Walter Scott Hendrixon.

A laboratory manual of general chemistry online

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is a mixture of equal volumes of two solutions, one containing 34.639
grams of copper sulfate to 500 c.c. and the other 173 grams sodium pot-
assium tartrate and 60 grams sodium hydroxide in 500 c.c. The cop-
per in 1 c.c. of the mixture is reduced by 0.005 gram of glucose.

To about 10 cc. Fehling's solution add enough glucose to precipi-
tate all the copper on boiling as cuprous oxide, Cu 2 O, which will settle,
leaving the clear liquid above. In a fresh mixture try a little milk
sugar. Try cane sugar.

Now dissolve about a gram of cane sugar in 100 c.c. of water, add a
few drops of con. HC1, heat to 70 degrees and let cool. Boil 10 c.c. of
this solution of invert sugar with 20 c.c. of Fehling's solution.

100. Esters, Acetic Ester: Acids act upon alcohols forming esters
and water much as they act upon bases forming salts


and water. Acetic acid and alcohol act thus: CH 3 CO 2 H+C2H 5 OH=
CH 3 CO 2 C 2 H S +H 2 O.

Use the apparatus shown in Fig. 6. With the delivery tube in the
test tube pour into test tube 20 c.c. water and mark the level with
gummed paper. Remove the water. Place in the flask 15 c.c. glacial
acetic acid, 15 c.c. alcohol and 5 c.c. con. sulfuric acid. Warm the flask
with a very small flame so that only a few drops shall distill over in
10 minutes. Now increase the heat and distill to the 20 clc. mark. In-
stead of measuring the distillate a two hole stopper and thermometer
may be used, and the distillation continued till it reads 95. Its bulb
should be in the vapor above the liquid.

The ester contains acetic acid and alcohol. To remove most of
these shake it persistently with twice its volume of water in a small
flask covered with the thumb. On standing the ester rises and forms
the top layer. Separate it from the water by using a separatory fun-
nel, or the funnel with pinch cock and exit tube as shown in Fig. 10.
Measure the purified ester.

In a test tube shake about 2 c.c. of your ester with an excess of
NaOH till it all disappears; that is, till it is all "saponified" forming
sodium acetate and alcohol.

Note odor of acetic ester. Upon this is based a good test for ace-
tic acid or its salts. To a little solid sodium acetate in a test tube add
a few drops of alcohol and a few of con. sulfuric acid and note odor.
Another test is this: To a neutral solution of an acetate add 3 drops
of ferric chloride. A red color should be obtained, and on boiling a
brown flocculent precipitate.

101. Saponification, Soap Making: In a dish or beaker place
five grams of olive oil or lard or tallow, add 15 c.c. of alcohol and 5 c.c.
of sodium hydroxide solution of concentration one to four. Do not use
a pipette for NaOH. With occasional stirring heat on a water or steam
bath, for at least half an hour, but better, an hour. All the alcohol
and most of the water should be evaporated. Examine the soap when
cold. If the saponification is complete a small portion should entire-
ly dissolve in water, save a slight milkiness. Dissolve more of the
soap and determine whether it will form a lather with soft water. Try
blowing soap bubbles with the solution. Determine whether the feel-
ing is that given by ordinary soap.

Filter some of the soap solution. Pour a little into distilled water,
into hydrant water, into solutions of calcium and magnesium sulfates.
Shake each tube and determine in which ones a persistent lather is
formed. To those in which it is not formed add more soap solution and
shake again, until a lather persists. Calcium and magnesium com-
pounds make water "hard." Why do such waters require more soap?


To a concentrated solution of soap add dilute hydrochloric acid in
excess. The precipitate consists of a mixture of organic acids of
which stearic acid, Cn H 33 COOH, is representative. The oils and fars
mentioned above consist of mixtures of what are known as esters of
these acids. The most common perhaps is stearin, (Ci-HusCOOhCsHs.
Write the equation for the saponification of stearin with sodium hy-
droxide; also, write equation for action of soap on solution of calcium


102. Silicic Acid, To 10 c.c. of a solution of sodium silicate in a
small beaker add con. HC1 drop by drop stirring for a few moments
after each addition. Silicic acid will separate as a jelly. Collect the
silicic acid on a filter, wash with water, transfer to a crucible, dry
over flame and finally ignite. When cold try to dissolve the residue
in water, in hydrochloric acid.

