Homogeneous and heterogeneous decomposition of hypochlorite - A study of the oxygen evolving side
reaction using mass spectrometry
Staan Sandin
Supervisors: Ann Cornell and Rasmus Karlsson
January 24, 2013
Abstract
Oxygen evolution from homogenous and heterogenous decomposition of hypochlorite is a small but nonetheless important side reaction in the electrolytic production of chlorate. In this diploma work, a method using mass spectrometry for analyzing the amount of oxygen formed in a hypochlorite containing electrolyte has been developed, and some preliminary experiments have been made. The method works satisfactory for initial screening, but for use in further studies, needs to be developed to include measurement of concentrations in the electrolyte and the ability to maintain a constant pH during experiments.
Based on results from the limited experiments made, some preliminary conclusions can be drawn. The amount of oxygen evolved was measured with the initial pH of 7, 8, and 9, and three dierent types of aqueous electrolytes at initial pH 7; NaOCl(0.19M), NaCl(1.8M/2.7M)+
NaOCl(0.19M), and NaClO4(1.8M) + NaOCl(0.19M). DSA (Dimensionally Stable Anode) particles, two types of cerium salts, and a cobalt salt were tried as catalysts, the concen- tration of the salts were 0.018 mM in all cases. The DSA particles and the cobalt used in this study catalyze the oxygen evolution reaction, while cerium does not. Both hypochlorous acid and hypochlorite ion seem to decompose separately into oxygen in the presence of cata- lyst, while the uncatalyzed decomposition mechanism require the presence of both species as no oxygen is detected outside of the pH range where they are both present (approximately 6<pH<10). The rate of oxygen formation has a maximum around neutral pH for both cat- alyzed and uncatalyzed decomposition, and the rate increases with a decrease in pH in the approximate interval 7<pH<10, below which it decreases. No clear eects of ionic medium or ionic strength were noticed in this study.
Contents
1 Introduction 3
1.1 The chlorate process . . . 3
1.2 Decomposition of hypochlorite . . . 4
1.2.1 The chlorate pathway . . . 5
1.2.2 The oxygen pathway . . . 8
1.2.3 In uence of experimental parameters . . . 8
1.2.4 Catalysis . . . 9
1.3 Aim of study . . . 10
2 Experimental 11 2.1 Apparatus . . . 11
2.1.1 Set-up . . . 11
2.1.2 Calibration . . . 12
2.1.2.1 Gas ow controllers . . . 12
2.1.2.2 Mass spectrometer . . . 15
2.1.2.3 pH-electrode . . . 16
2.2 Method . . . 17
2.2.1 Experimental procedure . . . 17
2.2.1.1 Calibration . . . 17
2.2.1.2 Procedure . . . 17
2.2.1.3 Di culties . . . 19
2.2.2 Source of errors . . . 20
2.3 Chemicals . . . 21
3 Results and discussion 22 3.1 Non catalyzed system . . . 22
3.1.1 In uence of pH . . . 22
3.1.2 In uence of ionic strength and ionic medium . . . 25
3.2 The catalyzed system . . . 26
3.3 Recommendations for future work . . . 29
4 Conclusions 31
Bibliography 32
1 Introduction
1.1 The chlorate process
The chlorate process is an important inorganic electrosynthesis process, producing approxi- mately 3 million tons worldwide[1]. Chlorate is produced by rst electrochemically forming chlorine at the anode, which then reacts between the electrodes with hydroxide formed at the cathode into hypochlorite. The hypochlorite then decomposes in the reactor bulk, forming chlorate. The e ciency of this process is lowered by, among others, unwanted reactions re- sulting in formation of oxygen. As hydrogen is also produced in the chlorate process, oxygen in too large amounts not only lower the e ciency of the process, but also creates a safety haz- ard. Even though the e ciency of the process is high, and the loss in e ciency as a result of oxygen formation is not that great, this still generates rather large economic losses. Making a rough calculation based on the worldwide production of chlorate, a current e ciency loss of 5 percent, an energy demand of 5000 kWh per ton of produced NaClO3, and an energy price of 0.5 SEK per kWh, the annual economic loss from the oxygen evolving side reactions amounts to 125 million SEK.
When developing new catalytic electrodes it is important to know whether corrosion products from the catalytic material can catalyze any side reactions. Permascand, who develops and manufactures electrodes for use in the chlorate process, is a company which could make use of the results from this study.
The main reactions in the chlorate electrolysis can be written as[2]:
Anodic reaction:
2Cl-→Cl2+ 2e- (1.1)
Cathodic reaction:
2H2O + 2e- →H2+ 2OH- (1.2)
Reaction in the electrolyte to form hypochlorite:
Cl2+ 2OH- →ClO-+ Cl-+ H2O (1.3) Hypochlorite is involved in an equilibrium reaction with water:
HOCl + H2O ClO-+ H3O+ (1.4)
Decomposition of hypochlorite in the bulk yields chlorate:
2HOCl + ClO- →ClO3+ 2Cl + 2H+ (1.5) Reaction (1.5) is an overall stoichiometric representation of the reaction, the mechanism of this reaction will be discussed further in section 1.2.1.
Reaction (1.4) is important as it aects the mechanisms involved in the decomposition of hypochlorite. This will also be described further in section 1.2.1.
