Exp. 34 - 4

Data Sheet

Name Henry Kissinger Date March 1, 1944

Hardness value of soap solution: 1 ml = 19.4 ppm CaCO3

## Blenk

<table><tr><td>Final burette reading</td><td>9.5 cm</td></tr><tr><td>Initial burette reading</td><td>6.5 cm</td></tr><tr><td>Volume of soap used</td><td>2.0 cm<sup>3</sup></td></tr></table>

## Artificial Hard Water

Unboiled water:

<table><tr><td>Final reading</td><td>30.8 cc</td></tr><tr><td>Initial reading</td><td>9.5 cc</td></tr><tr><td>Total volume used</td><td>27.3 cc</td></tr><tr><td>Total volume less blank</td><td>18.3 cc</td></tr><tr><td>Total hardness of water</td><td>355.02 ppm.</td></tr></table>

Boiled Water:

<table><tr><td>Final reading</td><td>26.9 cc</td></tr><tr><td>Initial reading</td><td>9.5 cc</td></tr><tr><td>Total volume used</td><td>18.4 cc</td></tr><tr><td>Total volume less blank</td><td>12.4 cc</td></tr><tr><td>Permanent hardness of water</td><td>240.56 ppm.</td></tr><tr><td>Temporary hardness of water</td><td>114.46 ppm.</td></tr></table>

## Tap Water

Unboiled water:

<table><tr><td>Final reading</td><td>10.4 cc</td></tr><tr><td>Initial reading</td><td>5 cc</td></tr><tr><td>Total volume used</td><td>5.4 cc</td></tr><tr><td>Total volume less blank</td><td>2.4 cc</td></tr><tr><td>Total hardness</td><td>46.56 ppm.</td></tr></table>

Boiled water:

<table><tr><td>Final reading</td><td>32.6 cc</td></tr><tr><td>Initial reading</td><td>27.9 cc</td></tr><tr><td>Total volume used</td><td>4.7 cc</td></tr><tr><td>Total volume less blank</td><td>1.2 cc</td></tr><tr><td>Permanent hardness of water</td><td>22.98 ppm.</td></tr><tr><td>Temporary hardness of water</td><td>13.58 ppm.</td></tr></table>

Supplementary Experiment II

THE ATMOSPHERE, NITROGEN AND NITROGEN COMPOUNDS

Directions:

Place 4 gms of potassium (or sodium) nitrite, 2 gms of ammonium chloride and about 20 ml of water in a small Erlenmeyer flask. Shake until the solid has dissolved and then clamp the flask to a ring stand. Equip the flask with a delivery tube and prepare to collect bottles of gas over water. Fill your 600 ml beaker with cold water and have it near by. The flask may be submerged in this water if the reaction becomes too rapid. Warm the flask gently with a very low flame. Control the rate of reaction by removing the flame when necessary. Collect and discard a half a bottle of gas. This would not be pure nitrogen. (1) Why? Collect one bottle of the gas, cover and set on the desk.

Place a short piece of magnesium ribbon in a small dry pyrex test tube. Hold the test tube in a horizontal position with the magnesium in the bottom. Insert the dry delivery tube from the generator half way in the test tube and after a few minutes heat the magnesium. What is the product formed? (2) Equation. Place a drop of water on the product and test for the odor of any gas escaping. (3) Equation accounting for this action. Thrust a burning splint into the bottle of gas which you collected. (4) List all the properties of nitrogen that you have been able to observe. Classify them as physical or chemical.

Obtain about 5 ml of lime water (calcium hydroxide) in a test tube and blow through it for a minute by means of a glass tube. (5) Account for the change observed. Pour 5 ml of lime water in a watch glass and set it on your desk, Observe the surface after about half an hour. What component of the atmosphere have you detected? Is this substance present in a fixed proportion? (6) Name two factors that tend to increase the proportion of this substance. One that decreases it.

You have observed two of the components of the atmosphere and now you will determine the percent of oxygen in the air. Fit your large test tube with a one hole rubber stopper which has a piece of glass tube in it. The tube should extend about 2 cm above the top of the stopper. Obtain 3 ml of pyrogallol solution and add to it 3 ml of sodium hydroxide from your desk. Quickly pour this into the test tube, put the stopper in place and place your finger firmly over the end of the glass tube. Shake the tube and contents for a couple of minutes and then invert the test tube in a beaker of water. After the mouth of the tube is well under the surface of the water, remove your finger. When the liquid no longer rises in the tube raise or lower it until the levels of the liquids inside and outside are the same. Place your finger over the glass tube and remove the test tube from the water. Measure the contents of the test tube. Since the pyrogallol absorbs the oxygen, the volume of the liquid represents the volume of this gas that was in the air. (7) Calculate the percent of oxygen in the air from these results and any additional measurements which you need to make. Tabulate all necessary information and label all calculations so that they are clearly understood.

Sup. II-2

Place one or two crystals of calcium chloride on a glass plate on your desk and observe it at the end of the laboratory period.
(8) Account for the results in terms of vapor pressures. Under what conditions would the calcium chloride not give these results? Observe the wet-dry thermometer demonstration in the front of the laboratory.
(9) Account for the difference in the readings. How would the readings compare in a room with hundred percent humidity? Explain. How would low humidity affect them?

## Ammonia:

Place five grams of ammonium chloride in your large test tube and add about five grams of calcium hydroxide to it. Rotate the tube until the two solids are mixed thoroughly. Fit the test tube with a one hole rubber stopper and a right angle glass tube. Clamp the tube so that the delivery tube points upward. Heat the contents of the test tube. Bring a bottle of hydrochloric acid close to the delivery tube.
(10) Account for the results. Equation. Place a small test tube over the delivery tube and after about a minute quickly place it mouth downward in a beaker of water.
(11) Write an ionic equation for the reaction that takes place. What substance acted as an acid? base? Hold a piece of wet litmus over the delivery tube.
(12) What ion is responsible for these results? Place about half a gram of ammonium chloride in a test tube and add 2 or 3 ml of sodium hydroxide solution. Warm the tube gently while a piece of wet red litmus is held above it.
(12) Write ionic equations for the reactions involved. Why is this a test for all ammonium salts?

## Oxides of Nitrogen:

Set up a generator after the model found in the front of the laboratory. Notice that the thistle tube should reach almost to the bottom of the bottle. Charge the generator with not more than 5 gms of copper turnings. Add water until the end of the thistle tube is covered (at least 10 ml). Now add an equal volume of concentrated nitric acid. Notice the color of the gas at first. What change occurs? (13) Account for the original color and this change. (14) Write a balanced equation for the production on nitric oxide by this method. Show method of balancing. Collect one full bottle of the gas over water. Collect a half a bottle and leave it inverted in the water. Thrust a burning splint into the full bottle of gas. Does nitric oxide support combustion? This test would distinguish it from what other oxide of nitrogen? Bring one of the oxygen generators to your desk. This consists of a tube filled with potassium chlorate and manganese dioxide. When this mixture is heated the potassium chlorate decomposes and oxygen is evolved. (15) Write the equation. What part does the manganese dioxide play? Heat the tube gently with the delivery tube in water. Allow a steady stream of bubbles to pass out for about 30 seconds and then place the end of the delivery tube under the bottle which is half full of nitric oxide. (16) Explain the two observations which you can make. When the water has risen nearly to the top of the bottle remove the delivery tube from the water and stop heating the potassium chlorate. What does the gas bottle now contain? (17) Write an equation which will account for its formation. In what commercial processes does this reaction play an important part?

Sup. II-3

Nitric acid:

This compound is one of the strong acids and as such has
properties which are characteristic of acids in general. (18) In what
way does a strong acid differ from a weak one? Write an equation for
the ionization of this acid. Place about 5 ml of water in your
evaporating dish and add 3 or 4 drops of sodium hydroxide. Then add
2 drops of phenolphthalein. What do the results tell you about the pH
of this solution? Add dilute nitric acid a drop at a time. Account
for the change. Write an ionic equation for the reaction. What is the
real acid in this reaction, the real base?

To 5 ml of nitric acid from your desk add 10 ml of water in a
large test tube. Now add a piece of zinc to this dilute acid and
warm gently to start action if necessary. Set the tube aside until
the action is complete and in the meantime continue with another part
of this experiment. Add sodium hydroxide solution until the mixture
in the test tube is basic. Now warm the test tube and hold a piece of
wet red litmus above the mouth. From the results of this test what
conclusion can you draw regarding the product formed when zinc reacts
with very dilute nitric acid? (19) Write a balanced equation. Show
the method. (20) What two factors determine the reduction product
obtained when nitric acid reacts with a metal? Illustrate with results
obtained in the experiment. What would you expect to obtain from
concentrated nitric acid on copper? Write the equation.

