As the first step, we studied the reactivities of NaOCl with different buffer components. NaOCl has a λmax of 292 nm with a molar extinction coefficient of 360 M–1 cm–1, (35) allowing for its monitoring at mM concentrations by ultraviolet–visible (UV–vis) spectrophotometry. Specifically, we used buffer components at 20 mM, which is at the lower end of the buffer concentrations commonly used. We started by studying the reaction of NaOCl (2 mM) at pH 7.4 and 37 °C with HEPES, which has no meaningful absorption in the region 250–400 nm. Upon incubation for 1 min, the peak at 292 nm disappeared to an undetectable level with concomitant appearance of a peak at 260 nm, indicating consumption of NaOCl by HEPES and the formation of a new chromophoric compound (Figure 2a). Looking at the structure of HEPES, one can readily see several oxidizable functional groups, including amino and hydroxyl groups. NaOCl is known to readily oxidize an amino group to chloramine. (36) Incidentally, chloramine is known to have a λmax close to 260 nm (λmax of 244 nm for the product of ammonia and NaOCl). (35) At the 5- and 30 min time points, the intensity at 260 nm also further decreased, presumably because of degradation or further reactions by chloramine. Indeed, chloramines are known to have stability issues. (37−40) HEPES has tertiary amines, which can form quaternary chloramine ions after reaction with NaOCl. (41) Such quaternary chloramine species can rapidly degrade, leading to a decrease in chloramine concentration. (40,41) Without the need to study the detailed chemistry, one outcome is unambiguous: there is strong interference from HEPES in hypochlorite-related studies.

After observing such a fast consumption of NaOCl by HEPES, we next studied the same with Tris-Cl, MES, ammonium acetate, citrate, and PBS. Indeed, NaOCl showed spectral changes upon the addition of each of the buffer compounds tested. Starting with Tris-Cl, the reaction seems to be very fast with the total disappearance of the peak at 292 nm (NaOCl) at the 1 min point and concomitant appearance of a peak at 250 nm, presumably corresponding to the chloramine product (Figure 2b). (35) Tris contains a primary amino and three hydroxyl groups, both of which are known to react with NaOCl, leading to the formation of chloramines and alkyl hypochlorite, respectively. (36,42) The spectrum of the Tris and NaOCl reaction mixture did not show any change within from 1 to 30 min (Figure 2b). This is different from that of the HEPES reaction and indicates a higher level of stability for the chloramine product from Tris. Such a difference in degradation propensity is consistent with literature stability reports of chloramines of primary and tertiary amines. (40,43)

As expected, MES, ammonium acetate, and citrate buffers all showed reactivity with NaOCl (Figure 2). All three buffers (MES, ammonium acetate, and citrate) showed NaOCl consumption within 1 min. In reactions with MES and ammonium acetate, a new peak at around 250 nm was observed, indicating the formation of the chloramine product. Further spectral changes were observed at the 5 and 30 min points, indicating degradation for the chloramine formed. Similar to HEPES, MES also has a tertiary amine, which can form quaternary chloramine ions after reaction with NaOCl (41) and rapidly degrade, leading to decreased concentration of the chloramine. (40,41) Looking at the structure of citrate, one would not expect to see chloramine formation because it lacks an amino group. Therefore, in the citrate and NaOCl reactions, we only observed a decrease in intensity at 292 nm without the concomitant formation of a new peak. The reaction of NaOCl with citrate has been reported to lead to the formation of alkyl hypochlorite, (42,44,45) which is still very reactive. (42) Further, the reactivity of alkyl hypochlorite can be different from that of NaOCl. (42,46−48) We emphasize that we do not intend to study the detailed reaction mechanism(s) and products from these buffer components. We are focused on showing the fact of buffer interference and raising a cautionary note so that others can pay particular attention to this issue when conducting their own ROS-related experiments, when applicable.

