Polarographic Determination of Phthalaldehyde N. HOWELL FURMAN and DANIEL R. NORTON' Princeton University, Princeton,
N. 1.
The present study was made in order to investigate the polarographic characteristics of phthalaldehyde and to develop a method for its determination. Phthalaldeh).de is reduced a t the dropping mercury electrode, giving two waves. The half-wave potentials of both waves become more negative with increasing pH. The reduction current constants of the first wave do not vary between pH 3 and 9 but increase between pH 9 and 12, while the constants for the second wave increase between pH 3 and 9 and decrease between pH 9 and 12. The limiting currents are controlled by the equilibria between reducible and nonreducible forms of phthalaldehyde, by their rates of diffusion, and by their rates of reaction. I t is postulated t h a t the reducible forms of phthalaldehyde are the normal and monohydrated forms, respectively. The reduction current of each wave in buffered solutions between pH 3 and 12 is proportional to the concentration of the aldehyde in the range of 0.40 to 2.00 millimolar. A t pH 14 phthalaldehyde is unstable and, as a result, the limiting currents diminish with time. The effects of changes in buffer capacity, solvent concentration, mercury pressure, and temperature are discussed. A method for the determination of low concentrations of phthalaldehyde is based on these studies.
benzaldehydes. The present work describes a polarographic investigation of phthalaldehyde upon which is based a method for the determination of this compound. EXPERIMENTAL
Apparatus. The polarograms were recorded with a Sargent Model S I 1 polarograph and with a photographic recording polarograph designed and constructed by Furman, Bricker, and Whitesell (9). The galvanometers were calibrated by the method described by Kolthoff and Lingane (12). Half-wave potentials referred to the saturated calomel electrode were measured graphically from the photographic record with a maximum deviation from the true value of 0.01 volt. Capillary constants were determined a t a mercury column height of 400 mm. with an open circuit a t a temperature of 25" C. The value of m (rate of flow of mercury) in air was 1.536 mg. sec.-l while the drop time i n 0.1.V potassium chloride was 3.68 seconds.
A
NUMBER of interesting studies of the polarography of aliphatic and aromatic afdehydes have appeared in the literature and have been summarized by Rawzonek ( 2 0 ) and by Kolthoff and Lingane ( 1 2 ) . Some of these studies illustrate the importance of considering the equilibrium between the free aldehyde and its hydrated form. I n the case of formaldehyde, because only the free formaldehyde is reducible, the reduction current is limited by the rate of dehydration of the nonreducible form ( 2 , 1 8 ) and also by the factors that influence the equilibrium, such as temperature, pH, solvents ( S ) , and buffers (19). The authors of the present xork show that similar considerations are important in describing the polarography of phthalaldehyde ( I ) arid its monohydrated form (11). Evidence for the existence of a hydrated form of phthalaldehyde has been given by Seekles ( 1 7 ) who studied the hydration of the aldehyde and isolated a stable monohydrate with a melting point of 45.3" C. In the case of acetaldehyde ( 4 ) and propionaldehyde ( 5 ) the polarographic waves are less dependent on external factors since these aldehydes are less hydrated than formaldehyde. 0-CHO
fi-CHO
u-CHO
~-cH(oH),
I
I1
The polarographic behavior of aromatic aldehydes is strongly dependent upon the pH of the medium and the electrolyte used. It has been observed that benzaldehyde gives two waves between pH 2.6 and 6.0 which merge into one wave above p H 6.2 ( I S ) . Substituted benzaldehydes show the same characteristics except for the pH a t which the two waves merge into one (6, 15, 16). In extending these studies to phthalaldehyde it was found that its polarography differsfrom that of benzaldehyde and substituted 1 Present address, U. S. Geological Survey, Department of t h e Interior, Washington 25, D. C.
POLAROGRA CP EH LL IC
THERMOSTAT VESSEL
-1 [---! ,- 4, -
----II k-
I I
_--
I
I
I
,----
I I
I I
I
I
I
1
I
I
, \
/
\
---
,-
/
I '-a'
Figure 1.