Place about half a gram of sulphur in your large test tube and
add about 5 ml of conc. nitric acid. Warm gently under the hood. What
is the gas evolved? How can you identify it? Pour the contents of the
tube into about 10 ml of distilled water and reserve for future tests.
Obtain 5 ml of magnesium sulphate solution. What ions does it contain?
Add some barium chloride solution. The precipitate is barium sulphate.
(21) Write an ionic equation for the reaction. This reaction is a test
for what ion? Test the solution which you reserved above for this ion.
(22) Account for its presence if found. Equation.

Obtain a small cube of egg white and add 1 ml of conc. nitric
acid to it in a pyrex test tube. Warm <u>gently</u> and note the color.
Cautiously add sodium hydroxide to the solution. The egg white is a
protein. Why will your fingers turn yellow when nitric acid is spilled
on them.

Dissolve a pinch of sodium nitrate in 3 or 4 ml of water and add about 2 ml of ferrous sulphate solution. Obtain about 4 ml of conc. sulphuric acid in a small test tube and pour this down the inside wall of the test tube containing the nitrate solution. (Your Instructor will demonstrate the method). Notice the ring formed at the junction of the acid and nitrate solution. The equation for this reaction will be found on the board. (23) Balance them.

Sup. III-2

the other five salts is derived. State whether each is weak or strong.
What kind of elements form hydroxides which act as strong bases?

### C. Test for the Alkali Metals:

Obtain about 1 ml (the rounded part of a small test tube full) of a solution of lithium chloride, potassium chloride and sodium chloride in three clean test tubes. In a fourth test tube obtain about 5 ml of dilute hydrochloric acid. Your Instructor will supply you with a platinum wire sealed in a glass tube. Clean this wire by heating it in an almost colorless flame (do not heat the glass handle) and then dipping the wire only in the dilute hydrochloric acid in your test tube. Repeat this dipping and heating until the wire imparts no color to the flame. Now dip the wire in the lithium chloride solution and hold in the flame. Note the color of the flame. Clean the wire as directed above and repeat the substituting sodium chloride solution for the lithium chloride. Look at the sodium flame through two thicknesses of cobalt glass. Is the light from this flame transmitted through the blue glass? Clean the wire and repeat the above test using potassium chloride solution in place of sodium chloride. What is the color of the potassium flame when seen normally? Is this light seen through the cobalt glass? Mix the solutions of sodium chloride and potassium chloride and make a flame test on the mixture. Which of the elements can you detect? Now view the flame through cobalt glass. What element is shown to be present? (9) Explain the use of cobalt glass in testing for the ions of the alkali metals.

## Alkaline Earth Metals

### D. The Elements:

Fill your large beaker a third full of water, place your index finger over the mouth of a small test tube which is completely full of water and lower the test tube in the beaker until the mouth is below the surface of the water. Remove your finger and leave the tube sitting mouth downward in the beaker. Obtain a shaving of calcium metal, drop it into the beaker and then move the test tube until the mouth is directly over the metal. What is the gas being collected? (10) In what way is this action similar to sodium? In what way is it different? (11) Write equation for the action of each of the following with water: calcium, strontium, barium. When the test tube is filled with gas quickly remove it from the beaker and hold its mouth close to a flame. (12) Explain. Drop a small piece of magnesium in 5 ml of water. Is a gas given off? Heat the tube and watch for signs of a gas forming near the metal. How does this metal compare with the others studied? (13) Write an ionic equation explaining the action. How can you explain the difference in activity?

### E. Hydrolysis of Salts:

Add a pinch of calcium acetate to 3 ml of distilled water and test the solution with litmus. This test shows that calcium acetate is the salt of a weak acid and a strong base. Since calcium hydroxide is a strong base, the element, calcium, is a distinct metal. What kind of reaction would you expect a barium acetate solution to give?

Sup. III-3

## F. Tests for the Alkaline Metals:

Obtain 1 ml of a solution of the chloride of each of the alkaline metals in separate test tubes. Make the flame tests on these three metals as directed for the alkali metals above.

## G. Common Useful Compounds of Calcium:

In Experiment 27 you heated calcium carbonate (marble chip) until it was partially decomposed. What were the products? What common name is given to the calcium compound formed? Obtain a small piece of calcium oxide and add two or three drops of water. What evidence is there of a chemical action? Add 5 ml of water and then test with litmus. What compound is formed? (14) Write a balanced equation. Complete the following statement: "A soluble metallic oxide and water produce a BASE." (15) Write two other equations illustrating this statement. What is the commercial name of the product formed when calcium oxide reacts with water? What use is made of this substance? (16) Explain why a solution of this substance would not have a high hydroxyl ion concentration. Why might you expect it to have a high concentration?

Heat a small lump of gypsum in a test tube. What evidence do you have that the original substance was a hydrate? (17) Write the equation for the reaction if you have driven off only part of the water. If you have completely dehydrated the compound. What commercial name is given to the product in each case? What disadvantage has the product in the last case? Obtain about three spoons of plaster of paris in an evaporating dish. Add water and stir to make a thick paste and then press a coin which has been lightly greased with vaseline firmly into the paste. Remove the coin after about ten minutes. (18) Write an equation for the reaction that has occurred. Give four important uses of plaster of paris.

Observe the dish of anhydrous calcium chloride which has been sitting in the laboratory so a considerable time. What has happened to it? What name is given to this process? (19) Explain why calcium chloride is so cheap. Give two important uses of this substance.

## H. Qualitative Tests:

In the above experiment you learned to test for the different alkali and alkaline earth metals and in previous experiments you have learned to identify ammonium compounds. In addition to these positive ions you have learned to distinguish the following negative ions: chloride, bromide, iodide, sulphate, nitrate, and carbonate. You will now be given an unknown, one substance (compound) and you will be expected to identify both the positive and negative ions present. Obtain your unknown substance from your Instructor and give him a clean dry, soft glass (not pyrex) test tube in exchange. Record the number found on the test tube (No __). The substance will be a solid and will be soluble in water. Dissolve a pinch of it in 5 ml of distilled water and make flame tests to determine the positive ion present. Use about 1/4 gm of the solid and make the ammonium ion test according to the directions in supplementary Exp. No. II. Dissolve half of the remaining solid in a test tube of distilled water and using small portions of this solution test for the negative ion present. Report the results of your test to your Instructor on a piece of paper.

Sup. I-1

Supplementary Experiment I

SULPHUR AND ITS COMPOUNDS

Directions:

You are familiar with the fact that sulphur is one of the
elements that exists in several different forms. (1) What name is
given to two or more forms of the same element? Your Instructor will
demonstrate the conversion of rhombic sulphur to the other forms.
Observe each form produced; note its characteristics and the conditions
under which it exists. Your Instructor will heat rhombic sulphur which
changes directly to a light yellow liquid. (2) What form of sulphur
predominates in this liquid? At what temperature does this change take
place? (3) Why is no monoclinic sulphur produced during this change?
(4) Under what conditions could monoclinic form be prepared from rhombic
(5) Describe all the changes that take place as rhombic sulphur is
heated from 20°C to 445°C. Notice how the monoclinic sulphur is pre-
pared; how the plastic sulphur is formed, What will happen to either
of these forms if allowed to stand at room temperature?

Obtain a piece of copper turning and examine the properties.
Drop the metal into a soft glass test tube and add about a gram of
powdered sulphur. Heat the tube moderately under the hood for a few
minutes and then heat strongly to boil off the excess sulphur. Examine
the material in the test tube. How does this copper sulphide differ
from the original copper? (6) Write an equation showing the valence
change involved. What was oxidized? Reduced? (7) Assuming that the
original copper weighed 283 mg. and that the product is pure cupric
sulphide, what is the weight of the final product? What would it weigh
if it were heated until it was completely converted to cuprous sulphide.
Show method clearly. (8) Explain the tarnishing of silver in the
presence of a "rubber band" or an egg.

The preparation of hydrogen sulphide from ferrous sulphide and
hydrochloric acid will be demonstrated. (9) Write an equation for this
reaction. What name is given to a water solution of this gas? Write
equations representing the equilibria that exist in such a solution.
(10) Write an equation representing the burning of hydrogen sulphide
in a limited supply of air. In an excess of air, Hydrogen sulphide
can be prepared also from a mixture of paraffin (a compound of carbon
and hydrogen) and sulphur. (11) Explain how such a mixture could
produce this gas. What other element have you studied that behaves
like sulphur in this reaction?