With the findings of strong reactivity of NaOCl with the various buffer compounds, next we were interested in understanding the reaction kinetics to put the issue of possible interferences by various buffer components into a proper context. Because of the fast reaction of NaOCl with buffer components, we decided to use stopped-flow spectrophotometry for subsequent reaction rate studies. Briefly, the reaction of 1 mM NaOCl and 20 mM each buffer compound was studied (Figure S1) by monitoring absorption changes at 325 nm, which was determined based on the observed signal-to-noise ratio. At 20 mM HEPES and 1 mM NaOCl, complete NaOCl consumption was observed in about 0.1 s (Figure 3a). The pseudo-first-order rate constant was determined to be 177 s–1 by a single exponential decay (Figure S2 and Table 1). The kinetics traces showed an initial plateau phase of about 1.73 ms. Though the reaction mechanism was not the focus of this study, we did consider various possibilities for the initial lag period. Since reaction progression was monitored by measuring the disappearance of hypochlorite at 325 nm, the lag period could mean the formation of a very reactive intermediate species with the same UV characteristics. The most likely species is chloramine of HEPES. However, chloramine is known to have a UV λmax at about 244 nm. (35) Further, chloramine is much less reactive than hypochlorite. (38) If one uses the reaction of an amine with chloramine or NaOCl as a gauge, the reactivity difference is on the order of 109 with the second-order rate constant being 0.32 M–1 s–1 for the reaction with chloramine (38) and 108 M–1 s–1 for the reaction with hypochlorite. (49) Therefore, chloramine is unlikely to be the contributing species. We did not further pursue other factors, because that would be beyond the scope of this study. The finding of the rapid and complete consumption of NaOCl within 100 ms by HEPES clearly indicates the need to avoid this buffer when studying hypochlorite concentrations.

In an attempt to determine the second-order rate constant, we first measured the pseudo-first-order rate constants by using HEPES in excess: 1 mM NaOCl with HEPES at 25, 30, 35, 40, 45, 50, and 60 mM, respectively (Figure S3). The plan was to plot the pseudo-first-order rate constants against HEPES concentrations to determine the second-order rate constants. Much to our surprise, the pseudo-first-order rate constant decreased with increasing HEPES concentration (Figures S3 and S4), contrary to our expectations. In our years of conducting similar kinetic experiments, we have never encountered this problem. One possible reason for this decreased pseudo-first order rate constant with increasing HEPES concentration could be due to self-association, which could change the reactivity of the nitrogen atom toward NaOCl. The piperazine moiety present in HEPES is known to be prone to self-association, though there is no report of self-association of HEPES. (50) This self-association could be the reason for the decreased pseudo-first-order rate constant with elevating HEPES concentration (Figure S4). Regardless of the specific rate constant, HEPES consumes NaOCl in less than 1 s at all HEPES concentrations (Figure S3). Additionally, a literature value determined using an indirect method also indicates the rapid consumption of NaOCl by HEPES (4400 M–1 s–1). (51) Without the need to study further details, one outcome is unambiguous: there is strong interference from the HEPES in NaOCl-related studies.

MES showed a slower reaction rate compared to HEPES but still consumed all of the NaOCl in about 1 s under the same conditions (Figure 3b). The kinetic traces for MES seem to show two phases, probably indicating secondary reaction(s). As a result, the pseudo-first-order rate constant was determined as 14 s–1 for the first phase by using a double-exponential method of 20 mM MES and 1 mM NaOCl (Figure S5 and Table 1). For the second phase, tertiary amines are known to form quaternary chloramine ions after reaction with NaOCl. (41) Such a quaternary chloramine species can undergo degradation, leading to the formation of an aldehyde and a secondary amine after hydrolysis. (40,41) The secondary amine formed in this reaction can also react with NaOCl; this could be a reason for the second phase when the NaOCl concentration changes. As discussed before, we did not further pursue the issues of degradation products, as that would be beyond the scope of this study. Regardless, MES consumed all of the NaOCl within 1 s (Figures 3 and S6). We determined the pseudo-first-order rate constants by using MES in excess: 1 mM NaOCl with MES at 25, 30, 35, 40, 45, 50, and 60 mM, respectively (Figure S6). Then, the pseudo-first-order rate constant was plotted against MES concentration to yield the second-order rate constant of 958 ± 30 M–1 s–1 based on the slope (Figure S7).

