I
I
N. 0. SCHMIDT and W. S. WISE Department of Sugar Chemistry and Technology, Imperial College ofTropical Agriculture, Trinidad,,B.W.I.
Chemical Solution of Evaporator Scale 1A
number of methods for chemical cleaning of evaporators are available; EDTA has the widest application, but the choice of method must be based on consideration of the chemical type of scale. Sodium citrate is recommended for the more soluble scales.
Tm
formation of scale in sugar factory evaporators is of considerable importance to the industry. By the end of a week's continuous operation, the scale that has accumulated has usually reduced the rate of heat transfer in the evaporators to such a n extent that further working of the scaled evaporator would be uneconomic. Detailed data on the effect of scale on evaporation rates are not available but in one factory (73) using a quadruple-effect evaporator without vapor bleed, in order to maintain the evaporation rate constant a t 9.0 pounds per hour per square foot of heating surface the steam pressure on the first calandria has to be steadily raised from 10 to 23 p s i . over a period of 7 days. As it is impracticable to work the calandrias a t much higher steam pressures, the only solution is to remove the scale. A popular method of descaling is to scrape the scale from the tubes with rotary wire brushes or cutters, varigus designs of which are commercially available; frequently a chemical treatment softens the scale before the mechanical cleaning. I t is difficult to obtain detailed information on the cost of mechanical cleaning, but a typical value in the British West Indies would be about $0.25 (BWI) per ton of sugar produced. This is not a n unduly high cost, and in many factories mechanical cleaning is preferred. There are, however, objections to mechanical cleaning. Some factory engineers consider that the life of evaporator tubes is shortened by abrasion, distortion, or fatigue of the tube material as the cleaning tools are passed up and down the tubes. No evidence appears to have been published on these points, although Schmidt ( 9 ) has shown that the abrasion of tube material caused by a mild wire brushing treatment depends on the composition of the scale and on average amounts to roughly 0: 1% of tube thickness per crop. In the bigger factories mechanical cleaning may present difficulties owing
to the time taken to clean the large number of tubes (upward of 1000 tubes per vessel), but this will depend on what other essential work is carried out during the week-end shutdown; the cleaning of the evaporators is not always a limiting factor. Perhaps the most important objection to mechanical cleaning is the fact that, in the British West Indies a t least, it is becoming increasingly difficult to obtain labor to carry out the cleaning, and some factories now using mechanical cleaning are considering 'chemical cleaning. An ideal solution of the scale problem would be some means of preventing the deposition of scale in the evaporator tubes; considerable effort, involving either chemical or physical treatment of the juice, has been made to achieve this end, but so far without complete success. Another possibility is to design evaporators which either do not scale ( 5 ) ,or can be descaled more conveniently-e.g., by using external calandrias. However, cost renders this solution unattractive as an immediate remedy. Under these circumstances the complete chemical cleaning of evaporators, already used in some factories, is becoming of increasing interest. Some scales contain a high proportion of silica and are readily soluble in hot, strong sodium hydroxide solution. Usually the scale consists of calcium salts with a small proportion of magnesium salts, the usual anions being sulfate, phosphate, silicate, and anions of organic acids. As such salts are soluble in acid solution, it has been common practice to use hydrochloric or sulfamic acid to dissolve evaporator scales. The disadvantage of using acid is that the iron or steel body of the evaporator is attacked, and cases have been recorded (7) of the collapse of evaporators due to corrosion by acid cleaning solutions I t has recently been found possible ( 7 7 ) to dissolve the calcium and magnesium salts of the scale in alkaline solutions of (ethylenedinitrilo) tetraacetic acid (ethylenediaminetetraacetic acid, EDTA). Factory trials have been carried out on an evaporator and a simple method of regenerating the spent EDTA solutions has been developed ( 8 ) . More detailed information was needed on the factors affecting the rate of solution of scale by EDTA.