The dierent side reactions contributing to the formation of oxygen in the chlorate process are[2, 3, 4]:
Decomposition of hypochlorite in the reactor bulk:
2OCl →O2+ 2Cl (1.6)
2HOCl → O2+ 2HCl (1.7)
Oxygen evolution from oxidation of water:
2H2O → O2+ 4H++ 4e (1.8)
Electrochemical chlorate formation:
6HOCl + 3H2O → 3
2O2+ 4Cl + 2ClO3 + 12H++ 6e (1.9) Anodic hypochlorite decomposition:
ClO+ H2O → O2+ 2H++ Cl + 2e (1.10) Although reaction (1.9) produce chlorate, it also produce oxygen. Therefore, reactor tanks are used where the homogeneous decomposition of chlorate in reaction (1.5) occurs to lower the concentration of hypochlorite in the electrolyzers, making the oxygen forming reactions slower. The opinions on which of the reactions (1.6)-(1.10) that contribute most to the formation of oxygen seem to dier. It was however not within the scope of this thesis to investigate further into this matter. The aim of this thesis was to develop a method for investigation of the formation of oxygen from homogeneous and heterogeneous decomposition of hypochlorite in the reactor bulk, reactions (1.6) and (1.7), and to do some initial screen- ing experiments to investigate the eects of ionic strength, ionic medium, pH, and possible catalytic compounds.
1.2 Decomposition of hypochlorite
The decomposition of hypochlorite in the reactor bulk (reactions (1.5), (1.6), and (1.7)) has been investigated by several workers. It seems widely agreed that the overall reaction (1.5) is of 3rd order at pH below 9, and of 2nd order at pH higher than 10[5], while the slow reaction towards oxygen is of 2nd order in hypochlorite ions. The eects of ionic strength, temperature, and pH on the rate of reactions (1.5), (1.6), and (1.7) will be discussed further below in sections 1.2.3 and 3.1. The catalyzed decomposition is discussed in sections 1.2.4
1.2.1 The chlorate pathway
The decomposition mechanisms for hypochlorite to chlorate depends on the pH of the elec- trolyte. In the literature found during this study, two main pH intervals were dominant. In the pH range 5 - 9, the decomposition reaction towards chlorate is of 3rd order. Above pH 10 it is of 2nd order[6, 7, 5] and between pH 9 -10 it is a mix of both. The reason for the change in mechanism is due to reaction (1.4) which, depending on the pH of the electrolyte, determines whether OCl or HOCl will be the dominating hypochlorite specie.
In the pH range 5 - 9 Adam et al.[6] suggested a mechanism for the stoichiometric total reaction (1.5), that explains the change in reaction rate when the electrolyte pH moves toward alkaline values. The suggested mechanism is based on observed rate laws, stoichiometry and several other experimental observations made of hypochlorite decomposition in the neutral pH region. The suggested mechanism is as follows:
2HOCl Cl2O · H2O (1.11)
This is the initiating step of the proposed mechanism where Cl2O is a short lived in- termediate which is directly consumed in reaction (1.12) or (1.14) (depending on pH) and (1.17).
OCl-+ Cl2O · H2O → HOCl + HCl2O2 (1.12)
HCl2O2 HClO2+ Cl- (1.13)
HOCl + Cl2O · H2O → HOCl + H2Cl2O2 (1.14) H2Cl2O2 HClO2+ Cl-+ H+ (1.15) Above pH 6 reactions (1.12) and (1.13) dominates the mechanism while below pH 6 the mechanism goes mainly via reactions (1.14) and (1.15). Reactions (1.12) and (1.14) are the rate determining steps in the proposed mechanism, and the path above pH 6 is the faster of the two. This is supposedly because of the higher negative charge of the oxygen atom in OCl compared to HOCl, which makes the interaction between OCl and the partially positive charged chlorine atom in Cl2O stronger. This is illustrated in gure (1.1) below.
Figure 1.1: Proposed visual reaction scheme for reaction (1.12)[6]
HClO2 ClO2 + H+ (1.16)
ClO2 + Cl2O · H2O → HOCl + HCl2O3 (1.17)
HCl2O3 HClO3+ Cl- (1.18)
HClO3 ClO3 + H+ (1.19)
HOCl OCl+ H+ (1.20)
In reaction (1.20), H+produced in steps (1.15), (1.16) and (1.19) converts OCl to HOCl.
Cl2+ H2O HOCl + Cl-+ H+ (1.21)
At levels of pH below 3, HOCl converts into Cl2 in the side reaction (1.21).
At higher pH levels (above 9), where the concentration of HOCl is low, and therefore the rate of reaction (1.11) is low, ClO2 begins to accumulate in the system[6] as reaction (1.17) becomes slower, (see gure 1.2).
Figure 1.2: Plot of the rate of disappearance of (HOCl + OCl) and the rate of formation of ClO2 vs. pH. Ri = (CiHOCl CtHOCl)/t, where CiHOCl is the initial concentration, and CtHOCl is the concentration at time t. The data point at 7.07 · 104M/s and pH 8 is o scale.
Conditions: CiHOCl= 0.15M, 0.5M borate buer, and 90oC.[6]
As reaction (1.11) becomes slower, ClO2 participates in the alternative reaction path described below. Above about pH 10 when HOCl is no longer present, the decomposition mechanism changes to one with a lower overall rate of reaction [5, 7]:
2OCl →ClO2 + Cl(slow) (1.22)
OCl + ClO ClO + Cl(fast) (1.23)
According to Adam et al.[6], the rate of reaction for hypochlorite decay to chlorate has two maxima, one below pH 8 and one at pH 13. It has a minimum in the range of pH 9 - 10. The rate of reaction at pH 8 is 10 times faster than the rate at pH 9 - 10, and four times faster than at pH 13 ( gure 1.2)[6].
