Kinetics and mechanism of the reaction of sodium azide with hypochlorite in aqueous solution
Introduction
The industrial production of sodium azide (NaN3) has surged over the past decade to meet the demand for automobile airbag inflator propellant. However it is highly toxic—even a cursory search of the Internet reveals that sodium azide toxicity is comparable to that of sodium cyanide when ingested [1]. The United States National Highway Traffic Administration (NHTSA) mandated that new passenger vehicles sold in the United States after 1996 have both driver-side and passenger-side air bags. As a result, sodium azide demand quickly rose to exceed 5 million kg in the U.S. by 1995 and has continued to rise, although the trend is now leveling off with the advent of azide-free inflators [2], [3]. An inflator module typically contains compressed disks or pellets of ≈60% (w/w) NaN3, blended with other ingredients, and packed into a sealed metal canister [4], [5]. Each driver-side inflator contains approximately 50–80 g NaN3, while the larger passenger-side inflator contains approximately 250 g [6]. The actual amount depends on the specific inflator, but the total installed in U.S. vehicles alone is now on the order of 50 million kg [7]. The smooth, rapid thermal decomposition characteristics, the high specific nitrogen content (65 mol%), and the long shelf life make sodium azide an attractive propellant material [6], [8].
Sodium azide has other uses. In the 1970s, it was used in some registered pesticide formulations, mainly for crops, and consequently discarded commercial products, off-specification product, container residues and spill residues were listed as hazardous waste by the United States Environmental Protection Agency (EPA) [9]. Lately, interest in sodium azide pesticide formulation has been revived as a replacement for methyl bromide [10]. Accordingly, sodium azide toxicity was recently on the agenda of an EPA Human Studies Internal Review Board [11]. Sodium azide is used as a preservative in certain laboratory reagents, samples and clinical fluids, and in 1989, there was an alert from the Food and Drug and Administration and the Center for Disease Control (CDC) concerning the need to rinse out sodium azide preservative in certain water filters prior to use [12]. Azide can also combine with metals to form explosive metal–azide complexes. The accumulation of azide in laboratory apparatus and drains, where it can react with lead or copper-containing fixtures, has caused explosions when routine maintenance work has been attempted [13].
The azide ion is readily protonated in aqueous solution (pKa = 4.65) to yield volatile hydrazoic acid (HN3), which is itself toxic, so the atmospheric fate of azide substance is also of interest and has recently been described [5], [14], [15], [16].
Sodium azide is highly soluble, which implies that releases into the environment could potentially migrate into sewers, streams, lakes and groundwater systems. And in fact, sodium azide has been found in groundwater at three manufacturing sites in three states, resulting in a multi-million dollar civil and criminal settlement [17].
Because of its ready availability and high toxicity, sodium azide has become a chemical of interest for the Department of Homeland Security [18], the CDC [19], and the EPA Water Supply Security Division [20]. Since it is possible that azide could be found in drinking water supplies it is of interest to know how it would behave in the presence of hypochlorite, which is commonly used in potable water treatment systems in the United States.
Oxidation by hypochlorite could also potentially be used as a treatment for much more concentrated azide-containing waste, but the results of this work lead us to recommend against such practice.
In a definitive series of mechanistic studies Margerum's group has shown that the reaction of hypochlorite with a range of nucleophiles, including CN−, I−, Br−, Cl−and SO32−, appears to proceed by way of Cl+ transfer, and not via oxygen atom transfer, as had long been thought [21], [22], [23], [24], [25], [26], [27]. Ignoring the existence of acid catalysis, the common rate law for reaction of hypochlorite with any of these nucleophiles, X−, can be represented by:
Thus, oxidation can proceed via HOCl and OCl−, and the reaction is first-order in both hypochlorite and X−. For all X−, except SO32−, the ratio is greater than 106. Therefore, the HOCl route will dominate in all natural waters (pKa HOCl = 7.31). In the case of SO32−, the ratio falls to 104 and so the OCl− route could become significant at pH ≥ ≈11.3. The value of kHOCl (M−1 s−1) decreases with the nucleophilicity of X−: CN− (1.22 × 109) > SO32− (7.6 × 108) > I− (1.4 × 108) > Br− (1.55 × 103) > Cl− (≤0.16) [28]. Since azide, the subject of this paper, is a pseudohalide with a nucleophilicity similar to that of bromide, it is of interest to compare the behavior of these two species in particular [29].
The first step of the HOCl route involves Cl+ transfer to an anion.whereas the first step of the much slower OCl− pathway formally involves an encounter between two anions, and so may be acid assisted.
Nevertheless, both pathways lead to a common intermediate, XCl, which may be quite stable, e.g., cyanogen chloride (CNCl) or bromochloride (BrCl) [28], [30]. XCl is subsequently lost through either alkaline hydrolysis:or by reaction with a second X− [30].
In the case of X− = N3−, this would lead to (N3)2, which would rearrange to 3N2. In fact, we will show later that this path is the dominant route for azide, unlike the halides, presumably because the production of N2 is thermodynamically so favorable.
Here we describe the kinetics and mechanism of the hypochlorite/azide reaction in aqueous solution and propose a mechanism to explain the observed rate laws. We use this information to estimate the lifetime of azide in chlorinated drinking water and to comment on the possible treatment of concentrated aqueous azide-containing waste with hypochlorite.
Section snippets
Materials
Water was first deionized (Calgon) and then distilled (Barnstead; conductivity <1 μS cm−1). Three sources of sodium azide were used (Aldrich 99.99+%; Mallinckrodt Practical; Acros 99%); no difference was observed in their electronic spectra or their kinetic behavior. We also tested two sources of sodium hypochlorite, which was either purchased (Clorox, ≈0.44 M) or prepared by bubbling chlorine (Matheson, 99.99%) into 0.2 M sodium hydroxide until Cl2 could first be seen in the headspace. The excess
Stoichiometry
A series of experiments described below support the following overall reaction stoichiometry.
N2 evolution was quantified by water displacement from two sealed vials connected in series. A mixture of 0.132 mmol OClT− and 1 mmol [N3T−] (40 mL; pH 7.0; 0.05 M phosphate) was allowed to react in a 40-mL septum-sealed vial (zero headspace), which was connected by a needle to an identical vial filled with water. As N2 accumulated in the second vial, water escaped through another
Discussion
It is of interest to use our newly acquired kinetic data to estimate the lifetime of azide in potable water treatment and distribution systems. This is governed by the chlorine concentration, for which two extremes exist. At the point of chlorination, the maximum free chlorine is dictated by its solubility (≈0.1 M) whereas further downstream, the residual chlorine level is normally about 1.5 μM [44], [45]. Assuming pH 7, and an infinite supply of chlorine (that upon dissolution yields ≈50% OClT−;
Acknowledgements
We would like to thank the EPA NEIC scientists and staff, particularly John Reschel, Beth Mishalanie, Eric Nottingham and the library staff for their help. One of us, EAB, would like to thank Ms. Diana A. Love, former Director, U.S. EPA NEIC, and Mr. Les Ogden for providing the financial and administrative support that made possible his sabbatical leave from the University of Arizona.
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