Environ. Sci. Technol. 2001, 35, 1127-1133
Equilibrium and Kinetics of Bromine Chloride Hydrolysis QIAN LIU AND DALE W. MARGERUM* Department of Chemistry, Purdue University, West Lafayette, Indiana 47907-1393
Aqueous-phase halogen reactions play an important role in tropospheric ozone depletion that is observed during Arctic sunrise where bromine chloride is a key intermediate. The temperature dependencies of BrCl(aq) equilibration with BrCl2-, HOBr(aq), Br2(aq), Cl2(aq), HOCl(aq), Br-, and other species (Br3-, Br2Cl-, Cl3-, OBr-, and OCl-) are determined as a function of Cl- concentration from pH 0 to pH 7. Values for K1 ()[BrCl2-]/([BrCl(aq)][Cl-])) at µ ) 1.0 M are 3.8 M-1 at 25.0 °C, 4.7 M-1 at 10.0 °C, and 5.5 M-1 at 0.0 °C, with ∆H1° ) -9.9 kJ mol-1 and ∆S1° ) -22 J K-1 mol-1. BrCl(aq) hydrolysis equilibria have little or no temperature dependence with Kh1 ()[HOBr(aq)][Cl-][H+]/[BrCl(aq)]) ) 1.3 × 10-4 M2 from 25.0 to 5.0 °C, µ ) 1.0 M. When conditions are adjusted to give a rapid partial hydrolysis of BrCl in equilibrium with HOBr and Cl- at p[H+] 4.31, a relatively slow reaction (kobsd ) 2.4 s-1) to form HOCl and Br- is observed. This takes place via BrCl reaction with Cl- to form Cl2, which hydrolyzes in the ratedetermining step to give HOCl. On the other hand, the rate of complete BrCl hydrolysis to form HOBr and Cl- at p[H+] 6.4 is extremely rapid with a first-order rate constant of 3.0 × 106 s-1 at 25.0 °C. The reverse reaction between HOBr, Cl-, and H+ has a rate constant of 2.3 × 1010 M-2 s-1, so that in seawater, where [Cl-]/[Br-] ) 700, the formation of BrCl is much faster than the formation of Br2 from HOBr, Br-, and H+. Rapid formation of BrCl(aq) and its subsequent reaction with Br- is a viable pathway to give Br2(aq). Photolysis of Br2(g) is believed to initiate the reactions associated with ozone depletion.
with ozone in the dark (15), and bromine release based on the diffusion of HBr(g) and HOBr(g) into the acidified aerosol (16). Kinetics data from Wang et al. (17) showed that the rate constant for BrCl(aq) hydrolysis to HOBr, Cl-, and H+ was greater than 105 s-1 at 25.0 °C. Vogt et al. (18) used this information to propose a mechanism in which HOBr(g) is scavenged by the aerosol and converted to BrCl(aq) and Br2(aq) through the following aqueous-phase reactions (eqs 1-3);
HOBr(aq) + Cl- + H+ h BrCl(aq) + H2O
(1)
BrCl(aq) + Br- h Br2Cl-
(2)
Br2Cl- h Br2(aq) + Cl-
(3)
This mechanism relies on aerosol acidity (19), the high abundance of Cl- in nascent sea-salt aerosols, and a very rapid reaction to generate BrCl(aq) in eq 1. Kirchner et al. (20) and Crowley et al. (21) presented experimental evidence for efficient halogen release by reactive uptake of HOBr onto alkali halide-doped ice surfaces and aqueous solutions containing Cl- and Br-. In short, bromine atoms have been implicated in the observed ozone depletion in the Arctic. However, the initiation step that converts sea salt bromide to photolyzable bromine remains ambiguous. Thermodynamic and kinetic data of the aqueous-phase reactions are needed to elucidate the mechanism of bromine release either from sea-salt aerosols, snowpack, or sea ice. In the present work, studies of two equilibria involving BrCl(aq) (eqs 1 and 4)
BrCl(aq) + Cl- h BrCl2-
(4)
and the equilibria between Br2(aq), Br-, and Br3- are reported as a function of temperature. Many other species, including Br2Cl-, Cl2(aq), Cl3-, Br2Cl-, HOCl, OCl-, and OBr-, must be considered in these equilibria. In addition, the rate of BrCl(aq) hydrolysis (measured by our pulsed acceleratedflow method) is shown to be even faster than the previous (17) conservative estimate. The reversible formation of BrCl2represents an additional reservoir of bromine that can be converted to Br2.