Calomel Electrode Vessel
The cell assembly used in this work is described in part by Furman et al. ( 9 ) , who used a Bakelite plate support, through which were inserted the electrodes and the tubes for the introduction of nitrogen. The present authors used the reference electrode vehsel and thermostat vessel illustrated in Figure 1. The design of the electrode vessel is similar to the one described by Hume and Harris (11). I t differs from the latter mainly in that the bridge is an all-glass unit attached directly to the calomel electrode vessel. A modification of this design has been reported by Mathers and Schoeneman (14). ,4s the temperature coefficient of the measured potential is very small, the cell does not require thermostating and can be placed directly upon the Bakelite support. The use of an agar-salt bridge without a porous plate ensures rapid and complete washing of the bridge between each polarogram. Since the entire cell assembly is easily cleaned and the reaction vessel is easily exchanged, the assembly gives efficient service where there are a large number of solutions to be polarographed. 1111
ANALYTICAL CHEMISTRY
1112 -4 Beckman Model G pH mctcr was used for H measurements. The resistances of the solutions were measurefwith a solar conductivity bridge, Model RC-lB, and it was found that corrections of potentials for IR drop were unnecessary in this study. Thc t,cmperatures of the solutions polarographed were controlled at 25.0" =k 0.1" C. Materials. Phthalaldehyde was prepared by the method of Thiele and Gunther (18). Recrvstallization from Detroleum ether gave long, pale >ellow needles with a melting range of 55.0' to 55.4" C. A carbon-hydrogen analysis of this compound gave 71.82% carbon and 4.77% hydrogen. The theoretical values are 71.63% carbon and 4.51% h drogen. The buffer components were rcagent g r d e and were used aithout further purification. The ethyl alcohol was checked polarographically and found to be free from reducible impurities. Buffer Solutions. The buffer solutions were prepared by neutralizing 800 ml. of a solution containing 0.4 mole of the appropriate weak acid with a solution of sodium or potassium hydroxide until the desired pH was obtained. The solution was subsequently diluted to 1 liter. The pH was selected to give Rtoek solutions that were approximately 0.2M in each buffer component. Solution of Phthalaldehyde. -4 stock solution was prepared by dissolving phthalaldehyde in absolute ethyl alcohol and then diluting to volume with distilled water to give a solution which WBR 2.00 X 10-2M in phthalaldehyde and 15'% by volume ethyl alcohol. Procedure. Except where otherwise specified the followiiig procedure was employed. Suitable aliquots of the buffer solution and ethyl alcohol were taken so that the final concentration of the supporting electrolyte was 0.1M in the buffer components and contained 1.5% by volume of ethyl alcohol. The saturated calomel electrode was placed in position, the mercury height raised to 400 mm., and nitrogen passed through the solution to he polarographed. Deaeration was continued for 10 minutes, during which time water from the thermostat was circulated around the reaction vessel. At the end of this period the nitrogen was passed over the surface of the solution and the polarogram recorded. Correction for the residual current of the hackground electrolyte was applied to the measurement of the reduction current. *
~~
Table I.
Stability of Phthalaldehyde Reduction Current",
Buffer NaHCOa-NaOH
PH 10.62
IGHPOrKOH
11.60
NaOH (1.OM)
14.0
w.
Time, Min. 27
1 92 1 89
Wave 2 5.89 5,77 5.90
2.08 2.16 2 16
4.82 4.77 4.85
25 100 287
0 ," t i
4.82 3.16 1.81
22 96 284
Wave 1
1 85
0 . I9 0.13
101
296
Concentration of phthalaldehyde, 1 . 2 0 millimolar.
reported that phthalaldehgde is transformed to phthalide upon the reaction of strong base (18). The results recorded in Table I demonstrate that phthalaldehgde is stable for 4.5 hours a t pH 10.6 and 11.6, while at pH 14 the reduction current diminishes with time, indicating its transfoimstion to phthalide which is reducible only a t potentials mol e negative than the discharge potential of the background electiolyte (21 ). I
I 14
.
12
.
10
.
I
rH
a -
6 -
4
-
- 0.7 1
I
'
2,
I
- 1.1 EL
1
91
-
1.5
S C E
I
Figure 3.
Effect of pH on Half-Wave Potentials
Data taken from Table I1
Figure 2.
Effect of Concentration of Phthalaldehyde upon Its Wave Form
Background electrol, te. Carbonate buffer, 15% by volume ethyl alcohol, adjusted t o pH 10.4 1. 2.