After you have been instructed in the preparation of hydro-
gen sulphide by this latter method, pass some of the gas through
solutions of each of the following salts: zinc chloride which has been
made basic with ammonium hydroxide, cadmium chloride, and copper sul-
phate. (12) Write an ionic equation representing the change in each
case. Dip a small piece of filter paper in lead acetate solution and
hold the wet paper at the outlet of a hydrogen sulphide generator.
This test is often used to detect the presence of hydrogen sulphide

Sup. I-2

and will be used for this purpose later in this experiment. Pass some of hydrogen sulphide through a dilute solution of potassium permanganate. The unbalanced equation representing this reaction will be written on the board. (13) Balance this equation and show the method clearly. What is the oxidizing agent, the reducing agent? What is oxidized? reduced? Repeat the test using potassium dichromate in place of potassium permanganate. (14) Answer question 13 with regard to this reaction.

Construct a gas generator according to the directions and figure found in Experiment No. . The thistle tube must reach within 1/4" from the bottom of the flask. Remove the stopper and pour about 10 grams of sodium bisulphite into the flask add water enough to cover the end of the thistle tube. Obtain in separate test tubes 5 ml. portions of each of the following solutions: potassium permanganate and potassium dichromate. Carry the generator, a gas collecting bottle and the test tubes to the hoods. Add dilute sulphuric acid through the thistle tube and collect a bottle of the gas by the upward displacement of air. Quickly invert the bottle in a 600 cc beaker half filled with water. (15) What do the results indicate? Test the solution in the beaker with litmus paper. What ion is responsible for this behavior? Account for its formation from sulphur dioxide and water.

Pass some of the sulphur dioxide gas through the permanganate solution and then the dichromate solution. (16) Obtain the unbalanced equations for these reactions from the board and balance them. Pass sulphur dioxide through half a test tube of distilled water for a few minutes and then pour the solution in an evaporating dish. Set the open dish on your desk and reserve the solution for a later part of the experiment.

Sulphuric acid is a very versatile compound. It exhibits three types of properties: those of a strong acid, an oxidizing agent and a dehydrating agent. These properties will be illustrated in this section.

Add about 1 ml of sodium hydroxide to a test tube and test with blue litmus. Now add 2 ml of dilute sulphuric acid to the tube and repeat the test with litmus. The reaction here is typical of what class of compounds? (17) Write an ionic equation representing the change. Add a piece of zinc to dilute sulphuric acid. What is the gas given off? Would any other acid act in this manner? (18) Write an ionic equation. The two reactions illustrated above depend upon the fact that sulphuric acid is a strong acid. What does this mean? In addition to being a strong acid, sulphuric acid has a high boiling point which makes it useful in the preparation of lower boiling acids. Wash out the gas generator used in the previous part of the experiment and charge it with about 5 g. of sodium chloride. Carry the apparatus to the hood and arrange it so that the delivery tube is about 1/4" above the surface of a small amount of water in a 100 cc beaker. Make some dilute sulphuric acid by adding 20 cc of the concentrated acid to 10 cc of water in a small beaker or a flask. (Add the acid to the water; never the reverse). Pour this acid through the thistle tube and watch for signs of a gas being given off. (19) What is the gas?

Sup. I-3

What does the acid supply in this reaction? Why is sulphuric acid better for this purpose than many other acids? (20) Write the equation representing the change that occurs when hydrogen chloride comes in contact with the water. (21) Apply the above method to the preparation of nitric acid. What physical property of sulphuric and nitric acids makes it possible to use this method? Explain.

Add two copper turnings to 5 ml. of concentrated sulphuric acid in a small pyrex test tube. Warm the tube *gently*. Test for a gas by waving some of it toward your nose with your hand. (Do not place the tube directly under your nose nor point it toward your face.)
(22) Why would you not get hydrogen? Account for the presence of the gas. Write an equation and show the method of balancing it. Place about three crystals of potassium iodide in a test tube and add conc. sulphuric acid. Hold a piece of filter paper which has been dipped in silver nitrate solution over the mouth of the tube. Account for the results on a basis of previous observation in this experiment.
What is the dark solid on the sides of the tube? (23) Write an equation for the change and show the method of balancing it. What is the role of sulphuric acid in the above reactions? Account for the difference in products.

Drop a wooden splint into a test tube containing about 3 cc (the round bottom of the test tube full) of concentrated sulphuric acid. Account for the results in this case and for the heat when you diluted the acid in above. (24) What is the role of the acid in this change?

Lafayette College
Department of Chemistry

GENERAL LABORATORY INSTRUCTIONS

A. Safety Precautions and First Aid.

None of the experiments assigned is hazardous when the directions for carrying it out are followed carefully. "Practical jokes" and failure to attend carefully to the work at hand have led on occasions to unfortunate results. Conduct which may lead to injury will not be tolerated.

In the event that any corrosive material, such as strong acids or alkalies, is spilled or spattered on your person or clothing, it should be washed off immediately with copious quantities of water. Never use any solution in an attempt to neutralize the effects of any such material; serious injuries may result. In case of any accident report to your Instructor.

Never taste any substance in the laboratory, regardless of how safe you may consider it, unless specifically directed to do so in your laboratory manual.

B. Reagents and Supplies.

Certain common reagents will be found in glass stoppered bottles on the shelf above your desk. Solid reagents will be found on shelves at one side of the laboratory. Containers are both labeled and numbered, and an alphabetical, keyed list hangs near. Liquids and solutions will be found on shelves on another side of the laboratory. Certain experiments may require Special Reagents or solutions, which will be made available temporarily and labeled "Special Reagent for Exp. No. __". When such reagents are to be used, this fact will be indicated in the directions of your laboratory manual by placing the symbol "(S.R.)" immediately after the name of the reagent.

In obtaining solids, take a watch glass with you to the reagent shelf and measure the desired material onto the watch glass while there. Never take reagent bottles to your desk. Your laboratory directions will generally specify in grams the amount of solid to be used. This means that you must weigh out this amount on the trip or platform balance, at least until you learn to estimate weights closely. The amounts taken should not vary from the specified amounts by more than about a half gram. Avoid taking more of the solid than you need from the bottle and do not return materials to the bottles if there is any possibility that they have been contaminated. Avoid waste of reagents; many of them are expensive and difficult to obtain at times.

Intro. -3

Waste solids must not be put in the sinks but disposed of in the stone jars at the ends of the tables. Waste liquids are to be poured into the sink and flushed down the drain with water. This latter is particularly important in the case of strong acids.

Table tops should be cleaned for a considerable distance on each side of your desk at the end of each laboratory period. This involves washing the top with a wet sponge, taking up the excess water with a sponge which has been squeezed out, and drying if necessary. In addition to the table top you are also responsible for the condition of the shelves and reagent bottles immediately above your desk.

Never leave the laboratory at the end of a period without checking with your Instructor.

Experiment 1

MANIPULATION OF GLASS

**Object:** To learn to make useful pieces of glass apparatus.

**Discussion:** Although the majority of pieces of chemical apparatus are made by expert glass blowers, it is sometimes necessary for the chemist to make apparatus for a special purpose. Glass tubes which are cut to the desired length and bent to form the proper angle are undoubtedly the type of apparatus which the chemist must make for himself most often. In this exercise you will make some apparatus, delivery tubes and stirring rods, which you shall use throughout the remainder of this course. It is therefore essential that you follow directions carefully and try to turn out as good products as possible.

**Directions:**

The first pieces of apparatus to be made are the stirring rods. You have in your equipment a piece of solid glass rod about 14 inches longs. Place this in your largest beaker so that it stands alone resting against one edge of the beaker. Mark a spot on the rod with your finger not more than two inches above the edge of the beaker; the rod is to be cut at this point. Place the rod flat on the table and by moving your file away from your body make a scratch on the rod at the point which you marked. Now grasp the rod firmly in both hands so that the thumbs almost meet at the marked point but on the opposite side of the glass rod from the scratch. Move your hands so as to bend the two ends of the glass rod in toward your body using your thumbs as a fulcrum.