Rate expressions for the above described mechanism in the pH 5 - 9 range presented by Adam et al. in the same study is:
d[ClO3]/dt = k1[HOCl]2(k2[OCl] + k4[HOCl])
k1+ k2[OCl] + k4[HOCl] (1.24) If pH ≥ 6, the expression can be simpli ed to:
d[ClO3]/dt = (k1k2/k1) [HOCl]2[OCl]
1 + (k2/k1) [OCl] (1.25) If the concentration of OClis low (in the study of Adam et al. less than 0.023 M), the expression can be further simpli ed to:
d[ClO3]/dt = (k1k2/k1)[HOCl]2[OCl] (1.26) The rate expression in equation (1.25) is identical to a rate law derived by Yokoyama and Takayasu[6], by which the following mechanism for the decomposition was presented:
HOCl + HOCl → H2Cl2O2 (1.27)
H2Cl2O2 + OCl →ClO3 + 2H++ 2Cl (1.28) Although this agrees well with the simpli ed rate expression in equation (1.26), it does not explain the experimentally observed rate expression in equation (1.25)[6].
The rate expression for the decomposition in the pH range 10 - 14 is[5]:
d[ClO3]/dt = kCl[OCl]² (1.29) In an early work by Lister[8] on the decomposition of hypochlorous acid, a mechanism similar to the one described above in reactions (1.22) and (1.23) was presented. The mech- anism describes the decomposition in electrolytes with pH 8 -11. The reaction mechanism, which has 2nd order kinetics, is:
2HOCl → Cl + 2H++ ClO2(slow) (1.30)
HOCl + ClO2 →Cl + H++ ClO3(fast) (1.31) A much slower reaction involving both HOCl and OCl- is also described:
HOCl + ClO- →Cl-+ H++ ClO2 (1.32) The experiments were made under conditions where OCl- is the dominant specie in order to distinguish between a mechanism involving HOCl and one that does not.
If Adam et al. is to be believed, the pH interval (8 - 11) for which this mechanism is derived, is a very mechanistically complex pH range since the mechanism presented by Adam et al. in the pH range 5 - 9, and the mechanism for pH ≥ 10, described above in reactions (1.22) and (1.23), are competing.
1.2.2 The oxygen pathway
In the literature found, no studies concerning the oxygen evolution at pH lower than 8 could be found. Adam et al.[6] studied the chlorate formation in the pH interval 5 - 8, and in this same study presented evidence that there are no reactions competing to that of decomposition to chlorate at neutral pH, and oxygen formation was never mentioned in the article. In a later study by the same[5], where the decomposition of hypochlorite in the pH 9 - 14 was studied, the formation of oxygen was considered. The reaction presented was (1.6) and the rate constant for the oxygen formation was found to decrease with increasing pH. Judging from these studies, oxygen formation by reaction (1.6) seems to start about at the same pH (around pH≈9) as where the decomposition mechanism towards chlorate begins to change towards the slower mechanism (reactions (1.22) and (1.23)).
Lister[8] examined the decomposition of hypochlorous acid in the pH range of 8 - 11. In this work he presented a mechanism for the oxygen formation where 1st order kinetics was observed:
HOCl + H2O → HCl + H2O2(slow) (1.33)
H2O2+ OCl →Cl + O2+ H2O(fast) (1.34) In later work by Lister[4] the oxygen formation was examined for more alkaline solutions (pH>11). The reaction is presented as a 2ndorder reaction with the rate law:
d[O2]/dt = kOx[OCl]2 (1.35)
Measurement of oxygen levels were not made in any of the studies by Adam et al. and any conclusions made about the formation of oxygen seem based on mass balance calculations.
Lister measured the oxygen formation by the use of a gas burette and based his calculations on the total amount of oxygen produced during the reaction.
1.2.3 In uence of experimental parameters
The rate of decomposition of hypochlorite is aected by temperature, pH, ionic strength, ionic media present, and the presence and intensity of UV light[4, 5].
The eect of pH on the mechanisms involved in the decomposition reaction was discussed above in sections 1.2.1 and 1.2.2. The maximum rate for the decomposition towards chlorate is reached around neutral pH[5, 6].
In a study by Adam and Gordon[5] where experiments were made in the 9 - 14 pH range, the eects of temperature, ionic strength, and chloride ion concentration were investigated.
In the same study, a model for the decomposition was presented over the entire pH range of
1 - 14. An equation describing the relationship between the decomposition rate constant, k2, ionic strength (µ), and temperature (T) was also presented:
log k2 = 0.149µ + logh2.083 · 1010T exp1.018 · 105/RT exp (56.5/R)i (1.36) As can be seen in the expression, the rate of reaction towards chlorate increases with increasing ionic strength and temperature, which according to Lister[4] is true also for oxygen formation.
It is also suggested in the same article by Adam et al. that the chloride ion can act as a catalyst, in addition to its contribution to the ionic strength, in the pH range 9 - 10 by the reactions:
HOCl + Cl HOCl2 (1.37)
HOCl2 + HOCl → Cl2O + Cl + H2O (1.38) It is not clear why the reaction is written with HOCl instead of OCl- which is the dom- inating specie of the two in this pH region. Experiments with the purpose of verifying this hypothesis were not made. However, by adding these two reactions to their model of the system, a better match with the experimental data was obtained.
1.2.4 Catalysis
In the literature found concerning the catalyzed decomposition of hypochlorite towards oxy- gen, catalysis by metal oxides, i.e. heterogeneous catalysis, is dominant. Homogeneous catalysis is only mentioned in the work by Adam et al. described above in section 1.2.3, where the chloride ion is considered as a potential catalyst. Furthermore, most of the work made on catalyzed decomposition is made in alkaline pH, which is not relevant for the chlo- rate process. In the literature found there seems to be a lack of knowledge concerning the species acting as catalysts as it is poorly described in what form the metal is under the dierent conditions, and it is mostly just assumed to be some kind of metal oxide.