Introduction
Experimental Section
Observations in the Arctic troposphere show that boundary layer ozone is episodically depleted from normal levels (∼40 ppb) to low levels ( 108 M-1 s-1 c
k6 ) 7.7 × 109 M-1 s-1 c kh1 ) (3.0 ( 0.4) × 106 s-1 b k-h4 ) 1.32 × 106 M-2 s-1 e k-7 ) 1.55 × 103 M-1 s-1 f
kh2 ) 97 s-1 g k-h2 ) 1.6 × 1010 M-2 s-1 g kh3 ) 22 s-1 h k-h3 ) 2.14 × 104 M-2 s-1 h
a All equilibrium constants and rate constants are reported at 25.0 °C and µ ) 1.0 M unless otherwise indicated. b This work. c Ref 17. d Bell, R. P.; Pring, M J. Chem. Soc. A 1966, 2, 1607-1609. e Ref 30. f Ref 31. g Ref 26 at µ ) 0.5 M (Kh2 ) 6.2 × 10-9 M2 corrected to µ ) 1.0 M). h Ref 27 at µ ) 0.5 M (Kh3 ) 0.96 × 10-3 M2 corrected to µ ) 1.0 M). i Ref 29. j Gerritsen, C. M.; Margerum, D. W. Inorg. Chem. 1990, 29, 2758-2762. k Gerritsen, C. M.; Gazda, M.; Margerum, D. W. Inorg. Chem. 1993, 32, 5740-5748.
TABLE 2. Spectral Characteristics of Halogen and Interhalogen Species species BrCl2BrCl(aq) Br2ClBr2(aq) Br3Cl2(aq) Cl3BrHOBr(aq) HOCl(aq) OBrOCl-
λ (nm)
E (M-1 cm-1)
343 232 343 232 381 245 232 390 232 186 362 266 232 325 232 290b 180 ( 20b ∼8800 10 400 ( 200b 12 100 ( 600b 100c ∼58 100d ∼100 332e ∼81 362f ∼31
a This work. b Ref 17. c Soulard, M.; Block, F.; Hatterer, A. J. Chem. Soc. Dalton Trans. 1981, 2300-2310. d Anbar, M.; Dostrovdky, I. J. Chem. Soc. 1954, 1105-1108. e Troy, R. C.; Margerum, D. W. Inorg. Chem. 1991, 30, 3538-3543. f Furman, C. S.; Margerum, D. W. Inorg. Chem.1998, 37, 4321-4327.
K7, Kh3, and our experimental data for Kh1. A plot of ln K1 vs (1/T) (Figure 3a) is linear and gives ∆H1° ) -9.9 ( 0.3 kJ mol-1 and ∆S1° ) -22 ( 1 J K-1 mol-1. This temperature dependence causes a small increase in the BrCl2-/BrCl(aq) ratio at lower temperatures.