103. Boric Acid: Place 15 g. borax in 50 c.c. water in a beaker.
Heat till dissolved and add 15 c.c. con. HC1 with stirring and allow the
liquid to cool thoroughly. Filter off and examine the boric acid. Dis-
solve a little of the substance in alcohol in a dish, set fire to alcohol
and observe color of the flame. Moisten a strip of turmeric paper with
the solution from the boric acid. Observe color, then treat the paper
with a solution of sodium carbonate, let dry, best on steam bath, and
note color.


106. The substance given for analysis may be in solution. If not,
dissolve in water, heating if needed. If not soluble in hot water use di-
lute nitric acid, and heat. If gas is given off one or more acids of
group (1) are present, and the special tests may be applied at once.

Group (1) : To a small portion of the solution add dilute HNO 3 in
excess if not already added, and heat. If gas is set free try to identify
it by odor, color anft tests. Only CO 2 , S0 2 , H 2 S, NO-, are likely to occur.
Make special tests for the following acids as given in the sections in-
dicated by the numbers: H 2 CO 3 (90), H 2 S0 3 (71), H 2 S 2 O 3 (73), H 2 S (68),
HN0 2 (54b).

Group (2) : To another portion of the solution add dilute nitric
acid in excess if not already present and if members of group (1) are
present boil to expel any gas. To a portion of the boiled solution add
barium chloride in excess. A white precipitate shows H 2 SO 4 (71).
Filter it off and make nitrate alkaline with ammonia. A yellow pre-
cipitate indicates H 2 CrO4 (140). A white precipitate indicates one or
more of the following: H 3 PO.,-(80), H 3 AsO 4 . (83), HsAs0 3 (83),
H 3 BO 3 (103), HF (38), H 2 C 2 O 4 (see following): To another portion of


the boiled solution add ammonia in excess, then CaCl 2 . An immediate
precipitate may mean any of the acids of this group save sulfuric. Add
an excess of acetic acid. If the precipitate is insoluble, oxalic acid or
hydrofluoric or both are present. To a third portion of the boiled solu-
tion add manganese dioxide, which will give C0 2 if oxalic acid is pres-
ent. It is well in any case to make the special test for boric acid, since
its barium and calcium salts are distinctly soluble.

Group (3) To the original solution add silver nitrate. If a pre-
cipitate is formed add an excess of dilute nitric acid. If the precipi-
tate all dissolves this indicates some acid in previous groups. If in-
soluble, one or more of the following are present: HC1 (37), HBr (37),
HI (37), H 4 Fe(CN) 6 (142), H s Fe(CN) 6 (142). A pure white precipitate
shows only HC1. If the precipitate is colored proceed to the special
tests for the others. For halogen acids see also 127b.

Group (4) Nitric and acetic acids must always be tested for if
the substance is soluble in water. For HNO 3 see (56b) , and for acetic,


107. Through a spectroscope suitably adjusted by the instructor
examine a flame colored by a sodium salt, and locate the sodium line on
the scale. Locate in the same way the red line given by potassium,
and that given by lithium. Draw a millimeter scale in your note book
and place the lines in their correct positions. If the spectroscope has
no scale, estimate as closely as possible the relative positions of the
three lines. Note that any chemical will show some sodium. Its line
is no evidence that sodium is present in considerable quantity.

108. What is the action of sodium on water? Standing well back
drop a bit of potassium into water in a bottle or beaker. How does the
result differ from that given by Na?

To reduce the violence of the action of the alkali metals on water
their alloys with some other metal are often used. Drop into water
"hydrone" which is an alloy of sodium and lead. Also try sodium
amalgam, an alloy of sodium and mercury. When the action of the
latter is over put the mercury in a dish provided for the purpose.
Never put mercury into sinks. Why? Test the solution in each case
with red litmus or turmeric paper. Test soapy feeling of the solution
between thumb and finger.

109. Preparation of Sodium Hydroxide: Dissolve 5 grams of so-
dium carbonate in 75 c.c. of water in a porcelain dish and reserve 5 c.c.
Heat the remainder nearly to boiling, and stir in a little at a time 5
grams of slaked lime. Boil gently for several minutes, replacing the
water evaporated, let settle and filter off the solution. If it destroys
the filter, let cool, add a little water and use another filter. To prove


the reaction, Na 2 CO 3 +Ca(OH) 2 =CaCO 3 +2NaOH, proceed as follows:
Wash the insoluble residue in the dish twice by filling nearly full of
water, stirring, letting the substance settle and pouring off the water.
Now put a little water on the insoluble substance in dish, and place
about the same amounts of lime and water in another dish and pour
upon each about 10 c.c. of dilute HC1. Compare the amounts of carbon
dioxide given off. Now treat the 5 c.c. reserved solution of sodium car-
bonate and the same volume of the filtrate with the same volumes of
HC1 and compare the gas given off. You started with a soluble car-
bonate and a very sparingly soluble hydroxide, "slaked lime." What
did you obtain by the reaction?