Notice the sharp edges on the broken ends of the glass rod. If left in this condition, these edges will cut your fingers and scratch your beakers. When a piece of glass rod or tubing is cut these sharp edges are removed by fire-polishing. Place the cut end of the glass rod in the bunsen flame just above the blue cone and rotate it slowly. Remove the rod once in a while and examine the end. Continue to heat the rod until the sharp edges have disappeared but do not heat so long that a heavy ball of glass forms or the end of the rod bends. If the other end of the rod has sharp edges, fire-polish it also. Place the hot glass on a wire gauze and allow it to cool. (Hot glass must never be placed directly on the table.) Using the same methods as above, measure off and prepare a stirring rod suitable for your 250 ml. beaker. Be sure that all the pieces of glass rod which you have are fire-polished.

Next you will make a delivery tube. Cut off two pieces of glass tubing about six inches each, using the same method described above for cutting the glass rod. Place the wing top on

Exp 3 -3

Questions and Problems:

1. What weight of 20% salt solution would be required to give the number of grams of pure salt corresponding to your desk number? What volume of this solution? Repeat these calculations for the diluted solution.

2. A reagent bottle of 250 ml capacity is to be filled with sulfuric acid, the density of which is 1.11 g/ml. What will be the weight of liquid contained in the full bottle? This solution of sulfuric acid contains 16% by weight of pure sulfuric acid. What weight of pure sulfuric acid will be contained in the full bottle?

3. Devise a method for determining the density of a solid of irregular shape. Make use of principles learned in both Experiments 2 and 3.

Experiment 4

CLASSIFICATION OF MATTER

**Object:** To learn the meaning of the term property and to distinguish between elements, compounds and mixtures by means of their properties.

**Discussion:** Most students, by the time they have reached college chemistry, could give a precise definition of the word property, but few realize the full significance of the term. To the chemist this term means the outstanding characteristics of a particular kind of matter. A knowledge of the properties of the common kinds of matter is extremely useful information. It is a means by which one substance is distinguished from another and also makes it possible to predict the uses to which a substance may be put. For example, in Experiment 2, you determined the density of a solid. If, upon comparing the results which you obtained with those obtained by your neighbor, you found any significant difference, you would be able to say immediately that the two of you were working with different substances. As an illustration of the second use of a knowledge of properties, consider the most outstanding characteristic of asbestos, the fact that it does not burn. Does this information alone suggest a large number of uses of this material? If more of its properties are known, then attempts to find uses for it will become more of a scientific problem and less a fumbling procedure of trial and error. The fact that it is white, fibrous, can be twisted into long cords all becomes valuable information. When you encounter a list of the properties of some substance, either as a result of direct observation in the laboratory or reading your text, think of them in terms of the values mentioned above; how can you identify this material if you see it again and what important uses are suggested by these properties?

It will be noticed that properties describe any particular kind of matter but not a particular piece of matter. Characteristics which apply to one bit of matter only are known as attributes. These of course are not nearly so important as the more general properties. Perhaps an illustration will distinguish between these two terms more clearly. If it were said that a piece of glass was 10 cm square, 1 cm thick, and weighed 250 grams, these would be attributes for they describe that particular piece of glass only. On the other hand, if it were said that glass is fragile, denser than water, and translucent, these would be properties. They apply to any piece of glass.

Properties are divided into two classes on a basis of the method of measuring and expressing them. Those which are measured accurately and expressed in terms which allow careful comparisons

Exp. 4 -2

are known as quantitative properties while the others are known as qualitative properties. The density of the solid which you measured in Experiment 2 is a quantitative property, as is the boiling point of a liquid. On the other hand, color, odor, taste, as they generally are observed and expressed are qualitative. It should be pointed out that there is no sharp division between these two. Although color is generally a qualitative property, it is possible to use a scale of colors with which this property may be measured and expressed accurately. It then would become a quantitative property

A discussion of elements, compounds, and mixtures will not be included in this introduction. This topic has been covered thoroughly in lecture, in your text, and the following experiments have been planned so that you can draw your own conclusions regarding the distinction between these forms of matter after you have carried out the described test.

Directions:

a) Examine some iron filings and note as many of the following properties as you can: Color, state, hardness, odor, and estimated density. Put a pinch of the filings in a test tube and add about a quarter of a test tube of hydrochloric acid. Is there any evidence of gas being evolved? If not, warm the tube gently. Smell the gas cautiously and describe its odor. Dip one of the poles of a magnet in paper and bring it near a small amount of the filings. Results? Place a pinch of the filings in a porcelain crucible and heat for about thirty seconds. Is there any change? Observe all the above properties of iron carefully; you will need this information later to determine whether or not a substance is iron. (1) Record in tabular form all the properties which you have observed above.

b) Repeat the above examination using powdered sulphur instead of iron. (2) Record your observations.

c) Obtain from your Instructor on a watchglass a small amount of material for examination. Is it homogenous? Use the tests which you applied to iron and sulphur. (3) How does this material compare with these elements? Give a reason for your answer. (4) Why would you classify the material as a mixture?

d) Mix thoroughly about 3 grams of iron filings and two grams of powdered sulphur. Pour this mixture into a soft glass test tube and heat gently under the hood. Remove the tube from the flame once in a while. Do the contents continue to glow after it has been removed from the source of the heat? After you have noticed this gow, continue the heating more strongly until no more dark fumes of sulphur can be seen in the tube. Allow the tube to cool, remove the contents and examine it to determine whether it is iron, sulphur, a mixture of these two, or a new substance.

Exp. 4 - 3

Try the effect of a magnet on it. Break off a small piece and test it with hydrochloric acid. Is a gas given off, and, if so, is this the same gas which was evolved when iron was treated with hydrochloric acid? (5) What property of the gas enabled you to answer this question? (6) Give a reason for your answer.

Questions and Problems:

I. List three properties of sugar which would distinguish it from pure white sand.

II. To what class of matter do iron and sulphur belong? Give a reason for your answer.

III. (a) Distinguish between a compound and an element.

(b) Between a compound and a mixture.

IV. It was found that 11 grams of iron combined with 6.3 grams of sulphur to form iron sulphide. How much sulphur would combine with 17 grams of iron?

V. It was found that 13.4 grams of carbon could be obtained from 34 grams of sugar. What is the percent of carbon in sugar?

VI. List as many properties as you can of glass which make it a good material for the manufacture of chemical apparatus. What are some of the properties which account for the disadvantages?

VII. In the first part of this experiment it was stated that a significant difference in the densities of two objects proved that they did not consist of the same material. Explain why the lack of such a difference in the density would not necessarily prove the materials to be the same. What would be necessary to prove them the same?

VIII. A regularly shaped block of sulphur measuring 5 cm x 3 cm x 2 cm has a mass of 62 g. What is the density? What mass of this substance would be needed to occupy a volume of 65 cc?

Exp. 5 -3

sulphur. Add two thirds of this quantity to the metal in the crucible and with a solid stirring rod carefully mix the metal and the sulphur. If any of the mixture is spilled here, an error will be introduced. Wipe the rod on the remaining one third of the sulphur to remove any pieces of the metal that may be adhering. Add this sulphur to the crucible, remove to the hood and continue the experiment as in the case of the copper.

Questions and Problems:

I. Examine the posted results obtained by the class. What conclusion can you draw regarding the quantity of copper and thus that combine? Do all the results support this conclusion? How do you account for this?

II. How do the quantities of the three metals which combine with one gram of sulphur compare? Can you answer question (b) in the discussion from these results?

III. Define in your own words a chemical law. Why would the result of this experiment be inadequate as a basis for a chemical law? What law does this experiment illustrate? Give a statement of it.

IV. Why was it not necessary to weigh the quantity of sulphur added to the crucible? Could you have added too much? Too little? Explain.

V. List all the evidence you can that the substance left in the crucible after heating was not the original metal.

VI. How many pounds of metal can be obtained from a ton of the metal sulphide on a basis of the results which you obtained in this experiment?

Experiment 6

OXYGEN

**Object:** To study the properties of oxygen, to learn about factors influencing rates of chemical reaction and to distinguish acidic and basic oxides.

**Discussion:**

Oxygen is the most abundant element on earth and without it life would cease to exist. It is, therefore, important to investigate some of its properties, the way in which it reacts and to learn from what it may be produced.

Oxygen is contained in very many different compounds. Some of them decompose and yield free oxygen when they are heated to moderate temperatures. In this experiment you will investigate several compounds containing oxygen to discover whether they decompose in this way at temperatures which you can reach with a bunsen burner.

