Walter Scott Hendrixon.

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dilute sulphuric acid. Turn the
small tube upward and determine
when the air is wholly expelled by
lighting the hydrogen with a test

tube as in Ex. 12. Now turn the tube downward, place it in a test tube
surrounded with water and heat the copper oxide, beginning at the end
of the column nearer the flask. Continue till the copper oxide is all
reduced, that is, all changed to red copper. Examine the liquid that
has collected in the test tube. What is it?

(b) For the glass tube and copper oxide in (a) substitute the
iron tube and contents from (13b) or use the iron tube and about 5
grams of ferric oxide. First fill the tube completely with hydrogen as
proved by lighting it with a test tube, heat the oxide strongly, protect-
ing the stoppers with the wet cotton as stated in 13b.

With a moderate flow of hydrogen continue the reduction 10 to 15
minutes. What is the source of the water in the test tube? When
cold and if ferric oxide was used test some of the contents of the tube
in con. HC1 warming. What gas is evolved? What evidence have you
that the reaction in 13b is reversed? What is a reversible reaction?
What governs the direction of this one? How could you produce a
state of equilibrium in it? What is the meaning of equilibrium?


15. Solubility of Solids: Some substances are very soluble in
cold, still more soluble in hot water. Shake 3 grams of powdered so-
dium nitrate or ammonium chloride with 5c.c. of water in a test tube
till all is dissolved. Is there any change in temperature of the water?
Now add 3 grams more and shake. Shake and determine whether it
all dissolves. If not, heat till dissolved. What occurs on cooling to
room temperature?

Some substances are very soluble in hot water, slightly in cold
water, and such may easily be purified by crystallization from hot so-
lution. By heating dissolve 4 grams of ordinary potassium chlorate
in lOc.c. distilled water. Filter boiling hot if there is any turbidity or


solid matter suspended in the liquid. When the solution is cold filter
off the crystals, setting filtrate aside and wash crystals with a little
cold water, dissolve a portion of them in distilled water and add sil-
ver nitrate to the solution and to filtrate. Compare the amount of
precipitate obtained with silver nitrate in the two tubes. It is due to
a chloride, commonly found in ordinary potassium chlorate.

Some substances are little more soluble in hot water than in cold
wau-r. Shake 5 grams of common salt with lOc.c. of water in a test
tube. Note the amount of salt undissolved. Boil the solution for a few
moments to saturate the water with salt. Note again, the amount of
salt. Filter the solution boiling hot and cool to room temperature.
Does much salt crystallize out? Why not? Compare the solubility 01
salt in hot and cold water with that of potassium chlorate. A few sub-
stances are even less soluble in hot water than in cold. Examples are
calcium sulfate (gypsum) and slaked lime.

16. Solubility of Liquids in Liquids, and the Separation of Solutes
between two non-miseible Solvents: Measure accurately in a cylinder
about 50c.c. of water reading at the lower surface of the meniscus. Now
carefully pipet into the cylinder 25c.c. of common alcohol. Mix thor-
oughly and read the volume. Is the volume now the sum of those of
the alcohol and water? Do water and alcohol dissolve each other

In a test tube place about lOc.c. of water and about 2c.c. of chloro-
form and shake. Let stand and note whether they are mixed. Add a
crystal of iodine and shake for some time and let stand. Which li-
quid takes up most of the iodine?

As above try to mix lOc.c. of \vater and 2c.c. of carbon disulfide.
Add about oc.c. of bromine water and shake. Which liquid takes up
most of the bromine?

To another tube add water and a little benzene, and a small crystal
of potassium permanganate and shake. Which liquid takes up the

Determine whether alcohol and chloroform, benzene and chloro-
form, carbon xdisulfide and benzene will mutually dissolve each other.

17 Solubility of Gases: Gases also vary widely in their solu-
bility in water. In text-book see solubility of oxygen, nitrogen and hy-
drogen sulfide which do not act chemically with the water, dissolve
only moderately and obey Henry's law, (which see). Others which
generally act chemically with water dissolve in very large amounts.

In a test tube or flask heat hydrant water and observe gas bub-
bles given off before the water begins to boil. Try the same with dis-
tilled water. Does it contain dissolved gases. Why does water when
drawn from the hydrant sometimes look milky and quickly become
clear on standing?


