Lafayette College
Department of Chemistry

Experiment 34

HARD WATER

Object: To study the causes of hard water and methods for softening it.

Discussion:

The control of the hardness of water is one of the practical problems of chemistry which is met in almost every home and industry. Wherever water is heated, soap is used, steam is generated or water is used in a manufacturing process such as a chemical industry, application of dyes, etc., the problem of hard water arises. The selection of the sites for many manufacturing plants is based upon the question of whether the water in that community is hard or soft. It is obvious, then, that this is a problem which is of vital importance to the layman, engineer and chemist.

No lengthy discussion of the causes and treatment of hard water will be given in the introduction since it is believed that a thorough understanding of these problems will be gained from the experimental part.

Directions:

a) Causes of hardness in water.

Obtain about 3 ml of solutions of calcium chloride, calcium nitrate and calcium sulfate in three test tubes and add several drops of soap solution to each. (You should obtain about 25 ml of soap solution in your small beaker and keep it at your desk for testing purposes.) (1) What ion is common to the three solutions? Assuming that soap is sodium stearate, $C_{17}H_{35}COONa$, write an equation for its reaction with this ion. Add several drops of soap solution to a few ml of distilled water. Shake this tube and compare the results with those obtained by shaking the soap with solutions of the calcium salts. (2) Explain the difference. Teast a solution of magnesium ions with soap. Repeat with a solution containing ferric ions. Add a pinch of pure calcium carbonate to a few ml of distilled water. Mix thoroughly and filter. Is the filtrate hard of soft? (3) What two conditions must be met by salts which cause hardness?

Add a pinch of powdered calcium carbonate to about 25 ml of distilled water in your small erlenmeyer flask and pass carbon dioxide from a Kipp generator through the suspension for two or three minutes. If you have not added too much of the solid you should be able to see that it has disappeared or at least that there is less present. (4) Remembering that carbon dioxide and water form carbonic acid, account for the change which has taken place. Equation. Filter the resulting solution and reserve for tests. Test a few ml of the filtrate with soap solution. (5) Account for the results, showing the origin of any ions which

Exp. 34 - 2

enter into the reaction.

b) - Removal of hardness.

Gently boil 5 ml of calcium sulfate solution for two minutes. Cool the solution and test with soap solution. Has boiling had any effect upon the hardness of this solution? Repeat this test using 5 ml of the filtrate from the last part of a) above. (6) Is there any sign of a change? Is the water hard after boiling? Account for the results. Water which acts like the filtrate from a) is known as temporary hard water, while that which acts like the calcium sulfate solution is known as permanent hard water. (7) Differentiate between the two.

To 3 ml of calcium sulfate solution add about 3 ml of sodium carbonate solution. (8) Write an equation accounting for the precipitate which is formed. What ions were removed from the solution? Would you expect the resulting filtrate to be hard or soft? Filter the suspension and test with soap. Repeat this test using tri-sodium phosphate solution in place of sodium carbonate. Repeat using sodium tetraborate solution (borax). (9) Write all equations. (10) Make a generalized statement about the method of softening permanent hard water.

To 3 ml of the filtrate from part a), temporary hard water, add about 1 ml of sodium hydroxide solution. Test the resulting solution with soap. (11) Account for the results by means of an equation. On an industrial scale, temporary hard water often is softened by adding calcium hydroxide. (12) Write the equation. It can be seen, however, that if an excess of this base is used, since it is soluble and ionized, the water still will be hard.

c) - Quantitative determination of hardness.

One of the methods most frequently used in determining the degree of hardness of water is to measure the amount of a standard soap solution which a given quantity of the water consumes before a permanent lather can be obtained. This method does not distinguish between the hardness caused by the different ions, it merely measures the total effect. Since this hardness may be the result of the presence of several ions from different compounds, the results are expressed in terms of the number of parts of calcium carbonate which must be present in a millions parts of water to produce the same degree of hardness. The concentration of the soap solution has been expressed in these terms and this value will be found on the label of the bottle.

All determinations will be made on 25 ml of hard water and thus it is necessary to know the amount of soap solution required to produce a lather with 25 ml of water which contains no hardness. This is known as the blank. Measure 25 ml of distilled water into a wide mouth bottle and arrange a solid or plugged rubber stopper to fit it. Rinse and fill a 50 ml burette with standard soap solution. Record the burette reading and then add the soap solution, a few tenths of a ml at a time. Stopper the bottle and shake between additions. The end point is reached when a lather is obtained which will persist for at least three minutes with the bottle lying on its side. Make the final reading and calculate the amount of soap solution used.

Exp. 34 - 3

The next determination is that of some artificial hard water which contains both permanent and temporary hardness. This water has been made up to contain a great deal more hardness than any natural water would be likely to contain. Obtain a measured 25 ml of the hard water in a clean, dry beaker and boil gently for five minutes. Allow the water to cool to room temperature, pour it into your graduated cylinder and add enough distilled water to make the volume up to 25 ml. Titrate with soap solution as directed in the case of the blank above. Be sure that the bottle which you use has been washed thoroughly with tap water and rinsed with distilled water. While the water is heating and cooling in the above part, titrate 25 ml of the unboiled water. From these two titrations it is possible to calculate the total hardness, temporary hardness and permanent hardness. Complete the information required on the data sheet.

Repeat the procedure described in the above paragraph using 25 ml of laboratory tap water in place of the artificial hard water. Record the data and make all calculations.

Questions and Problems:

I. Briefly discuss the effect of hard water on heating plants from the standpoint of heat lost and corrosion.

II. Devise a scheme for the removal of temporary and permanent hardness on a commercial scale.

III. What are the advantages of the permutite type of water softener?

IV. Explain in detail, writing all necessary ionic equations for equilibria involved, the formation of limestone caverns. How do stalagmites and stalactites form?

Experiment 35

IRON AND ITS COMPOUNDS

**Object:** To study some of the reactions of ferric and ferrous ions and their practical applications.

**Discussion:**

Iron is one of the first metals encountered in this course which only displays more than one valence. This experiment will be concerned chiefly with the means of identifying the ions in the different valence states and of converting them from one valence level to another. Many of the practical applications of the chemistry of this metal involve such changes.

**Directions:**

Prepare the blue print paper according to directions given on page 3 of this experiment. This will permit time for the paper to dry before the end of the period.

Obtain about one-half gram of ferrous ammonium sulphate and dissolve it in a test tube of distilled water which has been freshly boiled. Add an iron nail to the solution and stopper the tube. Obtain a test tube of ferric chloride from the reagent bottle. In each of two test tubes obtain 3 or 4 ml of potassium ferrocyanide and add a few drops of ferric chloride to one tube and to the other some of the ferrous ammonium sulphate solution. (1) The results of this and all the following tests are to be recorded in tabular form. Repeat the above test using each of the following in place of the potassium ferrocyanide: potassium ferricyanide, potassium thiocyanate, ammonium hydroxide, and sodium hydroxide. Record the color of the precipitate or solution formed in each case. (2) Write equations for all reactions which occur.