The most important type of reaction of oxygen is the one in which it enters into direct combination with other substances. This is called an oxidation reaction. One important kind of oxidation is combustion or ordinary burning, another is the breathing process of animals. When elements burn in oxygen oxides of the elements are formed. In general these oxides are either acidic or basic, that is, they form either acids or bases if they react with water. Oxides of metals form bases while oxides of non-metals form acids. This gives us an important method of classifying elements as either metallic or non-metallic. Litmus is a substance which has a red color in the presence of an acid but a blue color in the presence of a base. Paper which has been impregnated with a solution of litmus is called litmus paper. It may be used to determine whether a given solution is acidic or basic. Pure water, which is neutral, will not affect the color of litmus paper. In this experiment you will dissolve several oxides in water and test the solutions with litmus paper to learn whether they form acids or bases.

The rate at which a reaction takes place is influenced profoundly by several factors, such as temperature, presence of a catalyst and concentration of the reactants. In general, reactions proceed much more rapidly at higher than at lower temperatures. Some substances have the ability to speed a given reaction by being present in the reacting mixture, although they are not consumed finally in the reaction. These are called catalysts. Effects of catalysts are quite specific in most cases. Many of them are extremely important in modern industrial operations, such as catalytic cracking of gasoline, synthesis of ammonia and the manufacture of solid fats. Most reactions are speeded by increasing the concentration of the substances which react. You will study several of these factors in this experiment.

Exp. 6 - 2

Directions:

a) Laboratory sources of oxygen.

Oxygen causes substances to burn more brilliantly than does air. If a glowing splint on being thrust into a gas bursts into flame, the presence of oxygen is indicated. If it is extinguished a substance which does not support combustion is known to be present.

Place about half a gram of potassium chlorate in a dry test tube (pyrex or hard glass) and heat it at first gently until it is just melted to a clear liquid. Insert a glowing splint in the mouth of the test tube. (1) Is oxygen being evolved from the molten potassium chlorate? Now heat the tube more strongly until bubbles of gas are evolved. Test again with a glowing splint. (2) Is oxygen given off now? Heat separate portions of barium peroxide (dioxide), ferric oxide, mercuric oxide, silicon dioxide (sand) and lead dioxide. (Use a soft glass test tube in heating lead dioxide.) Test each for oxygen with a glowing splint. Not all of these substances will evolve oxygen at the temperature which you can reach with a bunsen burner. (3) Which did yield oxygen? (4) Which of these substances would you say are the most stable toward heat? Which are the least stable? Arrange your data and answers in tabular form.

b) Action of a catalyst.

Place about half a gram of potassium chlorate in a clean, dry pyrex test tube and support this in a vertical position with a burette clamp on a ring stand. It is advisable to remove any rubber tips from the clamp. Heat until the potassium chlorate has just melted. Is oxygen evolved? Remove the burner from under the test tube. With a spatula or horn spoon add a very small pinch of powdered manganese dioxide to the molten potassium chlorate, taking care to keep your face well away from the top of the test tube. (5) Describe what you observe. Is oxygen evolved? Contrast with the behavior of the molten potassium chlorate. Give an explanation of these phenomena.

c) Preparation and properties of oxygen.

Mix thoroughly 8 grams of potassium chlorate and 4 grams of powdered manganese dioxide. Test this mixture by heating rapidly a small portion of it in a test tube. If it decomposes quietly, without spattering, it may be used safely for the preparation of a quantity of oxygen. Place the mixture, after it has been shown to be satisfactory by the above test, in a large pyrex test tube and clamp this to the ring stand in a nearly horizontal position, placing the clamp close to the mouth of the test tube. Fit the test tube with a one hole rubber stopper and a delivery tube which reaches into the pneumatic trough as shown in the sketch in the laboratory. Heat the test tube gently and evenly, generating a slow, steady stream of gas. Collect 4 bottles of oxygen by the downward displacement of water, taking care to remove the delivery tube from the water when you interrupt heating.

Exp. 6 - 3

(6) Why? The bottles of oxygen may be preserved out of contact with air either by keeping them under water, mouth downward, or by placing a glass plate over the mouth of the bottle and removing it from the pneumatic trough.

Heat a glass stirring rod and dip it into a small quantity of powdered sulfur on a watch glass. (Perform under the hood!)
Some sulfur should adhere to the hot rod. Now heat the rod until the adhering sulfur burns. Note the color of the flame and the odor of the gas produced. Lower the burning sulfur into a bottle of oxygen which you have prepared. (7) Describe its behavior, comparing and contrasting the burning of sulfur in air and in pure oxygen. Is a different odor produced when it burns in oxygen? Does this experiment indicate a way in which reaction rates may be influenced? Explain your answer. Add about 5 ml of distilled water to the bottle after the sulfur has stopped burning and close the mouth of the bottle with the palm of your hand and shake.
(8) Do you observe anything? Test the water in this bottle with both red and blue litmus paper. Is there any evidence that the oxide of sulfur has reacted with the water? Test distilled water with red and blue litmus paper. (9) List as many properties of the oxide of sulfur as you have been able to determine from this experiment.

Place a small amount of charcoal in a deflagrating spoon and heat it in a flame until it glows or begins to burn. Lower this glowing charcoal into a bottle of oxygen and note the result. Add a little distilled water, close the bottle with the palm of your hand and shake. Test the liquid in the bottle with both colors of litmus paper. Add a few drops of lime water (calcium hydroxide solution). The white precipitate which forms is calcium carbonate and its appearance is a test for carbon dioxide.

Heat a small amount of "steel wool", held in a pair of tongs in the bunsen flame and plunge it while red hot into a bottle of oxygen. In order for this experiment to succeed it is necessary to heat the steel wool right next to the bottle of oxygen so that it will not cool too much before it can be inserted in the bottle of oxygen. Describe what you observe.

The combustion of other substances, such as a splint, lighted candle etc. may be tested in the remaining bottle of oxygen.

d) Reactions of oxides with water.

How did you observe solutions of sulfur dioxide and of carbon dioxide to react toward litmus paper? (10) Are these oxides acidic or basic? Classify sulfur and carbon as metals or non-metals on this basis.

Hold a 1 cm length of magnesium ribbon in your pair of tongs and heat it in the bunsen flame until it ignites. (Caution: Shield your eyes!) Hold the burning magnesium over a small, dry evaporating dish so as to catch the ash which forms. (11) Describe the flame of burning magnesium and the appearance of the ash which is formed. This ash is the oxide of magnesium. Add

Experiment 7

THE GAS LAWS

**Object:** To find the way the volume of a gas varies with the temperature.

**Discussion:**

It has been demonstrated to you that the volume of a gas decreases as the pressure increases. You will recognize this as a statement of Boyle's law. This typical behavior of gases is explained by the kinetic theory. It will be remembered that according to this theory gases consist of molecules which are a considerable distance apart as compared with those in liquids. When the pressure on the gas is increased, the molecules are pushed closer together and thus occupy a smaller volume.

You are familiar with a number of instances in which gases expand when heated. The weak automobile tire which explodes after running at a high speed in the summer is a familiar example. This behavior is also typical of all gases and is likewise explained by the kinetic theory. The molecules of the gas in addition to being widely separated are moving rapidly and the rate at which they move is proportional to their temperature. Thus, if the temperature increases, these molecules move faster, striking the walls of the vessel harder and oftener, all of which results in their exerting a greater force. If the force which these molecules exert is acting against a fixed external pressure, such as atmospheric pressure, the result is an increase in volume. There is a definite relationship between the change in volume with a given change in temperature when the temperature is expressed on the absolute scale. The statement of Charles' law expresses this relationship, and it is this law that you will confirm in this experiment.

**Directions:**

Half fill your 600 cc beaker with water and heat on a ring which is attached to a ring stand. While this is heating prepare a piece of glass tubing about 5 cm long and firepolished at both ends. When the glass is cool, insert it in a #4, 1-hole rubber stopper until the end is level with the bottom of the stopper. Place the stopper firmly in the neck of a dry, 125 ml erlenmeyer flask. Make a note of the position of the bottom of the stopper in the neck of the flask. This can usually be done with the aid of the green seal found on the flask. Later it will be necessary for you to replace this stopper at this same position. The stopper must be put in place firmly as you will lift the flask by the glass tube later. Set the flask in the beaker of hot water, forcing it

Exp. 7 -2

down until the water level is near the bottom of the stopper. It
should be held in this position by clamping the protruding glass
tube with a burette clamp. Continue to heat the water until it
has boiled for at least five minutes. Adjust the flame so that
the water boils gently. What is the temperature of the water?
Check it with a thermometer. What is the temperature of the gas
in the flask? (1) Why is it necessary to heat the water for at
least five minutes?