Note carefully the odor from a solution of ammonia due to am-
monia given off. Pour about Ic.c. of the solution into lOc.c. of water
in a dish and test with turmeric paper. Boil till the water is half
evaporated and test again with the paper. What is the effect of heat
on the solubility of gases. To a degree the following shows one of
many exceptions to the general rule:

Boil lOc.c. of dilute hydrochloric acid till three-fourths cf it has
evaporated. Boil the same volume of concentrated hydrochloric acid
till one half has evaporated. Add to each sample of boiled acid wheji
cold, the same amount of granulated zinc. Do they seem to act on zinc
at about the same rate? A more accurate determination would show
that they have the same concentration. By boiling long enough
samples of dilute and concentrated hydrochloric acids one arrives at
the same result; namely acids of concentration 20.2 per cent.

18. Chemical Action of Water: Refer back to 13 , for the action
of water on sodium and on iron, and to 9 for its action on sodium di-

Upon a piece of quick lime drop water slowly until the water is no
longer absorbed and the piece of lime looks wet. Place it in a dish
and note what occurs. This is the familiar "slaking" of lime.

10. Water in Combination: In a test tube heat a small amount of
copper sulfate, observing water given off and changed appearance. In
the same way try borax, alum, sodium phosphate. These are "hyd-
rates" and the water they contain is called "water of crystallization."
For contrast try potassium sulfate and common salt.

In a weighed porcelain crucible with lid weigh accurately about 4
g. of barium chloride, place the crucible on a triangle over a burner
and heat ten minutes. When cold weigh the crucible and contents
again and from the first weight of the barium chloride and the loss on
heating, calculate the per cent, of water.

Efflorescence: On a glass plate or watch-glass place a few crys-
tals of sodic sulphate, expose to air till the next laboratory period.

Deliquescence: Expose to air in dishes small pieces of calcium
chloride and caustic potash till the next laboratory period and record

20. Electrolytic Decomposition of Water: Support the U tube
shown in fig. 14 with a clamp and use the current described in 53.
Connect the side-arms of the tube with the trough of water by means
of short delivery tubes. Fill the U tube nearly to the side arms with
a 5th normal solution of sodium sulfate already made up and add a few
drops of litmus solution to each side of U tube. Start the current and
at the same time place over the ends of the delivery tubes two test
tubes the same size and full of water, thus collecting the hydrogen and
oxygen set free. Continue till the smaller volume equals about 5c.c.



and compare volumes. Note color of the solution at each electrode and
state cause. Name electrodes. From which comes the 0, and the H?

21. Purification of Water by Distillation: To a small volume of
hydrant water add a few drops of barium chloride solution. The white
precipitate shows carbonate and sulfate radicals present. Now add
dilute hydrochloric acid. The precipitate remaining shows sulfate
radical. To another portion add a few drops of dilute nitric acid and
a few of silver nitrate. A white precipitate shows chloride present.

Set up apparatus shown in fig. 6, fill flask one-third full of water,
distill enough to clean the tubes and reject it. Distill about lOc.c. of
the water and test it for sulfate and chloride radicals. None should
be obtained, all ordinary mineral matter being left in the boiler in dis-

Try the action of ammonia and so-
dium hydroxide on turmeric paper
and dilute siilfurio aoi^ on blue lit-
mus papei To one-fourth- of
flask full of water add about 5c.c. of
ammonia, distill and determine wheth-
er any ammonia distills over. Clean
apparatus carefully and try sodium
hydroxide, adding a few c.c. to one-
fourth flask of water. Distill from
a fresh portion of water containing
a little sulfuric acid and a little
potassium permanganate. Do they go
Fig. 6. over?


22. Preparation: Place a beaker with lOOc.c. of water in water,
preferably ice cold. Stir in little by little 5 grams of sodium dioxide,
NaO 2 the operation taking about 5 minutes. Even then some of the
peroxide will be decomposed giving off oxygen. Now add gradually
in the same way dilute hydrochloric acid till a drop of the liquid taken
out with the stirring rod just turns blue litmus paper red. This gives
a solution of hydrogen dioxide, but it contains also common salt. It
may be used where hydrogen peroxide is required below save where it
is used with silver and lead. There the commercial peroxide should be

Rub in a mortar with a little water about half a gram of starch,
transfer to a dish or beaker, add about lOOc.c. of water and a few crys-
tals of potassium iodide and heat to boiling. This is known as "starch-
iodide solution" and filter paper wet with it is called "starch-iodide pa-
per." Each is frequently required. To a part of the solution add a


few drops of the solution of hydrogen peroxide which sets free iodine
and this colors the starch blue. To see the color by transmitted light
dilute with much water. This is used as a test for iodine or hydrogen
peroxide or starch. That is, two being known to be present the pres-
ence or absence of the third can be determined by the test.

