Supporting Information to the Article
PHOTOCHEMICAL OXIDATION OF CHLORIDE-ION BY OZONE IN ACID AQUEOUS SOLUTION
Alexander V. Levanov,*,a Oksana Ya. Isaykina,b Nazrin K. Amirova,a Ewald E. Antipenko,a Valerii V. Lunina,b
a Department of Chemistry, M.V. Lomonosov Moscow State University, Leninskiye Gory 1, building 3, 119991 Moscow (Russia)
b A.V. Topchiev Institute of Petrochemical Synthesis, Russian Academy of Sciences, Leninsky prospect 29, 119991 Moscow (Russia)
* Corresponding Author. E-mail address: ; Fax: (+7) 495-939-4575; Phone: (+7) 495-939-3685
Fig. S1. Effect of ozone concentration in initial gases on chlorine emission rate. Points are the experimental data. Lines represent the results of model calculations: black solid line – reactions (R1-R17), HO3 = 0.24, kLa = 0.3 s–1; red solid line – reactions (R1-R17), HO3 = 0.16, kLa = 0.4 s–1; black dashed line – reactions (R1-R10, R12-R16, R18), HO3 = 0.24, kLa = 0.3 s–1, k18 = 7.15 × 105 s–1, n18 = 1.0725; red dashed line – reactions (R1-R10, R12-R16, R18), HO3 = 0.16, kLa = 0.4 s–1, k18 = 2.26 × 106 s–1, n18 = 1.107. Concentrations in reaction solution [H+] = 0.1 М, [Cl–] = 1 М, [Na+] = 0.9 М.
Fig. S2. Effect of ozone concentration in initial gases on chlorate formation rate. Points are the experimental data. Lines represent the results of model calculations: black solid line – reactions (R1-R17), HO3 = 0.24, kLa = 0.3 s–1; red solid line – reactions (R1-R17), HO3 = 0.16, kLa = 0.4 s–1; black dashed line – reactions (R1-R10, R12-R16, R18), HO3 = 0.24, kLa = 0.3 s–1, k18 = 7.15 × 105 s–1, n18 = 1.0725; red dashed line – reactions (R1-R10, R12-R16, R18), HO3 = 0.16, kLa = 0.4 s–1, k18 = 2.26 × 106 s–1, n18 = 1.107. Concentrations in reaction solution [H+] = 0.1 М, [Cl–] = 1 М, [Na+] = 0.9 М.
Fig. S3. Effect of H+ concentration in reaction solution on chlorine emission rate. Points are the experimental data. Lines represent the results of model calculations: black solid line – reactions (R1-R17), HO3 = 0.24, kLa = 0.3 s–1; red solid line – reactions (R1-R17), HO3 = 0.16, kLa = 0.4 s–1; black dashed line – reactions (R1-R10, R12-R16, R18), HO3 = 0.24, kLa = 0.3 s–1, k18 = 7.15 × 105 s–1, n18 = 1.0725; red dashed line – reactions (R1-R10, R12-R16, R18), HO3 = 0.16, kLa = 0.4 s–1, k18 = 2.26 × 106 s–1, n18 = 1.107. Ozone concentration in initial gases 10.4 g/m3; concentrations in reaction solution [Cl–] = 1 М, [H+] + [Na+] = 1 М.
Fig. S4. Effect of H+ concentration in reaction solution on chlorate formation rate. Points are the experimental data. Lines represent the results of model calculations: black solid line – reactions (R1-R17), HO3 = 0.24, kLa = 0.3 s–1; red solid line – reactions (R1-R17), HO3 = 0.16, kLa = 0.4 s–1; black dashed line – reactions (R1-R10, R12-R16, R18), HO3 = 0.24, kLa = 0.3 s–1, k18 = 7.15 × 105 s–1, n18 = 1.0725; red dashed line – reactions (R1-R10, R12-R16, R18), HO3 = 0.16, kLa = 0.4 s–1, k18 = 2.26 × 106 s–1, n18 = 1.107. Ozone concentration in initial gases 10.4 g/m3; concentrations in reaction solution [Cl–] = 1 М, [H+] + [Na+] = 1 М.
Fig. S5. Effect of the magnitude of volumetric mass transfer coefficient kLa on calculated rates of chlorine emission and chlorate formation. Kinetic calculations have been performed with the set of reactions (R1 – R17). Experimental conditions: concentrations in reaction solution [H+] = 0.8 М, [Cl–] = 1 М, [Na+] = 0.2 М, ozone concentration in initial gases 10.4 g/m3, the number of photons absorbed per unit volume of reaction solution (4.34 ± 0.06) × 1020 photons L−1 min−1.