Obtain two 5 ml portions of a solution containing ferrous ions (from side shelf). Make one of the solutions acidic with dilute sulphuric acid and pass hydrogen sulphide through both of them for about half a minute. (3) How do you account for the results? How would it be possible to separate cupric and ferrous ions if both were in the same solution?

Add an iron nail to about 3 or 4 ml of dilute hydrochloric acid and allow it to stand for several minutes (it may be necessary to warm the tube in order to start the reaction). Dilute with about twice the volume of distilled water and make tests on portions to determine the valence state to which the iron has been oxidized by the hydrogen ion. (4) Write an equation using only those ions involved in the change. What role does the iron play in this reaction? Dilute about 2 ml of ferric chloride solution with an equal volume of distilled water and add two iron nails. Acidify the solution with a few drops of dilute hydrochloric acid. Feat the mixture and allow it to stand for several minutes. Mix thoroughly by shaking while holding your

Exp. 35 --2

thumb over the mouth of the tube and then test for ferric ions with potassium thiocyanate. (5) What has happened to the ferric ions? In which valence state will this element exist in the presence of an excess of the metal? Write an equation for the change. Show the method of balancing.

The ferrous ion is an active reducing agent. Acidify about 2 ml of ferrous sulphate with dilute sulphuric acid and then add potassium permanganate solution a few drops at a time until a faint color persists. Test for the presence of ferrous and ferric ions. (6) Balance the following equation:

$$Fe^{++} + MnO_4^{-} + H^{+} \rightarrow Fe^{+++} + Mn^{++} + H_2O$$

Repeat the test using potassium dichromate in place of potassium permanganate:

$$Fe^{++} + Cr_2O_7^{2-} + H^{+} \rightarrow Fe^{+++} + Cr^{+++} + H_2O$$

Both of these reactions are used in the quantitative determination of the iron in a solution. All of the iron is reduced to the ferrous state and then either permanganate or dichromate solution is added in sufficient quantity to just oxidize it to the ferric state. The quantity of iron may be calculated if the concentration and the volume of the oxidizing solution is known.

The reaction of the ferric ion with the thiocyanate ion offers an excellent opportunity to illustrate the effect of concentration on an equilibrium. The reaction, as is indicated below, is reversible:

$$Fe^{+++} + 3 CNS^{-} \rightleftharpoons Fe(CNS)_3$$

The ferric thiocyanate which is formed is a soluble compound with a deep red color whereas the ions on the left side of the equation do not have an appreciable color. Thus the relative concentrations of the ions and the undissociated molecule at equilibrium can be estimated from the color of the solution. Add 4 ml portions of ferric chloride solution (0.1N) and Potassium thiocyanate solution (0.1N) to 35 ml of distilled water. Mix thoroughly and divide so that you have about 5 ml in each of four test tubes. To the first tube add 4 ml of ferric chloride solution, to the second 4 ml of potassium thiocyanate, and to the third about 2 g. of solid potassium chloride and distilled water to make the volume equal that in the other two tubes. Compare the color of each tube with that of the standard, i.e., the fourth, and determine in which direction the equilibrium has been shifted. Account for the change in each case. (7) Record your results. Give a statement of the law involved.

One of the practical applications of the reactions of iron is in blue printing processes. The actual preparation of the blue print paper should be carried out at the beginning of the period so that the paper will have time to dry. The following test should be performed at this time in order to understand the process. In a darkened room prepare two test tubes each containing 3 ml of ferric chloride solution and 3 ml of oxalic acid solution. Leave one of the test tubes in the darkened room and expose the other to bright sun

Exp. 35 --3

light for 5 to 10 minutes. Add potassium ferricyanide to each tube.
(8) What has happened to the ferric ion in each case? What is the
role of oxalic acid in this reaction?

**Preparation of blue print paper:** Use precautions to prevent the
solutions in this part of the experiment from coming in contact with
your clothes. They will cause stains which are difficult to remove.

In a darkened room mix 15 ml of ferric ammonium citrate solution
with an equal volume of potassium ferricyanide. Place several layers
of newspaper on the desk top and place on it about three sheets of
white paper (supplied). Coat these evenly with the solution using
absorbent cotton as a sponge. Set them aside on fresh newspaper in
the dark room and allow them to dry. In the meantime prepare or
obtain some object to be used in the printing. Any relatively flat,
opaque object such as the type of key used in a "Yale" lock may be
used. Other objects which may be used are photographic negatives
or a drawing which you make on tracing paper or any other thin trans-
parent paper. The drawing must be made with either India ink or a
very soft 'lead' pencil.

When the prepared paper is dry place it on a book or a piece of
cardboard, coated side up. Place the object to be printed on the
paper (drawings and negatives should be covered with a dry clean glass
plate) and expose it to the bright sunlight for several minutes.
Return to the darkened room, remove the object, wash the paper first
in dilute hydrochloric acid solution (1 part dilute HCl to 25 parts
of water), and finally in water. The print may now be dried in
ordinary light. It can be labeled by using a 0.1N sodium hydroxide
solution as an 'ink'. (9) Explain the chemical changes involved in
the printing of the blue print paper.

**Questions:**

I. What weight of potassium permanganate would be required to oxidize 3.72 gms of ferrous sulphate to ferric sulphate in an acid solution?

II. What weight of hydrogen chloride would be in the acid necessary to react with 4.75 gms of iron? What weight of acid would be used if it is 27% HCl? What volume of hydrogen could be collected if it were measured over water at 19°C and 754 mm?

Experiment 36

PHOTOCHEMISTRY

Object: To observe chemical action of light and basic reactions used in photography.

Discussion:

Light is a form of energy and as such has power to cause certain chemical reactions to take place. Visable light, especially in the red end of the spectrum, is not very energy rich and so is able to bring about only those reactions which take place quite readily. On the other hand light in the blue part of the spectrum and particularly in the ultra-violet (invisable) is much more rich in energy and can cause reactions which require considerable energy. It is a fundamental law in photochemistry that only light which is absorbed can cause a chemical reaction.

You are familiar with photochemistry in your everyday life. In the summer on the beach the sunlight causes a complicated reaction to take place in your body and you acquire the envied deep smooth tan or a lobster red depending upon your type of skin and upon your moderation. You know that light is essential to normal healthy growth of plants and animals. Everyone has used a camera and knows that chemical action of light is the basis of image formation on the negative.

