D. A. Davenporl and Aflf N. Sbba Purdue University Lafayette, Indiana
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Hypodermic Syringes in Quantitative Elementary Chemistry Experiments
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2.
General chemistry experiments
In the first1 of these two papers we described ways in which hypodermic syringes can be used in simple, quantitative experiments on the gas laws. We now wish to demonstrate the versatility of syringe techniques in other types of experiments involving measurement of gaseous volume.2 Gas Densities
Equal volumes of various gases are introduced into the syringe which is then capped (a broken needlemount sealed with solder is convenient) and weighed. The known density of air is used to calculate the true weight of the syringe. As may be seen from Table 1, a volume of 25 ml is a particularly good choice under ordinary conditions. Table 1. Gas Densities Weieht in millierams of 25 ml of eas at 26'C and 750 mm Ha 2 . 3 He 4 . 1 Nx 28.0
Air 29.3 OB 32.3 COX44.5
Changes in buoyancy must be allowed for (especially with hydrogen and helium) if the sample volume differs appreciably from 25 ml. In such case a simple proportional correction of the observed weight will not suffice. Indeed, the fact that an empty syringe weighs the same as one filled with 25 ml of air comes as a salutary reminder to many students. Analysis of Air
A small volume of alkaline pyrogallol is introduced into a syringe containing a measured volume of air. The syringe is capped and shaken until no further gas absorbtion occurs. Typical results fall in the range 20.5-22.0 per cent oxygen by volume. The use of a 100-ml syringe yields an effectivelecture demonstration ss the per cent by volume of nitrogen remaining may be read directly. Analysis of Gaseous Mixtures
By proper choice of absorbent solutions a variety of simple gaseous analyses'may be carried out. Suitable 'DAVENPORT, D. A,, J. CHEM.EDUC.,39, 252(1962). R e ferred to as Part I hereafter. %Forsatisfactory results with glass syringes, lubrication must he such as to allow free response to slight pressure differentials. A qualitative test may he made hy ascertaining that the piston responds readily to gentle blowing and sucking. We haw found that plastic syringes w o ~ kbadly, or not at all, althoughFrigerio and Trotter, (J. CHEM.EDUC.,39, 594(1962)) have recently de~cribedtheir use in a simple gas law experiment using a Bourdon gauge.
unknowns may be made up by blowing air, oxygen, nitrogen, ammonia, alkylamines, hydrogen chloride, sulfur dioxide, etc., into large balloons which are then sealed with a short length of glass tubing closed a t one end with a serum-stopple. The small pressure differential (-30 mm) between the gas inside and the air outside the balloon enables the unknown to be readily sampled. Analyses may be checked by the gas density method described above or by effusion measurements as described in Part 1.' To illustrate the quality of a typical analysis some results from a related experiment are given in Table 2. Table 2.
Analysis of Gases Dissolved in Water Sample A, ml
Sample B, ml
5.0 6.7 13.3 1.99
4.5 6.8 13.7 2.02
GO2
0,
I% NsOs
These gaseous mixtures were obtained by degassing tap water by boiling. The carbon dioxide is estimated using aqueous alkali as absorbent and the sum of carbon dioxide and oxygen using alkaline pyrogallol. The amount of carbon dioxide may, of course, vary but the N2-O2ratio should be 2.03 (at 2 5 T ) if the dissolved gases are in equilibrium with air. Solubilities of Gases in Liquids
The solubilities of gases such as carbon dioxide and hydrogen sulfide in a variety of solvents a t various temperatures are readily measured. A 25 ml sample of the gas is shaken with 5 ml of distilled (or a t least degassed) solvent. Determinations a t other than room temperature may be made by immersing the syringe in ice and water baths. Typical data are given in Table 3. Table 3.
Solubility of COz in Various Solvents
Solvent
Temp ("C)
H20 Hz0 GHsOH C3H60H CaHsNH2 C.H,NH,
25 0 25 0 25
0
-Solubility in g/100 mlFound Literature" 0.14 0.34 0.42 0.68 0.26 0 37
0.141 0.327 0.44 0.70 0.23
SEIDELL, A., AND LINKE, W.F., "Solubilites of Inorganic and Metal-OrganicCompounds," Vol. 1,4thed., 0.Van Nostrand Co., Inc., Princeton, N.J. 1958, pp. 460, 476, 489.