FIGURE 1. Spectrophotometric data for the determination of K1 ) [BrCl2-]/([BrCl(aq)][Cl-]) as a function of temperature. Conditions: 1.00 cm cell, [BrCl]T ) 3.84 × 10-5 M, [H+] ) 1.00 M, and µ ) 1.0 M. Open circles are experimental data, and solid lines are curvefitting results for EBrCl2 ) (3.72 ( 0.07) × 104 M-1 cm-1 and EBrCl ) 650 ( 700 M-1 cm-1 after taking into consideration contributions from 10 other species in equilibrium. Temperature Dependence of Br2 + Br- h Br3- Equilibrium. As the pH of the BrCl solutions increases, more Brforms and the possibility of absorbance contributions due to trace levels of Br3- becomes a concern. Values of the equilibrium constant K4 ) [Br3-]/([Br2(aq)][Br-]) could not be found as a function of temperature, so the dependence was measured. A 9.30 × 10-6 M [Br2]T solution was made by adding a small amount of Br2(l) into an acidic solution containing excess bromide ([Br-] ) 0.0585 M, p[H+] ) 2.0, µ ) 1.0 M). The absorbance at 266 nm (A266) is given in eq VOL. 35, NO. 6, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 2. Species distribution at 25.0 °C in accord with Scheme 1 under conditions of [BrCl]T ) 4.88 × 10-5 M, [Cl-] ) 1.0 M, and µ ) 1.0 M. 8, and values of K4 at each temperature are calculated by use of
A266 )
Br3K4[Br-][Br2]T 1 + K4[Br-]
FIGURE 3. (a) Temperature dependence of BrCl/BrCl2- equilibrium. ∆H1° ) -9.9 ( 0.3 kJ mol-1 and ∆S1° ) -22 ( 1 J K-1 mol-1. Conditions: [BrCl]T ) 3.84 × 10-5 M, [H+] ) 1.00 M, and µ ) 1.0 M. (b) Temperature dependence of Br2/Br3- equilibrium. Conditions: [Br2]T ) 9.30 × 10-6 M, [Br-] ) 0.0585 M, p[H+] ) 2.0, and µ ) 1.0 M. Dark circles are experimental data, and the solid line is curve-fitting result for ∆H4° ) -5.7 ( 0.7 kJ mol-1 and ∆S4° ) 4 ( 3 J K-1 mol-1 at 25 °C, and ∆C°p4 ) (-3.1 ( 0.6) × 102 J K-1 mol-1. The plot is not linear indicating that ∆H4° and ∆S4° change with temperature because of the existence of two hydrates of bromine: Br2(aq) and H2OBr2(aq). Their relative concentrations vary with temperature.
(8)
As shown in Figure 3b, ln K4 vs (1/T) is not linear, indicating that a nonzero ∆C°p4 term needs to be incorporated (eq 9). The curve in Figure 3b is the best fit of the data to
ln K4 )
-∆H°4 ∆S°4 ∆C°p4 298.15 T 1+ + - ln RT R R T 298.15
(
) (9)
The result gives ∆H4° ) -5.7 ( 0.7 kJ mol-1 and ∆S4° ) 4 ( 3 J K-1 mol-1 at 25.0 °C, and ∆C°p4 ) (-3.1 ( 0.6) × 102 J K-1 mol-1. Negative ∆C°p values have also been observed for Br2 and Cl2 hydrolysis equilibrium constants, Kh2 and Kh3 (2628). This has been attributed to the existence of two different hydrated forms of these halogens in water, where the relative concentrations change with temperature. By contrast, the ∆C°p value for bromine chloride is approximately zero. We attribute this to a more constant and larger degree of hydration of the more polar BrCl molecule (29). Temperature Dependence of Bromine Chloride Hydrolysis Equilibrium and Species Distribution with pH. The BrCl(aq) hydrolysis equilibrium is interconnected with many other equilibria in accordance with Scheme 1. As pH increases, BrCl(aq) hydrolyzes, and concentrations of HOBr, HOCl, Br-, and OCl- as well as other species such as Br2 and Br2Cl- increase significantly as shown in Figure 2. Absorbance measurements at 232 nm were taken for each BrCl2-/BrCl(aq) solution from p[H+] 3.2 to p[H+] 6.5 in 1.0 M [Cl-] with an initial [BrCl]T concentration of 4.88 × 10-5 M. The hydrolysis equilibrium constant Kh1 ()[HOBr(aq)][Cl-][H+]/[BrCl(aq)]) was evaluated by solving the 12 simultaneous equations (10 equilibrium constants and eqs 6 and 7) with use of a SigmaPlot program (25) to fit the experimental data. All 12 species in eqs 6 and 7 were considered. Values for K3, K4, Kh2, K5, K7, 1130