Fill your pneumatic trough with water and have it near the
remainder of the apparatus. Crasp the neck of the flask firmly
with your crucible tongs held in your left hand. Hold the flask
at its present level in the water while you remove the burette
clamps. Catch the glass tube with the thumb and second finger;
place the index finger firmly over the opening of the tube. Lift
the flask out of the water and invert it in your trough making
certain that your finger is held tightly on the tube until it is
under water. The water must be kept boiling until the flask has
been removed. Keep the flask in this position and pour water from
the trough over it for three or four minutes. (2) Why did the
water enter the flask? How can you tell when the flask has been
held in the water long enough? Adjust the position of the flask
so that the water level inside and outside are the same. Place
your finger over the opening in the glass tube and set the flask
upright on your desk. Remove your finger, dry the neck of the
flask and the rubber stopper with a towel and remove the stopper,
being careful to let the water which is held in the glass tube
enter the flask. Measure the volume of the water with your gradu-
ated cylinder. (3) What does this volume represent? Record it on
the data sheet. Determine the temperature of the water in the
trough and record this in its proper place.

Fill the flask completely with water and replace the rubber
stopper at the same position it previously occupied. Notice that
the glass tube is filled. This is important. Wipe the outside of
the flask as before. Carefully remove the stopper without losing
water and determine the volume. Record as the initial volume of
the gas. Use a *dry* flask and repeat the experiment as a check.
Fill in all the information requested on the data sheet.

Questions and Problems

I. What is the barometric pressure today? Why is it not necessary
to consider pressure in this experiment?

II. What would be the effect of starting this experiment with a wet flask? Explain why the error would be so large.

Exp. 7 - 3

III. A volume of gas which is the same as your desk number expressed in ml is collected over water at 757 mm pressure. What volume would this gas occupy if it were dry and measured at 765 mm? Temperature is assumed constant at 20°C.

IV. What is the volume of dry air measured under standard conditions which is in your flask just after it has been removed from the trough? Show clearly all steps involved in solving this problem. Why is it not necessary to subtract the vapor pressure of water in the experiment which you performed?

V. Explain the following in terms of the kinetic theory:

(a) The volume of a gas in a vessel with a movable top increases as the temperature increases.

(b) The pressure exerted by a gas held in a container of fixed volume increases as the temperature.

(c) Why does not the pressure increase in (a)?

VI. 250 ml of dry hydrogen was collected in three bottles by the displacement of water, alcohol and mercury under the same atmospheric pressure. Arrange in order of the final volume, indicating which volume is the greatest. Explain why the final volumes would vary.

Experiment 8

HYDROGEN

**Object:** To prepare some hydrogen, study certain of its physical and chemical properties and to observe diffusion of gases.

**Discussion:**

Hydrogen is widespread in many naturally occurring compounds. About 11% of water is hydrogen and appreciable amounts of hydrogen are contained in many forms of vegetable and animal matter. Most of us are familiar with the fact that petroleum contains compounds of hydrogen. Relatively only small amounts of hydrogen are found uncombined in nature. This may be attributed to the reactivity of hydrogen. Because of its importance, it is appropriate to learn how it may be prepared and to study some of its outstanding properties.

All solutions of acids contained combined hydrogen and under proper circumstances this may be liberated to yield a quantity of free hydrogen. Certain metals* will react with solutions of acids and give a salt of the metal plus free hydrogen. This is often a convenient method for the production of small amounts of hydrogen in the laboratory. Another source of hydrogen is water, from which it may be prepared in pure condition by electrolysis. This is both a commercial and a laboratory method.

Molecules of hydrogen are lighter than molecules of any other substance. According to the Kinetic Molecular Theory, if several gases are compared at the same temperature, the lightest molecules will on the average travel fastest. Hence molecules of hydrogen should diffuse more rapidly than molecules of all other gases. Certain chemical properties of hydrogen make it an appropriate gas for one to use in a study of gaseous diffusion.

**Directions:**

Set up a gas generating apparatus as shown in the sketch. Place about 25 g of mossy zinc in the flask. Just cover the zinc with water and insert the stopper containing the thistle tube and delivery tube in the flask. Be sure that the thistle tube is below the level of the water in the flask. (1) Why? Make a reduction tube shown in the sketch, using a pyrex test tube. The tube for carrying hydrogen into the reduction tube should reach nearly to the bottom of the test tube. Place about 2 g of black copper oxide (cupric oxide) in the bottom of the dry test tube and assemble the reduction tube.

* Only those metals above hydrogen in the E.M.F. series (a list of metals according to their activity) liberate hydrogen. The significance of this fact will be emphasized later.

Exp. 8 -2

Have the delivery tube from the generator reach into the pneumatic trough in such a way that gas may be collected by the downward displacement of water. Add to the generator about 50 ml of dilute sulfuric acid by pouring it down the thistle tube.
(2) Do you note a particularly vigorous reaction? Add about 1 ml of copper sulfate solution through the thistle tube.
(3) Does gas evolution increase? What might copper sulfate be termed?

Fig. 1.
Generator

Tip for flame.

Fig. 2.
Reduction Tube

Discard the first two bottles of gas. (4) Why? Allow one bottle to fill a little more than half full of air by tipping the mouth above the water level. Displace the remaining water with hydrogen. Collect two additional bottles of pure hydrogen and cover the three bottles with glass plates. Place the mouth downward. (5) Why? Then disconnect the delivery tube in the pneumatic trough from the generator and substitute the reduction tube which you made previously. Pinch the rubber connection, thereby shutting off the flow of hydrogen until the pressure causes the liquid in the generator to back up the thistle tube. Now add about 25 ml of dilute sulfuric acid through the thistle tube taking care to prevent air being aspirated into the generator. Allow the action to continue causing air to be swept out of the reduction tube. Wrap a towel firmly around your generator at this time.

Carry the closed bottle of hydrogen which contains a little over half air to a burner under the hood and quickly remove the glass plate and present the mouth of the bottle to the flame. (6) What do you observe? Is a mixture of air and hydrogen explosive? Light a wooden splint and thrust it well into an inverted bottle of hydrogen. Then withdraw the splint slowly. Be sure to observe everything. (7) Does hydrogen support combustion? Does it burn? Place an inverted bottle of air tightly mouth to mouth over the remaining bottle of hydrogen, allowing them to

Exp. 8 - 3

stand in this position for two or three minutes. Test the in-
flammability of the contents of each bottle. (8) What do you
observe? How do you explain this behavior on the basis of the
Kinetic Molecular Theory?

By this time air should be displaced from your reduction
tube. Ignite a burner which is placed about three feet from the
generator, after making sure that the generator is wrapped in a
towel. Invert a small, dry test tube over the jet of the re-
duction tube and hold it there for about two minutes until hydro-
gen has displaced the air in it. Remove the test tube and close
its mouth with your thumb immediately, carry it to the bunsen
burner, remove your thumb and quickly plunge the mouth of the
still inverted test tube into the bunsen flame. If the hydrogen
is impure it will explode with a "bark". If this happens collect
another test tube full and repeat this procedure until hydrogen
is collected which burns quietly with a practically colorless
flame. Use this flame to ignite the hydrogen which issues from
the tip of the reduction tube. DO NOT IGNITE THE HYDROGEN JET
DIRECTLY, but in the above manner. Hold an inverted, small, dry
beaker over the burning jet of hydrogen. (9) What do you observe?
Explain. Hydrogen when pure burns with a practically colorless
flame. The yellow color of the flame in this experiment is
caused by sodium from the glass tip. Your Instructor will demon-
strate a hydrogen flame to you.

When the hydrogen jet is burning steadily, begin to heat the
contents of the reduction tube carefully with your bunsen burner,
gradually increasing the temperature. (10) Describe everything
which you observe as the heating progresses. There are at least
three things which you will note. After you are convinced of
what happens, allow the reduction tube to cool while a stream of
hydrogen passes through it. Then remove the test tube and ex-
amine the substances in it. (11) Explain in some detail your
observations and the reactions which you infer from them.

Questions and Problems:

I. If 3.83 l of hydrogen, measured at N.T.P., were generated what would the volume of this hydrogen be if collected over water at a pressure of 751 mm and at 25°C.?

II. How would you pour a bottle of hydrogen into another con-
tainer? Explain your answer.

III. What is the difference between a substance which is combus-
tible and one which supports combustion? Is there a very sharp
distinction? Explain your answer.

