ROQERF. BARTHOLOMEW
3442
A Study of the Equilibrium KNO,(l)
Jc KNO,(l) + 1/202(g)over the
Temperature Range 550-750"
by Roger F. Bartholomew Research and Devebpment Laboratories, Corning G h s s works, Corning, New York
(Received March 90, 1966)
+
Equilibrium constants for the reaction KN03(1) KNOz(1) '/zOz(g) were determined. A value of 27.6 f 1.2 kcal mole-' was found experimentally for the heat of reaction over the temperature range 550-750". Heat content data were obtained for liquid potassium nitrate and nitrite, which led to a calculated value of 27.5 f 0.3 kcal mole-' for the heat of reaction a t 650". The equilibrium constant could be expressed by the equation log K =
-27.6 f 1.2 kcal mole-' 2.303RT
1.2 + 24.4 2.303B f
These results are discussed in the light of the discrepancy existing in the literature.
The thermal stability of alkali metal nitrates has been the subject of much interest because of the applications of these salts at high temperatures in the treatment of steel,' the strengthening of glass by ion exchange,2 and in the field of munitions. The mechanism for decomposition is complex, involving the initial production of nitrite, which further decomposes to the metal oxide and oxides of n i t r ~ g e n . ~In , ~the case of potassium nitrate, it was found that a quasi-equilibrium is set up below 800°.6*6 This equilibrium reaction can be written KNW)
__
KNOz(1)
+ '/zOz(g)
(1)
The equilibrium constant, K , can be expressed by the equation
where N * refers to the mole fraction of the subscript salt present in the melt at equilibrium, y is the activity coefficient, and Pol is the pressure of oxygen above the melt. There is evidence from surface tension data' KNOz system is ideal, and from that the KN03 electrical conductivity that the NaN03 NaNO2 systems is ideal, so the y terms can be dropped. Two papers have been published dealing with the measurement of the equilibrium constant defined in eq 2. Freeman6 followed the change in volume of a closed
+
The Journal of P h y s k d Chemistry
+
system, a t a constant pressure of 1 atm of oxygen, during the decomposition of potassium nitrate and oxidation of potassium nitrite. Sirotkin6 determined K by analyzing the melt, during decomposition of potassium nitrate, for nitrite as a function of time until it became constant. To ach.ieve constant partial pressure of oxygen, air was continuously bubbled through the melt. From the temperature dependence of the equilibrium constant, Freeman obtained 31 kcal mole-' for the heat of reaction, and Sirotkin found 26 kcal mole-'. The present investigation involved following both the change in volume of the system at constant oxygen pressure and the nitrite content, by analyzing the resulting equilibrium mixture. The heat capacities of molten potassium nitrate and potassium nitrite N. P. Popovshaya, Chem.Abstr., 54, 15177d (1960). (2) 8. D . Stookev. J. 8. Olcott. H. M. Garfinkel. and D. L. Rothermel, "Advances jn Glass Technology," Plenum' Press, New York, N. Y., 1962,pp 397-411. (3) 8. Gordon and C. Campbell, Anal. Chem., 27, 1102 (1955). (4) J. C. Casanova, Bull. SOC.China. France, 429 (1959). (5) E.S.Freeman, J. Am. C h m . SOC.,79, 839 (1957). (6) G. D . Sirotkin, Russ. J . I o r g . Chem., 4, 2538 (1959). (7) H.Bloom, F. G. Davis, and D. W. Jones, Trans. Faraday SOC., 56, 1179 (1960). (8) H.Bloom, I. W. Knaggs, J. J. Molloy, and D . Welch, ibid., 49, 458 (1953). ( 1 ) P. I. Protsenko and
STUDY OF
THE
EQUILIBRIUM KNOB(1) + KNOZ(1)
+ l/tOz(g)
were measured as a function of temperature in order to calculate the heat of reaction.