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FIGURE 4. Spectrophotometric data for the determination of Kh1 ) [HOBr(aq)][Cl-][H+]/[BrCl(aq)] as a function of temperature. Postmixing conditions: 0.962 cm cell, [BrCl]T ) 4.88 × 10-5 M, [phosphate]T ) 0.0063 M, [Cl-] ) 1.0 M, and µ ) 1.0 M. Open circles are experimental data and solid lines are curve-fitting results. BrCl(aq) hydrolysis has little or no temperature dependence with Kh1 ) (1.3 ( 0.1) × 10-4 M2 from 25.0 to 5.0 °C. , and KHOCl are known, and Khl was obtained from 19 KHOBr a a data sets with variable p[H+] by use of the previously determined values of K1, BrCl2, and BrCl. The 19 data sets plus the 25 data sets with variable [Cl-] were then combined and iterated until values for K1, BrCl2, BrCl, and Kh1 converged. This procedure gives Kh1 ) (1.34 ( 0.09) × 10-4 M2 at 25.0 °C and the same values obtained previously for K1 and the molar absorptivities. Figure 4a shows the fit of these data to experimental absorbance values obtained at 232 nm. Figure
TABLE 3. Data for Observed Slow Reaction from 1 to 20 s in the BrCl(aq) Hydrolysis Experimentsa p[H+]
∆A (mAU)b
kobsd (s-1)
kpred (s-1)
3.64 3.85 4.31 4.58 4.91 5.29 6.01
29 37 42 52 41 24 6
9.5 3.9 2.4 1.4 1.1 0.8 0.5
5.6 3.7 1.5 0.90 0.5 0.4 0.3
a Post-mixing [BrCl] ) 4.88 × 10-5 M, [Cl-] ) 1.0 M, [phosphate] T T ) 0.0063 M, µ ) 1.0 M, 232 nm, 25.0 °C, 0.962 cm cell. b ∆A ) A02 - Afinal.
SCHEME 2. Cl2(aq) Pathway for Conversion of HOBr and BrCl to HOCl and Br-
FIGURE 5. Two-step process in BrCl(aq) hydrolysis experiment. Premixing conditions: solution 1: [BrCl]T ) 9.76 × 10-5 M, [Cl-] ) 1.0 M, p[H+] ) 2.84, µ ) 1.0 M; solution 2: [phosphate]T ) 0.0125 M, [Cl-] ) 1.0 M, µ ) 1.0 M, p[H+] ≈ 5.7. Post-mixing conditions: [BrCl]T ) 4.88 × 10-5 M, [phosphate]T ) 0.0063 M, [Cl-] ) 1.0 M, p[H+] ) 4.31, µ ) 1.0 M, 25 °C, 0.962 cm cell. The fast step (absorbance change from A01 to A02) corresponds to the hydrolysis of BrCl(aq) to HOBr(aq) and rapid equilibria adjustment among BrCl/BrCl2-, HOBr, and Cl2/Cl3-. The slow step (absorbance change from A02 to Afinal) corresponds to the shift of above equilibria toward HOCl(aq) formation through Cl2(aq) hydrolysis. 2 is the species distribution from pH 0 to pH 7 calculated using the 12 equations and the evaluated constants. At low pH, BrCl(aq) and BrCl2- are the dominant species. As pH increases, BrCl(aq) and BrCl2- hydrolyze, and significant amounts of HOBr, HOCl, Br-, and OCl- form. As the temperature decreases, the values of Kh1 are nearly constant [(1.34 ( 0.09) × 10-4 M2 at 25.0 °C, (1.44 ( 0.08) × 10-4 M2 at 15.0 °C, (1.24 ( 0.09) × 10-4 M2 at 5.0 °C)] as shown for the data in Figure 4, panels b and c. We conclude that BrCl(aq) hydrolysis has little or no temperature dependence with Kh1 ) (1.3 ( 0.1) × 10-4 M2 from 25.0 to 5.0 °C. The fit of the data