Sensitivity Sensitivitj
L/mo, '/pa,
7 X 10-7\1 phthalaldehyde 7 X 10-4.M phthalaldehyde
Stability of Phthalaldehyde. Thiele and Giinthcr ( 1 8 ) reported that phthalaldehyde is stable in the dark but is transformed to a powdery mass upon exposure to light. The present studies show that the aldehyde undergoes no apparent transformation when stored in the darkness for a period of a year. By measuring the reduction current per unit concentration of the aldehyde it wafi ascertained that its stock solution was stable during a period of onc week that it was observed. It has been
Wave Form. Phthalaldehyde is reduced at the dropping mercury electrode, giving two waves. The first wave is well formed over a wide range of pH and aldehyde concentration. The second wave is well formed a t low pH but a rounded maximum appears a t high pH. At an aldehyde concentration of 10-3M the maximum of the second n-ave is very small at pH 7.3, while a t pH 10.4 it is 6.7y0 of the height of the wave. The ratio of the height of the maximum to the height of the wave decreases with decreasing aldehyde concentration as shown in Figure 2. An unsuccessful attempt was made to suppress the maximum of the second wave. Using a carbonate buffer a t pH 10.4, the following compounds were studied to determine theii effectiveness in suppressing the maximum: lanthanum chloride, 1naphthol, methyl red, methylcellulose, phenolphthalein, mcresolsulfonphthalein, and dimethylaminoazobenzene. These compounds were studied a t concentrations of 0.001, 0.01, and O.lyoin solutions which were 7.0 X 10-3.V in phthalaldehyde. None of these compounds were effective in suppressing the maximum. Gelatin cannot be used as a maximum suppressor as it reacts with phthalaldehyde under the conditions specified here.
1113
V O L U M E 26, NO. 7, J U L Y 1 9 5 4 Table 11. Polarographic Data of Phthalaldebydc in I'arious Buffer Media Buffer Tartaric acidNaOH CHsCOOH-NaOH
3 . 10 4.73 7.26 . 9.15 10,30 10 65 11.60" 14.0
KHzPOrKOH
HaBOs-NaOH NaHCOi-NaOH NaHCOa-NaOH KiHPOrKOH X a O H (1. OM)
-0.64 -0.72 -0.89 -0.97 -0.98 -1.03 1.03 -1.12
-
- I .07
0.47 0.49 0.36 0.59 1.19 1.56 1.76
-1.09 -1.29 -1.40 -1.40 -1.44 -1.46 -1.54
. . , b
Third wave appeared a t thiR pII with an El z of -1.74 IR of 2.62. b Wave height decreases with time.
0 0 3 6
71 87 34 20
5 04 5 00
4.OOb
...
volts and a n
Since maximuin suppressors werc ineff ective, the niasinium was diminished by limit,ing the conrrntration of the aldehyde t o values less than 2 X 10-3.22. 1 - d e r these ronditions t>hrmaximum did not interfere with the measuremrnt of t,he reduction current. Effect of p H upon Half-Wave Potentials. The experimental results are shown in Tahlc I1 and in Figure 3. The half-wavc potentials were measured in carti supporting elertrolytr a t phthalaldehyde concentrat,ions oi 0.40, 0.80, 1.20, 1.60, and 2.00 millimolar. There were no signifimnt variations of El*.>with a change in aldehydc caoncentration. The of bot,h waves become more negat,ive with an increase in pH. The relationship is nonlinear and the points do not fall on a smooth eurvr. -41though a n interpretat,ion of t,hesc results is difficult because of the complexity of the systrins involved, it is apparent that each of the two waves represents one or more reducible forms and that the El12 values of both waves are affert,ed in the samc qualitat,ive way by a changc in the naturc and pH of the huffcr system.
4, the reduction current constants for the first wave do not, vary between p H 3 and 9 hut increase between p H 9 and 12, while the constants for the second wave increase between pH 3 and 9 and decrease between pII 9 and 12. Above p H 9 the decrease in IR of the second wave is approximately equal t o the increase in ZR of the first wave. This seems t o indicate that R change in p H shifts the equilibrium betareen two reducible species. Below p H 9 this relationship does not exist, indicating that, an equilibrium between two reducible species is not the controlling factor under these conditions. Relationship between Reduction Current and Concentration. The reduction currents of the watvcs are directly proport,ionril to thc total concentration of thc aldehyde between 0.40 and 2.00 X I O - 3 M over a p H range of 3.10 to 10.65. While it is possihlc that the relationship holds a t p H 11.6 where the aldehyde was found to be stable, it has been definitely established (hilt thc relationship no Iongrr holds a t pH 14. Representative data given in Figure 5 show the lincar relationship betneoi reduction current, and concentration at 1111 7.26, 9.15, a n d 10.65. I n these studies it was ohscrvcd that, the half-wave potrntials arc1 independent of t h r aldchytfc conrrntration.