Fig. S6. Optimized values of the kinetic parameters of process (R18), k18 and n18, as functions of Henry’s law constant of ozone HO3 and volumetric mass transfer coefficient kLa. Kinetic calculations have been performed with the set of reactions (R1 – R10, R12 – R16, R18). Experimental conditions: concentrations in reaction solution [H+] = 0.8 М, [Cl–] = 1 М, [Na+] = 0.2 М, ozone concentration in initial gases 10.4 g/m3, the number of photons absorbed per unit volume of reaction solution (4.34 ± 0.06) × 1020 photons L−1 min−1.For kLa ≤ 0.1 s–1 at HO3 = 0.24 and kLa ≤ 0.2 s–1 at HO3 = 0.16, it is impossible to get the agreement between the experimental and calculated rates by increasing the parameter k18.
Table S1. Reactions R1 – R17 included in the mechanism of photochemical oxidation of aqueous chloride ion by ozone.
# / Irreversible reactions / Rate constantR1 / O3 + hν → O(1D) + O2 / 4.57 × 10−2 s−1 (see footnote #3 in the main text)
R2 / O(1D) + H2O → 2OH / k = φOH × 1.0 × 1012 s−1; φOH = 0.05 (Reisz et al. 2003)
R3 / O(1D) + H2O → H2O2 / k = φH2O2 × 1.0 × 1012 s−1; φH2O2 = 0.9 (Reisz et al. 2003)
R4 / Cl– + O3 → ClO– + O2 / (1.80 × 10–3+1.56 × 10–2[H+])/(1+9.07 × 10–2[H+][Cl–]) L mol−1 s−1
(Levanov et al. 2012; Levanov et al. 2003) (see footnote #4 in the main text)
R10 / Cl2 + H2O2 → 2Cl– + O2 + 2H+ / 183.3/(1+2.27[H+][Cl–]) L mol−1 s−1 (Connick 1947)
R11 / Cl2– + H2O2 → 2Cl– + HO2 + H+ / 6.50 × 105 L mol−1 s−1 (Yu 2004)
R12 / Cl2 + HO2 → Cl2– + H+ + O2 / 1.00 × 109 L mol−1 s−1 (Bjergbakke et al. 1981)
R13 / Cl2O2 + H2O → Cl– + ClO3– + 2H+ / 1 × 104 s−1 (Quiroga and Perissinotti 2005)
R14 / ClO + ClO → Cl2O2 / 2.50 × 109 L mol−1 s−1 (Klaning and Wolff 1985)
R15 / Cl2– + O3 → ClO +Cl– + O2 / 9.00 × 107 L mol−1 s−1 (Bielski 1993)
R16 / ClO + H2O2 → HOCl + HO2 / 3 × 108 L mol−1 s−1 (Su et al. 1979)
R17 / ClO + HO2 → HOCl + O2 / 4.2 × 109 L mol−1 s−1 (Atkinson et al. 2007)
# / Reversible reactions / Equilibrium constant / Forward reaction rate constant / Reverse reaction rate constant
R5 / HOCl ⇄ H+ + ClO– (Adam et al. 1992) / 3.98 × 108 М / 1.99 × 103 s−1 / 5.00 × 1010 L mol−1 s−1
R6 / Cl2+H2O⇄H++Cl–+HOCl (Wang and Margerum 1994) / 1 × 10–3 М2 / 22 s−1 / 2.14 × 104 L2 mol−2 s−1
R7 / Cl– + OH ⇄ ClOH– (Yu 2004) / 0.70 М−1 / 4.2 × 109 L mol−1 s−1 / 6.0 × 109 s−1
R8 / ClOH– + H+ ⇄ Cl + H2O (Yu 2004) / 7.4 × 106 М−1 / 2.4 × 1010 L mol−1 s−1 / 1.8 × 105 s−1
R9 / Cl + Cl– ⇄ Cl2– (Yu 2004) / 1.4 × 105 М−1 / 7.8 × 109 L mol−1 s−1 / 5.7 × 104 s−1
EFFECT OF ADDITION TO THE REACTION SET (R1 – R17) OF VARIOUS REACTIONS, AND THE PROCESSES OF HO2 DISAPPEARANCE AND OH GENERATION, ON THE CALCULATED RATES OF CHLORINE EMISSION AND CHLORATE FORMATION
Kinetic calculations have been performed with the set of reactions (R1 – R17), with the addition of some other processes which were expected to increase the calculated rates. Henry’s law constant for ozone and volumetric mass transfer coefficient were taken to be HO3 = 0.24 and kLa = 0.3 s–1.
Chlorine emission rate, (dnCl2/dt)/Vliq, and chlorate formation rate, (dnClO3−/dt)/Vliq, are expressed in the units μmol L−1 min−1.
Experimental conditions: concentrations in reaction solution [H+] = 0.8 М, [Cl–] = 1 М, [Na+] = 0.2 М, ozone concentration in initial gases 10.4 g/m3, the number of photons absorbed per unit volume of reaction solution (4.34 ± 0.06) × 1020 photons L−1 min−1.
Table S2. Effect of addition of some chemical reactions to the reaction set (R1 – R17) on calculated rates of chlorine emission and chlorate formation.