Experimental Section
Materials. Crucibles of 1-in. diameter were manufactured out of both 309 stainless steel and fine silver (99.99% pure). The crucible material was found to have no effect on the equilibrium position. The potassium nitrate used was Baker Analyzed reagent grade (99.5% pure). The potassium nitrite was supplied by Baker and Adamson and was 95.0% minimum purity. Both salts were dried in an air oven a t 120" for several days. The potassium nitrite was further dried by a method proposed by Rhodes and Ubbelohde. The melting points of potassium nitrate (334") and potassium nitrite (428") were determined by measuring cooling curves. Apparatus. Constant Pressure System. The apparatus built was of the same basic design as described by Freeman? However, two major modifications were made. The crucible was suspended in the furnace by a length of silver wire attached to a bar of iron contained in a long glass tube. A magnet was used to move the iron bar and, therefore, the crucible in and out of the reaction vessel. A sample was weighed and placed on a glass spoon, which was then put in the system. The furnace was brought to temperature with the empty crucible in the reaction zone and the sample in the cool part of the tube. After the system was evacuated for several hours, it was flushed with oxygen and finally adjusted so that the pressure in the system was 1 atm. The crucible was quickly raised to the top of the furnace; the spoon was rotated to drop the salt into the crucible, which was then lowered into the furnace. The change in volume of the closed system was followed as a function of time. By this technique, the number of moles of salt present at time t = 0 was known accurately. The other major change was in the control system for the level of fluid in the buret to maintain constant pressure. A needle valve having an adjustable arm at right angles to the valve shaft was constructed out of stainless steel. One side of this arm was attached to a fixed point by a spring, while the other side was connected to a solenoid. The latter was activated by m electronic relay when the mercury level in the Utube made contact in the circuit. When the valve was open, mercury entered the buret from a reservoir. The system could be reversed to have mercury flowing out of the buret when contact was broken in the Utube. The use of this valve has been described by Bartholomew and Flood.10
3443
Calorimeter. A drop calorimeter was used to determine the heat contents of the salts. I n brief, this consisted of a tube furnace with a trap door in the bottom of the tube suspended over a copper calorimeter, which was surrounded by a water bath in a dewar flask. The sample was contained in a tightly sealed silver crucible. Runs were made over a temperature range through which decomposition or oxidation of the salt was negligible. This was checked by noting the change in weight of the crucible and contents after any run; no weight change was discernible. Analysis. The melt obtained at equilibrium was dissolved in deionized water. Nitrite contents were determined by titrating aliquots with KMn04. Both direct titration and back titration with a known excess of sodium oxalate were used. Several melts were analyzed for their nitrate content by a potentiometric titration with iron(I1) in a concentrated sulfuric acid medium." The presence of oxide in the melt was also investigated using a standard solution of hydrochloric acid and phenolphthalein as the indicator. On no occasion was oxide detected. Results and Discussion The volume of oxygen evolved per initial gram of potassium nitrate is shown as a function of time in Figure 1. Runs were made at four different temperatures; the run at 750" reached equilibrium in 60 min, while at 600" it took 300 min to be reached. Figure 2 shows the volume of oxygen taken up per initial gram of potassium nitrite against time. From these two sets of data the volume evolved or taken up at equilibrium was determined. The second column in Table I lists the values obtained. To obtain the mole fraction of nitrite present at equilibrium, the stoichiometry indicated in eq 1 was used; that is, for every mole of oxygen evolved or taken up, 2 moles of nitrite is produced or removed from the system. Analytical figures obtained for the weight percentage of nitrite present in the melt at equilibrium are shown in the fourth column of Table I, while the sixth column shows the nitrate content of the melt in the cases in which it was determined. The sum of nitrate plus nitrite was found to be, within experimental error, representative of eq 1. The mole fraction of potassium nitrite obtained from the analytical data is given
(9) E. Rhodes and A. R. Ubbelohde, Proc. Roy. SOC. (London), A251, 156 (1959). (10) R. F. Bartholomew and E. A. Flood, Can. J. Chenz., 43, 1968 (1965). (11) C.D.McKinney, Jr., W. H. Rogers, and W. M. McNabb, Ind. Eng. Chem., Anal. Ed., 19, 1041 (1947).