-
0
2.0
c >
-
I. 2
-I
2
0.4
c a I
0.0
Figure 5 . WAVE I
W4'rE
1
2 3 4 5
ti
i
I
1.5
3.5
I
5.5
I R
Figure 4.
Effect of plI on Reduction Current Constants Data taken from Table I1
Effect of p H upon the Reduction Current Constants. The data are recorded in Table I1 and Figure 4. The reduction currents were measured in each buffer system at phthalaldehyde concentrations of 0.4, 0.8, 1.2, 1.6, and 2.0 millimolar. Since the relationship between reduction current and aldehyde concentration is a linear one, the data are presented in the form of reduction current, i,, per unit concentration, C . The value ZR is defined as the reduction current constant, i,/C, with the units microamperes pcr millimole per liter. -4s illustrated in Figure
CURRENT
12 0
(MICROAMPERES)
Concentration of Phthalaldehyde bcrbtls Reduction Current
Curl e
2
8.0
4 0
REDUCiICN
Wa\ e 1:irst Fir+t First Serond Second Second
Buffer Phosphate Aorate Carbonate Phosnhatn Carbonate Borate
pH 7.2R 9.1.5 10.50 7.26 10.50 9 1.5
Buffer Capacity Study. The effect of buffer capacity upon thc half-~mvepotential and reduct,ion current was observed by recnortiing polarograme of 1.60 millimolar solutions of pllthrilaldehyde in huff er solutions which were adjusted to approxiniately 0.05, 0.10, and 0.15-11 in both acid and salt components. Studies wcre made of the folloxing buffer systemR: -tartrate buffer at, p1-I 3.10, acetjatc buffer a t pII 4.73, phosphate buffer at pFI 7.26, borate buffer at pII 9.15, and carbonate buffer a t pII 10.30. I n these experiments, the buffer capacity had no measurable effect on the half-wave potentials and it was concluded that the solutions were adequately buffered. The effect of change in buffer capacity upon the reduction current was negligible except, in two cases. These are shown in Table 111. The measurable increases arc not necessarily related to buffer capacity but may be duc t o the effect of the background electrolyte concentration upon the equilibrium and the rate of transformation of nonrctluc.ihlc t o reducible forms of phthalaldehyde. Effect of Solvent Concentration upon and IR. Increasing
ANALYTICAL CHEMISTRY
1114
the ethyl alcohol concentration caused a measurable shift in E112 toward a more negative value and a decrease in ZR. I n a borate buffer a t p H 9.26 an increase of ethyl alcohol content from 1.5 to 6.5% caused a shift in El12of 0.03 volt toward a more negative value and a decrease of I R of 9%. The effects on both waves were the same. It is important, therefore, to control the solvent concentration in the solutions polarographed.
Table 111. Effect of Buffer Capacity on Reduction Current Buffei CHGOOH-Na0H
pH 4.73
KHIPO~KOH
7.26
Change in Increase, Molarity of yo>of i, for Buffer Components Wave 2 0.05t00.10 14 0 . 1 0 to 0 . 1 5 12 0 . 0 5 t o o 10 O.lOtoO.15
7 2
Relationship between Reduction Current and Pressure on the Dropping Mercury. The reduction currents of 1.00 X 10-3M phthalaldehyde were measured a t different heights of the mercury column in a carbonate buffer adjusted to p H 10.50 and the results are shown in Table IV. If the reduction currents were true diffusion currents, they should follow the relationship i, = k(k,,,, ) l j z , where z7 is the reduction current, k is a constant, and h,,,, is the pressure of the mercury column corrected for back pressure (12). The constant, k , was determined for both waves over a wide range of pressure of the dropping mercury. Reduction currents that are determined by a rate constant are found to be independent of the height of the mercury in the reservoir. If this were the case, a 31% decrease in k would have resulted between h,,, of 26.9 and 56.9 em. The reduction current of wave 1 showed a 21% decrease while the reduction current of wave 2 showed a 5% increase over the previously mentioned range in mercury column height. From these results it was concluded that the reduction current of the first wave is largely rate controlled while the reduction current of the second wave is largely diffusion controlled under the conditions employed in this study.