110. Comparative Tests of Crude and Pure Sodium Hydroxides:
Dissolve 1-2 grams of crude and pure NaOH, each in 10 c.c. of distilled
water. Make each acid throughout with dilute nitric acid and
warm. If either gives off gas test for carbon dioxide as in 90. Test
small portions of each acidified solution for Cl with silver nitrate, and
other portions for SO 4 with BaCl 2 . Test other portions for iron by add-
ing to each a few drops of potassium ferrocyanide and potassium ferri-
cyanide, which will give a blue color if iron is present.

HI. Purification of Common Salt: If practicable use the crude
rock salt of the feed store, but ordinary salt will do. Dissolve about 5
grams in 25 c.c. of water. To one third add dilute HC1 and test for
SO 4 , and preserve the tube and contents. To another portion add a
little ammonium chloride, make alkaline with ammonium hydroxide
and add ammonium carbonate. The precipitate is calcium carbonate.
Filter it off and to the filtrate add sodium phosphate which will give a
precipitate on standing if magnesium is present. See 80. Save this
tube and contents.

Make a fully saturated solution of the crude salt, first reducing it
to powder if rock salt is used, and shaking a long time with water.
Filter if necessary and to the solution add an equal volume of pure
con. HC1. This will precipitate most of the salt. If gasous HC1 were
added more salt would precipitate. Filter off the salt and wash with
three small portions of water. Dissolve some of this salt in water,
test for SO 4 , calcium, magnesium, as above and compare with the
results of these tests with crude salt.

112. An Acid Salt, Acid Potassium Tartrate, Cream of Tartar:
This is one of the few slightly soluble salts of potassium. As the term
is commonly used there are no insoluble salts of Na or K. Dissolve
about 4 grams of pure, dry potassium carbonate in 25 c.c. of water and
10 grams of tartaric acid in 50 c.c. of water, measuring the latter so-
lution. To the carbonate solution add one drop of methyl orange, then
add cautiously the acid solution till a faint red color is obtained. This
gives the soluble normal salt, K 2 C4H 4 O 6 . Note what volume of the acid


solution was added, then add as much more, stir and let stand some
time. The precipitate .is the acid tartrate, cream of tartar. Which
salt is least soluble?

Filter off the cream of tartar, let it dry on the filter and weigh
with a balanced filter. Calculate the weight of acid sodium carbonate
to mix with it to make one kind of baking powder which when wet acts
thus: HKC4H 4 O 6 +HNaCO3=NaKC i H 4 O c +H 2 O+CO 2 . What makes the
bread rise when baking power is used? Try your baking powder in a
tube with water and test for COz.

113. Potassium Nitrate from Sodium Nitrate: Dissolve in 50 c.c.
H 2 O, 25 g. sodium nitrate and the calculated amount of potassium chlo-
ride required in the reaction

KCl+NaN0 3 =KNO 3 +NaCl.

Evaporate to one-half the volume, let the separated salt settle, and de-
cant the clear, hot liquid into a beaker, press solid with spatula and
let liquid run into beaker. This solution should turn solid when cold.
Transfer this to a filter and let drain. Press solid between folds of
filter paper, dissolve in least hot water and let crystallize. Transfer
crystals to filter and let dry.

114. Qualitative Analysis: Determine first the presence of the
ammonium radical and the alkali metals in known substances, and
then their presence or absence in "unknowns," using the scheme as
given 152, Group V. In the initial work fresh substances will be ex-
amined, and of course what is said of filtrates from previous groups
and their preparation for analysis does not apply.


115. Upon a piece of quick lime drop water as Icng as it is taken
up, place it in a dish and observe from time to time. After it has be-
come powdery place some of the "slaked" lime in a jar of water, shake
it thoroughly and let settle. Filter a portion of the nearly clear lime
water placing the funnel in a flask to protect from the carbon dioxide
of the air. Test the clear solution with turmeric paper. To portions
of the lime water add one-third of their volumes of ferric chloride and
magnesium chloride respectively. The precipitates are hydroxides
of the metals. Test the alkalinity of barium hydroxide and its action
on solutions of the same metals.

116. Pass carbon dioxide into 25 c.c. of clear lime water until the
precipitate ot calcium carbonate, CaCOs, dissolves, forming the acid
carbonate, H.CaCCOaK The latter is the chief substance that gives
"temporary hard water." To a small portion of the solution add clear
lime water. How may lime soften temporary hard water? Will lime
also remove magnesium from water? Boil another portion of the so-


lution which will reverse to the left the reaction which occurred with
CO 2 in excess.