at 15.0 and 5.0 °C include the known effect of temperature on K3, K7, Kh2, Kh3, and our experimental data for K1 and K4. Temperature dependencies of K5, KHOBr , a and KHOCl have little effect on the evaluations. a Partial Bromine Chloride Hydrolysis. A concentration jump experiment was performed by mixing BrCl2-/BrCl(aq) solutions with phosphate buffer solutions in the APPSF instrument. Under post-mixing conditions of 1.0 M [Cl-] and p[H+] > 6.5, BrCl(aq) hydrolyzes completely, and 100% of the BrCl2- signal is lost within the dead time (3 ms) of the instrument. This observation is in good agreement with the previous report (17) that BrCl(aq) hydrolysis is a very rapid process with kh1 greater than 105 s-1. However, under conditions where BrCl hydrolysis is not complete (1.0 M [Cl-] and p[H+] 3-6), a two-step process is observed. In Figure 5 at p[H+] 4.31, 50% of the initial BrCl2-/BrCl(aq) is hydrolyzed at equilibrium (g10 s), but a 46% drop in absorbance occurs within a few milliseconds while the remaining 4% change takes several seconds. Spectral scans (220-400 nm) taken from 125 ms to 5 s confirm that the absorbance decay of the slower reaction is due to the loss of BrCl2-, which serves as an indicator for the progress of other reactions. The slow decay is p[H+] dependent with observed first-order rate constants that decrease as the p[H+] increases (Table 3). The observed final absorbance also decreases as the p[H+]
increases because the extent of hydrolysis is greater. To evaluate accurately the BrCl(aq) hydrolysis equilibrium and rate constants, it is important to understand the nature of the slow reaction. Equilibrium calculations of the system in accordance with Scheme 1 are given as a function of p[H+] in Figure 2. At p[H+] 4.31, it is clear that significant amounts of HOCl(aq) and Br- as well as HOBr(aq) form at equilibrium. Therefore, after the rapid BrCl(aq) hydrolysis to HOBr(aq), there must be other pathways by which HOBr(aq) and [BrCl]T are slowly converted to HOCl(aq) and Br-. One possibility would be the direct reaction in eq 10 where water reacts at the chlorine atom rather than the bromine atom:
BrCl + H2O h BrClOH2 h Br- + HOCl + H+ (10) The equilibrium constant for this reaction (Kh4) can be evaluated from Kh1 and K7 as given in Table 1, where K7 is obtained from thermodynamic values for each species (30). Earlier kinetic studies (31) measured rate constants at 25.0 °C for the production of HOBr from HOCl + Br- (1.55 × 103 M-1 s-1), from H+ + HOCl + Br- (1.32 × 106 M-2 s-1), and from H2PO4- + HOCl + Br- (3.0 × 103 M-2 s-1). Hence, we can show that the sum of the rate constants for the reverse reaction in eq 10, under the conditions used for Figure 5, gives a value of 1.63 × 103 M-1 s-1. From this value and Kh4, it can be shown that the rate constant for the forward reaction in eq 10 would be 2.9 × 10-2 s-1. The ratio of [BrCl]/([BrCl2-] + [BrCl]) is 1/4.8 so that the rate constant for the conversion of BrCl2- to HOCl + Br- + H+ cannot exceed 6.0 × 10-3 s-1 for this pathway whereas the observed rate constant is 2.4 s-1. Instead, an indirect conversion of BrCl(aq) to Cl2(aq) and then to HOCl(aq) is proposed (Scheme 2) to explain the two-step process. The high concentration of Cl- used favors this pathway. In Scheme 2, the