23. Oxidation with Hydrogen Peroxide: Moisten a strip of filter
paper with very dilute lead acetate and expose it to hydrogen sulfide
a little of which may be made in a test tube as in 68. The black sub-
stance is lead sulfide, PbS. Pour upon the paper a few drops -of hy-
drogen dioxide, which will change the lead sulfide to white lead sul-
fate, PbSO*. Why is this called oxidation?

To a solution of silver nitrate add NaOH and then carefully add
just enough of a solution of ammonium hydroxide to dissolve the pre-
cipitate at first formed. Now add commercial hydrogen peroxide. The
gray precipitate is finely divided metallic silver, and the escaping gas
is oxygen which may be tested by trying in the tube a glowing splint-
er. This action appears to be one of reduction, but it is not primarily.
Probably a higher oxide of silver is formed and at once decomposes
into silver and oxygen. The next two cases are of the same sort.

To about a gram of manganese dioxide in a test tube add hydrogen
dioxide and test the gas with a glowing splinter. Repeat using a con-
centrated solution or a few crystals of potassium permanganate in-
stead of the manganese dioxide.


24. Preparation of Chlorine: Chlorine is dangerous if breathed.
Experiments with it should be conducted in hoods. If carried out on the
students' desks only half the usual number should work in the room at
one time and windows should be freely opened. They should stand to
windward of the apparatus when collecting the gas. When through
collecting at once place the delivery tube in a test tube nearly full of
concentrated sodium hydroxide, and remove the source of heat. Be-
fore beginning the experiments students should read them through,
the right plan in any case, provide everything necessary and thus re-
duce the time to the minimum. Make a solution of starch-iodide as di-
directed in 22. ,

Chlorine is made commercially by the electrolysis of fused com-
mon salt or a solution of common salt. All other methods depend up-
on the oxidation of the hydrogen of hydrochloric acid. Several
methods may be illustrated on a very small scale.



Upon a few crystals of potas-
sium permanganate in a test tube
pour about Ic.c. of con. hydrochlo-
ric acid and pour a little of the
heavy, greenish yellow gas into a
little of the starch solution in an-
other tube. In the same way treat
a little potassium dichromate heat
and test for chlorine. Try in the
same way lead dioxide and con.
hydrochloric acid. Use a very
little potassium chlorate and con-
centrated hydrochlpric acid, HC1;
also, about 1 cc. con. HC1 and a
few drops of con. nitric acid.

Fig. 7.

25. Set up the apparatus as in fig. 7, with the water bath one-
fourth full of water. A copper can is best, but a tin can or a beaker will
serve for a water bath. See that the thistle tube reaches nearly to the
bottom of the flask. In the flask place 25 grams of manganese diox-
ide MnO 2 , and add 40c.c. of concentrated HC1 diluted with lOc.C. of
water. Heat the flask and collect five jars or bottles of chlorine. Tlie
green color will show when they are full. They should be well filled,
but a large excess should not over-flow into the room. One jar should
have the bottom wet with con. sulfuric acid to dry the gas for use in

When through collecting place the delivery tube in NaOH in a
test tube. Let the flask cool, removing water bath, and proceed with the
next experiment.

26. Properties of Chlorine: To show bleaching action and the
need of water, suspend in the jar of dry chlorine strips of colored
cotton cloth and litmus paper. After a few moments note any fading
of the colors, then moisten the strips and suspend again in the jar,
and note effect. Refer to a text-book for information on hypochlorous
acid and bleaching with chlorine. What does the bleaching?

Into a jar of chlorine pour successively with shaking, small vol-
umes of solutions of litmus, cochineal, much diluted ink.

From a piece of antimony scrape with a knife a very little of the
metal letting it fall into a jar of chlorine. Using forceps or tongs heat
to redness a strip of copper foil and lower it into the same jar of chlo-
rine. If "Dutch metal" is used it need not be heated.


Burn the laboratory gas at the end of a glass tube and lower the
small flame into a jar of chlorine. What is the black substance? When
the green color has disappeared blow breath over mouth of jar, which
will give a fog, consisting of droplets of water containing HC1. Will
carbon burn in Cl? Try an ignited piece of charcoal.