For the reactions between NaOCl and Tris-Cl or ammonium acetate, respectively, the reaction was found to be complete, even before the first spectral scan by the stopped-flow instrument (Figure S1). The observation suggests that the reaction occurred within the mixing time of the stopped-flow instrument, which is about 2.2 ms (Figure 3), giving a pseudo-first-order rate constant of 2092 s–1 as the lower limit. This provides the lower end of the second-order reaction rate being about 1.0 × 105 M–1 s–1. For comparison, the reported second-order rate constant for reactions between a methyl amine and NaOCl is about 1.9–3.6 × 108 M–1 s–1. (49) Regardless of the specific number for the rate constant, all of the indications are that the reaction between NaOCl and Tris-Cl or ammonium acetate is very fast. For normal ROS-related experiments on time scales of min to h, there is no meaningful difference when the reactions between the buffer component and NaOCl are this fast because these buffer components can consume NaOCl within a few seconds.

The reaction kinetic information will inform the degree of interference, depending on the reaction in question. For example, the second-order rate constant of methionine oxidation by NaOCl is 3.7 × 108 M–1 s–1, (52,53) which is significantly (5 orders of magnitude) faster than the oxidation of HEPES. Then, one would not expect to see meaningful competitive reaction of HEPES with NaOCl unless HEPES is in excess by more than 5 orders of magnitude. On the other hand, boronate oxidation by NaOCl has a second-order rate constant 5.7 × 103 M–1 s–1, (54) which is similar to the second-order rate constant of HEPES with NaOCl (4.4 × 103 M–1 s–1). (51) Then, one would expect significant interference of boronate oxidation by HEPES because HEPES and boronate have comparable reactivity with NaOCl and yet HEPES is present in excess. Naturally, for a probe having a reaction rate with NaOCl that is slower than that with HEPES, one would expect severe interference by this buffer component.

The oxidation of a boronic acid moiety by peroxide is a commonly used approach for designing fluorescent probes for various ROS. For example, Chang and colleagues developed a number of boronate-based probes for the detection of ROS. (25) Such a reaction involves a peroxy anion attacking the boron atom with an open shell, followed by rearrangements leading to the formation of a hydroxyl group in the position of the boronic acid moiety (Scheme 1). Because of the widespread use of such boronic acid chemistry, we started by examining the effect of buffer components on the reaction of a boronate with NaOCl. We chose 4-acetylphenylboronic pinacolate ester (APBE) 1 as a model compound for this study. This selection was made because of its known rate constant 5.7 × 103 M–1 s–1 in its reaction with NaOCl and its simplicity. (54) Briefly, we examined the effect by adding a stoichiometric amount of NaOCl to a solution of APBE 1 (100 μM) in different buffers and analyzed the products using HPLC. We first conducted a control reaction by incubating APBE 1 with NaOCl in PBS without DMSO and observed the quick and quantitative conversion of APBE 1 to a new peak corresponding to the oxidized product 2 (Figure 4). Such results also help confirm that the decreased UV absorbance of NaOCl in PBS (relative to NaOCl alone; Figure 2f) was due to pH changes, not NaOCl consumption. Interestingly, when the same reaction was conducted in HEPES, MES, Tris-Cl, or ammonium acetate buffer by adding NaOCl to a solution of APBE 1 in the respective buffer, no product 2 formation was observed (Figure 4). Obviously, HEPES, MES, Tris-Cl, and ammonium acetate were able to react with NaOCl fast enough to prevent APBE 1 oxidation. Furthermore, the chloramine products from such reactions do not react with boronate within the time scale studied. Because HPLC was used to study reaction profiles, there is the question as to whether HPLC mobile phase (H2O, ACN, and trifluoroacetic acid (TFA)) might interfere with the outcome (Figure 4) due to the NaOCl reaction with mobile phase. Therefore, we conducted experiments with ACN and TFA and were able to rule out the interference from ACN or TFA during HPLC analysis. First, NaOCl consumption was complete within 1 s when Tris, HEPES, MES, or ammonium acetate was used (Figure 3). Therefore, by the time of HPLC injection, no NaOCl is expected to remain to interfere with reactions. The results in Figure 4c indicate so. For the reaction in PBS, a second-order rate constant is 5.7 × 103 M–1 s–1 for the reaction of boronate with NaOCl, (54) giving a first calculated t1/2 of 1.7 s at 100 μM each. The reaction kinetics indicate that 95% of the reaction will be completed in less than 60 s, which should lead to an almost complete consumption of APBE 1 by the time of the HPLC injection. Indeed, the reaction in PBS buffer showed complete conversion of APBE 1 to the product (Figure 4). Third, we incubated ACN with hypochlorite and saw no consumption of NaOCl (Figure S8). Fourth, we observed the lack of impact on APBE 1 integrity by a TFA-hypochlorite combination (Figure S9). We should note that the reaction kinetics of NaOCl are similar for APBE 1 and HEPES with the second-order rate constant being 5.7 × 103 M–1 s–1 and 4.4 × 103 M–1 s–1, respectively. (51,54) Incidentally, both are faster than many click reactions. (65,66) Presumably, because HEPES is present in large excess of the probe (boronate), the majority of the NaOCl was consumed by the buffer component, even if NaOCl was added after APBE 1 addition, leading to skewed results for ROS concentration determination.