hydrolysis of BrCl(aq) to HOBr(aq) is very rapid when BrCl2-/BrCl(aq) solutions are mixed with VOL. 35, NO. 6, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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phosphate buffer solution. In a subsequent reaction, the equilibrium between BrCl(aq), Cl-, and Cl2(aq) adjusts rapidly (k8 ) 7.0 × 103 M-1 s-1 and k-8 ) 7.7 × 109 M-1 s-1) (17). These steps are complete within the dead time of the stopped-flow instrument and cause the large absorbance drop shown in Figure 5 due to the conversion of BrCl2-/BrCl(aq) to HOBr(aq), with small amounts of Cl2(aq) and Br- also present. After this, a slow and reversible Cl2 hydrolysis (kh3 ) 22 s-1, k-h3 ) 2.14 × 104 M-2 s-1) (27) takes place. Phosphate buffer accelerates this reaction in both directions. (Values of kp ) 2.0 × 104 M-1 s-1 and k-p ) 17 M-2 s-1 are calculated from the Brønsted-Pedersen relationship reported in ref 27.) Hydrolysis of Cl2(aq) results in the formation of HOCl(aq) and the further loss of BrCl2-/BrCl(aq) as the equilibria are shifted such that the initial [BrCl]T is now distributed among BrCl(aq), BrCl2-, HOBr(aq), Cl2(aq), and HOCl(aq). As a consequence, the absorbance drops further. At p[H+] 4.31, this corresponds to the observed 4% signal change from 0.1 to 10 s. The Cl2(aq) hydrolysis step is rate-determining. The rate constant (kpred) for the loss of BrCl2- via the Cl2(aq) hydrolysis step under reversible conditions is expressed by
kpred )
(
-
K8[Cl-]
KaH2PO4 kh3 + kp [H2PO4-]T H2PO4+ [Br ] Ka + [H ]
(
-
Kh1
K8[Cl-]
)
)
K8K5[Cl-]2
+
+ + [H+][Cl-] [Br-] [Br-] [H+] k-h3[H+][Cl-] + k-p [H2PO4-]T[Cl-] (11) H2PO4Ka + [H+] 1 + K1[Cl ] +
In the process of Cl2(aq) hydrolysis and the equilibria readjustment, [Br-] increases and kpred decreases as the reaction progresses. Our kinetic measurements were made under conditions where [Br-] is close to its equilibrium level so that we essentially have a concentration-jump relaxation experiment. The relaxation rate constants (kobsd in Table 3) increase with decrease in p[H+] because the fraction [BrCl]/ ([BrCl2-] + [BrCl] + [HOBr]) increases and the rate constants for the reverse reactions increase due to [H+] and [H2PO4-] interaction with [HOCl] and [Cl-]. We selected the observed absorbance at the midpoint of the relaxation curve to calculate [Br-] in order to solve for kpred in eq 11. The predicted rate constants (Table 3) are somewhat larger than kobsd, but are sufficiently close to support the mechanism via Cl2 that is given in Scheme 2. Many of the rate constants used for kpred were measured at µ ) 0.50 M rather than µ )1.00 M used in this work. This may account for some of the differences in kpred and kobsd. It should be noted that above p[H+] 6 the interconversion from HOBr + Cl- to HOCl + Bris still necessary and should be even slower. However, the absorbance of BrCl2- was used as the indicator for these reactions and it is no longer present at higher pH. BrCl(aq) Hydrolysis Kinetics by PAF-IV. The kinetics of BrCl(aq) hydrolysis to HOBr(aq) at 25.0 °C were measured on the PAF-IV instrument by following the loss of [BrCl]T at 232 nm when BrCl solution was mixed with phosphate buffer solution. The following conditions were used: premixing BrCl solution: [BrCl]T ) 5.48 × 10-5 M, [Cl-] ) 2.0 