Evaluation of the Chloroisocyanurate Hydrolysis Constantst Michael L. Pinsky” and Hua-Ching Hu FMC Corporation, Research and Development Laboratories, Princeton, New Jersey 08540
Equilibrium free chlorine measurements based on a new linear sweep voltammetric technique for free chlorine are reported for the first time at typical low use concentrations found in swimming-pool water. Analyses were performed on buffered solutions over a wide range of parent compound (trichloroisocyanuric acid and sodium dichloroisocyanurate) concentrations as well as at fixed concentrations with various amounts of isocyanuric acid present as would be expected in “chlorine-stabilized” pool water. The measured values were combined with appropriate physical-chemical expressions derived from the chemical equilibrium equations so that the two key chlorine hydrolysis constants could be evaluated. The overall equilibrium was observed to be established within the time of measurement (1min). The pH and temperature dependence of the equilibria were also studied. With isocyanuric acid present the free chlorine residual is expected to be longer lived because some of the most photoactive component, the OC1- ion, is converted to the more photostable chloroisocyanurate form in rapid and reversible equilibria.
Introduction The chlorinated isocyanurates, also called s- triazinetriones, are dry sources of free chlorine. They have been used for disinfection of swimming-pool water over the past 20 yr. They are also used in other sanitizing applications such as laundry detergents and food processing equipment. The two active chlorine compounds of interest are the soluble sodium dichloroisocyanurate available as the dihydrate and the slightly soluble trichloroisocyanuric acid. Their structures along with the parent compound, isocyanuric acid, are shown in Figure 1. One major problem in studying the aqueous chemistry of these compounds a t their normal use concentrations (ca. 1.5 mg/L free available chlorine) has been the inability to distinguish the free chlorine as hypochlorous acid and hypochlorite ion from the chlorine bound to the nitrogen atom. The public health community normally requires measurements of hypochlorous acid (HOC1) because it is considered to be the most effective sanitizing agent. In this report free chlorine refers to the sum of HOCl + 061-. All measurements, both chemical and electrochemical, determine the sum. The distribution, however, is pH dependent. HOCl is normally regarded as more germicidal than OC1-, and, as an example, for E. coli bacteria, OC1- is l / ~ oas effective as HOCl. Both field and laboratory studies (1-6) have suggested that free chlorine and isocyanurate species exist in dynamic equilibrium with one another. This means that chemical or destructive test methods are able to determine only the total free chlorine because the equilibria controlling the free chlorine concentration shift during the test reaction. Therefore, nondestructive measurements of HOCl and OC1- ion are required for equilibrium studies of chlorine-isocyanurate mixtures.
+ Portions of the experimental and theoretical work were presented at the ACS Anaheim Meeting t o the Environmental Chemistry Division’s Disinfection Symposium.
0013-936X/81/0915-0423$01.25/0 @ 1981 American Chemical Society
The equilibrium constants of primary interest in this study are the two which govern hydrolysis of the nitrogen-chlorine bonds on the isocyanurate ring to form hypochlorous acid, hypochlorite ion, and the protonated form of isocyanurate. Following the convention of O’Brien ( 5 ) ,reactions 1and 2 are written for convenience as alkaline hydrolyses because the optimum pH range in swimming-pool water is 7.2-7.8. C12CA-
+ OH- s HClCA- + OC1-
HCICA-
+ OH-
H2CA-
+ OC1-
(K1)
(1)
(Kz)
(2)
Most of our experiments were performed at pH 7.5. Here, C12CA- is the dichloroisocyanurate ion and H&A- is the anion of isocyanuric acid, H&A. K1 and Kz refer to equilibrium 1 and 2, respectively. The acid dissociation constants which control hydrolysis of hypochlorous acid (HOCl) (7) and isocyanuric acid over this pH range are as follows:
Previous studies of the complex equilibrium chemistry of free chlorine-isocyanurate mixtures were based on spectrophotometric measurements (3-5) of solutions containing millimolar concentrations of free chlorine and H&A, which are up to 1000 times greater than those used in practice or in this study. The spectra of these solutions exhibit maxima attributable to OC1- ion and diffuse characteristics of some isocyanurate species. At any particular pH, however, there are no features which can be uniquely assigned to any single component because all spectra are seriously overlapped ( 3 , 5 ) .In fact, according to the published constants ( 5 ) )there can be at least three and perhaps five different isocyanurate compounds responsible for the absorbances between 200 and 240 nm. In O’Brien’s study, Kz was evaluated over the pH range 9.4-10. However, the problem of having too many isocyanurate species at significant concentrations is most serious in that pH range because of contributions from reactions 3 and 4 leading to the isocyanurate anion, HCA2-, and the mono-
HClCA-
H+t ClCA2-
(4)
chloroisocyanurate anion, ClCA2-. Curve-resolving techniques are not yet advanced sufficiently to deconvolute such composite spectra with an acceptable degree of certainty. Consequently, the absorbance at 292 nm contained unresolved contributions from the isocyanurates, so that at best only an upper limit of the OC1- ion concentration could be determined. In O’Brien’s study, K1 was established by combining three other constants. A summary of published constants is given in Table I. The majority of our experiments were carried out at ambient conditions with temperatures of 22 f 2 “C. For calculations utilizing these data, we choose a PKHoCl of 7.50 and a P K H ~ Cof A 7.0 after Gardiner ( 3 ) .The later experiments performed with controlled temperatures were evaluated by using PKHoCl values from Morris (7) and P K H ~ Cvalues A from our Iaboratory. Volume 15, Number 4, April 1981
423
,.