Boronate-based probes are also being used for the detection of ONOO–, (60,61,67) because of the rapid reaction of boronate with ONOO–. (54) The second-order rate constant of boronate reaction with ONOO– reaction has been reported as 1.2 × 106 M–1 s–1. (54) However, there are the following additional complexities we needed to consider while working with ONOO–. First, the ONOO– is known to show pH-dependent decomposition. (68) Specifically, at pH 7.4 the ONOO– has a t1/2 of 1.9 s. (68) Second, ONOOH and ONOO– have significantly different reaction kinetics. (52,69) For example, the second-order rate constants of the reaction of ONOOH and ONOO– with methionine have been reported to be 2 × 103 M–1 s–1 and 2 × 10–1 M–1 s–1, respectively. (52,69) Third, boronate reacts 103-fold faster with ONOO– compared to NaOCl. (54) To analyze the interference from buffer molecules in ONOO– detection without the competitive reaction of boronate with ONOO–, we first added the ONOO– to the buffer solution followed by a stoichiometric amount of APBE 1 (100 μM). The reaction was analyzed using HPLC. As expected, quantitative product formation was observed only in PBS, whereas in other organic buffers, no significant product formation was observed (Figure 5). The quantitative product formation in PBS and the lack of significant product formation in HEPES and Tris indicate the ONOO– consumption by the buffer molecules. These results are consistent with the literature reports of the ONOO– reactivity with the functional groups present in these organic buffers. (70−72) Overall, the results demonstrate significant interference from organic buffers in the ROS concentration determination.

Because boronate-based probes are also commonly used for detecting H2O2, (25,55−58) we are interested in examining buffer effects in such studies. Though they are commonly referred as being selective for H2O2, boronate compounds react faster with other ROS such as NaOCl and ONOOH than with H2O2, by more than 2000-fold and a million-fold, respectively. (54) To study the effect of the buffer component on the boronate reaction with H2O2, we chose three buffer solutions, PBS, HEPES, and Tris-Cl at the same pH for ease of comparison. APBE 1 has a second-order rate constant of 2.2 M–1 s–1 for its reaction with H2O2, giving a first t1/2 of 13 h at 10 μM each. (54) Because of this slow reaction kinetics, we studied the reaction with a 10-fold excess of H2O2. Briefly, we incubated 10 μM of APBE 1 with 100 μM of H2O2 in different buffer solutions at 37 °C. Then, 20 μL aliquots were sampled every 15 min for HPLC analyses. As expected, the rates of product formation in each buffer were different (Figure 6). The reaction carried out in PBS reached near completion in 2 h (Figure 6), whereas in HEPES or Tris-Cl, the reaction did not complete in 2 h (Figure 6). At the 15 min time point, almost 40% of product 2 was formed in PBS and only ∼22% was formed in HEPES or Tris (Figure 6). At the 2 h time point, almost 100% of product 2 formed in PBS and only about 73 and 66% of product formed in HEPES and Tris-Cl buffer, respectively (Figure 6). The results may seem surprising since H2O2 is not known to be a strong enough oxidizing agent to oxidize HEPES or Tris within the time frame of the experiments. (73) The rate constants for reactions between a tertiary amine and H2O2 have been reported to be on the scale of 10–5 M–1 s–1, too slow to be an interference of the reaction between a boronate and H2O2. (74) Then, what could be the reason for the observed interference by such buffer components? In order to explain the buffer effect, it is important to recognize the Lewis acid nature of the boron atom. With its open shell, boron is prone to reaction with an electron-rich species, which is an essential step in oxidation by a peroxide (Scheme 1). In the case of APBE, pinacol ester 4 at low mM concentration is known to quickly reverse (hydrolyze) to its boronic acid 8 form in aqueous solution (Scheme 1). (75) This free boronic acid 8 species has an open shell and is prone to oxidative cleavage (Scheme 1). In the case of APBE 1 at low μM concentrations as in the experiments conducted, hydrolysis is expected to be rapid. Because the pKa of 4-acylboronic acid 8 is expected to be at least 7.8, (76) it exists mostly in the oxidation-prone free boronic acid 8 form at pH 7.4. (77−79) However, boronic acid 8 is known to have fairly high affinities for polyols and amino alcohols, including diethanolamine (80) and sugars. (77,78) It is conceivable that the structural properties of HEPES and Tris afford them the ability to chelate to boronic acid and keep a portion in the anionic tetrahedral form, which is not prone to oxidative degradation. All of these could be the reason for the observed interference by HEPES and Tris. Overall, these results indicate slowed product formation from H2O2-mediated oxidation of boronate in HEPES and Tris-Cl buffers than in PBS.