M, p[H+] ) 2.3, and µ ) 2.0 M; premixing phosphate buffer solution: [phosphate]T ) 0.0125 M, [Cl-] ) 2.0 M, p[H+] ) 6.6, and µ ) 2.0 M; post-mixing solution: p[H+] ) 6.4, [Cl-] ) 2.0 M, and µ ) 2.0 M; cell path ) 2.05 cm. High chloride concentration was used to suppress the reaction rate. A very rapid BrCl(aq) hydrolysis to HOBr(aq) was observed under the above conditions. Data analysis using eq 5 yielded an observed first-order rate constant (kr) equal to (3.5 ( 0.5) × 105 s-1 (Figure 6). This large value causes the reaction to 1132
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FIGURE 6. PAF-IV kinetic data for the determination of the BrCl(aq) hydrolysis rate constant, kh1. Open circles are experimental data, and the solid line is curve-fitting result. kh1 ) kr/0.116 ) (3.0 ( 0.4) × 106 s-1 because only 11.6% of [BrCl]T is present as BrCl(aq) under the experimental conditions. Premixing: [BrCl]T ) 5.48 × 10-5 M, [Cl-] ) 2.0 M, p[H+] ) 2.3, µ ) 2.0 M; [phosphate]T ) 0.0125 M, [Cl-] ) 2 M, p[H+] ) 6.6, and µ ) 2.0 M. Post-mixing: [BrCl]T ) 2.74 × 10-5 M, [phosphate]T ) 0.0063 M, [Cl-] ) 2.0 M, p[H+] ) 6.4, µ ) 2.0 M. occur very close to the inlet jets at the center of the cell. A 14% error is associated with this rate constant as the performance limit of the PAF-IV instrument is approached. Under the experimental condition of 2.0 M [Cl-], only 11.6% [BrCl]T is present as BrCl(aq). Taking this into account, the actual rate constant kh1 of BrCl(aq) hydrolysis is (3.0 ( 0.4) × 106 s-1 at 25 °C, which is 30 times larger than the minimum value of 105 s-1 estimated by Wang et al. (17). Hence, eq 1 is even more strongly preferred than previously thought as a source of BrCl(aq) because the rate constant of BrCl(aq) formation from HOBr, Cl-, and H+ (k-h1) is directly related to BrCl(aq) hydrolysis rate constant kh1 with k-h1) kh1/Kh1. The reverse reaction between HOBr, Cl-, and H+ has a rate constant of 2.3 × 1010 M-2 s-1, so that in seawater, where [Cl-]/[Br-] ) 700, the formation of BrCl is 1000 times faster than the formation of Br2 from HOBr, Br-, and H+ at 25.0 °C. Recent studies (32) with a new PAF instrument, designed to operate over a wider range of temperatures, gives a value of kh1) (1.75 ( 0.05) × 106 s-1 at 0.0 °C for the BrCl(aq) hydrolysis rate constant. We conclude that BrCl(aq) formation is an important pathway in the uptake of HOBr at lower temperatures as well as at room temperature. Subsequent very rapid reactions of BrCl(aq) with Br- will form Br2(aq). Henry’s law constants for Br2 and BrCl (29) and the equilibrium constants for Br2Cl- and BrCl2- indicate that Br2 is 4.0 times more volatile than BrCl from a 1.0 M Cl- solution at 0 °C. Photolysis of Br2(g) can initiate chain reactions (8-13) that lead to loss of tropospheric O3 as observed at Arctic sunrise.
Acknowledgments We thank the National Science Foundation for financial support (ATM-9631572 and CHE-9818214).
Supporting Information Available Listings of data (7 pages). This material is available free of charge via the Internet at http://pubs.acs.org.
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Received for review June 14, 2000. Revised manuscript received January 3, 2001. Accepted January 11, 2001. ES001380R
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