Exper imental The first experiments were carried out by circulating test solutions through short
lengths of discarded scaled evaporator tubes ( 6 ) , but it was difficult to obtain reproducible quantitative results in this way. The technique was therefore developed of casting plaster of Paris cylinders which were rotated a t constant speed in the test solutions? Smooth cylinders, 8 cm. long and 2 cm. in diameter, were cast from dental-grade plaster of Paris in glass specimen tubes, a brass rod being clamped axially in the specimen tube before the plaster was added. The glass tube was subsequently removed by cracking it and peeling from the cast. The flat ends were protected from the action of the test solution by coating with cellulose acetate cement. Cylinders were discarded when diameter was decreased by 2 mm. When not in use, they were stored in distilled water. For carrying out a run the brass rod was held in a chuck connected by a flexible drive to a reduction gear driven by a */3-hp. induction motor to ensure that the cylinders rotated a t constant speed (140 r.p.m.). The brass rod passed through a bearing in a large rubber bung which was fitted into a wide glass tube of volume about 500 ml. The glass tube was contained in a water bath, the temperature of which was controlled to within 0.5' C. In starting a run the test solution was brought u p to temperature and then the cylinder, previously warmed in distilled water to the temperature of the bath, was introduced. The motor was then started and thereafter 2-ml. samples of the solution were removed a t regular intervals and analyzed for calcium by the Schwarzenbach titration method ( 3 ) , using Eriochrome Black T as indicator. Immediately after the removal of the sample, 2 ml. of a solution identical with the test solution, but not containing EDTA, were added to the test solution to maintain a constant volume. The test solutions consisted of buffers in which EDTA and sometimes other substances such as sodium fluoride were dissolved. The buffers were hydrochloric acid for p H values less'than 2, acetic acid-sodium acetate for p H 3 to 6, ammonia-ammonium chloride for p H 8 to 11, and sodium hydroxide solutions for p H above 12. All p H values were determined using a Beckman laboratory model p H meter. The solutions of magnesium-EDTA chelate were prepared by adding the calculated amount of magnesium sulfate to a 6OmM EDTA solution. T o avoid large variations in sulfate concentration during the run, all VOl'. 50, NO. 5
MAY 1958
81 1
solutions, except in the citrate runs, contained lLM sodium sulfate. I n the experiments with sodium citrate solutions, a fine precipitate of calcium citrate appeared during the run. By blowing compressed air through the solution for a few seconds it was possible to obtain a uniform suspension of the precipitate and the samples were then removed by pipet in the usual way. The calcium citrate precipitate dissolved immediately in the EDTA solution used for the Schwarzenbach method of calcium estimation, so that the subsequent titration gave the total amount of calcium sulfate removed from the cylinder.
Figure 1. Reaction rate varies linearly with concentration Initial
concentrations
of EDTA (mM): 0 100 8 50 @ 40
.
0 0
Resu Its I t was thought likely that a t constant temperature, speed of rotation, and radius of cylinder, the solution process would be first-order, giving a n equation: kA - (C, - C ) V
dC/dt = (1 1 where k is a rate constant, cm. m i n . 7 A is the surface area of the cylinder C is the calcium concentration a t time t C, is the calcium concentration a t infinite time Vis the volume of solution
The value of C m is the sum of two terms: the nonionic calcium chelated by the EDTA, and C, the saturation concentration of calcium sulfate in the buffer solution. If E is the concentration of total EDTA present, the final equilibrium concentration of calcium chelate can be put equal to aE, where cy tends to 1 a t high p H values and to zero a t low p H values. Using this notation we have
+ C,
(2 1
C, = CYE
-4 slight correction term must be added to Equation 1 to allow for the fact that samples were removed from the solution and replaced by buffer solution. I n this case the total EDTA4concentration, E, will itself decrease with time. For convenience we now introduce Y , the concentration of unchelated EDTALe., E - C-which was in fact what was experimentally determined. Toward the end of a run, when all the EDTA has been consumed and the solution contains free calcium ions, the value of Y may become negative. Assuming that the sampling is carried out continuously a t u ml. per minute, we can easily derive the equation : - d Y / d t = -kA/V [(l - a ) E - C,] f
Y ( k A / V 4- v / V ) ( 3 ) I n the case where a = 1-Le.: a t p H values where the EDTA chelates virtually a stoichiometrically equivalent amount of calcium ions-this equation becomes - d Y / d t = k A / V X C, Y W / V v/V) (4) A test of Equation 4 is shown in Figure 1.