Table IV.
Relationship between Reduction Current and Pressure on Dropping 3Iercury"
Drop Time at Reduction Current, Constant, El, 2 Sec. pa. k = ir/h:Ar:, hcorr.. Wave Wave Wave Wave FTaves FTave Wave Waves Cm. 1 2 1 2 1 + 2 1 2 l f 2 7.0 16.9 6.2 1 73 6.96 0 422 1.27 5 23 1.70 26.9 3.9 1.92 4.4 6.82 8.74 0.371 1 . 3 2 1.69 36.9 3.3 2.9 8.33 10.37 0.336 2.04 1.37 1.71 46.9 2.6 2.3 2.15 9.34 11.49 0.314 157 1.68 56.9 2.2 1.5 2.22 10 40 12.62 0.295 1.38 1.68 73.9 1.7 1.5 2.28 12.04 14.32 0.266 1 . 4 0 1.67 a Polarographic measurements were recorded in a 10-3.M solution of phthalaldehyde in a bicarbonate-carbonate buffer adjusted to pH 10.50. Measurements were made at 23' C. using a Leeds 8: Northrup Electroohemograph. The capillary coqstants at a mercury column height of 28.3 cm. were: m = 2.356 mg. set.-' in air w~ithan open circuit, drop time first wave 4.4 seconds, drop time second wave 3.9 seconds. Using these values, an average correction for back pressure of 1.4 cm. was calculated. The correction was applied to obtain the figures under the column hoarr. in the , table.
The constancy of k shown for the relationship between the sum of the reduction currents of the two phthalaldehyde waves and hoifi, is probably significant. As shown in Table IV, the average deviation of the constants for the sum of the two waves was only 0.7%. A4ssuming that the concentration of the electrode surface of the reducible form corresponding to wave 2 is controlled by a diffusion process, that the reducible form corresponding to wave 2 is in equilibrium with that corresponding t o wave 1, and that reduction processes for wave 1 and wave 2 involve an equal number of electrons, it follows that the sum of the reduction currents of the two waves should bear a constant
relationship to h,',::. The half-wave potentials were also measured and found to be independent of the height of the mercury column. Temperature Study. The temperature coefficient of reduction current controlled by the rate of diffusion of a reducible substance is on the order of 1 to 2y0 per degree centigrade, while that controlled b y the rate of reaction has a much larger temperature coefficient. The temperature coeficients recorded in Table V indicate that kinetic currents were operative in a tartrate buffer at p H 3.02 and in a phosphate buffer at p H 7.21. In the borate buffer a t p H 9.26 the reduction current of the first wave was a p parently rate controlled while that of the second wave was diffusion controlled. The concentration of phthalaldehyde in this study was 2.00 X 10-3M. The temperature had no effect upon half-wave potentials.