reaction 1;therefore, the equilibrium C12CA- anion concentration can be neglected for these calculations without serious errors. The constants will be shown to be well-behaved in that they will have predictive value in solutions of different compositions.
V
CI
II / \
\
Experimental Section
ci Tr ic h lor o i so c yo nu r i c Ac i d
0
II
cI\
7
/\
N /cI
\o-
04 c\N//c
NO+
Sodium Dichloroisocyonurate
0
II
0// C \ N / c \ o
I
H
lsocyonuric Acid
Flgure 1. Molecular structures of the three commercial compounds of interest. The isocyanurates are also called s-triazinetriones.
In order to calculate Kz a t three controlled temperatures, we found the values for K , by solving the equation log K, = -4470.99/T
+ 6.0875
-
0.017060T
where T is in kelvins (9).The identical values were used to convert the alkaline hydrolysis constants K Z to hydrolysis constants in water, where
Kzb = K w K 2 / K ~ o c i At pH 7.5, the important isocyanurate species are H3CA and HzCA- anion. Hydrolysis of ClZCA-anion to HClCA- anion is most likely while significant concentrations of the CICA2anion ( 5 ) do not exist, but this will be verified later. The low concentrations (10-5-10-4 M) used in this study simplify the calculations because an assumption can be made, and tested later, that the ClZCA-anion is more than 90% dissociated in
Reagents. Standard solutions of chlorine of concentration 3X M were prepared by diluting 4-6% sodium hypochlorite solution (Fischer Scientific Co.) with pH 7.5 phosphate buffer solution. The solutions were standardized iodometrically against a 0.1 N standard sodium thiosulfate solution. The standard solutions were freshly prepared daily and stored in a dark bottle to prevent photodecomposition. Sodium dichloroisocyanurate solutions were prepared by dissolving 0.1 g of NaClzCA-2HzO (FMC Corp., trade name is CDBClearon) in 20 mL of phosphate buffer solution. The total available chlorine was determined iodometrically against standard sodium thiosulfate solution. The solution was freshly prepared daily and stored in a dark bottle. Isocyanuric acid solutions were prepared by dissolving 0.05 g of 99.6% H3CA (FMC Corp.) in 25 mL of water with 1-2 mL of 0.1 N sodium hydroxide solution added to accelerate the otherwise very slow dissolution process. Supporting electrolyte solutions were prepared by using a mixed water system consisting of 90% deionized distilled and 10% distilled water. We subsequently found that this water contained less than 0.1 mg/L of cupric ion which facilitated the measurements (10).The supporting electrolyte solutions were prepared by dissolving 13.8 g of sodium phosphate, monobasic, in -950 mL of the mixed water system. The solutions were adjusted to pH 7.5 by adding concentrated sodium hydroxide before adjustment to final concentration. Linear Sweep Voltammetry. Linear sweep voltammetry is a technique in which a potential is rapidly applied to an electrode in an unstirred solution. The resulting current is thus a function of potential. The determination of free chlorine was performed with a PAR (Princeton Applied Research Corp.) Model 170 electrochemistry system using a three-electrode system (IO).The working electrode was a PAR wax-impregnated graphite electrode (Catalogue No. 9319); the counter electrode was a platinum foil electrode; and the reference electrode was a calomel saturated electrode. The cell was a PAR jacketed polarographic cell (Catalogue No. 9300 and 9301) covered with black tape to protect the chlorine from photodecomposition. All current-potential curves were recorded from f0.5 V to -0.5 V vs. SCE at a rate of 10 mV/s at room temperature. The peak potential is located a t --0.2 V vs. SCE. The current was measured between the base line and the peak. In some cases, a plateau was obtained in the current-potential curve instead of a peak. The current was then measured between the base line and the plateau. The measured current is proportional to the surface area of a particular electrode. For a typical solution of 5 X M sodium hypo-
~
Table 1. Summary of Equilibrium Constants and Experimental Methods ref
4 3
5 this work
KI
31.6 44.2 98.0 235
K2
9.6 3.2 7.7