+
8 12
+
Y mM
For these runs a 1.OM ammonia-ammonium chloride buffer was used, containing EDTA at initial concentrations of 0, 40, 50, and 100mM. The temperature of the solution was 40' C. The solution also contained 1.OM sodium sulfate, so that the increase in sulfate concentration due to the dissolving of calcium sulfate did not markedly affect the total sulfate concentration. Values of -dY/dt were estimated from successive experimental values of Y and t and Figure 1 shows the relation of -dY/dt to Y , which was taken as the mean concentration over which -dY/dt was measured. It can be seen from Figure 1 that Equation 4 is obeyed. The rate -dY/dt varies linearly with Y, and the results from different initial concentrations of EDTA all lie on the same straight line, which also includes the points obtained when no EDTA was present. The negative intercept on the Y axis gives the values of C,. I n the work reported below, values of k and C, were obtained by drawing graphs similar to Figure 1 and applying Equation 4. I n the few cases where a differed sufficiently from unity, Equation 3 was used; under the conditions of the experiments straight-line plots were obtained, as a slow sampling rate ensured that E was effectively constant over the time of the experiments. In some cases, values for C m were obtained by rotating the cylinder for some hours in the test solution, no replacement solution being added and no attempt being made to obrain rate measurements. A series of experiments was carried out at different temperatures, to find the activation energy for the dissolving process (Table I). From these results, by plotting log k against 1 / T , a n activation energy of 4.0 kcal. was found for the process of dissolving calcium sulfate in EDTA solutions. The variation of k with pH is shown in Figure 2 where each point is the mean of at least three determinations. Results
INDUSTRIAL AND ENGINEERING CHEMISTRY
were obtained under different conditions: buffer solutions containing no EDTA; solutions containing 60m.M EDTA; solutions containing 6Omltl EDTA and 0.5M sodium oxalate, fluoride. carbonate, or phosphate, which form sparingly soluble calcium salts; and solutions containing 60mA.Mmagnesium-EDTA chelate. As no specific effects were found. the average of the results obtained in the presence of oxalate, fluoride, carbonate, or phosphate is plotted in Figure 2. Values of C m under a variety of conditions are shown in Figure 3, CCObeing the total concentration of calcium which can dissolve in a given solution of EDTA. C m is the sum of the chelated calcium and the saturation concentration of calcium sulfate in the buffer solution. Discussion of Results
Although there is a certain degree of scatter, Figure 2 shows that the rate constant for the solution of calcium sulfate in EDTA solutions is independent of p H and is not influenced by the presence of fluoride, phosphate. carbonate. or oxalate, all of which form sparingly soluble calcium salts. The same value for the rate constant is obtained in buffer solutions containing no added EDTA, except for very acid solutions, when the rate constant increases slightly. In the experiments using magnesiumEDTA4chelate a lower rate constant was obtained than for pure EDTA solutions, although equilibrium studies show that calcium ions form almost the stoichio-
Table 1.
Variation of Rate Constant with Temperature
Temp., C.
k , Cm./Min.