IV. Explain why a hydrogen jet will burn in air, but a mixture of air and hydrogen explodes.

V. 6.54 g of zinc treated with enough acid yield 2.24 l of hydrogen at N.T.P. If you wish the number of liters of hydrogen corresponding to your desk number, measured over water at 20°C. and 735 mm, how much zinc should you use?

Experiment 9

EQUIVALENT WEIGHT

**Object:** To become familiar with the concept of equivalent weight and to determine the equivalent weights of two metals.

**Discussion:**

Throughout your study of chemistry, you will find that there are many important quantitative relationships. It is of great importance that one understands these and knows how to apply them, for chemistry is an exact science. Often it is required to know what weight of a given element will react with completely or displace completely a certain weight of another substance. That is to say, one may wish to know how much of some material is equivalent to a fixed quantity of something else. For example, it might be desired to know what weight of magnesium would be exactly equivalent to 1 g of hydrogen, such information being needed to calculate the weight of magnesium which would displace one gram of hydrogen from an acid.

An equivalent weight of any substance is defined as the weight in grams of that substance which is equivalent to one atomic weight (1.008 g) of hydrogen. The equivalent weight is, accordingly, the weight of a substance which just will displace completely or react completely with 1.008 g of hydrogen. It is self-evident that one equivalent weight of any substance is equivalent exactly to one equivalent weight of any other substance. At N.T.P. 1.008 g of hydrogen (i.e. one atomic weight) occupies 11.2 l, which is, therefore, termed the equivalent volume of hydrogen. It is clear that if one determines the weight of a particular substance which is required to liberate 11.2 l of hydrogen at N.T.P. he will have found the equivalent weight of that substance.

It should be noted clearly that it makes no difference to the definition or determination of the equivalent weight whether the particular substance in question actually will react with hydrogen or not. A series of substances and reactions can be found through which it will be possible to relate a weight of hydrogen to an equivalent weight of any other substance.

In this experiment you will determine the equivalent weights of two metals, zinc and aluminum, by finding the weights of these metals required to displace from hydrochloric acid 11.2 l of hydrogen at N.T.P.

Exp. 9 -2

Directions:

This is a quantitative experiment and will require careful,
accurate work in order to obtain satisfactory results. Record on
the data sheet all of your measurements as soon as they are made.

Fit a stopper and delivery tube into one of your 250 ml bottle
to make a gas generator so that gas generated in the bottle can be
collected in the pneumatic trough by downward displacement of water
The delivery tube must contain a section of rubber tubing so that
the apparatus will be flexible. Place about 40 ml of dilute hydro-
chloric acid in the bottle.

Weigh to the nearest milligram a piece of zinc which shall be
within the limits 0.5 - 0.6 g. Tie a piece of silk thread securely
about this piece of metal and suspend in through the stopper in the
generating bottle above the surface of the hydrochloric acid.
Great care must be taken to prevent the acid from coming in contact
with the metal at this time. Stopper the bottle tightly. Adjust
the delivery tube so that the gas generated in the generator will
be collected in a bottle in the pneumatic trough.

Now lay the generator bottle on its side in such a way that
the hydrochloric acid comes in contact with the piece of zinc, but
exercise the utmost care to prevent any liquid from entering the
delivery tube. Hydrogen will be formed by the action of the acid
on the metal and displaced gas will be collected in the pneumatic
trough. Of course, if liquid escapes from the generator bottle,
the true volume of hydrogen generated will not be measured. When
the action has ceased, all of the metal should be dissolved. Remove
the delivery tube from the bottle of gas which you have collected,
level the water within this bottle with that in the trough,
place a glass plate across the mouth of the bottle (while the water
levels are the same) and remove the bottle, taking care to prevent
any change in the amount of water therein. Measure the volume of
water contained in the bottle. Then determine the volume of the
bottle. Measure the temperature of the water in the pneumatic
trough and assume that this is the temperature of the gas which
you have collected. Vapor pressures will be found in a Handbook
or tabulated on the board.

Repeat the experiment using this time sheet aluminum, but ob-
serving the following changes. The piece of aluminum should be
0.14 - 0.17 g. Instead of using dilute hydrochloric acid, prepare
40 ml of a special concentration by pouring carefully 20 ml of con-
centrated hydrochloric acid into 20 ml of dilute hydrochloric acid.
This is to be used to generate the hydrogen.

Calculate the equivalent weights of zinc and aluminum from
the data which you have obtained. What fraction of the atomic
weight is the equivalent weight in the two cases?

Exp. 3 - 3

Questions and Problems:

I. 13.5 g of aluminum just reacts with 53.2 g of chlorine. On the basis of your determination of the equivalent weight of aluminum, calculate the equivalent weight of chlorine. How many grams of hydrogen would react with 53.2 g of chlorine?

II. Calculate the volume of gas you would obtain over water at 20°C and 755 mm pressure if you used excess hydrochloric acid and the number of grams of aluminum corresponding to your desk number.

III. The valence of an element is defined as a number which represents the number of atoms of hydrogen which will react with or be displaced by one atom of the element in question. Compare this with the definition of equivalent weight given in this experiment. On the basis of these two definitions calculate the valence of zinc and of aluminum. (Recall that one atomic weight of an element contains the same number of atoms as one atomic weight of any other element.)

IV. A certain non-metal has an atomic weight of 30.98. It is found that 2.582 g of this element will react completely with 0.252 g of hydrogen. Calculate its equivalent weight. What is its valence? Calculate how many grams of oxygen would react with one atomic weight of this element. Write a formula for this compound.

Exp. 9 -4

Data Sheet.

<table><thead><tr><th></th><th>Zinc</th><th>Aluminum</th></tr></thead><tbody><tr><td>1. Weight of Metal used.</td><td>g</td><td>g</td></tr><tr><td>2. Volume of bottle</td><td>ml</td><td>ml</td></tr><tr><td>3. Volume of water in bottle at end.</td><td>ml</td><td>ml</td></tr><tr><td>4. Volume of gas collected</td><td>ml</td><td>ml</td></tr><tr><td>5. Temperature of water</td><td>°C</td><td>°C</td></tr><tr><td>°K</td><td></td><td>°K</td></tr><tr><td>6. Atmospheric pressure (Instructor)</td><td>mm</td><td>mm</td></tr><tr><td>7. Vapor pressure of water at the temperature of expt. (Handbook)</td><td>mm</td><td>mm</td></tr><tr><td>8. Partial pressure of hydrogen</td><td>mm</td><td>mm</td></tr><tr><td>9. Corrected volume at N.T.P.</td><td>ml</td><td>ml</td></tr><tr><td>10. Fraction of equivalent volume (" weight)</td><td></td><td></td></tr><tr><td>11. Equivalent weight of metal (calc'd)</td><td></td><td></td></tr><tr><td>12. Atomic weight of metal (Text)</td><td></td><td></td></tr><tr><td>13. Fraction of atomic weight equal to equivalent weight.</td><td></td><td></td></tr></tbody></table>

Show your calculations in this space.

Exp. 3 -4

Data Sheet.

<table><thead><tr><th></th><th>Zinc</th><th>Aluminum</th></tr></thead><tbody><tr><td>1. Weight of Metal used.</td><td>.5385 g</td><td>.18 g</td></tr><tr><td>2. Volume of bottle</td><td>269 ml</td><td>269 ml</td></tr><tr><td>3. Volume of water in bottle at end.</td><td>50 ml</td><td>8 ml</td></tr><tr><td>4. Volume of gas collected</td><td>219 ml</td><td>261 ml</td></tr><tr><td>5. Temperature of water</td><td>13° C</td><td>13° C</td></tr><tr><td></td><td>296 °K</td><td>296 °K</td></tr><tr><td>6. Atmospheric pressure (Instructor)</td><td>752 mm</td><td>752 mm</td></tr><tr><td>7. Vapor pressure of water at the temperature of expt. (Handbook)</td><td>11.2 mm</td><td>11.2 mm</td></tr><tr><td>8. Partial pressure of hydrogen</td><td>740.8 mm</td><td>740.8 mm</td></tr><tr><td>9. Corrected volume at N.T.P.</td><td>ml</td><td>ml</td></tr><tr><td>10. Fraction of equivalent volume ( " " weight)</td><td></td><td></td></tr><tr><td>11. Equivalent weight of metal (calc'd)</td><td></td><td></td></tr><tr><td>12. Atomic weight of metal (Text)</td><td></td><td></td></tr><tr><td>13. Fraction of atomic weight equal to equivalent weight.</td><td></td><td></td></tr></tbody></table>

Experiment 10

WATER

**Object:** To observe some of the characteristics of water and to find how this substance may be purified.

**Discussion:**

The importance of water is far too great and its uses too numerous to discuss completely in the limited space available. Almost any of us could list several hundred uses of this compound, many of which are so essential that life could not continue long without it. Our body itself consists of from 60 to 65% water, and we must constantly take in additional amounts of this substance to maintain this percentage. Did you ever stop to realize that when you are looking at the average individual, weight 150 lbs., that you are looking at about 90 lbs. or over ten gallons of water?