Lister[9] investigated the catalytic eects of oxides of the transition metals Cu, Ni, Co, Mn, and Fe on the decomposition of hypochlorite. He found that neither Mn nor Fe catalyze either of the above described decomposition pathways to chlorate and oxygen, respectively, and that Cu, Ni, and Co catalyze only the decomposition towards oxygen. He presented a mechanism for the catalyzed decomposition through the oxidation of the metals, M, to unstable oxides:
2MO + ClO →M2O3+ Cl (1.39)
M2O3 + ClO →M2O3·ClO(adsorbed) (1.40)
M2O3 ·ClO(adsorbed) → 2MO + Cl + O2 (1.41)
According to Lister, the reason that Mn and Fe do not catalyze the reaction is in part because of their oxidation to stable oxides. This mechanism through oxidation, loss of oxygen, and re-oxidation, according to Lister, also gives an explanation as to why these metals do not catalyze the decomposition towards chlorate.
The catalytic eects of iridium compounds on the decomposition reaction was investigated in studies made by Ayres and Booth[10, 11]. They found the catalyzed decomposition to be of the rst order, and that the formation of oxygen increased with decreasing pH. This indicating HOCl as well as OClas sources of oxygen in the catalyzed decomposition.
They also mentioned the results of earlier work by Homan and Ritter, that stated that ruthenium and rhodium had little eect on the decomposition reaction, experimental data from this study has however not been found.
Other work concerning the catalytic eects of the materials used in dimensionally stable anodes (DSA), which are used in the chlorate process, could not be found. However, it is mentioned by Kotowski and Busse[12] that the DSA-material does not catalyze the oxygen formation, and that any observed oxygen formation on the DSA surface would be the result of an internal electrochemical process where the electrode support and the surface material acts as anode and cathode, respectively.
In Peters[13] master's thesis, trials with dierent potential catalysts for chlorate formation were made. Addition of silver chloride and ruthenium dioxide to the electrolyte showed increases in the decomposition rate of hypochlorite. Although only the concentration of hypochlorite and pH were measured, bubbles were not observed in any substantial amount in the former case, while they were to a large degree in the latter, this indicating that silver chloride acts as a catalyst for chlorate production, and ruthenium dioxide for oxygen evolution.
1.3 Aim of study
The aim of this study was to develop a method for measuring oxygen formation from hetero- geneous and homogeneous decomposition of hypochlorite, and also to use this method to do preliminary studies of potential catalysts for the reaction. The potential catalysts would be compounds that can be present in electrode surfaces, and exist in the electrolyte as an eect of corrosion and long-term use.
2 Experimental
2.1 Apparatus
The set-up of the apparatus took quite some time of the thesis period with work such as control measuring of gas ows, test calibration of the mass spectrometer and leakage control.
When all of this and more problems were nally solved, time was short and the number of experiments was therefore limited.
2.1.1 Set-up
The apparatus consisted of a mantled glass reactor sealed with a te on lid in which there were holes made for gas inlet, gas outlet, pH-meter, and an inlet for adding acid without letting air into the system. The gas was controlled with two gas ow regulators, one with a high capacity of about 5 l/min (FC2 in gure), and the other one with a capacity of around 200 ml/min (FC1 in gure), depending on type of gas. A schematic gure of the set-up can be seen below in gure 2.1.
Figure 2.1: Experimental set-up
After making its way through the reaction mixture, the carrier gas (in this case argon) and formed gaseous products continued through a cooler, partly lled with drying pearls, and into the mass spectrometer. The pH was measured at 2 minute intervals by a pH-electrode, which also measured the temperature. Using argon as a carrier gas with a known gas ow
made it possible to calculate the amount of oxygen produced from the output-values given in percent by the software used to control the mass spectrometer. A table of the system components can be seen in table 2.1 below.
Table 2.1: System components
Component Maker Model No.
MS Hiden Analytical HPR20
FC controller Brooks Instruments 0154
FC1 Brooks Instruments 5850S
FC2 Brooks Instruments 5850E
Heater Julabo Labortechnik GMBH VC 3 BASIS
pH-electrode Metrohm pH Lab 827
The pipe leading the carrier gas into the reaction vessel seen in the gure above was a ground glass diuser of porosity 3. The cooler used in the set-up was a 160mm Allihn condenser.
2.1.2 Calibration
2.1.2.1 Gas ow controllers
Careful calibration of the gas ow controllers was made as the precision and reliability of the controllers are important. Below, calibration curves for the dierent gas ow controllers tested are presented. The gas ow rate was measured with a soap bubble gas ow meter where you, as the name suggests, use a known cylindrical volume and the time it takes for a soap membrane to travel through this volume to calculate the ow rate.
The gas ow controllers rst tested were analogue rotameters, and these were calibrated for air and nitrogen. The results of these calibrations can be seen below in gure 2.2.
0 10 20 30 40 50 0
200 400 600 800 1000 1200
position of float (mm)
ml/min
Brooks R−2−15−A, N
2 meter
v = 40.0202 + 20.2948x
80 90 100 110 120 130 140 150
100 120 140 160 180 200 220 240 260
position of float (mm)
ml/min
Brooks R−2−15−AAA, N
2 meter
v = −50.4766 + 1.9821x
Figure 2.2: Calibration curves for analogue rotameters
Also, a digital mass ow controller with internal gas control, dierent from the ones nally used in the set-up was tested with nitrogen ( gure 2.3).