32 40 50 60 70 76
0.110 0.115 0.117 0.166 0.196 0.223
E VA P 0 R A T 0 R D E S C A L I N 0
0. 5
T
0.05
L
@ ' \
t I
2
4
6
8
I I
IO
I2
PH
Figure 2. Rate constant of calcium sulfate in EDTA solutions is indePendent of PH
Figure 3. Relation of calcium concentration to pH a t infinite time:
@ Buffer alone
A. B.
0 Buffey and EDTA c) Buffer, EDTA, and sodium oxalate, fluoride,carbonate,or phosphate 0 Buffer and sodium citrate 0 Buffer and magnesium-EDTA chelate
'
Buffer alone Buffer and 60 mM EDTA C. Buffer, 60 m M EDTA, and 0.5M sodium oxalate D. Buffer, 60 m,M EDTA, and 0.5M sodium fluoride E. Buffer, 60 mM EDTA, and O.5M sodium carbonate Ail buffer solutions contain 1.OM sodium sulfate
metric amount of calcium-EDTA chelate by displacing magnesium from its EDTA chelate. These results can be interpreted in terms of simple film theory. The fact that a physical rather than a chemical process is controlling the rate of reaction is indicated by the low activation energy found and by the fact that the rate d solution is increased by increasing the flow of liquid past the scale (6). If it is assumed that the reaction between calcium ions and EDTA is rapid and takes place in an infinitely thin reaction zone situated within the stagnant film (72), the following equation is readily deduced : dN/dt = DEA/d ( E -k C, Dc/DE - C) ( 5 ) where DE and Dc are diffusion constants for EDTA and calcium sulfate, respectively, and d is the film thickness. Equation 5 is of the same form as Equation 1. The value of k is given by D8/d and is independent of the rate of chelate formation from calcium ions and EDTA, if this is a sufficiently fast reaction. The experiments with the magnesium-EDTA chelate in which a lower value of k was found suggest that the displacement of magnesium ions from the EDTA chelate by calcium ions is a relatively slow processi I n the case of buffers containing no EDTA, a similar approach gives dN/dt = DcA/d (C, - C) (6) The value of k is given by Dc/d and, assuming that the diffusion constants for EDTA and calcium sulfate are not mark-
F
edly different, the conclusion from Equation 6 is that the rate constant, k, for buffer solutions alone should be the same as that obtained with the buffer solution containing EDTA. This conclusion is supported by Figures 1 and 2. The variation with p H of values found for Cm-i.e., LYE C,--shown in Figure 3, presents a number of interesting features which can be explained in terms of simple ionic theory. Curve A was obtained in buffer solutions containing no added EDTA. The value of C, is roughly constant from pH I to 11. Below pH l the value of C, increases, while above pH 11 it decreases to zero. The decrease is presumably due to the fact that the surface of the cylinder becomes coated with calcium hydroxide in very alkaline solutions; the solubility of calcium hydroxide decreases with increasing pH, owing to the common-ion effect of hydroxyl ions. Below p H 1, a significant amount of sulfate exists as bisulfate ions, so that to maintain the solubility product [Ca++] [Sod--] constant, the calcium ion concentration becomes larger-Le., the value of C, increases. Curve B was obtaiqed with solutions containing 60mM EDTA. Over the p H range of 6 to 1I the value of C is roughly constant. Above pH 11 Cm decreases and tends to the value 6OmM. Below pH 6 CWdecreases markedly and by p H 3 curve B joins curve A . The difference between curves A and B is due to the term LYE. At high pH
+
values, when a = 1, the EDTA is able to chelate the stoichiometric amount of calcium, here 60mM, while at pH values below 4 the EDTA binds calcium less strongly owing to the competition of hydrogen ions for the EDTA anions and LY decreases eventually to zero. The result is that curve B lies parallel to and 6OmM above curve A at pH values above 6, but below p H 6 the difference between curves A and B diminishes, until by pH 3, the curves join together. Curve C, obtained with EDTA in the presence of 0.5M sodium oxalate, is similar to curve B but joins curve A at a higher pH value than does curve B. This is essentially due to the fact that calcium oxalate is very much less soluble than calcium sulfate. The surface of the cylinder becomes coated with calcium oxalate and solution of calcium from the cylinder is possible only when the concentration of calcium