The method is limited to gasel of intermediate solubility but the scope of the exqeriment may be extended somewhat by studying the d e c t of increasing pH on solubility. Volume 39, Number 12, December 1962
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617
Henry's Law
Vapor Pressures of Liquids
As has been described in Part 1,' the pressure of a gas in a syringe may be varied by piling on, or suspending from, the piston varying numbers of "calihrated" (by Boyle's Law) textbooks. Thus it should he an easy matter to verify Henry's Law but because of the slowness of the approach to equilibrium the experiment is rather tedious. However as may be seen from Figure 1
Three methods have been used to study the vapor pressures of volatile liquids from zero to 2.5 at,mospheres. (a) A measured volume of dry air is sealed in a syringe by slipping a small serum stopple over the Luer-Lok fitting of the syringe (see Part 1,' Fig. 4). small amounts of a relatively volatile liquid are injected by means of a second syringe and the resulting gaseous volume noted. The vapor pressure is calculated by assuming Dalton's law of partial pressures to be valid. The procedure is repeated at various temperatures. (b) A small volume of the liquid is drawn into a syringe containing 5 ml of dry air a t 0% The syringe is capped and the total vapor volume is recorded as the syringe is heated. The contribution of the air to the total vapor volume at each temperature is readily calculated (providing air is essentially insoluble in the liquid), and the vapor pressure may again be found using Dalton's law of partial pressures. (c) Vapor pressures above the standard boiling point may be obtained by the following somewhat hazardous method. A sample of degassed liquid is drawn into a syringe, all air is expelled, and the syringe capped. The syringe is placed in a water bath which is slowly heated. Simultaneous temperature and volume readings are taken. At the boiling point a marked acceleration in the motion of the piston is easily noticed. Successive books are piled on the pistons as soon as a temperature rise causes the piston to "take off ." In the later stages of the experiment leaks tend to develop but they do not seem to affectthe results. Figure 3 compares some results obtained for diethyl
Figure 1.
Henry's Law plot for the solubility of Con.
(which gives data for the C02-H,O system a t 25'C) satisfactory results are eventually obtainable. Because the pressure varies, the ideal gas law should be used to determine the amount of gas remaining undissolved. Liquefaction of Gases
The pressure/volume relationships of easily liquefiable gases such as ammonia, sulfur dioxide and methylamine are readily studied at gradually diminishing temperature by the method described for Boyle's Law in Part 1.' Glycerol must not be used as a lubricant but silicone and fluorocarbon oils work well. Typical data for ammonia are shown in Figure 2. The liquecan be studied a t room temperature faction of NO2-N20n but the results are complicated by the pressure dependence of the equilibrium. (See below.)
~b
10
A 3b 40 TempemtuR
Figure 3.
io ('GI
$0 t
Vapor pressure curve for ether.
ether using methods (a) and (c) with accepted values. The vapor pressure of ether cannot be studied by method (b) as air appears to be surprisingly soluble. Table 4 gives some data obtained for carbon tetrachloride using method ( b ) . Stoichiometry Figure 2.
618
Prelrure-volume isotherms for ommania.
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Any reaction which quantitatively yields an insoluble gas can be used to illustrate stoichiometric behavior.
Various procedures are possible but the following two will serve most purposes. (a) A small glass (or copper) cup is glued to the bottom of the piston. Measured lengths (I/% in., 1 in. and 1'/2 in., for example) of magnesium ribbon are placed in the cup and dilute acid is sucked up beneath. The syringe is capped and shaken. The volumes of Hp are read directly and values of 9.7 ml, 19.5 ml, and 29.4 are typical. If a good balance is available the equivalent weight of magnesium is, of course, readily found. Aluminum foil, iron wire, etc., may conveniently replace the magnesium. This equipment is also suitable for the gas-volumetric analysis of aqueous hydrogen peroxide. Thus manganese dioxide may be placed in the cup and a sample of peroxide drawn in beneath. The results may be compared (reasonably favorably) with those found by permanganate titration. Table 4.
Second, the reaction is extremely sensilzoe to pH and a high-capacity buffer is required. Figure 4 illustrates a lecture demonstration version of this experiment which utilizes 100 ml syringes. The buffer used here is a 1 M solution of (XHJ~HPOIwhich has a pH of 7.90 a t 20°C. Fifty ml of buffer solution are placed in each of four 100 ml Erlenmeyer flasks containing a magnetic stirring bar and maintained a t 20°, lo0, 0' and -10' respectively, The flasks are
Vapor Pressure of CCla (Atmospheric Pressure 748 mm)
Cdculttted vo!urne of t,mp ( T ) s.17(ml)
Tempera-
Calculated Actual Total vapor vapor pres- vapor presV O I I I I ~ P (ml) sure (mm) sure (mm)
Tima
Figure 4.