CaC0 3 +H 2 0+C0 2 - =(reversibly)H 2 Ca(C0 3 ) 2 .

117. Pure Calcium Chloride: Pour off the liquid from flask in
which carbon dioxide was made, make it alkaline with milk of lime ob-
tained by shaking slaked lime and water in a jar and pouring at once.
Let settle and filter. How does this remove iron and magnesium?
Pass into the filtrate C0 2 and boil. Why? Filter again if necessary
and evaporate to dryness in a porcelain dish and heat. Expose a little
of the calcium chloride to air till next period and observe again. Is it
deliquescent? Dissolve the remainder in about 10 times it weight of
water, and use as calcium chloride solution.

118 Comparative Solubilities of Salts of Ba, Sr, Ca, Mg: Carbon-
ates: To solutions of Ba, Sr, Ca and Mg chlorides from shelf add equal
volumes of water then to each about one-fifth of its volume of ammo-
nium chloride, and finally to each, ammonium carbonate. What com-
pounds are precipitated? How could magnesium be separated from
the other three metals? Add to its solution sodium phosphate, and
see the tests for phosphoric acid (80) arid Mg (128).

119. Chromates: To solutions as under carbonates, but omitting
Mg, add a little acetic acid then a solution of pure potassium chro-
rhate, or dichromate. How could barium be separated from strontium
and calcium?

120. Suit* ates : To solutions of Ba, Sr, and Ca chlorides add a so-
lution of magnesium sulfate and let stand for a few moments. How
does this prove that Ba, Sr, and Ca sulfates are less soluble that mag-
nesium sulfate? To fresh solutions of Ba and Sr chlorides add a solu-
tion of calcium sulfate. How do the results show that Ba and Sr sul-
fates are less soluble than calcium sulfate? To a solution of barium
chloride add a solution of strontium sulfate and let stand a short time.
How do we know that barium sulfate is less soluble than strontium
sulfate? Arrange the sulfates in the order of their increasing solu-
bility in water.

121. Calcium Oxalate: To a solution of calcium chloride add an'
excess of ammonium carbonate, let stand a few moments and filter. To
the filtrate and also to a solution of calcium sulfate add a solution of
ammonium oxalate and let stand half an hour. State how you know
that calcium oxalate is less soluble than calcium carbonate or calcium
sulfate. Devise a scheme for the separation of Ba, Ca, Sr and Mg.

Calcium sulfate is present in many natural waters and causes
"permanent hardness" in the sense that it is not precipitated by boil-
ing, though on concentration by evaporation, it forms a hard deposit
on the boiler. To calcium sulfate solution add a solution of sodium
carbonate. What two chemicals may be added to soften water showing


both temporary and permanent hardness?

Test the water of the laboratory for both sorts of hardness.

122. Make analyses of a solution containing barium, strontium,
calcium and magnesium according to the directions of Group IV; also,
analyses of unknown solutions or solids which may contain these me-
tals, and solutions or solids which may contain also metals of Group
V and the ammonium radical.


123. Weigh accurately a small dish or beaker, add 2.65 grams to
the weights and exactly balance with pure sodium carbonate. Dis-
solve the carbonate in water, transfer with rinsings of dish to a half
liter flask, using a funnel and taking care that none of the solution is
lost. Do not even lose some by removal on the stirring rod. Make up
the volume to the mark with water, and mix by placing the thumb
over mouth of flask and inverting several times. This is a decinormal
solution of Na 2 CO 3 . Why? Transfer this solution to a bottle or larger
flask and wash the graduated flask.

Measure in a small cylinder 7.5 c.c. pure con. HC1 and dilute it 10
700 c.c. and mix well. Fill a buret with the acid solution to well above
the zero, fill the tip of buret and bring the surface to or below the
zero. Read accurately at the lowest point of the meniscus. With a pipet
(see 4) place 20 or 25 c.c. according to capacity of the pipet, of the so-
dium carbonate solution in a dish or beaker, add to it 2 drops of methyl
orange. For comparison it is well to place beside it about 50 c.c. of
water and add- to it two drops of the indicator. From the buret run in
the acid as rapidly as you wish to about 15 c.c. then a few drops at a time
with stirring till the solution becomes faintly red as shown by compar-
ison with the indicator in water. Make another titration, which should
agree within a few tenths with the first.