DCFH is known to show rapid response toward NaOCl. (29) We started by determining the second-order rate constant of the DCFH reaction with NaOCl, using stopped-flow spectrophotometry. Briefly, we determined the pseudo-first-order rate constant by using NaOCl in excess. Specifically, we first determined the pseudo-first-order rate constant of the reaction between 20 μM DCFH and NaOCl at 250, 300, 350, 400, 450, and 500 μM (Figure S12). All of the solutions were prepared in 10 mM PBS at pH 7.4. Then, the pseudo-first-order rate constants were plotted against NaOCl concentration (Figures 7, S12, and S13), yielding a second-order rate constant of 580 ± 18 M–1 s–1 (Figures 7, S12, and S13). Such a number is smaller than that for the same reaction of NaOCl with HEPES, MES, Tris, and ammonium acetate. Therefore, we anticipate interference issues using these buffer molecules.

We studied the effect of buffer reactivity on DCFH oxidation by incubating 10 μM DCFH with 100 μM NaOCl in different buffer solutions at 37 °C. Ninety-six-well plates were used for this study. We first used an experimental protocol designed to examine the effect of the buffer without having the competitive reaction of DCFH with NaOCl. Specifically, 50 μL of 400 μM NaOCl was first incubated with 100 μL of a 10 mM buffer at 37 °C for 5 min. Subsequently, 50 μL of 40 μM DCFH was added to the solution before recording of the fluorescence using a plate reader (λex 495 nm and λem 530 nm). As expected, NaOCl led to no fluorescent turn-on effect of the probe when the experiments were conducted in HEPES, MES, Tris-Cl, or ammonium acetate buffer, indicating consumption of NaOCl by the buffer molecules (Figure 8) and the inability of the chloramine products to react with DCFH, whereas the same experiments in PBS and citrate buffer showed fluorescence turn on (Figure 8). Next, a similar experiment was conducted without the incubation step of the buffer and ROS (Figure S14). Briefly, 50 μL of 400 μM NaOCl was added first to 100 μL of 10 mM buffer. This was immediately followed by the addition of 50 μL of 40 μM DCFH and then recording of the fluorescence using a plate reader (λex = 495 nm and λem = 530 nm). The results were such that the highest fluorescence signal was observed in PBS followed by citrate buffer (Figure 8). In MES buffer, the fluorescence intensity was about 50% of that in PBS. In other buffers, including HEPES, Tris-Cl, and ammonium acetate, no significant fluorescence turn-on of the probe by NaOCl was observed, indicating consumption of NaOCl by the buffer molecules (Figure 8). These results are in-line with the reaction kinetics as MES showed a slower reaction rate with NaOCl compared to HEPES, Tris-Cl, and ammonium acetate (Figure 3 and Table 1). Furthermore, the results indicate that these buffer molecules can cause interference of variable degrees depending on the experimental protocols. Overall, the results demonstrate strong interference from four buffers: HEPES, MES, Tris-Cl, and ammonium acetate.