h"1
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Rater of hydrolysis of sodium borohydride atpH 7.9.
closed with serum stopples with hypodermic needles through them. The syringes are scremed into place only after temperature equilibrium has been established. Then a t a given instant 2.8 ml of a 0.1 d l solution of sodium borohydride in slightly basic (pH 10) aqueous solution is injected into each vigorously stirred flask. The volume of Hz is read a t suitable time intervals. The data of Figure 4 are self-explanatory. The values for -lO°C are only approximate as some of the (NH4)yHP04crystallizes a t this temperature. The usual logarithmic plots give satisfactory straight lines. An apparent energy of activation may be calculated but is practically meaningless in view of the change in pH of the buffer as the temperature is lowered from 2ooc to - 10°C.
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(b) If the reactants can be mixed in solution the apparatus illustrated in Part 1,' Figure 4 (and used for the first of the vapor pressure methods above) is convenient. Thus a sulfamic acid solution may be placed in a 30 ml syringe and a nitrite solution injected with a second (small) ~ y r i n g e . ~I t is perhaps pertinent to point out here that because of their compactness and calibration syringes make very convenient weight burets.' Rates of Reaction
Fermentation of sucrose and the catalyzed decomposition of hydrogen peroxide are two reactions which yield gaseous products a t a conveniently measurable rate. Both of these are well described in the literature, and they could no doubt be adapted to syringe techniques. For the sake of variety we wish to describe some measurements on the rates of hydrolysis of sodium borohydride in alkaline buffer solutions. The kinetics of this reaction have been extensively described in the literature5 but its pedagogic utility is not widely knoum6 Two warnings are imperitive. First, many gases do not readily leave even hugely supersaturated solutions. Consequently, vigorous and continuous shaking is necessary and the addition of a little sand helps. WASER,J., "Quantitative Chemistry," R7. A. Benjamin, J
Inc., New York, 1961, p. 87. 'THOBURN, J. M., J. CHEM.E D U C . , GlG(1959). ~~, 6 Dams, R. E., BROMELS, E., AND KIBBY,C. L., J. Am. Chem. Soe., 84,885(1962). 6 MARGERUM, D. W., MARTIN,F. D., AND DAVIS,R. E., "Laboratory Manual for General Chemistry," Tri-State Onset Co., Cincinnati, O., 1961, p. 76.
The N02-Ns01 Equilibrium
The qualitative effects of temperature and pressure changes on the nitrogen dioxide/dinitrogen tetroxide equilibrium are very easily demonstrated by syringe techniques. The syringe should be lubricated with the minimum workable amount of Fluorolube. A sample of gas is then introduced. It is now a simple matter to show that as the temperature is raised the gas is more expansive than would be predicted on the basis of Charles' law. Also, a t any one temperature, the gas may be shown to be more compressible than would be predicted on the basis of Boyle's law. This second effectis more readily apparent a t elevated temperatures where the equilibrium proportion of nitrogen dioxide is . . higher. Now if these qualitative observations could be made quantitative, we would have a very simple technique for studying the thermodynamic functions of this equilibrium. The situation is complicated by the fact that neither nitrogen dioxide or dinitrogen tetroxide behaves as an ideal gas,? but nonetheless, t,he data in ~DANIELS F., , AND VERHOEK,F., J. Am. Chem. Sac., 53, 1250 (1931). Volume 39, Number 12, December 1962
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Tables 5 and 6 have been calculated assuming ideality. In order to calculate the average molecular weight, M.,, it is necessary to know the weight of the sample. This was found by direct weighing a t room temperature for the experiment summarized in Table 5 and it was calculated from an interpolated gas density for the experiment summarized in Table 6. The value of a, the degree of dissociation, is calculated from the formula: a=-
Thermal Expansion of 0.0434 grams of NOINlOh at a Constant Pressure of 7 5 3 mm
Table 5. Temperature ("C)
Volume (ml)
Ma,
Literature ('1
Calculated a KP
KP
a
92 - Ma,
M,
and K,, the equilibrium constant, from: Table 6.
I t is seen that agreement with accepted values7 is surprisingly good in view of the extreme simplicity of the equipment and the unjustified assumption of gas ideality. I n conclusion, it must he admitted that of all the syringe experiments described, this is the one most susceptible to small experimental errors.
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Pressure (ah) 10 1.5 2.0
Compression of 0.0730 g of NOI-NIO, 51°C Volume (ml) 29.9 19.1 13.9
Msv 65.7 68.6 70.7
a
K*
0.40 0.34 0.30
0.76 0.78 0.79
at