0 5 10 15 20 25
0 5 10 15 20 25
value on flowmeter (ml/min)
measured value (ml/min)
Digital, N
2
data
v = 0.27324 + 0.87542x
Figure 2.3: Calibration curve for digital gas ow controller
The ow controllers nally used in the set-up were two Brooks controllers, the calibration curves of which can be seen below in gure 2.4. These were controlled externally with a control unit, where the percentage of maximum ow was entered for desired gas ow. At this point, the gases had been changed to air and argon, both because this made it possible to detect air entering the system, and that the nitrogen gas contained CO2 which has a peak in the mass-spectrum that overlaps with nitrogen. The calibration was only made once and tests of the ow controllers was not made afterwards.
0 1 2 3 4 5 6 7
50 100 150 200 250 300 350 400 450
% of max (input)
gas flow (ml/min)
FC1 − Argon
data
v = 95.902 + 63.3093x
0 20 40 60 80 100
0 50 100 150
% of max (input)
gas flow (ml/min)
FC2 − Argon and air
Argon
vAr = −4.0607 + 2.0067x AirvAir = −2.9531 + 1.433x
Figure 2.4: Calibration curves for digital gas ow controllers used in set-up
These calibration curves were then used to calculate compositions of gas mixtures of air and argon used for calibration of the mass spectrometer. It might seem like unnecessary work to measure the gas ows in this way since the controllers do have speci ed conversion factors for ow rates of dierent types of gases. However, if one compares the data from these measurements with the curves calculated from coe cients in the gas controller manuals, the importance of these measurements are obvious. See gure 2.5 below for comparison.
0 1 2 3 4 5 6 7 50
100 150 200 250 300 350 400 450
% of max (input)
gas flow (ml/min)
FC1 − Argon
Calibration Specified
0 20 40 60 80 100
0 50 100 150
% of max (input)
gas flow (ml/min)
FC2 − Argon and air
Calibration, Ar Specified, Ar Calibration, Air Specified, Air
Figure 2.5: Comparison of calibration curves from calibration data with curves generated with speci ed gas conversion factors.
The digital and externally controlled mass ow controllers from Brooks were chosen be- cause of the external digital control, making the control of the gas ows more exact, repeat- able, and simple. Moreover, as can be seen in the gures above, the fact that they have a more linear ow to input relationship than the analogue controllers makes the gas control more precise. The ow used for the argon carrier gas during the experiments were 16.15 ml/min.
2.1.2.2 Mass spectrometer
Before initiating an experiment, the mass spectrometer software, QGA professional, was calibrated using a known mixture of air and argon at a single point. By using air as a calibration gas, the nitrogen levels could be monitored which made it possible to identify air leakage into the system. A test of this calibration can be seen below in gure 2.6.
0 1 2 3 4 5 6 7 8 0
1 2 3 4 5 6 7 8
Calculated % O2
% O2
Calibration test QGA professional
Calculated % Measured %
Figure 2.6: Test of mass spectrometer software calibration. Calibration gas composition:
86.09% Ar, 11.00% N2, 2.91% O2.
The line in the gure represents the known oxygen percentage by volume, which was easily calculated as the composition of the air used in this study (technical air; 21% O2, 79% N2), as well as the gas ows ( gure 2.4), were known. The circles represent the output data from the mass spectrometer software, which were measured at known gas compositions (data points taken when conducting the calibration of ow controllers). The raw data from the mass spectrometer was given in partial pressure, but was automatically converted into percentage of volume by the software.
2.1.2.3 pH-electrode
The pH-electrode was calibrated using two calibration solutions of pH 4 and 7 at the desired temperature at least once every day of testing. If two or more trials were made in the same day, the calibration was always checked against the calibration solutions before the initiation of a new run. If the values varied more than pH ±0.1, the pH-meter was re-calibrated. The eventual change of pH output values when checked against calibration solutions before and after experiments were regrettably not noted.
Because of the time it takes for the gas mixture to travel through the system, the pH value and the oxygen percentage value from the MS have to be corrected when analyzing the data. To match these values, the time it takes for oxygen to travel through the system was therefore subtracted from the time values of the mass spectrometer output data. This time was measured by emptying the system of air, lling it with argon, and injecting an approximate amount of 1 ml of air by the acid inlet, noting the time it took from injection to the rise in oxygen and nitrogen values in the mass spectrometer software. The value, or
2.2 Method
2.2.1 Experimental procedure
2.2.1.1 Calibration
The mass spectrometer was calibrated before an experiment if a control of the calibration showed large deviation in measured values from calculated values. This control was made simply by running a known gas composition through the system. Usually, the calibration had to be re-made if more than one day had past since the last calibration. Before taking calibration measurements, a background was taken when argon was run through the system at the same ow as during experiments. This because at the low ow of argon that was used during the experiments, a very small amount of air was entering the system just before the mass spectrometer capillary, where the system was open to the atmosphere. Taking a background measurement eliminates any background interference during a run, making the results easier to interpret.
2.2.1.2 Procedure
The reactor was lled with 140 ml milliQ water and connected to the system. The heating and stirrer was turned on, and argon at high ows (248 ml/min) was used to vent the system of air. When catalysts were used, they were added at this stage. The venting of the system was monitored with the mass spectrometer, and when the temperature stabilized at about 68◦C, and no air remained in the system, the measurement was re-started and the run initiated.
