygen from the oxides. The process of reduction is more fully explained on a following page.

Because hydrogen is a component of all acids, it is sometimes known as an "acid-forming element".

The usual method of preparing this gas in the laboratory is by using a metal such as zinc, and sulfuric or hydrochloric acid.

EXPERIMENT No. 79 Hydrogen From Iron

(CL-44, CL-55, CL-66, CL-77)

APPARATUS: Sodium bisulfate, copper sulfate, powdered iron, test tube, candle or alcohol lamp.

PROCEDURE: Dissolve five measures of sodium bisulfate and one half measure of copper sulfate in a test tube one quarter filled with water. Heat slowly, then add one measure of powdered iron.

SUMMARY: Iron frees hydrogen from sodium bisulfate. The products are ferrous sulfate and free hydrogen. As in the previous experiment, the copper sulfate acts as a catalyst.


EXPERIMENT No. 80 Hydrogen From Zinc

(CL-44, CL-55, CL-66, CL-77)

APPARATUS: Sodium bisulfate, zinc, test tube, candle or alcohol lamp, copper sulfate, wooden splint.

PROCEDURE: Dissolve four measures of sodium bisulfate and one half measure of copper sulfate in a test tube one quarter filled with water. Add two or three small pieces of zinc. Heat carefully and note the reaction. Place a lighted wooden splint at the mouth of the test tube.

SUMMARY: Zinc displaces hydrogen from sodium bisulfate to form zinc sulfate and free hydrogen. (The sodium bisulfate dissolved in water forms dilute sulfuric acid.) The copper sulfate in this experiment serves to hasten the reaction and is called a catalytic agent in the language of chemistry.

generating hydrogen gas


EXPERIMENT No. 81 Test for Hydrogen Peroxide


APPARATUS: Sodium chromate, drinking glass, hydrogen peroxide solution (drug store), hydrochloric acid, test tube, and stirring rod.

PROCEDURE: Dissolve one half measure of sodium chromate in a test tube one quarter filled with water and add four drops of hydrochloric acid. Note the light orange color. Pour this solution into a drinking glass and add five drops of hydrogen peroxide and stir.  Note that the solution becomes blue and then rapidly becomes light brown.

SUMMARY: This test is often made to demonstrate the presence of hydrogen peroxide in rain water.


Chemists studied the atmosphere for a great many years before the elements helium, neon, argon, krypton and xenon were discovered. With the exception of helium, which is obtained from natural gas, all the other members of this group are obtained from liquid air.


There is very little helium in the atmosphere, only about four parts in every million. Next to hydrogen it is the lightest of all gases, but still is about twice as heavy as hydrogen. The United States is able to produce more helium than any other country because of our natural gas wells in Texas and Utah. It has thus supplanted hydrogen in U. S. Navy’s lighter-than-air craft for some years. It does not have quite as much lifting power as hydrogen, but its safety factor more than outweighs this small disadvantage - it being non-inflammable.

Of the other rare gases in the helium group, neon and argon are the only ones having any commercial importance. We first applied neon gas in a commercial way, about fifteen years ago when neon signs came into existence. These glass tube signs contain a very small amount of neon gas and when a current of electricity is passed through the tube, the characteristic red glow of neon appears. Recent developments in the field of electric signs have now made it possible to improve greatly on the original design of neon signs so that a great variety of colors and effects are obtainable today, many without the use of neon.


We have already seen that the name given to the process when oxygen combines with an element to form an oxide, is oxidation.

The reverse of this process can also take place, that is, the oxide can give up its oxygen and the name given to the process is reduction.  Various substances such as sodium bisulfite or hydrogen may be used as reducing agents or we can use our blowpipe to accomplish this purpose.

EXPERIMENT No. 82 The Reduction Of An Iron Compound

(CL-44, CL-55, CL-66, CL-77)

APPARATUS: Sodium carbonate, ferric ammonium sulfate, magnet, blowpipe, charcoal block, candle.

PROCEDURE: Place on your charcoal block a mixture of one measure of sodium carbonate and a half measure of ferric ammonium sulfate.  Direct a reducing flame at this mixture by means of the  blowpipe.  Heat it for two or three minutes (see instructions on use of the blow-pipe). Pass a magnet over the tiny pieces and observe what happens.

SUMMARY: When the flame is directed at this mixture, we are able to reduce the ferric or iron compound to its metallic state, proved by the fact that it is attracted by the magnet.

EXPERIMENT No. 83 Reduction Of Iron Salicylate

(CL-55, CL-66, CL-77)

APPARATUS: Ferric ammonium sulfate, sodium salicylate, test tube, sodium bisulfite.


thermit welding

Metal and Thermit Corp.  

The upper photograph shows the heat reaction resulting from the ignition of aluminum and iron oxide used in Thermit welding. Lower left is an example of a typical casting fracture requiring the thermit weld. Lower right is a diagram of the actual welding process wherein the molten metal is poured into the form built around the fractured casting.


PROCEDURE: Put one eighth of a measure of ferric ammonium sulfate and an equal amount of sodium salicylate in a test tube and add water up to the three quarter mark. Note the color of the solution.  Add five measures of sodium bisulfite and shake well. Again note whether any change in color occurs.

SUMMARY: The deep red solution is ferric salicylate. Sodium bisulfite acts as a reducing agent and readily reduces the ferric salicylate to ferrous salicylate, a colorless compound. The reducing action takes place on the iron which changes from a ferric to a ferrous state.

EXPERIMENT No. 84 Reduction Of Logwood

(CL-55, CL-66, CL-77)

APPARATUS: Logwood, test tube, candle or alcohol lamp, sodium bisulfite.

PROCEDURE: Fill a test tube three quarters full of water. Add two measures of logwood and heat gently. Stop heating when the solution becomes a deep red and pour this off into another test tube. Add five measures of sodium bisulfite. Shake well and note any changes.

SUMMARY: Sodium bisulfite again acts as a reducing agent and removes oxygen from the logwood. When this takes place, the solution becomes colorless.


Thermit welding has often been called industry’s master maintenance tool as this form of welding makes it possible to repair broken parts on heavy machinery, locomotives, ships’ rudder frames, anchors and other large castings which, due to their large size, otherwise could not be repaired.

The thermit reaction is essentially a reduction of iron oxide with aluminum used as the reducing agent. Thermit itself is a mechanical mixture of finely divided aluminum and iron oxide. When this mixture is ignited, a reaction of tremendous heat occurs. The aluminum, having a high affinity for oxygen, leaves the iron in a pure metallic form which, because of its weight, falls to the bottom of the mixture. The molten iron is then directed into a form built around the portion to be welded so that the area surrounding the broken part itself is heated to a very high temperature. The whole mass then integrates or fuses and the seam or crack is entirely obliterated. When the casting is cool, the form is removed, the surplus material cleaned off and the casting is as good as new.

In times of war, thermit incendiary bombs are widely used. The tremendous heat generated by these bombs makes it exceedingly difficult to combat them. Dry sand is considered the most effective extinguisher to use.

ContentsLionel Chem-lab - Chapter 4    or   Back to the Experiments Page

"The Science Notebook"  Copyright 2008-2018 - Norman Young