In this experiment you will make one or two tests which illustrate fundamental reactions employed in photography. When light strikes silver halides a chemical reaction occurs which liberates metallic silver, which forms a latent image. Silver chloride reacts with sodium thiosulfate (hypo) to produce the complex sodium silver thiosulfate:

$$ \underline{\text{AgCl}} + \text{Na}_2\text{S}_2\text{O}_3 \rightarrow \text{NaAgS}_2\text{O}_3 $$

Directions:

Obtain about 2 ml of silver nitrate solution in each of two test tubes. In that which follows one is to be exposed to light, but the other must be kept in the dark as much as possible. Add 2 ml of sodium chloride solution to each tube. (1) Describe the precipitate, noting especially its color, Write ionic equations for its formation. Expose the one tube and precipitate to strong sun light for 15 - 25 minutes. Keep the other dark. (2) What do you observe to happen to the precipitate exposed to light? What caused the action? When you are convinced of the action, treat both precipitates with sodium thiosulfate solution. (3) What happens in each case? (4) Are both precipitates completely soluble in excess sodium thiosulfate? Describe the appearance of any insoluble material. (5) Explain how use may be made of this reaction in developing a photographic negative. (6) Name at least two complex ions of silver aside from any which may have been formed in this experiment. Write equations showing their formation.

Supplementary Experiment IV
WHY REACTIONS GO TO COMPLETION

**Object:** To study the various types of chemical reactions and determine why each type goes to completion.

**Discussion:**

Chemical reactions can be classified in five different groups: 1. simple combination, 2. simple decomposition, 3. simple replacement, 4. double decomposition, and 5. oxidation-reduction. Reactions of the first three types are governed by principles embodied in the electro-chemical series. Elements high in the series are likely to unite vigorously with almost all nonmetals, conversely, their compounds are usually stable: elements high in the series will replace those less active. (1) On the side of your report sheet list the electrochemical series, and the halogen activity series. Double decomposition reactions will go to completion if one of the products is removed from the reaction mixture through insolubility, volatility, non-ionization, or complex ion formation.

**Method:**

A. Simple combination reactions:

Take about ½ inch of iron fillings in a test tube and mix with an equal volume of powdered sulfur. Heat strongly in a pyrex test tube until a reaction occurs. (2) Is this an exothermic or an endothermic reaction? Write a balanced equation for the reaction. Reserve a portion of the product for part 4. Heat a small piece of steel wool in a hot flame. (3) What evidence of a reaction is there? Write a balanced equation. (4) Which of the following reactions would you expect to take place? Write equations for those that will:

Simple replacement reactions need not involve two elements. Whenever two substances react to produce a third compound, the reaction has been one of simple combination. Pass carbon dioxide from the Kipp generator through 5 ml of distilled water for three minutes. Test the resulting solution with litmus. (5) Write a balanced equation for the reaction. Place sufficient phosphorus pentoxide to cover the tip of your horn spoon in a test tube and add 10 ml of water. **Caution!** Do not permit phosphorus pentoxide to touch the skin. (6) What evidence of reaction was there? Test the solution with litmus. (7) Write an equation for the reaction. Add a small piece of calcium oxide to 5 ml of water in a test tube and test with litmus. (8) Equation. (9) Write equations for the reactions between sulfur dioxide and water; sulphur trioxide and water; nitrogen pentoxide and water. Bring the stopper from a bottle of ammonia water to the mouth of an open bottle of hydrochloric acid. (Ammonia gas is given off by ammonia water. (10) What evidence of reaction is there?

Sup. Exp. IV - 2

B. Simple decomposition.

Heat about three grams of mercuric oxide in a test tube with a strong flame until a reaction is evident. Test for oxygen by inserting a glowing splinter into the tube. Repeat with calcium oxide, ferric oxide, copper oxide, and magnesium oxide. (12) Do your results agree with the results you would expect from the electrochemical series? Write balanced equations for all decompositions that take place. What form of energy could be used to decompose compounds stable to heat? (13) Write balanced equations for the decomposition of calcium hydroxide, ammonium hydroxide, sulfurous acid, sulfuric acid, ammonium nitrate, and ammonium nitrite.

C. Simple replacement.

Take three 5 ml samples each of cupric sulfate, mercuric nitrate, ferric chloride, silver nitrate, magnesium sulfate, HCl and calcium chloride solutions. Add a small piece of granulated zinc to one tube of each solution, a small iron nail to the second tube of each solution, and a copper turning to the third tube of each solution. (14) Write balanced equations for each reaction that takes place, and explain why the other substances did not react. To five ml of sodium bromide solution add 1 ml of chlorine water; add 1 ml of carbon tetrachloride and shake. Repeat with sodium iodide solution. Add 1 ml of bromine water to 5 ml of sodium chloride and sodium iodide solutions in separate tubes, add carbon tetrachloride and shake. (15) Write balanced equations for the reactions that take place.

D. Double decomposition reactions.

(16) Write balanced equations for all reactions that go to completion in this section, underlining all precipitates.

1. Reactions that take place through insolubility.

To 3 ml of silver nitrate, mercurous nitrate, lead nitrate, and mercuric nitrate in separate tubes, add 1 ml of hydrochloric acid.

To 3 ml of cupric sulfate, magnesium sulfate, sodium carbonate, silver nitrate, and calcium chloride in separate tubes, add 1 ml of barium chloride solution.

To 3 ml of ferric chloride, barium chloride, silver nitrate, and copper sulfate solutions, add 1 ml of sodium hydroxide solution.

2. Reactions that take place through volatility.

To a small piece of ferrous sulfide (prepared in part A.) add 5 ml of hydrochloric acid *under the hood*.

Add hydrochloric acid to a pinch of sodium bicarbonate, and repeat with sodium carbonate, and with sodium bisulfite.

To a pinch of ammonium chloride, add 3 ml of sodium hydroxide solution and warm. Smell cautiously.

3. Reactions that take place due to the formation of water (Neutralization).

Take 5 ml of sodium hydroxide in an evaporating dish and add three drops of phenolphthalein solution. Add hydrochloric acid drop by drop until the pink color just disappears. Evaporate to dryness, dissolve the residue in 5 ml of distilled water, and test with litmus. Repeat with limewater and nitric acid. Write

Sup. Exp. IV - 3

Aluminum hydroxide and sulfuric acid

Barium hydroxide and phosphoric acid

Ferric hydroxide and nitric acid.

### E. Complex ion formation.

To 5 ml of copper sulfate solution in a test tube add ammonia water drop by drop until a precipitate forms. Add 5 ml additional ammonia water. (17) Write a balanced equation for the reaction. Repeat with silver nitrate and ammonia water. (18) Equation.

F. To 2 ml of potassium permanganate solution, add 5 ml of sodium bisulfite solution. (19) What evidence for reaction is there? Write a balanced equation for the reaction. Repeat with potassium dichromate and sodium bisulfite. (20) Equation. Show the electron changes in each case. Without repeating these reactions, write balanced equations for the following reactions showing all electron changes:

Copper and conc. nitric acid

Copper and dilute nitric acid

Copper and hot conc. sulfuric acid

Sulfur and conc. nitric acid

Phosphorus and conc. nitric acid

(21) Which of the reactions in parts A and B are oxidation-reduction reactions? Show the electron changes that take place.