The logging of pH was turned on 10 seconds after the initiation of a run. Just before initiating the run, 80 ml of hypochlorite solution, (stored in a refrigerator at 5◦C), was lled into a measuring vessel. The hypochlorite solution was added through the hole for the pH- electrode at two minutes after initiation, during which the argon ow was high to quickly vent any air that entered during the addition. Upon addition of the hypochlorite solution, the pH rose to around 11, the temperature of the solution fell to about 55◦C, and it took approximately 10 minutes for it to heat up again. After 10 minutes, the argon ow was lowered to that which would be held during the reaction. After 13.5 minutes, acid was added to lower the pH to the desired level. This took about 1 - 3 minutes depending on pH and type of electrolyte. Upon acid addition, the reaction was started. The run was usually terminated after 100 minutes.
The QGA professional data, and the pH electrode output data from a run is illustrated below in gure 2.7, where percent of O2 and pH is plotted as a function of time.
0 10 20 30 40 50 60 70 80 90 100 0
0.5 1 1.5 2 2.5
H2O−NaOCl (14g/l), pH 7, 70oC − % O2 vs. time
time (min)
% O2
0 10 20 30 40 50 60 70 80 90 100
4 5 6 7 8 9 10 11
time (min)
pH
Figure 2.7: Example plot of output data
The data illustrated in this gure is not so easily interpreted as it is hard to say exactly what the momentaneous value represent because of the oxygen build-up within the system, both in the electrolyte, in the gas volume above the liquid, and in tubes and connections between reactor and mass spectrometer. If a constant value was achieved, which would be the case at constant pH, the value could easily be converted into the rate of oxygen formation.
As can be seen in the above gure, no constant value is achieved, and to compare dierent runs it is more bene cial to integrate the total amount of oxygen produced, as in gure 2.8 below.
0 10 20 30 40 50 60 70 80 90 100 0
1 2 3 4 5 6 7 8 9
NaOCl (14g/l), pH 7, 70oC − ml O2 (accumulative) vs. time
time (min) ml O2
Figure 2.8: Amount of oxygen generated over time, same data set as in gure 2.7.
Assuming only argon and oxygen were present in the gas, the volumetric ow rate of oxygen could be calculated using the known ow rate of argon and the measured oxygen percentage in the gas by equation (2.1).
vO2 = yO2 ·vAr
1 yO2 (2.1)
Where yO2 is the percentage of oxygen, and v is the volumetric ow rate of gas in ml/min.
This equation, and integration by the trapezoidal rule, were used to produce plots like the one seen in gure (2.8).
2.2.1.3 Di culties
The addition of acid was probably the hardest part of the procedure. The time it took to get to the desired level of pH was dierent each trial, and if care was not taken, the pH could drop below the target value, or air could get into the system.
When catalysts were used (see section 2.3), these were added with the water, before any hypochlorite was added. Upon addition of hypochlorite, the catalyst precipitated because of the rise in pH, and adsorbed on the diuser. Upon addition of acid, the catalyst particles in the bulk dissolved (indicated by change of color), however the particles adsorbed on the diuser did not (indicated by diuser discoloration). This makes it di cult to know the manner of the catalyst reaction since the amount of catalyst in the bulk is not known. It is furthermore impossible to know if the reaction is catalyzed by the particles adsorbed on the diuser, the catalyst dissolved in the bulk, or both. This problem would be solved by
adding the catalyst after the addition of acid. However, to do this the set-up would have to be altered by adding a separate inlet for the addition of catalyst.
2.2.2 Source of errors
The measurement of amount and rate of addition of both acid and hypochlorite solution were made manually in all trials. This could result in variation of amount and rate of generated oxygen in the dierent trials. Because of the limited amount of time available for experiments, no trial was repeated except that of H2O NaOCl with the initial pH set to ∼ 7. An example of the dierence in measurement results in between trials can be seen below in gures 2.9 and 2.10.
0 10 20 30 40 50 60 70 80 90
0 0.5 1 1.5 2 2.5 3
H2O−NaOCl (14g/l), 70oC − % O
2 vs. time
time (min)
% O 2
0 10 20 30 40 50 60 70 80 90
3 4 5 6 7 8 9 10 11
time (min)
pH
Figure 2.9: Dierence in between trials shown with H2O NaOCl electrolyte with initial pH set to ∼ 7.
0 10 20 30 40 50 60 70 80 90 0
1 2 3 4 5 6 7 8 9 10
NaOCl (14g/l), pH 7, 70oC − ml O
2 (accumulative) vs. time
time (min) ml O 2
Figure 2.10: Dierence in between trials shown with H2O NaOCl electrolyte with initial pH set to ∼ 7.
2.3 Chemicals
The chemicals used in the experiments are presented in table 2.2 below.
Table 2.2: Chemicals used in the experiments
Compound Chemical formula Maker Article number
Sodium perchlorate monohydrate NaClO4·H2O AnalaR 103134Y
Sodium chloride NaCl Merck 1.06404.1000
Hydrochloric acid HCl, 37% Fuming Merck 1.00317.2501
Cerium nitrate hexahydrate Ce (NO3)3·6H2O Alfa Aesar 2332972 Cerium chloride heptahydrate CeCl3·7H2O Aldrich Chemistry 18618-55-8
Cobalt nitrate hexahydrate Co (NO3)2·6H2O Merck 1.02536.0100 Drying pearls, Orange, Heavy metal free - Aldrich Chemistry 94098-500G
pH calibration buer solution, pH 4 - Metrohm 6.2307.100
pH calibration buer solution, pH 7 - Metrohm 6.2307.110
Hypochlorite solution 1N NaOCl in 0.1N NaOH Prolabo 23039 5N The hypochlorite solution used in the study was assumed to be free of metal impurities and was always stored in a refrigerator at about 5◦C. The DSA particles were of 30% RuO2 + 70% TiO2 and produced by Permascand. The size of the particles was not measured.
RO- ltered water (MilliQ) was used in all experiments.