ions in solution is less than &,ox/ [Ox- -1, where is the solubility product of calcium oxalate. The calcium EDTA chelate is in equilibrium with a small concentration of calcium ions, which, as the pH is lowered, increases because of the competition of hydrogen ions for the EDTA anions. When the equilibrium concentration of the calcium ions reaches the value &ox/ [Ox-- 1, no solution can take place from the calcium oxalate surface of the cylinder-Le, a = 0. Owing to the low value of the solubility product of calcium oxalate, the calcium ion concentra[Ox--] a t a tion becomes equal to comparatively high pH. Curve D was obtained in the presence of 0.5M sodium fluoride; as calcium fluoride is less soluble than calcium oxalate, curve D is displaced to a slightly higher pH than curve C. Curve E was obtained in the presence of 0.5Msodium carbonate. At high p H VOL. 50, NO. 5
MAY 1958
813
Table II. Solution of Calcium Sulfate in 0.100M Sodium Citrate Solutions at 40” C. and pH 6 Na2S04
Concn., iM
C m , *Vl
0 0.1
0.083
0.110
0.5 1.o
0.069
0.061
values, because of the low solubility ofcalcium carbonate, curve E is similar to the curves obtained for iolutions containing fluoride and oxalate. As the p H is lowered, however, carbonate ions are converted to bicarbonate ions, and because calcium bicarbonate is much more soluble than calcium sulfate, E joins B a t a p H of about 7. Results Obtained with Sodium Citrate
While the results obtained from both laboratory and factory experiments have shown that EDTA can be successfully used to dissolve evaporator scales, a possible objection is the comparatively high cost. I t was therefore decided to investigate the possibility of using cheaper complexing agents. Sodium citrate appeared to be worth investigating. as it is known to form a complex (7) with calcium ions : Ca+Cit3- = CaCit-
-+
although in more concentrated solutions calcium citrate is precipitated. Experiments were carried out in which the plaster of paris cylinders were rotated in sodium citrate solutions. As expected, the calcium sulfate was removed from the cylinder and eventually a precipitate of calcium citrate appeared. The formation of the precipitate did not alter the rate a t which calcium sulfate was subsequently removed from the cylinder. Good first-order plots were obtained for the solution process and the values of k (Figure 2) are very close to those found for EDTA solutions. The disadvantage in using sodium citrate is that the complex CaCit- has a much lower formation constant than the Ca-EDTA chelate and sodium citrate does not dissolve the less soluble calcium salts. This was found experimentally, in that the sodium citrate had no effect on the calcium sulfate cylinder in the presence of added fluoride. The effect was further emphasized by using sodium citrate solutions containing varying amounts of sodium sulfate. The C m values found are shown in Table 11. The Cm value, which is a measure of the amount of calcium sulfate removed from the cylinder. is reduced by the presence of sulfate ion, because the concentration of calcium ions in equilibrium with the CaCit- complex is comparatively high. M’hen the calcium ions
8 14
have reached the concentration corresponding to the solubility product of calcium sulfate, no further solution from the cylinder is possible (cf. the explanation of curve C in Figure 3). I t is obvious that sodium citrate is of much more limited application than EDTA in dissolving scale. However. in experiments (70) on a sugar factory evaporator in which the scale consisted of a high proportion of calcium sulfate. the evaporator tubes were completely cleaned by the treatment. Application of Results These results are of interest in relation to the practical problems of evaporator cleaning. Essentially the rate of cleaning will depend on the twofactors, k and Cm. Apart from the case of magnesiumEDTA chelate, the only variables affecting k are the rate of stirring and temperature. Although an increase in temperature corresponds to an increase in k , owing to the low activation energy of the dissolving process. the rate is not greatly sensitive to temperature variations. It is possible to increase the value of k by increasing the rate of stirring, but this cannot usually be varied under the conditions of cleaning sugar factory