### Exercises:

Complete the following general equations:

Active metal and active non-metal

Metal oxide and water

Non-metal oxide and water

Metal and acid

Metal and salt of less active metal

Halogen and salt of less active halogen

Acid and base.

ONLY A BOY

I remember the first time I tried it--
I was only a kid of fifteen
And even tho' she was much younger than me
She was far more compos'd and serene.

I was eager, yet awkwardly backward;
Uncertain of how to proceed,
But she seemed not to notice the hesitancy
With which I prepared for the deed.

It was out in the barn, I remember,
At the close of a lush, summer day
And the evening was scented with clover in bloom
And fragrant of freshly mown hay.

I remember I spoke to her softly,
As I cuddled her face in my hands,
And I saw in the depths of her wide eyes the look
Of a loved one who understands.

I remember she moved a bit closer,
How I stood, while my head seemed to spin
With the thought of the thing I was going to do--
Yet reluctant, somehow, to begin.

And her eyes seemed, I thought, to rebuke me,
For waiting--for being afraid,
And even old Nellie, our ancient plow horse,
Looked over her manger and neighed.

Long later, I stood up, uncertain
Of whether to stay or to run,
A-tingle with pride, yet shaken and awed
As I know that, at last, it was done.

I remember (it seemed hours later)
How my heart hammered under my blouse
With the joy of a boy that's turned into a man
As I made my way back to the house.

Twenty years have gone by since that evening
But I've never forgotten our vow,
The thrill and joy that I felt as a boy
On that day when I milked my first cow.

Experiment 20

MOLECULAR WEIGHT

**Object:** To learn to use one of the methods of determining the molecular weight of a substance.

**Discussion:**

There are several methods of determining the molecular weight of a substance, but not all are applicable to every substance. In Experiment 15 you used one of these methods, depression of freezing point, which can be used only in the case of soluble un-ionized substances. The method which will be illustrated in this experiment can be applied to gases or liquids which can be converted into gas without decomposing. You will determine the molecular weight of the gas oxygen.

It is well known to you at this stage of the course that one gram molecular weight of any gas, measured under standard conditions, has a volume of 22.4 liters. It is obvious therefore that in order to find the molecular weight, all that is needed is sufficient data from which to find the weight of this volume of gas. Often this information is obtained by acutely weighing a known volume of gas but in this experiment you will obtain the weight by indirect methods.

**Directions:**

Obtain a rubber stopper for your large pyrex tube and fit it with a delivery tube for the collection of gas over water. Cut a piece of glass tubing to act as an extension on the delivery tube. This should be long enough so that the extended delivery tube almost reaches the bottom of an inverted gas bottle which is sitting on the bridge of your pneumatic trough.

Add about one gram of powdered manganese dioxide to the ignition tube and heat it in the flame until no more moisture is given off. Be certain to heat the tube its full length so as to drive all the moisture from the walls. Allow the tube to cool to room temperature, add about 2 g of potassium chlorate (S.R.) and weigh on the analytical balances. The potassium chlorate has been dried. Rotate the tube carefully, holding it at a slant so that the manganese dioxide and potassium chlorate are mixed but none of the mixture must be spilled. A loss of any of this material would introduce a serious error in your results. Clamp the tube with a buret clamp at the extreme top so that it rests at about a 45° angle with the table top. It is important that the following operations be done in the order listed. Fill a gas collecting bottle and invert it on the bridge of your trough. Insert the rubber stopper in the ignition tube. Pass the delivery tube extension through the hole in the bridge and into the bottle, then connect it to the end of the delivery tube. Heat the mixture in the ignition tube very gently, starting from the top and moving down until you have collected about three quarters of a bottle of the gas. Do not try to collect a full bottle for if any gas escapes from the bottle the experiment must be started over.

Exp. 20 -2

After you have collected the required amount of gas, be certain
that the end of the delivery tube extension is above the level of
the water in the gas bottle and remains so until the ignition tube
and its contents have cooled to room temperature. The delivery tube
must be removed from the gas bottle *before* the rubber stopper is
disconnected from the ignition tube. (1) Why?

Adjust the bottle so that the levels of the water inside and outside are the same, cover the mouth with a glass plate and set up right on your desk. By means of your graduated cylinder determine the volume of water needed to refill the bottle. What does this volume represent? weigh the tube but do not empty it afterward. Record all the necessary data on the sheet supplied. Since you used only a small portion of the oxygen from the potassium chlorate, it is not necessary to start with a new mixture. The last weighing of the first trial becomes the first weighing for the second trial. Complete the information on the data sheet and report the results your Instructor.

Questions:

I. Briefly describe the steps involved in applying this general method of determining the molecular weight of a liquid. Could the molecular weight of all liquids be determined in this way? Explain

II. 215 ml. of a gas, collected over water at 21°C and 752 mm. wa.
found to weigh 0.5375 g. What is the molecular weight of the gas?
If it is known that its molecule contains equal numbers of carbon
and hydrogen atoms, what is the formula?

III. Why was it not necessary to know the weight of the manganese dioxide or the potassium chlorate in the experiment? How would moisture in the manganese dioxide have affected the results?

Experiment 22

NITRIC AND SULPHURIC ACIDS

**Object:** To study the chemical characteristics of nitric and sulphuric acids and their salts.

**Discussion:**

Usually when compounds are studied simultaneously in the laboratory it is because they are derived from elements which appear in the same group in the periodic table. In such cases the elements usually exhibit the same valence and often marked similarities in their chemical properties. Nitrogen and sulphur are not in the same group but are in adjacent groups and therefore do not resemble each other as much as each resembles the elements in its group. The acids derived from these non-metals are, however, very similar in many respects and therefore are studied together in this exercise.

Both of these compounds are classified as strong acids but this refers to one characteristic only, the complete ionization of the molecules into hydrogen ions and negative ions. (1) According to the Bronsted theory what ions would be present in a solution of each of these acids? (Remember that sulphuric acid is dibasic.) According to this more modern theory what is the real acid in these solutions? Explain. In this experiment you will study many types of behavior that depend upon the presence of these hydrogen ions, such as turning litmus red or neutralizing bases. In such cases the compounds are behaving as acids.

Both of the compounds discussed above have other characteristic types of behavior which are not typical of acids and are in no way concerned with the presence of the hydrogen ions. Thus they act as strong oxidizing agents and when they react in this way, the fact that they are acids have no influence on the reaction. It is common to hear the beginning student make the statement that "nitric acid is a strong acid because it dissolves copper." Both parts of his statement are true and if he would omit the word "because", the statement as a whole would be true. Nitric acid reacts with copper because it is a strong oxidizing agent and not because it is a strong acid. In this experiment you will study some of the reactions which take place when these compounds act as oxidizing agents.