On a deflagration spoon lower a small bit of white phosphorus in-
to a jar of Cl and avoid inhaling the Cl or fumes of PC1 3 . It should
soon melt, then take fire.

Now take the flask and all jars to the sink best under hood and
standing well back fill them with water. Wash well any remaining
MnOi. and place it in a vessel provided for that purpose.


27. Preparation : Read the experiment through and have all the
necessary materials at hand so that once begun the experiment may
be carried through rapidly. ^

All soluble chlorides give hydrochloric acid when treated with
concentrated sulfuric acid. For many reasons sodium chloride, com-
mon salt, is to be preferred.

Set up the apparatus as in fig. 7, omitting the water bath. In the
flask place 20 grams of sodium chloride. Dilute 35 grams (20c.c.) of
concentrated sulfuric acid by pouring it slowly with stirring into 8c.c,
of water in a beaker. Pour slowly into the flask and let stand a few
moments till acid and salt are in contact throughout then apply a low
heat, best using a burner with crown top. Collect the gas in dry jars
or bottles in the same way as chlorine. Abundant fumes will indicate
when the jars are full. After collecting two jars, fill a dry bottle
which has been fitted with a stopper and a short piece of tubing one
end flush with the large end of stopper and the other drawn out and
cut off so as to leave a small orifice. It should reach at least to the
middle of the bottle's length. Insert stopper when full and place the
bottle mouth downward in water. Press down the bottle and pour cold
water over it to start the absorption. A fountain will result.

Into a test-tube of water insert the delivery- tube so that it reaches
a very little below the surface of the water. Note ready absorption of
the gas. Is the solution lighter or heavier than water? Lower the
tube as necessary to absorb all the gas. Why does the water become
warm? Continue till the water is nearly saturated, then try the action
of small portions of the solution on a little zinc, marble, sodium car-
bonate. Compare its action with that of the dilute HC1 from the shelf.
Test its action on blue litmus paper. To a little of the solution add a
few drops of silver nitrate then a few drops of dilute nitric acid. The
precipitate is silver chloride. This is a test for hydrochloric acid or
a chloride. Repeat using instead of HC1 a few drops of a solution of


salt or of any other chloride. With a known chloride could it be used
as a test for silver? Save flask and its contents for 29.

28. Acid, Base, Salt, Neutralization: Wet the inside of a jar with
a solution of concentrated ammonia and pour out excess of liquid.
Cover jar with a glass plate. Place^ it over a jar of the HC1 gas from
the last experiment, bringing the jars mouth to mouth and remove
plate. The ammonia unites with acid forming ammonium chloride, a

Standing well away drop one or two small bits of sodium into a
little distilled water in a bottle. When the action is over test the wa-
ter with turmeric paper. The action of sodium on water gives sodium
hydroxide, a base, and what gas? (see 13).

Test sodkim hydroxide from shelf bottle with the papers. Pour
about 5c.c. of the alkali into a dish and add dilute hydrochloric acid till
the solution turns litmus paj>er red, testing by taking out a drop of the
solution with a stirring rod and touching the paper ; never place pa-
pers in the solution or dip them into it. Now add a few drops of NaOH
or dilute HC1 as may be necessary with stirring, till the solution
changes neither turmeric nor blue litmus paper. It is now neutral.
Evaporate to dryness, taste the residue. Place upon it a few drops of
con. sulfuric acid and note odor of the gas. What was the solid resi-

29. The preparation of HC1 from salt and sulfuric acid gives a
good example of a reversible reaction:

NaCl+H 2 SO 4 = (reversibly) HNaSO 4 +HCl.

HC1 is a stronger acid than H 2 SO4. In the cold or in a small closed
space even on heating the reaction would not complete itself to the
right. But, the HC1 is easily volatile and heating drives it out of
solution and away so that it cannot react with HNaSO 4 to the left. Sul-
furic acid is volatile only at very high temperature. Try restoring
HC1 thus: Dissolve with the least volume of water and heating, the
contents of the flask used in preparing HC1. Cool some of the solution
of HNaSO4 in a test tube and add a few c.c. of con. HC1 from shelf,
which will precipitate sodium chloride. Why is sulfuric acid used to
prepare easily volatile acids from their salts? Why are the reactions
completed when heat is used?

30. Preparation of Bleaching Powder, Potassium Hypochlorite
and Potassium Chlorate: Study these subjects in a text-book and read
the experiment through.