evaporators. The smaller values of k obtained with the magnesium-EDTA chelate are of interest in connection with the use of regenerated EDTA solutions for evaporator cleaning (8). Magnesium salts are not in general removed by the regeneration process and gradually accumulate in the EDTA solution as the magnesiumEDTA chelate. The formation of this chelate has two adverse effects: the smaller rate constant and the fact that the magnesium-EDTA chelate cannot dissolve magnesium salts from the scale. Therefore. difficulties may be encountered with EDTA solutions which have been used on scales of high magnesium content. This difficulty would be overcome by using a different method of regenerating EDTA based on precipitation of the pure acid (4, 8). However. this process is unsuitable for regular use in sugar factories (8),as calcium sulfate has to be precipitated before the EDTA is precipitated. The variations in Cm are much larger than the variations in k and their effect on the rate of cleaning is of importance. The results in Figure 3. B, show that for calcium sulfate scales the rate of cleaning becomes independent of p H above p H 6. It has been the practice to use very alkaline EDTA solutions for cleaning evaporators. but the above results show that this is not necessary. a conclusion confirmed by a factory trial (2). The importance of this result is in connection with the regeneration process. which involves acidification of the spent solution with sulfuric acid to precipitate calcium sulfate and decantation of the supernatant layer, the p H of which
INDUSTRlAL AND ENGINEERING CHEMISTRY
is then raised by the addition of caustic soda. By not having to raise this p H to a very high value. there is a considerable saving in caustic soda. with a corresponding saving in the amount of sulfuric acid required for acidification. The results obtained emphasize the difficulties of using acids to dissolve calcium sulfate scales. The ability of acids to dissolve calcium sulfate depends on the increase of C, a t low p H values due to the formation of bisulfate ions in solution. I t can be seen from Figure 3, that the increase in C, is found only a t rather high concentrations of hydrochloric acid and is comparatively small, These results correspond to solutions containing 1M sodium sulfate, so that the solubility of the calcium sulfate is somewhat depressed by the common ion effect. This does not essentially alter the conclusion that acid is less suitable than EDTA for dissolving sulfate scales. Hydrochloric acid, or even weak acids, can, however, be used to dissolve other scales such as calcium carbonate and this is the basis of the soda ash process of cleaning evaporators, in which the scale is first boiled with soda ash, and thus converted to calcium carbonate, which is then dissolved in acid. The results with sodium citrate indicate that it is a satisfactory cleaning agent for the more soluble scales, but it may have a restricted application. I t should be of use in descaling rum stills, where the scale is essentially calcium sulfate. literature Cited
(1) Avalos, M., Keller, A. G., Proc. Intern. Sac. Sugar Cane Technologists,
9th Congr., in press.
(2) Bennett, M. C., Schmidt, N. O., Wiggins, L. F., Wise, W.S.,Intern. Sugar J . 58, 249-52 (1956). (3) Biedermann, W., Schwarzenbach, G., C h k i a 2, 56-9 (1948). (4) Buckley, G. D., Thurston, E. F., Chem. &Y Ind. (London) 1956, p. 493. (5) Chandler, J. L., Proc. Brit. W e s t Indies Sugar Technologists, 1957 meeting,
in press. (6) Connolley, F. H., A.I.C.T.A. thesis, Trinidad, 1955. (7) Hastings, A. B., McLean, F. C., Eichelberger, L., Hall, J. L., da Costa, E.: J . Bid. Chem. 107, 35170 (1934). (8) Holland, I . D., Massiah, B. V.; Meyers, J. C.: Schmidt, N. O., W’iggins, L. F., Wise, W. S., Proc. Brit. W e s t Indies Sugar Technologists 1954, pp.
155-61.
(9) Schmidt, N. 0.:Ibid., pp. 141-9. (10) Schmidt, N. O., unpublished experi-
ments. (11) Schmidt, N. O.?Wiggins, L. F.?IND. ENG.CHEM.46, 867-70 (1954). 112) Sherwood. T. K.. “Adsorution and Extraction,” p. 194, Mckraw-Hill, New York, 1937. (13) Springer, H. B., Proc. Intern. Sac. Sugai Cane Technoiogists, 8 t h Congr., 754-65 (1953). RECEIVED for review April 15, 1957 ACCEPTED October 5, 1957 Part of the research program of the British West Indies Sugar Research Scheme.