Sulphuric acid exhibits a third type of behavior which is in no way connected with the two types previously described. As you will see from some of the tests which you perform in this exercise, it is an active dehydrating agent. It does this by the formation of several hydrates, the most important of which results from the following change:

$$H_2SO_4 + H_2O \longrightarrow H_2SO_4 \cdot H_2O$$

Exp. 22 - 2

The compounds of nitrogen as a whole are rather unstable. Many of them decompose readily and produce gases which occupy volumes many times those of the compounds from which they were produced. This is one of the very important properties which make nitrogen compounds so valuable as explosives. The salts of nitric acid do not decompose readily enough for this purpose. They are, however, broken up with comparative ease by heating. The completeness of this decomposition is determined by the metal which forms the positive part of the salt.

### Directions:

Neither of the acids studied here will be prepared during this experiment. (2) Write equations for the preparation of sulphuric acid. What are the conditions used? (3) Would pressure influence the second reaction? Explain. (4) Write an equation for the preparation of nitric acid from sodium nitrate. Explain why the other compound used in this reaction is particularly adapted to this purpose. (5) What is the modern method of manufacturing nitric acid? (6) What is the anhydride of nitric acid? Write the equation for the reaction of the compound with water. Why is this method not used in the production of nitric acid on a commercial scale?

Test dilute nitric and sulphuric acids with litmus paper. What particle is responsible for the results? (7) Write ionic equations for the reaction of each of the acids with a soluble hydroxide of a metal. Write equations for the reaction of the acids with a basic oxide.

Add a piece of zinc to about 5 ml of dilute sulphuric acid in a test tube. (8) Describe the results and write an ionic equation. Rewrite the equation using only the molecules, atoms or ions which changed. (9) Would these same results be obtained with copper? (If in doubt make a test.) Explain in terms of electrons. Repeat the above test using zinc and dilute nitric acid under the hood. (10) Is there any evidence of hydrogen being given off? Explain fully making reference to any pertinent information which you gained by observing the above tests.

Prepare dilute sulphuric acid by adding about 5 ml of the concentrated acid to about 3 ml of water in a small beaker. (Notice the order in which these liquids are mixed. It is important.) Add the dilute sulphuric acid which you have prepared to about one gram of sodium chloride in a test tube. Is there any evidence of chlorine being given off? Repeat the test using concentrated nitric acid and sodium chloride. Warm the tube gently under the hood. Is chlorine given off? Interpret these results. What do they indicate with regard to the relative oxidizing power of nitric and sulphuric acids?

Add a few pieces of copper turnings to about 5 ml of dilute nitric acid from your reagent shelf. What evidence is there that the acid is reduced? (12) Explain why it is impossible in the test as you have made it, to determine whether nitrogen dioxide

Exp. 22 - 3

or nitric oxide was the reduction product. (13) Write balanced equations in which each is formed. Show the method of balancing. What conditions are necessary in each case? Repeat the above test using concentrated sulphuric acid and copper turnings. Warm the tube *gently* and wave some of the evolved gas toward your nose. (14) What is it? Write a balanced equation accounting for its formation. (Allow the sulphuric acid to cool and then discard it in a sink into which water is running. Do not pour hot sulphuric acid into water and never pour any acid into the sink without flushing it down the drain with plenty of water.) Notice that the reducing agent used above was a metal. To what class of compounds do the oxidation products belong?

Mix about 1 ml of dilute sulphuric acid with half a test tube of water. Add to this 3 or 4 ml of barium chloride solution. The white precipitate which is formed is barium sulphate. Write an ionic equation for the reaction. Rewrite the equation using only those ions which entered the reaction. Test the solubility of the precipitate in dilute hydrochloric acid. The test with hydrochloric acid distinguishes barium sulphate from other white precipitates which might be formed. Repeat the test using barium chloride and sodium sulphate solution. (16) Write the ionic equation.

Add 1 ml of dilute nitric acid to 5 ml of distilled water in a test tube. To this add 2 or 3 ml of freshly prepared ferrous sulphate solution and mix thoroughly. Hold the tube at a slant and pour concentrated sulphuric acid down the inside wall of the tube so that it sinks and forms a layer under the water solution. If you have difficulty, your Instructor will demonstrate the method. Observe the results. This is a test for the nitrate ion provided no iodide nor other interfering ions are present. (17) Account for the formation of the brown ring by means of one or more equations.

To about half a gram of sulphur in your large pyrex tube add 2 or 3 ml of concentrated nitric acid. Warm gently under the hood for about two minutes and dilute with about 5 ml of distilled water. Filter and test with barium chloride solution. (18) Write an equation accounting for the ion which you found to be present. (19) To what class of compounds does the oxidation product belong when the *non-metal*, sulphur, reacts with nitric acid?

Slowly and with constant stirring add 5 ml of concentrated sulphuric acid to 10 ml of water in a small beaker. (20) What evidence is there that a chemical change has taken place? What type of compound has been formed? (21) Why is it essential that sulphuric acid be added to water rather than the reverse process? Is this true of all acids? Explain. Drop a wood splint into about 5 ml of concentrated sulphuric acid. (22) If the composition of wood is represented by $(C_6H_{10}O_5)_x$, account for the change which you have observed.

Exp. 22 - 4

Questions:

I. How many liters of SO₃, measured at 40°C and 745 mm, would be needed to make 7 liters of 0.82 N sulphuric acid?

II. What factors determine the reduction products obtained when nitric acid acts as an oxidizing agent? Use the metals in the electrochemical series to explain one of these factors.

III. Write equations showing three practical ways in which nitric acid may be obtained from atmospheric nitrogen.

IV. Explain why no hydrogen is obtained when metals which are above hydrogen in the electromotive series react with moderately concentrated nitric acid. Is it possible to obtain hydrogen by the reaction of nitric acid on a metal? Explain.

V. What is aqua regia? Write equations for its formation. Explain why it is used instead of nitric acid in dissolving the noble metals.

VI. A solution of sulphuric acid with a density of 1.1394 gms per ml contains 20% of the acid. What volume of this would be needed to make 5 liters of 0.25 N acid?

Exp. 24 -2

secondary sodium ortho-phosphate (Na₂HPO₄) or disodium phosphate,
tertiary sodium ortho-phosphate (Na₃PO₄) or trisodium phosphate.