Volume 70,Number 11
November 1966
ROGERF. BARTHOLOMEW
3444
Table I
Decomposition of Potassium Nitrate
T ,O
Equilibrium vol, om8 (STP) of Oa/ g of KNOs
C
I?* KNOa (gas)
27.6 14.3 6.4 3.10
750 700 648 600 550
% KNOi (analysis)
N*KNO~
0.248 0.129 0.0577 0.0282
16.42 8.44 5.47 1.97 0.733
...
...
% KNOs
N*KNO* (analysis)
N*KNO~ (analysis)
(analysis)
0.189 0.099 0.0643 0.0233 0.00869
...
...
92.88 94.42 97.23
0.914 0.934 0.967
...
+
N*KNOI (analysis)
1.013 0.998 0.990
...
*..
Oxidation of Potassium Nitrite
T," C
Equilibrium vol, oms (STP) of 021 g of KNOz
750 700 650
98.8 117.0 124.0
I
I
I
7% KNOn (analysis)
N*KNOZ (gas)
0.250 0.113 0.058
I
I
I
I
15.6 8.66 4.59
I
I
I
N*KNO~ (analysis)
0.180 0.101 0.054
I
A-
L f IOI
> m
IW
--C
Figure 1. Volume of oxygen (cc g-1 (STP)) evolved us. time during the decomposition of KNOa.
in column five. As the mole fraction of potassium nitrite was defined by NKNOa
=
I x)
I
40
I
I
I
M w 70 TIYE Ill MINUTES
I
m
I
w
I
100
I 110
I
im
1
Figure 2. Volume of oxygen (cc g-' (STP)) taken up vs. time during the oxidation of KN02.
400
200
TIME IN MINUTES
I x)
nitrite is the starting material. Table I1 shows the value of the equilibrium constants determined under the headings of the method used to obtain them.
Table
II : Equilibrium Constants
nKNOa
+ nmoa
%moa
where n refers to the number of moles of the subscript salt, the mole fraction of potassium nitrate a t equilibrium should be given by 1 - N*KNo~.The equilibrium constant, K , is given by the relationship
if the starting point of the reaction is potassium nitrate, or by K i n v e r s e , the reciprocal relationship, if potassium The Journal of Physical Chemistry
750 700 650
600 550
0.333 0.148 0.061 (648") 0.029
...
0.233 0.110 0.064 (648') 0.024 0.0088
0.333 0.125 0.062
0.220 0.112 0.057
...
... ...
...
The reciprocal of Kinverse is given so that comparison can be made between the two sets of data. It was noted that agreement between the two techniques
STUDY OF THE EQUILIBRIUM KNOs(1)
+
KNO&
+ l/z02(g)
3445
using the same starting material was not as good as agreement between the same technique using different starting materials. These results do indicate that the same equilibrium is attained from either direction. Thermodynamic equations relating the equilibrium constant with temperature can be written
AFO
:=
-RT In K
=
AH" - TASO
where the symbols have their usual meanings. Therefore, a plot of log K against 1/T OK-' should be linear with a slope of -AH0/2.30R and an intercept of AS"/ 2.3QR. This plot is shown in Figure 3, where values of log K and log K-linverseare plotted against 1/T OK-'. The heat of reaction, AH", was found to be 27.6 f 1.2 kcal mole-', and the entropy change, AS", associated with the reaction was found to be 24.4 f 1.2 cal mole-' deg-'. Table I11 compares the values obtained in this study with previously reported data. Unfortunately, Freeman and Sirotkin gave no estimate of the errors in their values. Table 111: Values of AH" and AS" ASo,
AHO,
koa1 mole-'
This work Freemans Sirotkina
cal mole-1 deg-1
27.6 f 1 . 2 31 26.0
24.4 f 1 . 2 31.1 22.9
12.
I1 ' 0.1
%
f
8
0 KIANALYSISI
m
3
X
8
K-llANALYSl51 INVERSE
0.0
L
INVERSE 10.0
I 11.0
.+~ - ' , 1 0 4
-
Figure 3. Equilibrium constant as a function of temperature.