In addition to the use of water to maintain life, the chemist finds many uses for this substance in the laboratory. Many times it is used as a reactant, a substance actually entering into a chemical reaction, but more often it is used as a solvent and a medium in which reactions between other substances take place. Both the layman and the chemist then are very much interested in the purity of the water which they are using. However, when these two speak of "pure water," they have entirely different requirements in mind.

The study of the purification of water for drinking purposes is a very important topic, but unfortunately it is not easily studied in the chemistry laboratory. During this experiment we shall consider the purification of water chiefly from the chemist's point of view.

The most important cases in which water is used as a reactant have been studied under hydrogen or oxygen. It will be recalled that this compound was used as a source of both of these elements. About the only other important use of water as a reactant is its reaction with the oxides of the elements.

**Directions:**

Test some tap water for impurities by evaporating 2 or 3 ml on a clean watch glass over a water bath. (A water bath is made by setting the watch glass over your 250 ml beaker about half full of boiling water.) While this tap water is evaporating set up the apparatus according to the sketch on the board. (1) What evidence do you find of impurities in the tap water? Charge the flask with about 50 ml of tap water and boil gently until you have distilled 5 or 6 ml of water. (2) How should the purity of this water compare with that of the tap water? How could you test it? Carry out the test. (3) Describe the results. Add about a gram of solid copper sulphate to the remaining water in your distilling flask and continue the distillation as before. (4) What evidence do you now

Exp. 10 -2

have that the impurity (copper sulphate) is not being distilled over. Such impurities are known as non-volatile ones. Empty the distilling flask, wash out with water but do not completely disassemble the apparatus.

Obtain about 5 ml of water in each of two clean test tubes. Add about one drop of ammonium hydroxide to one of them and then add two drops of phenolphthalein to each tube. (5) How can you detect the presence of ammonium hydroxide in water? Fill your distilling flask with 50 ml of fresh water and add 2 drops of ammonium hydroxide to it. Distill the water, catching the distillate in clean test tubes. Change the tubes so that not more than 5 ml of the distilled water is caught in one tube. Place the tubes with the distillate in your test tube rack in the order in which they were obtained. Add a drop of phenolphthalein to each of the tubes and continue the process until a definite change is observed. (6) Explain the results which you have obtained. Ammonium hydroxide is an example of a volatile impurity. (7) Can water which contains such impurities be purified? Explain.

Add a few drops of soap solution to 5 ml of distilled water in a test tube and shake. Repeat using hard water from the shelf near the liquid reagents. (8) How do the results compare? This behavior of the water is due to the presence of certain impurities which react with the soap. The chemistry of these changes cannot be fully understood by you at this time but will be taken up in detail later in the course. While this type of impurity does no harm as far as drinking water is concerned, it is very important from the standpoint of household and industrial uses of water.

Review the reaction of the oxides of elements with water. These tests were performed in Experiment 6. (9) Describe the results and make a generalization based on them.

Questions:

I. Contrast the requirements for pure water from the standpoint of drinking purposes and laboratory purposes. List three essential steps in the process of preparing drinking water and give the purpose of each.

II. Describe a chemical test which might enable you to distinguish a metallic element from a nonmetallic one. Why cannot this test be applied to all elements?

III. Explain the difficulty of cooking potatoes on a high mountain. How can this difficulty be overcome?

IV. A bottle of dry gas which consisted of 1/4 nitrogen and 3/4 oxygen was placed mouth down over a pan of water the temperature of which was 27° C. The bottle was placed in the water so the levels inside and outside were the same, and it was found that the atmospheric pressure was 754 mm. Find the pressure exerted by each of the gases in the bottle.

Exp. 11 -3

that a small opening is left. Heat gently with a low flame for about five minutes then increase the flame and heat strongly for about fifteen minutes. At the end of this time remove the flame, place the cover on the crucible so that it is completely covered and allow to cool. Avoid the opening of the crucible as much as possible; carry it to the balance and weigh. Repeat the heating for about five minutes, cool and weigh again. (13) What is the object of this second heating and weighing? Repeat the experiment as a check if directed to do so by your Instructor. When you have completed the data sheet you will know the percent of water in your hydrate and the number of molecules of water associated with each molecule of the anhydrous salt.

Questions:

I. Explain what would happen if you put some anhydrous copper sulphate and hydrated sodium carbonate under the same bell jar.

II. Several crystals of hydrated sodium carbonate were placed in a closed jar and an equal amount was placed in an open dish. Those in the jar effloresced for only a short time while those in the dish continued to do so. Explain fully.

III. How much water do you buy when purchasing five pounds of washing soda (sodium carbonate decahydrate, $Na_2CO_3·10H_2O$)?

IV. Give three important characteristics of a chemical change. Which of these three did you use in this experiment to determine whether or not a chemical change had occurred?

V. Give the important characteristics of a compound as contrasted with those of a mixture. Which of these were used to determine whether a hydrate is a mixture or a compound?

VI. What do the results of the class as a whole indicate regarding the composition of hydrates? Explain.

Experiment 12

LAW OF MULTIPLE PROPORTIONS

**Object:** To examine the relationships between composition of two compounds which are composed of the same elements.

**Discussion:**

One of the first laws studied in this course and verified in the laboratory was the law of definite proportions. It will be recalled that this law dealt with the composition of compounds and stated that this composition does not vary. In stating the above law, it is necessary to specify that two elements combine in a definite proportion when forming the same compound. Often the same two elements combine in different proportions, but, when they do so, different compounds are formed. Thus there is a red solid and a black solid, obviously different compounds, both containing only the elements copper and oxygen.

An examination of the quantitative relationships of pairs of such compounds shows that the law of definite proportions still applies to each individual compound. Exactly the same proportions of copper and oxygen are always found in the red compound. The proportions in the black compound are just as definite, though different from those in the red. A further consideration of these two compounds shows that if you started with two samples of copper weighing 10 g each and made the red oxide from one and the black from the other, the weights of oxygen involved would differ, and an interesting relationship would exist between them.

Unfortunately, the two compounds referred to above cannot be used in the laboratory for this study. Instead, you will use the compounds potassium chlorate and potassium perchlorate. These two differ from the oxides of copper in that they consist of three elements. A similarity can be seen, however, if the compounds use here are considered to consist of the two substances, potassium chloride and oxygen just as the oxides of copper consist of the two substances, copper and oxygen. You should not get the impression that either potassium chlorate or potassium perchlorate can be made by combining potassium chloride and oxygen directly.

The information desired from this experiment is: What is the relationship between the amounts of oxygen associated with 1 g of potassium chloride to form each of the two compounds studied?

**Directions:**

Weigh your *clean*, *dry* ignition tube carefully on the chemical balance and record the weight on the data sheet. Add an estimated 1 g or 1 1/2 g of potassium chlorate (S.R.) to the tube and weigh again. Clamp the tube at the extreme top end with a burette clamp so that it makes about a 60° angle with the table top. Start heating the contents of the tube very gently until it melts. Continue the heating until a clear liquid remains in the tube and there is no evidence of a gas being evolved. Remember this is a quantitative experiment and if the tube is heated too rapidly,

Exp. 12 -2

the oxygen will carry off liquid and solid particles. A serious error would thus be introduced. Allow the tube to cool to room temperature and weigh.

Wash your tube, dry thoroughly in the flame and repeat the above experiment substituting potassium perchlorate (S.R.) for potassium chlorate.

Questions:

I. An experiment similar to this one with two different compounds gave the following results:

<table><thead><tr><th></th><th>potassium chloride</th><th>oxygen</th></tr></thead><tbody><tr><td>Compound A</td><td>1.742</td><td>0.374</td></tr><tr><td>Compound B</td><td>1.542</td><td>0.662</td></tr></tbody></table>

Calculate items 7, 8, 9, 10 of the data sheet from this information

II. How does the atomic theory explain the law of multiple proportions?

III. If a liter of oxygen weighs 1.429 g at N.T.P., what would be the volume of oxygen, measured at existing conditions, obtained in each experiment which you performed today? All information needed must be listed and the method of your solution shown clearly.