Among other points to be covered in this experiment are tests
which will distinguish the three types of phosphates and also the
three types of ortho-phosphate salts. In most of these tests, as
in the majority of qualitative tests, use will be made of the form
ation of precipitates as identifying characteristics of unknown
substances. Although it is not possible to consider the formation
of these precipitates from the quantitative point of view at this
time, it would be well to consider it from the qualitative side.
A precipitate is the result of a chemical change in which an in-
soluble compound is one of the products. Another concept of this
phenomenon and one which will prove more useful later is that a
precipitate will form as a result of mixing two solutions which
contain more of two kinds of ions than can exist together in a
stable solution. As an illustration we might use the well known
test for chlorides, the precipitation of silver chloride. Silver
chloride is a very slightly soluble substance and thus only a very
small concentration of silver ions and chloride ions can exist to-
gether in solution. The product of the concentrations of these
ions which can exist together in solution is known as the solu-
bility product. At a given temperature the solubility product of
a slightly soluble compound is a constant which is characteristic
of that compound. If relatively concentrated solutions of sodium
chloride and silver nitrate are mixed, there will be in the solu-
tion *momentarily* more silver and chloride ions than can remain to-
gether in a stable solution. Another way of saying this is that
the solubility product of this compound, silver chloride, has been
exceeded. The excess ions form a precipitate which separates from
the solution until the number of ions remaining just satisfies the
solubility product of the compound. An understanding of the above
principle will aid in interpreting some of the tests for the phos-
phates.

As mentioned above, phosphoric acid is a tri-basic acid and
among the first of the poly-basic acids which you have studied.
In discussing the ionization of acids we have been concerned
chiefly with the mono-basic acids and these are relatively simple
since each molecule, if it is ionized at all, is ionized entirely.
This is not true of acids such as sulphuric and phosphoric acid but
instead these poly-basic acids ionize in steps. This can be
illustrated best by an equation using phosphoric acid as an example.
Pay particular attention to the size of the arrows as these in-
dicate the relative degrees of ionization.

$$
\mathrm{H_3PO_4 \rightleftharpoons H^+ + H_2PO_4^-}
$$

$$
\text{[Chemical structure of } H^+ + HPO_4^{2-} \text{ ---} \text{ [Chemical structure of } H^+ + PO_4^{3-} \text{ ---}]
$$

$$
\text{[Chemical structure of } H^+ + PO_4^{3-} \text{ ---} \text{ [Chemical structure of } H^+ + H_2PO_4^- \text{ ---}]
$$

Exp. 24 -3

(4) Explain why you would not expect to find many phosphate ions in a solution of phosphoric acid. (5) In the second equilibrium above, H₂PO₄⁻ belongs to what class of compounds? Why? HPO₄²⁻ is acting as what type of compound in this same equilibrium? Explain. (6) Does it always act as this type of compound? Explain. (7) List all the acids found in the above equilibria. All of the bases. Which is the strongest acid? Base? (8) Write equations representing the ionization of sulphuric acid.

Some of the most important uses of the phosphates depend upon the fact that the sodium salts hydrolyze. At this time you should review Experiment 17, which deals with hydrolysis. The following ionic equations represent the hydrolysis of the three sodium ortho-phosphates. Study these equations carefully now and again after you have performed the part of this experiment which deals with the hydrolysis of these salts. (Note that the heavy arrow indicates the preponderant concentrations at equilibrium.)

$$ \text{NaH}_2\text{PO}_4 \rightleftharpoons \text{Na}^+ + \text{H}_2\text{PO}_4^- \quad \uparrow \quad \text{H}^+ + \text{HPO}_4^{2-} $$

$$ \text{Na}_2\text{HPO}_4 \rightleftharpoons 2\text{Na}^+ + \text{HPO}_4^{2-} $$

$$ \text{Na}_3\text{PO}_4 \rightleftharpoons 3\text{Na}^+ + \text{PO}_4^{3-} $$

$$ \text{HOH} \rightleftharpoons \text{OH}^- + \text{H}^+ \quad \downarrow \quad \text{HPO}_4^{2-} $$

$$ \text{HOH} \rightleftharpoons \text{OH}^- + \text{H}^+ \quad \downarrow \quad \text{H}_2\text{PO}_4^- $$

Exp. 24 -4

(9) Would you expect a solution of tertiary sodium phosphate to give an acidic or basic reaction? Why? (10) How would the reaction given by a solution of this salt compare with that given by a solution of secondary sodium phosphate? Reason. (11) In the hydrolysis of trisodium phosphate, would you expect a higher concentration of HPO₄²⁻ or H₂PO₄²⁻? Explain.

### Directions:

**Oxidation of phosphorus:** Place a piece of red phosphorus about the size of a pea in a small pyrex test tube and add one ml of concentrated nitric acid. Warm this gently under the hood. Treat the sample with two additional separate portions of the acid and warm between additions. Finally heat to drive off all of the nitric acid, allow the tube to cool and add about 15 ml of distilled water. Stir to dissolve as much of the solid as possible and, if the solution is still cloudy, filter it. Test the filtrate with litmus. What type of compound has been formed?

(12) What does this tell you regarding the element phosphorus?

(13) Write a balanced equation for the reaction which took place. Show valence changes clearly. Reserve the filtrate for future tests.

**Tests for phosphates:** Add about 2 ml of the filtrate obtained above to 5 ml of ammonium molybdate solution. If a precipitate does not form at once warm the tube and set aside for a few minutes. The precipitate is ammonium phosphomolybdate, and its formation indicates the presence of the phosphate ion.

Another standard test for the phosphate ion is the formation of a precipitate when "magnesia mixture" is added. Add 3 or 4 drops of conc. ammonium hydroxide to 3 ml of magnesium chloride. The white precipitate which forms is magnesium hydroxide, and it was formed because there were too many magnesium and hydroxyl ions in the same solution; the solubility product of magnesium hydroxide was exceeded.

$$Mg^{++} + 2OH^{-} \rightleftharpoons Mg(OH)_{2}$$

Add about 3 ml of ammonium chloride solution to the tube and shake it.

In order to understand what has happened it is necessary to recall why the precipitate formed and the source of the ions which formed it.

$$\text{a) } \text{MgCl}_2 \rightarrow \text{Mg}^{++} + 2\text{Cl}^{-}$$

$$\text{b) } \text{NH}_4\text{OH} \rightleftharpoons \text{NH}_4^+ + \text{OH}^{-}$$

Magnesium chloride is a very soluble salt and is completely ionized, and thus supplies a high concentration of Mg⁺⁺. Ammonium hydroxide is a weak base, ionizes only slightly, and thus supplies few hydroxyl ions. This low concentration, however, is too great to remain in solution with the Mg⁺⁺. When ammonium chloride is added, the large concentration of ammonium ions which it supplies to the solution affects the ionization of the ammonium hydroxide and shifts equilibrium (b) to the left. The result of this is that fewer hydroxyl ions are left in solution; in fact,

Exp. 24 -5

not enough to exceed the solubility product of magnesium hydroxide. Thus the "magnesia mixture" has a high concentration of magnesium and ammonium ions and a small concentration of hydroxyl ions.

Add 1 or 2 ml of the filtrate of phosphoric acid obtained by the oxidation of phosphorus to 3 ml of the "magnesia mixture." The precipitate, which is ammonium magnesium phosphate, often forms slowly. Notice its characteristics. How does it compare with magnesium hydroxide?