1
I
rso
roo TEMPERATURE O K
Figure 4. Heat content vs. temperatura for liquid KNOI and KN02.
necessary quantities. A drop calorimeter was used, as described in the Experimental Section, to determine the heat contents of both potassium nitrate and nitrite above their melting points. The data obtained, shown in Figure 4, can be fitted by the equation HT - H ~ 8 . 2= aT b, where a and b are constants, T is the absolute temperature, and HT - Hzss.z is the heat content at T"K relative to 298.2"K expressed in kcal mole-'. Table IV lists the values of a and b for KNOl and KNOZfound in this work.
+
Table IV:
I n order to calculate the heat of reaction, it is necessary to know the heat contents of liquid potassium nitrate and nitrite as a function of temperature. As these data are not available for potassium nitrite, calorimetric measurements were made to obtain the
I 650
HT
Values of a and b in the Equation
- H Z S J = aT + b
Salt
Reference
loa,, kcal mole -1 deg-1
KNOa KNOa KNOn
Thiswork Kelley12 This work
29.3 29.5 24.5
b, koa1 mole-]
Temp range,
-5.45 -5.25 -4.72
611-677 611-700 701-780
O K
The agreement between the values determined in this investigation and those given by Kelley12is within experimental error. The heat of reaction at 298.2"K can be found from the known heats of formation of potassium nitrate and nitrite.13 A calculated value of 29.2 kcal mole-' was ~~
_____
(12) K. K. Kelley, "Contributions to the Data on Theoretical Metallurgy," Part XIII, Bureau of Mines Bulletin 584, U. S. Government Printing Office, Washington, D. C., 1960.
(13) F. D. Rossini, D. D. Wagman, W. H. Evans, L. Levine, and I. Jaffe, "Selected Values of Thermodynamics Properties," National Bureau of Standards Circular 500, U. 8.Government Printing Office, Washington, D. C., 1952.
Volume 70, Number 11
November 1966
ROGERF. BARTHOLOMEW
3446
obtained. The assumption must then be made that extrapolation from the temperature range, over which the heat contents of the salts were determined to that at which decomposition takes place, is valid. A further effect must be taken into account before AH" at 650" can be determined. The change in heat content of oxygen gas must be known. This can be allowed for, as the quantity HT - H2ga.zfor oxygen gas has been tabulated by Kelley12 as a function of temperature. The calculated value of AH" was found to be 27.6 kcal mole-' with a possible error of *0.3 kcal mole-'. Comparing this value with that obtained experimentally, it is seen that they are in excellent agreement. The heat of reaction was determined using the equilibrium constants obtained separately from the gas data and analysis data. I n the case where the equilibrium constant was found from analysis of the melt, AH" was 26.4 i 1.1 kcal mole-'. From the gas data, a value of 29.1 1.7 kcal mole-' was obtained. The agreement is within experimental error, but shows the same trend as found in the previously reported data5*6 (see Table 11). That is, the analysis of the melt gave a lower value of AH' than that found from the gas data. The explanation for this difference appears to he inherent in the method. In the cooling of the melt for analysis, some of the nitrite may be oxidized to nitrate. From kinetic datal6 the rate of oxidation
*
The Journal of Physieal Chemistry
of potassium nitrite is very rapid at 750" but quite slow at 600". This fact is in agreement with experiment as can be seen from Figure 3. A t 750" there is quite a spread between the equilibrium constant obtained by the two techniques, indicating that the nitrite concentration in the melt, on analysis, is less than that obtained from the gas data. The possibility of side reactions taking place can be ruled out on the grounds that the gas data at 750" gave the same value of the equilibrium constant, using either potassium nitrate or nitrite as the starting material. Any side reaction would involve the evolution of oxides of nitrogen, which would increase the volume of gas evolved in the decomposition runs and decrease the volume of gas taken up in the oxidation runs. It would therefore appear that the nitrite concentration of the melt at equilibrium, as determined from the gas data, is inherently more accurate then that found by analysis.
Aclcnotvledgments. The author wishes to thank Stanley Lewek and Dr. T. R. Kozlowski for their help in constructing the system and making some of the measurements. He also wishes to express his appreciation to Dr. W. A. Plummer for making available the calorimeter for the heat-content measurements and to Dr. H. M. Garfinkel and Professor B. R. Sundheim for many helpful discussions.