3 Results and discussion
3.1 Non catalyzed system
The in uence of pH, ionic strength, and ionic medium was investigated by varying the pH in three dierent types of aqueous electrolytes:
Electrolyte 1 NaOCl(0.19M or 14g/l)
Electrolyte 2 NaCl(1.8M) : NaOCl(0.19M or 14g/l) NaCl(2.7M) : NaOCl(0.19M or 14g/l) Electrolyte 3 NaClO4(1.8M) : NaOCl(0.19M or 14g/l)
3.1.1 In uence of pH
Both in the case of catalysed and uncatalyzed decomposition, the oxygen evolution seemed to have a maximum around neutral pH. Trials without catalysts were made at three dierent initial pH; 9, 8, and 7. As written, these are only the initial pH as the pH decrease during the reaction for all cases except when the initial pH is set to 9. As will be seen, this leads to preliminary conclusions about the system in the more acidic pH region as well. In gures 3.1 and 3.2, data from these trials are illustrated.
0 10 20 30 40 50 60 70 80 90 0
0.5 1 1.5 2 2.5 3
H2O−NaOCl (14g/l), 70oC − % O
2 vs. time
time (min)
% O 2
pH0 7 pH0 8 pH0 9
0 10 20 30 40 50 60 70 80 90
3 4 5 6 7 8 9 10 11
time (min)
pH
Figure 3.1: In uence of pH on the formation of oxygen
0 10 20 30 40 50 60 70 80 90
0 2 4 6 8 10 12
time (min) ml O 2
NaOCl (14g/l), pH 7, 70oC − ml O
2 (accumulative) vs. time
pH0 7 pH0 8 pH0 9
Figure 3.2: In uence of pH on the formation of oxygen, integrated values
As can be seen in gure 3.1, steady state was achieved only when the initial pH was set to 9. This is mainly believed to be a result of the reactions forming chlorate which result in a
lowering of pH. These reactions are very slow at alkaline pH. How much of the change in pH is caused by the oxygen formation cannot be determined because of the lack of measurement data on chlorate and hypochlorite concentration.
A quick look at the line representing the run made around neutral pH (denoted pH 7) in gure 3.1, might give the impression that oxygen formation decrease and ultimately cease because of depletion of hypochlorite. This is however not believed to be the case, the decrease is a result of change in pH. This would best be veri ed by conducting experiments at constant pH, but the following was noticed in the present experiments. After a run was nished, the pH was lowered to a level where chlorine starts to form by reaction (1.21), the result of this can be seen below in gure 3.3. Since chlorine was formed, hypochlorite must have been present in the solution.
0 50 100 150 200 250 300
0 1 2 3 4 5 6 7 8 9
NaOCl (14g/l),pH 7, 70oC − % O
2 and Cl
2 vs. time
time (min)
%
% O2
% Cl2
0 50 100 150 200 250 300
2 3 4 5 6 7 8 9 10 11
time (min)
pH
Figure 3.3: Veri cation of hypochlorite presence by lowering of pH and measuring chlorine As can be seen in gures 3.1 and 3.2 above, the uncatalysed oxygen formation has a maximum around neutral pH, which is not in agreement with the ndings of Adam et al.[6]
who stated that there are no competing side reactions to that of chlorate formation at neutral pH. The experiments determining this were made at 50◦C and pH 7.1 using a buer solution.
The statement is based on the ratio of reacted hypochlorous acid and formed chlorate being 2.91±0.09 throughout the decomposition. Gas measurement equipment of any kind was not used during the experiments.
The deviation could be a result of the dierence in temperature, electrolyte composition, or contaminants present in the solution that catalyzed the decomposition. It could also be that the amounts of oxygen formed during the experiments of Adam et al. were so small that they were simply not noticed as any gas possibly generated was not measured.
The results presented above is however in agreement with Lister's proposed mechanism involving both OCl and HOCl as seen in section 1.2.2, since oxygen is only detected within the pH range where both species are present in signi cant amounts (6<pH<10).
3.1.2 In uence of ionic strength and ionic medium
Results from trials with dierent electrolyte compositions at an initial pH of about 7.3 can be seen below in gures 3.4 and 3.5.
0 10 20 30 40 50 60 70 80 90
0 0.5 1 1.5 2 2.5 3
NaOCl (14g/l),pH 7, 70oC − % O
2 vs. time
time (min)
% O 2
H2O NaCl(2.7M) NaClO4(1.8M) NaCl(1.8M)
0 10 20 30 40 50 60 70 80 90
3 4 5 6 7 8 9 10 11
time (min)
pH
Figure 3.4: In uence of ionic medium and ionic strength on the evolution of oxygen
0 10 20 30 40 50 60 70 80 90 0
1 2 3 4 5 6 7 8 9 10 11
time (min) ml O 2
NaOCl (14g/l), pH 7, 70oC − ml O
2 (accumulative) vs. time
H2O NaCl(2.7M) NaClO4(1.8M) NaCl(1.8M)
Figure 3.5: In uence of ionic medium and ionic strength on the evolution of oxygen, inte- grated values
Although there was a dierence in the amount of oxygen produced during the time of measurement, it was not signi cantly higher than the dierence between runs, as discussed above in section 2.2.2.
3.2 The catalyzed system
The reason for doing this research was to investigate if corrosion products from electrodes might have any catalytic eects on the formation of oxygen from decomposition of hypochlo- rite. Although the eects can clearly be seen in the plots made from the data collected during these experiments, the particle size and active surface is not known.