Divide the volume of the alkali by that of the acid which gives the
decinormal concentration factor of the acid. Measure 500 c.c. of the
acid and multiply by the factor, which will give the total volume to
which the 500 must be made up with water to become decinormal.
Why? Transfer the 500 c.c. to a larger vessel, add the necessary water,
mix and titrate again against the alkali. They should neutralize each
other volume for volume.

With these two standard deci-normal solutions the concentration
of any other acid or alkali may be determined, or solutions of desired
concentration may be made by the same method used in making the
acid solution.

124. In titrating weak acids methyl orange cannot be used ; see
next section. One must use a very weakly acid or neutral indica-
tor, such as litmus, phenoltalein, congo red. But these are affected by


the carbonic acid from the carbonate and one must titrate the solution
at the boiling point or use NaOH free from carbonate. Find the per
cent of acid in vinegar by running it from a buret into a measured
volume of the carbonate boiling and containing a few drops of phenol-
talein, or use cold NaOH supplied by the instructor, instead of Na 2 COs.
Methyl orange and phenolphthalein are weak, complicated organic
acids. The following will give a correct, general idea of their action:
Any strong acid sets free any weak acid from its salt. Let NaR be
such a salt where R is the negative radical. Then,

In the case of phenolphthalein R" gives the red color, while HR is
colorless. Phenolphthalein is such a weak acid that even carbonic
acid is strong enough to act in the same way as HC1 in the equation
and form HR. Hence when the stage in the titration represented by
HNaCOs is passed, and H 2 CO a is formed this acts as a relatively strong
acid, forms HR and thus destroys the color. On the other hand methyl
orange is a stronger acid than carbonic and the yellow color of its
negative ion in Na + +R~ persists till there is a slight excess of HC1,
when red HR is formed.

125. The facts stated in 124 are well illustrated by titrating in the
same solution both normal and bicarbonate by using different indica-
tors, as fallows:

Dissolve about 0.2 gram of normal sodium carbonate in 25 c.c. of
water, without heating, add a few drops of phenoltalein, fill a buret
with your deci-normal acid, read and run into the carbonate solution
drop by drop near the end, till the pink color just disappears. The
carbonate is now all HNaCOs. Add a few drops of methyl orange, read
the buret and run in the acid till the solution takes on a faint tinge of
red, using the indicator in water for comparison as in 123. Read again
and compare the volumes required to change the normal to the bicar-
bonate, and to neutralize the latter. ' Why are they approximately
equal ? Why are these carbonates alkaline, having no ion OH ? See 6 1.



126. (a) Dissolve 5 grams copper chloride in 10 c.c. of con. HC1
and 10 c.c. water; or (b) prepare a solution of the copper chloride by
dissolving 5 grams copper sulfate and 2.5 grams common salt by heat-
ing with 10 c.c. of water in a test tube. When dissolved set the tube
in cold water for several minutes. Pour off the solution from the sep-
arated sodium sulfate. Why is this formed? Now add to the solution
10 cc. con. HC1, let stand a few moments and filter off the salt. Why
is salt thus formed? Whether (a) or (b) boil very gently the solu-


tion with 5 grams finely divided copper in a small flask, replacing
evaporated liquid if necessary with dil. HC1. Continue till colorless
or till a few drops poured into water gives no blue color. You now
have HCuCl 2 . Pour a part into water which decomposes it giving in-
soluble CuCl. Does this dissolve in ammonia? Compare with silver
chloride. Put the ammonia solution in a white dish, stir and note that
the cuprous ion is rapidly oxidized to cupric ion as shown by increas-
ing blue color.

To a portion of the liquid from flask add an excess of NaOH. What
is the red compound? See Fehling's solution (99). Heat a little of the
white precipitate with Br water and give result in terms of ion formed.
Expose some of the white CuCl on the filter to sunlight and note result
after an hour.

There is no cupric but only cuprous iodide. To a few c.c. copper
sulfate solution add a little potassium iodide solution, and test the so-
lution for free iodine with starch paper. Add Ic.c. carbon disulfide
shake and let settle. Note color of the carbon disulfide.

Determine whether dilute sulfuric, hydrochloric and nitric acids
give H by their action on copper and cadmium and explain results.

(c) Tests for Copper: To half a test tube of water add a drop of
solution of any cupric salt, and make alkaline with ammonia. The
blue color is due to the complex ion Cu(NH 3 )4 ++ . To a like dilute solu-
tion of copper add a few drops of acid and a little potassium ferrocy-
anide. Dilute equal parts of each solution till the colors are just vis-
ible and state which test is the more delicate.

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Online LibraryWalter Scott HendrixonA laboratory manual of general chemistry → online text (page 6 of 8)