Arrange the apparatus as shown in fig. 8. The wash bottle con-
tains diluted sulfuric acid, made by adding 40c.c. of the con. acid with
stirring, to lOc.c. of water. Its purpose is to remove the greater part
of the HC1 and water from the Cl. The horizontal test tube contains
about 2 grams of dry slaked lime spread evenly throughout i r s




length. The upright test tube should contain a solution of 4 grams of
potassium hydroxide dissolved in 12c.c. of water. It must be cooled
and kept cold by surrounding with water as shown.

Charge the flask with MnO 2 and diluted HC1 as in 25. Maintain a
moderate stream of chlorine for about 15 minutes or until the larger

portion of it seems to pass
through the solution in the second
test tube. Now preserve half the
contents of this tube as potassium
hypochlorite, and heat the re-

j, |>| mainder to boiling and without

( cooling continue to pass chlorine

into it for about five minutes or
till a drop taken out with a stir-
ring rod does not feel soapy to the
fingers. Now heat to boiling and
filter the solution. Cool by stand-
ing the tube in water, when crys-
tals of KC1O 3 form. When quite
cold filter them off and wash with
a little cold water. Test the ni-
trate with silver nitrate. What
was formed * besides potassium
chlorate? Dissolve the chlorate
Fig. 8. by passing about 3 cc. of boil-

ing water through the filter several times, cool, let crystallize, pour
off the water from the crystals, dissolve them in water and add
silver nitrate. Compare the first and second precipitate formed by sil-
ver nitrate. Pure chlorate would give no precipitate.

Into a little of the hypochlorite solution place a drop of diluted
ink, into another .portion a bit of colored cloth. After noting any
bleaching add a little dilute acid to each solution and note effect. To
a third portion add a few drops of strong solution of ammonium hyd-
roxide. What gas is given off?

Try the action of dilute acid on a little of the dry bleaching pow-
der from the horizontal tube. Dissolve a portion of it so far as pos-
sible, filter and try the action of the filtrate on ink, colored cloth, lit-
mus paper, before and after adding dilute acid.

To show the instability and oxidizing power of potassium chlorate
mix on paper 5 grams of the salt and 5 grams of powdered sugar,
but do not grind them together in a mortar. Place the mixture on an
iron plate, take out a drop of con. H 2 SO 4 with the stirring rod, stand
well away and drop the acid upon the mixture.




31. Preparation of Bromine: Place 5 grams of manganese diox-
ide and 5 grams of sodium bromide on paper, hold the neck of the re-
tort pointing slightly upward and slide the mixture in at the tubulus
without letting it fall into the neck of the retort. Support the retort
as in fig. 9, add 50c.c. of dilute sulfuric acid through a funnel. The
flask should contain about lOOc.c. of water and the tip of the retort
neck should dip under its surface. Apply heat and continue till all
the bromine has distilled over. Move the flask away so as to bring
the neck of the retort above the surface and then remove the burner.
Is bromine soluble in water? Is it heavier than water? Set aside the

flask containing bromine for lat-
er use.

Give two reasons why we do
not prepare Br in the same way
as Cl; that is, by the action of
hydrobromic acid on manganese

May Cl be prepared by the
same method used for Br; that
is, by use of NaCl. MnO 2 and di-
luted H 2 SO 4 ? Try it in a small
way in a test tube, but strength-
en the acid by adding about Ic.c.
con. H,S0 4 to 2c.c. of dilute acid.
Write equations for preparation
of eaph.

~~ Why is it not best to use con.

Fig. 9. sulfuric acid in preparing either

Br or Cl by this method?

31. By the same method prepare a little iodine, using about O.r>
gram of sodium iodide, a gram of MnO 2 and 5c.c. of a mixture of di-
lute and con. suifuric acid. Note sublimed iodine near the mouth
of the test tube. Compare the reaction with that in the preparation
of Cl and Br by the same method.

Heat a few crystals of iodine in a dry test tube. Does it form
a liquid before subliming? What is sublimation? Note crystals
higher up on walls of the tube.

33. Hydriodic Acid and Comparison of HCI, HHr, HI: Treat
very small amounts of sodium chloride, sodium bromide and sodium
iodide with a few dropi of con. sulfuric acid. How did you prepare
HCI? Could HBr be prepared in the same way? What is the black
substance set free by the action of the acid on sodium iodide?


When salt and sulfuric acid are heated together hydrochloric

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Online LibraryWalter Scott HendrixonA laboratory manual of general chemistry → online text (page 2 of 8)