The compounds tried were particles from DSA electrodes, cerium chloride CeCl3·7H2O, cerium nitrate Ce (NO3)3·6H2O, and cobalt nitrate Co (NO3)2·H2O. The concentration of the DSA-particles was unknown, while the concentration of the metals was 0.018 mM in all cases. This concentration was chosen based upon concentrations used in the literature found during the study.
The reason for using both cerium nitrate and cerium chloride, was to eliminate any eects that might arise due to the counter ion. In gures 3.6 the eects of the dierent compounds can be seen.
0 10 20 30 40 50 60 70 80 90 100 0
1 2 3 4 5
H2O−NaOCl (14g/l), 70oC − % O
2 & pH vs. time
time (min)
% O 2
No catalyst DSA Ce(NO3)3 CeCl3
0 10 20 30 40 50 60 70 80 90 100
3 4 5 6 7 8 9 10 11
time (min)
pH
Figure 3.6: Comparison of dierent catalyst compounds
As can be seen in the gures, cerium does not seem to act as a catalyst for oxygen formation, while particles from DSA do. When running a catalyst trial on cobalt, the oxygen levels increased immediately upon addition of the hypochlorite, which had not occurred in any other case. As a consequence, the procedure of the cobalt trial varied from that of the others, the results of this run can be seen in gure 3.7.
0 10 20 30 40 50 60 70 80 90 100 0
2 4 6 8 10 12
H2O−NaOCl (14g/l) − Co(NO
3)
2, 70oC − % O
2 & pH vs. time
time (min)
% O 2 % O2
% Cl2
0 10 20 30 40 50 60 70 80 90 100
3 4 5 6 7 8 9 10 11
time (min)
pH
Figure 3.7: Trial using cobalt as catalyst
As cobalt catalyzed the oxygen formation immediately upon addition of the hypochlorite solution, the argon ow was high (248 ml/min) during the rst 10 minutes of the run, before being lowered to the ow normally used during measurements. This resulted in the low values seen in the plot of % O2 vs. time during the rst 10 minutes. This was corrected in the plot below ( gure 3.8). After the rst 30 minutes of the run, the pH was lowered to approximately 3.5 to see how the oxygen formation was aected. This however, resulted in large amounts of Cl2 being formed which makes the measurement data unreliable as the system was not calibrated for chlorine. Although comparison with the other trials at the same pH are not possible, it is clear that cobalt is a very potent catalyst for the oxygen evolution reaction compared to the other compounds. In gure 3.8, a comparison between integrated values of trials with the dierent compounds is illustrated.
0 10 20 30 40 50 60 70 80 90 100 0
5 10 15 20 25 30
time (min) ml O 2
NaOCl (14g/l), 70oC − ml O
2 (accumulative) vs. time
No catalyst DSACe(NO
3)
3
CeCl3
Co(NO
3)
2
Figure 3.8: Comparison of integrated values of catalyst trials
As can be seen in gures 3.6, 3.7, and 3.8, the oxygen formation does not cease completely when pH decreases below 6.5, as it does when no catalyst is used. For the cobalt trial, the formation is relatively high even at pH above 10. This indicates that the catalyzed oxygen formation is active for both hypochlorous acid and hypochlorite ion separately.
According to Pourbaix diagrams[14], for the pH used, cobalt seems to exist as CoO2 over the entire range. Ruthenium, which in the particles exists as RuO2, could in this kind of solution be oxidized (by formed oxygen) to RuO4²-. Cerium exists as Ce (OH)2²+ up to a pH of around 10, above which it changes into CeO2.
3.3 Recommendations for future work
If this type of method is to be used in further similar studies, the apparatus and procedure should be improved. To get a better understanding of the in uence of pH on the formation of oxygen, the apparatus should include a pH-stat so that the pH could be kept constant during the reaction. Doing this would enable determination of kinetic parameters.
The gas volume of the system (volume above electrolyte surface in the reactor, hoses, etc.) should be reduced as much as possible to get a faster gas build-up and thereby reduce the time dierence between formation and detection of oxygen.
The addition of catalyst should be made after the acid addition to avoid precipitation of the metal ion catalyst, if such is used. This change would also eliminate the problems which occurred when using cobalt as a catalyst (see section 3.2).
The addition of any substance during the measurement period should be made using a glass burette or other device enabling addition without letting air into the reactor vessel.
Measuring the concentrations of hypochlorite, chlorate, and chloride is also recommended as this, together with measurement of change in pH, would enable comparison of the con- tribution from the dierent decomposition paths to the change in pH. Which in turn would give a better understanding of the mechanisms involved in the oxygen formation. It would also make it possible to see whether the compounds added catalyze not only the oxygen formation, but also the formation of chlorate. Finally, to make the results more relevant for the chlorate industry, experiments should be made under chlorate process conditions.
4 Conclusions
The ionic medium used, and its concentration, does not have any substantial eect on the decomposition reactions towards oxygen within the concentration ranges used in this study.
The uncatalyzed oxygen formation has a maximum around neutral pH, and involves both the hypochlorite ion (OCl-), and hypochlorous acid (HOCl) in the pH range of approximately 6<pH<10. The catalyzed decomposition of hypochlorite forming oxygen however is active for both hypochlorous acid and hypochlorite ion separately.
DSA particles and cobalt catalyze the oxygen formation while cerium does not.
The method developed during this study is considered successful, but should however be developed further for use in future work as described above in section 3.3.
Acknowledgments
I would like to thank my supervisors, Ann Cornell and Rasmus Karlsson for their support during the entire project. I would also like to thank the people at Permascand and Eka Chemicals who let me participate in their meetings and shared their ideas with me. Lastly, a thanks to all of the co-workers at Applied Electrochemistry for contributing to a stimulating and entertaining working atmosphere.
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