34 6.0 28 60
424
Environmental Science & Technology
expll methods
pH titration, spectrophotometry spectrophotometry pH titration, spectrophotometry linear sweep voltammetry at carbon electrode
lonlc strength, M
0.5 0.3 0.02 0.1 0.2 0.2 0.2 0.2
t$mmp. C
23 23 25 22 f 2 15.5 25.0 30.0
60
50
r4,
0
+ 0.5
t 0.3
to1
1.5
-0 3
I
POTENTIAL (VOLTS VS SCE)
chlorite, an average current of 30.2 units (each unit on the chart paper is equivalent to 0.1 PA) was observed with a standard deviation of 0.78 unit. Total available chlorine was determined by iodometric titration. An aliquot of the solution was diluted with sulfuric acid followed by the addition of excess potassium iodide. The liberated iodine was then titrated with standard sodium thiosulfate. Procedures.Free chlorine calibration curves were prepared by adding predetermined amounts (10-100 pL) of freshly prepared, diluted sodium hypochlorite solution into the cell already filled with 20 mL of oxygen-free supporting electrolyte solution. The solution was scanned immediately from +0.5 V to -0.5 V vs. SCE. To determine the dissociation of NaClZCA, we performed measurements on solutions prepared by adding a predetermined amount of freshly prepared NaClzCA solution into the cell already filled with 20 mL of oxygen-free supporting electrolyte solution. The current-potential curves exhibited the same shape and potential values as in the pure NaOCl solution. The free chlorine was measured after a lapse of 1min. The degree of dissociation at a given concentration level is defined and calculated as the ratio of free chlorine to the total available chlorine measured iodometrically. T o determine the equilibrium free chlorine concentration in an aqueous chlorine-cyanuric acid system, we added various amounts of H&A and a fixed concentration of sodium hypochlorite solution into the cell already filled with 20 mL of oxygen-free supporting electrolyte solution. The free chlorine concentration was measured after stirring the solution for 1 min. Solutions of isocyanuric acid containing 0.1 M NaCl were titrated with standardized 0.1 N NaOH. The pH at one-half the volume added to reach the equivalence point is equal to the pK; pKs were found for H&A at 15.5,25, and 30 "C (f0.2 "C).
Results and Discussion The current-potential curves of NaOCl standards are shown in Figure 2. All solutions gave well-defined current-
M NaOCi; (d) 5.65
M NaOCI; (c)3.77 X
Figure 2. Current-potential curves of free chlorine from NaOCI: (a)residual current;(b) 1.88 X X M NaOCI; (e) 7.53 X M NaOCI; (f) 9.42 X M NaOCI.
/
30
E
2c
I
0
N I1
K V
I =02 pHs75
v
w
16
K K
a V
IC
!
C
/
/ 3 7
3
NaOCl
565
7
5
CONCENTRATION,MOLES/L
9
2
X IO5
Figure 3. Calibration curve of NaOCI.
potential curves as well as satisfactorily low residual currents. The calibration curve of free chlorine with concentrations to 9.42 X M is linear, as shown in Figure from 1.88 X 3. The current values are proportional to the free chlorine concentration based on the following electron-transfer reaction: C10- HzO 2e- C1- + 20H-. Table I1 contains data from the current-potential curves of C12CA- solutions a t room temperature. The measured free chlorine concentrations are always smaller than the total free chlorine, indicating that the C12CA- anion does not dissociate completely a t very low concentrations.
+
+
-
Volume 15, Number 4, April 1981 425
Table II. Experimental Free Chlorine Concentrations from Sodium Dichloroisocyanurate 10S[C12CA-], M
0.76 1.05 1.52 1.86 2.46 3.02 3.04 3.10 6.08 9.03 11.4 15.1
105( [HOCI] M
+ [OCl-I), 1.26 1.70 2.34 2.72 3.49 4.23 4.26 4.28 8.15 11.9 13.5 18.7
a
( 1 - a)/ (2a2- a)
0.83 0.81 0.77 0.73 0.71 0.70 0.70 0.69 0.67 0.66 0.59 0.62
0.310 0.378 0.553 0.804 0.973 1.07 1.07 1.18 1.45 1.61 3.86 2.55
The derivation of eq I11 which satisfies these needs is given in the Appendix. (1-a) c -- - [C12CA-]o 2d-a
K2
At higher C12CA- ion concentrations where the fraction of dissociated free chlorine is smaller, the assumptions required to derive eq I11 (Appendix) are not valid. One consequence of this condition is the discontinuity in the left-hand-side term at a = 0.50, and the negative values for a