The silver ion is used as the precipitating agent in another test for the phosphate radical. This test works best when the solution is nearly neutral so disodium phosphate (secondary sodium ortho-phosphate) will be used. Dissolve a small amount of this salt in 1/4 of a test tube of water and add silver nitrate solution. Notice the color and other characteristics of the precipitate formed. Add dilute nitric acid to the tube and note the results. (14) Write an ionic equation for the formation of the precipitate. Would this distinguish a phosphate from an iodide? Explain. Add silver nitrate to 3 ml of the phosphoric acid filtrate obtained in the first part of the experiment. If no precipitate is formed add dilute sodium hydroxide one drop at a time until the solution is faintly basic to litmus.

Record in the table below the observations made in the following tests: Add one or two crystals of sodium pyro-phosphate to 3 ml of water in a test tube. Add silver nitrate solution. The precipitate is silver pyro-phosphate (Ag₄P₂O₇). Add dilute nitric acid. Test the pyro-phosphate with a mixture of equal parts of dilute nitric acid and ammonium molybdate. Acidify another portion of the pyro-phosphate with acetic acid and test with egg albumen solution. Repeat all the above tests using sodium metaphosphate (NaPO₃) is place of sodium pyro-phosphate.

<table><thead><tr><th></th><th>AgNO<sub>3</sub></th><th>Ammonium Molybdate</th><th>Egg Albumen</th></tr></thead><tbody><tr><td>Ortho-phosphates</td><td></td><td></td><td></td></tr><tr><td>Meta-phosphates</td><td></td><td></td><td></td></tr><tr><td>Pyro-phosphates</td><td></td><td></td><td></td></tr></tbody></table>

Preparation of the sodium salts of pyro- and meta-phosphate acids: Heat one gram of disodium phosphate (Na₂HPO₄·12H₂O) in your pyrex ignition tube. This should be heated gently and rotated until the water of hydration is driven off, then clamp the test tube near its tip and heat it strongly for fifteen minutes. Allow the tube to cool and dissolve the contents in 10 ml of water. (15) Make tests to determine what compound has been formed. Write equations for the changes that take place.

In the same way as directed above, heat 1 g of sodium ammonium phosphate or some mono-sodium phosphate. Cool, dissolve in water and make tests to identify the salt formed. (16) What is given off during this change? Devise a test to identify it,

Exp. 24 -6

Write the equation for the reaction which has taken place during
the heating.

Add about 0.1 g of phosphorus pentoxide to 15 ml of distilled water in a small beaker and stir the mixture. (Do not allow any of the P<sub>2</sub>O<sub>5</sub> to come in contact with your hands) Test the solution with litmus. What does this indicate about the nature of the element phosphorus? Apply test to determine what was formed. (17) Write an equation for the reaction.

By means of the indicators listed in Experiment 17, determine
as closely as possible the pH of the dilute solutions of the three
sodium ortho-phosphates and the solution of phosphoric acid. The
process can be speeded considerably by using litmus paper to de-
termine which indicators are to be used on each solution. Obvious-
ly, if the solution is acidic to litmus, the use of phenolphthalein
would be a waste of time. (18) Record your results in tabular
form. Do your results agree with those expected after a study of
the hydrolysis equations given in the Introduction.

Qualitative analysis: In the experiments which you have per-
formed you have learned to test for many of the negative ions. You
will now be given an unknown substance and asked to identify the
negative ion which is present. The substance may contain any of
the following ions: nitrate, sulphate, and phosphate, iodide,
chloride or bromide. All of these, except the phosphate, will be
present either as the normal sodium salt or as the acid. The
phosphate ion may be derived from any one of the three sodium
ortho-phosphates or from phosphoric acid. In addition to report-
ing the negative ion which is present, you are to determine also
the compound from which it is derived. Report your unknown num-
ber and the results of your tests to your Instructor before leav-
ing the laboratory.

Questions:

I. How many ml of 0.5 N KOH would be required to neutralize the acid which could be prepared by completely oxidizing 1 g of phosphorus?

II. What phosphorus compound is most widely used as a fertilizer?
Why? How is it prepared? Give equations.

III. Explain the fact that, although few phosphate ions exist in
a phosphoric acid solution, the addition of ammonium molybdate
practically completely precipitates this substance.

IV. By use of equations and a brief explanation, show how the
sodium salts of phosphoric acid can control the acidity of a
solution.

Experiment 26

BISMUTH

**Object:** To complete the study of the Nitrogen Family of elements.

**Discussion:**

With this experiment you will complete your study of the Nitrogen Family of Group V of the Periodic Table. Bismuth is the heaviest of these elements and, therefore, it would be expected to show the most metallic properties of this family. However, bismuth salts are hydrolyzed more or less strongly in aqueous solution much in the same manner as antimony salts. For example, it is necessary to have a considerable amount of free hydrochloric acid in solutions of both antimony trichloride and bismuth chloride to prevent hydrolysis from occurring to such an extent that the oxy-chloride is precipitated. Bismuth, like the other members of this family shows valences of +3 and +5, all though +3 is the more common.

**Directions:**

Dissolve a few granules of metallic bismuth in about 2 ml of concentrated hydrochloric acid, adding a few drops of concentrated nitric acid and warming to aid solution. (1) What is the function of the nitric acid? What do you observe which proves this? Dilute the solution of $BiCl_3$ which you have obtained in this manner with about half a test tube of water. The white precipitate which you obtain is $BiOCl$. (2) Write ionic equations to account for its formation. (3) What is the name given to this type of reaction? Add concentrated hydrochloric acid to the precipitate and observe whether it is soluble in excess acid. (4) Explain this behavior on the basis of ionic equilibria.

Obtain 2 ml of bismuth chloride solution from the side shelf and add sodium hydroxide solution at first slowly, then in large excess. (5) Describe your observations and write ionic equations for them. (6) Is bismuth hydroxide amphoteric? Explain.

Place a bright iron nail or piece of iron in a few drops of bismuth chloride solution. (7) Explain and write ionic equations for the reactions involved.

Bubble hydrogen sulfide through a few ml of bismuth chloride solution. (8) Describe and write equations for the reaction.

Add sodium hydroxide solution slowly to 2 ml of stannous chloride solution ($SnCl_2$). When the permanent precipitate which first forms has been re-dissolved by excess sodium hydroxide you have a solution of sodium stannite ($Na_2SnO_2$). (9) Write ionic

Exp. 26 - 2

equations to explain its formation. Is stannous hydroxide amphoteric? Explain. Sodium stannite is a fairly strong reducing agent, it being oxidized to sodium stannate ($Na_2SnO_3$). Use this solution of sodium stannite for the following test for bismuth. Neutralize about 3 ml of bismuth chloride with sodium hydroxide. Filter and wash with about 5 ml of distilled water. Discard the filtrate. Now pour about 3 ml of the sodium stannite solution which you prepared onto the precipitate of bismuth hydroxide. (10) Write equations and explain your observations. This constitutes a test for bismuth which is employed rather frequently in analytical procedures.