Table Y. Effect of Temperature on Reduction Current
Tartaric acidNaOH
3.02
15-25 25-35
8 60 8.41
6.23 5 50
KHIPOI-KOH
7.21
15-25 25-35
7.92 8.60
7.48 6.32
Boric acidNaOH
9.26
15-28 23-38
1l.d 12.7
1.03 1.32
DISCUSSIOR
Reducible Forms. The only example of a diffusion current, the second wave in a borate buffer a t p H 9.26, is found under the conditions where the reduction current constant reaches a maximum. -4comparison of the polarographic characteristics of the second phthalaldehyde wave with the single wave of benzaldehyde show striking similarities as shown in Table VI. It is postulated, therefore, that the compound being reduced a t the potential of the second wave is the monohydrate of phthalaldehyde (11). On the basis of known relationships between the polarographic half-wave potentials and structures of carbonyl compounds, Coulson and Crowell ( 8 ) showed that the relative ease of reduction of a series of aldehydes was largely a measure of resonance and tautomeric effects. Assuming that the compound being reduced a t the potential of the second wave is the monohydrate of phthalaldehyde (11), it is postulated on the basis of resonance and tautomeric effects that the compound being reduced a t the potential of the first wave is the free form of phthalaldehyde (I). Hydrated Forms. The temperature coefficients of phthalaldehyde reduction currents are indicative of hydrated species. From the authors' interpretation it IT ould appear that both reducible and nonreducible hydrates are present in solution. The existence of a monohydrate vias repoited earlier by Seekles (IT), who showed that the rate of solution of phthalaldehyde in water is slow and that its extraction from aqueous solutions ~ i t h organic solvents is difficult. I n a study of the freezing point curves of the phthalaldehyde-water system, Seekles isolated a monohydrate with a melting point of 45.3' C. The existence of a nonreducible hydrate stable only in solution is postulated on the grounds that the reduction currents of both aldehyde waves are far less than would correspond t o the normal reduction current for the carbonyl groups. This rules out the existence of substantial amounts of free phthalaldehyde in aqueous solution and requires that there be a nonreducible form of phthalaldehyde in solution. The existence of a nonreducible hydrate of phthalaldehyde is further suggested in reports that acetaldehyde ( 1 ) and diacetyl(10) are present partly in their hydrated forms in aqueous solutions. I t is postulated that the nonreducible hydrate of phthalaldehyde is the cyclic isomeric monohydrate or the dihydrate.
1115
V O L U M E 26, NO. 7, J U L Y 1 9 5 4
presence of a substance that r e a c t s IV i t h phthalaldehyde and (2) the presence of a subE1 2, Temp. stance that is itself reducible ir/Cm.2 / s t 1 / 6 Coeff , a Aldehyde Buffer PH Vojts a t the dropping mercury elecBenzaldehyde IiHzPOh-KOH G.33 -1.24b 3.50b ... trode. I t was observed that (single wave) 7.0 -1.31 ... 1.04e NHICI-NHIOH 10.0 -1.40 ... ... ammonia and amino acids react Phthalaldehyde KH?POI;KOH 7.26 -1.29 ... ... (second w a r e ) Boric acid9.15 -1.40 3.75 ... with phthalaldehyde in basic NaOH 9.26 ... 1 03 solutions. By selecting a buffer NaHCOn-NaOH 10.50 --'i.'48 ... ... in the acid range these interRange of 16O t o 2 3 O C.: % increase of ir,per C . b D a t a of Coulson a n d Crone11 ( 7 ) . Their buffer contained aninioniurn, acetate, chloride, and phosphate ions. ferences can be avoided. If D a t a of Mathers and Schoeneman ( 2 4 ) . ieducible impurities are present, it may be possible to select one of the two phthalalThat the limiting ruirents are controlled in moqt cases by the dehyde waves that orcur at half-wave potentials wherr there rate of transformation of the hydrated forms is shown by the is no interference. The change of El with pH can alqo be effects of temperature and mercury pressure. Only onc case of used to avoid interferences by reducible impurities. a true diffusion cuirent was noted. The effects of pH, solvent, The average deviation of a series of determinations was lee.; than 2%. While the method was used to determine milligram and buffer type upon the reduction characteristics of phthalaldehyde indicate that 1 he reacfions are probably subject to acid-base quantities of phthalaldehyde it is capable of measuring concnentrations as small as a microgram per milliliter. catalysis. Slopes of Polarographic Waves. Experimentally, it has heen The polarographic method is reliable and accurate for the determination of low concentrations of phthalaldehyde and it should shown that irreversible electrode reactions often follon the 0.0% prove useful in extending the knowledge of other reactions of relationship, E = El 2 log; where E1 2 is constant 01 Zd - Z phthalaldehyde. and independent of the concentration and where CY is less than one ( 1 2 ) . This relationship is found in the case of phthalalde4CKVOWLEDGMEYT hyde. 1-alues of CY weie calculated by determining the slopes The authors wish to thank Dean Resford for the preparation of the polarographic waves and they are recorded in Table VII. of phthalaldehyde, Robert I>. Caswell for the carbon-hydrogen An approximation was made in the use of reduction currents for analysis, and Charles K. Mann for contributing to the maximum diffusion currents. suppression studies.
Table VI.