(11) On the basis of the various experiments which you have performed to date, devise a scheme of analysis, whereby you could detect any one of the following substances:

$H_3PO_4$, $H_2PO_4^-$, $HPO_4^{2-}$, $PO_4^{3-}$, $As^{3+}$, $AsO_3^{--}$ (or the meta form $AsO_2^-$), $H_3AsO_3$, $Sb^{3+}$, $SbO_3^{--}$ (or as meta form $SbO_2^-$), $Bi^{3+}$.

Obtain from your Instructor an unknown containing one or more members of the above list and analyze it to determine which are present. Report those substances which you have determined to be present, giving the observations on which you have based your decisions.

Questions and Problems:

I. Theoretically what volume of $H_2S$ at 735 mm and 26°C would be required to precipitate all of the bismuth from a solution containing 0.552 g of $BiCl_3$? What experimental facts make it necessary to use larger quantities?

II. What volume of 2.79 normal sodium hydroxide is required to convert 91.5 ml of 1.372 molar $BiCl_3$ into $Bi(OH)_3$?

III. Devise and describe tests by which you could prove whether any given element has an hydroxide which shows amphoteric properties. What other tests are there by which you could distinguish metals, non-metals and metalloids. Write equations which are illustrative of the reactions which you mention in answering all parts of the question.

Experiment 27

CARBON AND CARBON DIOXIDE

**Object:** To study the properties of carbon, carbon dioxide and the compounds derived from this oxide.

**Discussion:**

Carbon is one of the most important of all the elements. Its compounds are more numerous than those of any other element and the uses of both the element and its compounds are found in all fields of human activity. The compounds which carbon forms with hydrogen, oxygen, nitrogen and a few other elements are so numerous that the study of them constitutes a separate branch of chemistry, "organic chemistry". In this present experiment you will be concerned chiefly with the element in its uncombined condition and with its most important oxide carbon dioxide. This latter substance, and the compounds derived from it, carbonates, are among the inorganic compounds of carbon.

The element itself has three major uses and innumerable ones of less importance. In the form of coke and charcoal it is used as a fuel; in these same forms but principally in the form of coke it is used as a reducing agent; in the form of charcoal it is used as an adsorbing agent. These three uses will be illustrated in this experiment.

The fact that carbon dioxide is relatively unreactive and has no harmful physiological effect except at very high concentrations makes it one of the most widely used gases. The behavior of this gas and the compounds derived from it illustrate some of the most important principles of chemistry, the majority of which have been met before in this course and will be reviewed here in connection with the study of this element.

It was discovered early in the course that carbon dioxide is an acid anhydride. (1) Describe the test and results in Experiment 6 from which this information was gained. The acid which is formed when this oxide reacts with water is both weak and unstable and these facts explain a large part of the behavior of carbon dioxide and the carbonates. The following represents the equilibria which exist in a solution of this gas in water:

Experiment 29

ANALYSIS OF COAL

**Object:** To determine the approximate composition of a sample of coal.

**Discussion:**

The importance of coal as a fuel is demonstrated by the fact that the output of this product in the United States is approximately 500,000,000 tons per year and that 53% of all the energy used is obtained from coal. This energy is obtained in the form of heat through the process of burning and then may be converted into any one of several other forms such as electrical or mechanical energy.

When a consumer purchases coal he is purchasing potential energy and one of the most important things which he wishes to know about the product is the fuel value or the amount of heat which can be obtained when a definite weight of the material is burned. This value can be expressed in terms of calories per gram but in practice is more often expressed in terms of British Thermal Units (B.T.U.) per pound. The latter unit is defined as the quantity of heat required to raise one pound of water one degree F. The most direct method of obtaining this value is by means of a calorimeter which is an instrument used to measure the quantity of heat evolved during a chemical reaction such as burning. If possible such an instrument will be demonstrated and explained in the laboratory.

Coal is a mixture rather than a compound and is the result of a slow and gradual change in the vegetative matter which grew during the carboniferous period in the earth's history. The various types of coal represent the steps in this process but it must be emphasized that even within one type such as bituminous wide variations in composition are found. The composition of coal, which can be determined by an analysis, is another indicator of its value as a fuel. It is even possible to approximate the fuel value from the information obtained in an analysis. The following table indicates the way in which the fuel value varies with the composition of the material. Remember that the information represents the analysis of a typical sample of each kind of coal and should by no means be taken to mean that all samples of anthracite coal would have the composition given here:

Exp. 29 - 2

<table><thead><tr><th>Sample</th><th>Water</th><th>Volatile Matter</th><th>Fixed Carbon</th><th>Ash</th><th>Fuel Value B.T.U. per lb.</th></tr></thead><tbody><tr><td>Wood</td><td>7.4%</td><td>----</td><td>----</td><td>0.4%</td><td>4,500</td></tr><tr><td>Peat</td><td>10.0</td><td>50.0</td><td>25.0</td><td>10.0</td><td>4,300</td></tr><tr><td>Lignite</td><td>33.0</td><td>29.0</td><td>29.0</td><td>7.0</td><td>6,500</td></tr><tr><td>Bituminous Coal</td><td>3.4</td><td>35.5</td><td>51.3</td><td>9.7</td><td>13,000</td></tr><tr><td>Anthracite Coal</td><td>3.3</td><td>3.3</td><td>85.0</td><td>9.1</td><td>13,350</td></tr></tbody></table>

## Directions:

(a) Moisture: The percent of moisture in a sample of coal is determined by drying a weighed sample in an oven which is held at 105 degrees C. and determining the loss in weight. Since the necessary equipment is not available for use in this laboratory, this value has been determined for each of the samples supplied and will be found on the label of the container. Obtain a sample and record its number and the percent of moisture on the data sheet.

(b) Volatile matter other than moisture: Weigh a clean crucible and cover and record the weight on the data sheet. Powder the sample of coal and weigh about two grams in your crucible by difference. Set the covered crucible on a triangle and heat with the full flame of your burner. This heating should continue for seven minutes.

Cool to room temperature and weigh. (1) Explain why it is important to have the cover firmly on the crucible during the above heating. Calculate the total percent of volatile matter in the sample and then the volatile matter other than moisture. (2) Name as many of the materials as you can which are evolved in this procedure. Are these materials present in the original coal as such or are they formed as a result of the heating?

(c) Fixed carbon: Place the uncovered crucible with the remains from the above determination on the triangle. Heat strongly until only a white ash remains. The process can be speeded by tilting the crucible slightly and rotating it from time to time. Care must be taken that no flames nor drafts enter the crucible and that none of the contents is spilled.