Comparison of Polarographic Characteristics of Phthalaldehyde and Renzalde hyde
+
Table VII.
t
.4lpha Buffer Tartaric acid-NaOH CHKOOH-NaOH HaB03-NaOH NaHCO3-NaOH K~HPOI-KOH NaOH (1.OM)
LITERATURE CITED
Alpha Values of Phthalaldehyde PH 3.10 4.73 9.15 10.62 11.60 14.0
Wave 1 0.96 0.80 0.54 0.76 0.86
...
Wave 2 0.53 0. ti9 0.84 0.83 0.73 0 58
(1)
Bell, R. P., and Darwent, B. deB., Trans. Faraday SOC.,46, 34 (1950).
(2) Bieber, R., and Trumpler, G., Helo. Chim. Acta, 30, 706 (1947). ( 3 ) Ibid.. D. 1534. lbid., 2000.
i.
lbid., 31, 5 (1948). Brdicka, R., Arkiv Kemi Mzizeral. Geol., 26B, KO,19 (1948). Coulson, D. &I., and Crowell, W. R., J . A m . Chem. Soc., 74, 1290 (1952).
In general the magnitude of the values of cy of the two waves bear a reciprocal relationship to each other. The value of CY for one xave increases while that for the other n a v e decreases as the pH is changed. T h r corresponding data for the phosphate buffer a t p H 7.26 nap excluded from the table since its wave is unsymmetrical. The average slopes calculated for both phthalaldehyde waves between pH 5 and 10 are 0.043 volt per pH unit. Coulson and Crowrll ( 7 ) rrported a slope of the same v h e for benzaldehyde which i- a two-electron reduction. Using the Ilkov~i.equation, an estimation as made of the number of electrons per molecule in the reduction of the specie identified with wave 2 in a borate buffer a t p H 9.15. The calculated value of 2.3 gives further evidence of similarity hetween the reduction characteristirs of this specie postulated to be the monohydrated form of phthalaldehyde and benzaldehyde. Analytical Method. The polarographic determination of low concentrations of phthalaldehyde is based upon the linear relationship betn een ieduction current and concentration. The accuiacy of the detrrmination depend6 upon a number of factors. The solvent concentration of the solutions polarographed should be held constant. Care should be taken to control the temperature, as the temperature coefficients of the kinetic currents are on the order of ten times those of diffusion currents. When a fresh buffer solution is prepared, a new calibration curve should be constructed. Two types of inteiferences might be encountered. (1) the
Ibid., p. 1294. Furman, Pi. H., Bricker, C. E., and Whitesell, E. B., IND. ESG. CHEM.,Ax.4~.ED., 14, 333 (1942). Harrison, S., Collection Czechoslov. Chem. Communs., 15, 818 (1950).
Hume, D. N., and Harris, IT.E., IND. Esc. CHEM..ANAL.ED., 15, 465 (1943). Kolthoff, I. >I., and Lingane, ,J. J., "Polarography," Vols. I and 11, Kew York, Interscience Publishers, 1952. Korshunov, I. .4., and Sazanova, L. N., Zhur. Fiz. Khiin., 23, 202 (1949).
Mathers, A. P., and Schoeneman, R. L., J . Assoc. Ofic.Agr. Chemists, 35, 830 (1952).
Portillo, R., and Tarela, G., Analcs lis.
2/
qctim. ( M a d r i d ) , 40,
839 (1944).
Santavy, F., Collcctzoii Czcchoclor. Chem. Communs., 14, 145 (1949).
Seekles, L., Rec. trais. cIi7i~z.,42, 706 (1923). Thiele, J., and Gunther, O., A n n . , 347, 100 (1906). Vesely, K., and Brdicka, R., Collection Czechosloz. Chem. Communs., 12, 313 (194T).
Wawzonek, S . , A X I L . CHEM.,21, 61 (1949): 22, 30 (1950); 24, 32 (19.52).
Wawzonek, S., and Fo-sum, J. H , J . Electrochem. Soc., 96, 234 (1949). RECEIVED f o r review, December 14, 1953. Accec>ted Aipril 8, 1954. Presented before t h e Analytical Cheniistry Section of the X I I t h International Congress of Pure and Applied Chemistry, New York, N . Y., 1961. T h e study was supported by a fellowship provided by Wallace Laboratories, Inc., Llonmouth Junction, 1v. J. Taken in p a r t from a thesis presented by Daniel R. Norton t o t h e Faculty of Princet,on University in partial fulfiilnient of t h e requirements for the P1i.D. degree, J u n e 1948.