Reactions / (dnCl2/dt)/Vliq / (dnClO3−/dt)/Vliq(Experiment) / 47.5 / 3.55
(R1 - R17) / 13.5 / 2.30
(R1 – R17) +
a) Cl2– + HO2 → 2Cl– + O2 + H+, k = 3.1 × 109 L mol−1 s−1 (Yu 2004) / 13.5 / 2.30
(R1 – R17) +
b) HO2 + HO2 → H2O2 + O2, k = 7.7 × 105 L mol−1 s−1 (Elliot and Bartels 2009);
c) H2O2 → 2OH, k = 2.77 × 10–4 s−1, the photolysis rate constant was calculated with the formula
k = jH2O2,254·NΦ·e H2O2,254·ln10/(60·NA), where jH2O2,254 = 0.5 is the primary quantum yield (Goldstein et al. 2007; Yu and Barker 2003), NΦ = (4.34 ± 0.06) × 1020 photons L−1 min−1 is the rate of UV photons absorption by the reaction solution in the experiments of this work, e H2O2,254 = 20 L mol–1 cm–1 is the molar absorptivity and hydrogen peroxide in aqueous solution (Chu and Anastasio 2005). / 13.5 / 2.31
(R1 – R17) +
d) HO2 + O3 → OH + 2O2, k = 1.2 × 106 L mol−1 s−1 (Nizkorodov et al. 2000). This is the rate constant in the gas phase. In aqueous solution k < 1 × 104 L mol−1 s−1 (Sehested et al. 1984). / 13.5 / 2.30
(R1 – R17) +
e) Cl– + O(1D) → ClO–, k = 1 × 1010 L mol−1 s−1, diffusion-controlled rate constant (Caldin 2001) / 14.0 / 2.28
(R1 – R17) +
f) HO2 + O(1D) → OH + O2, k = 1 × 1010 L mol−1 s−1, diffusion-controlled rate constant (Caldin 2001) / 13.5 / 2.30
(R1 – R17) +
g) H2O2 + O(1D) → OH + HO2, k = 1 × 1010 L mol−1 s−1, diffusion-controlled rate constant (Caldin 2001);
b) HO2 + HO2 → H2O2 + O2, k = 7.7 × 105 L mol−1 s−1 (Elliot and Bartels 2009);
c) H2O2 → 2OH, k = 2.77 × 10–4 s−1. / 13.5 / 2.30
(R1 – R17) +
b) HO2 + HO2 → H2O2 + O2, k = 7.7 × 105 L mol−1 s−1 (Elliot and Bartels 2009);
c) H2O2 → 2OH, k = 2.77 × 10–4 s−1;
d) HO2 + O3 → OH + 2O2, k = 1.2 × 106 L mol−1 s−1 (Nizkorodov et al. 2000);
e) Cl– + O(1D) → ClO–, k = 1 × 1010 L mol−1 s−1, diffusion-controlled rate constant (Caldin 2001);
f) HO2 + O(1D) → OH + O2, k = 1 × 1010 L mol−1 s−1, diffusion-controlled rate constant (Caldin 2001) / 14.0 / 2.30
(R1 – R17) +
h) O3 → O(3P) + O2, k = 1 × 10–5 s−1 (Ignatiev et al. 2008; Sehested et al. 1991);
i) H2O2 + O(3P) → OH + HO2, k = 1.6 × 109 L mol−1 s−1 (Sauer et al. 1984);
j) HO2 + O(3P) → OH + O2, k = 1.6 × 109 L mol−1 s−1 (assumed value);
k) Cl– + O(3P) → ClO–, k = 1 × 108 L mol−1 s−1 (assumed value);
b) HO2 + HO2 → H2O2 + O2, k = 7.7 × 105 L mol−1 s−1 (Elliot and Bartels 2009);
c) H2O2 → 2OH, k = 2.77 × 10–4 s−1. / 13.5 / 2.31
Table S3. Effect of intensity of additional sink of HO2 on calculated rates of chlorine emission and chlorate formation. Kinetic calculations have been performed with the set of reactions (R1 – R17), plus a process of HO2 disappearance, HO2 → … .
Rate of HO2 disappearance, mol L–1s–1 / (dnCl2/dt)/Vliq / (dnClO3−/dt)/Vliq(Experiment) / 47.5 / 3.55
0 / 13.5 / 2.30
10 / 13.5 / 2.29
100 / 13.6 / 2.24
1 × 103 / 14.5 / 1.77
4 × 103 / 16.9 / 0.71
5 × 103 / The steady state is not established in the calculations.
Table S4. Effect of intensity of additional source of OH on calculated rates of chlorine emission and chlorate formation. Kinetic calculations have been performed with the set of reactions (R1 – R17), plus reactions
b) HO2 + HO2 → H2O2 + O2, k = 7.7 × 105 L mol−1 s−1 (Elliot and Bartels 2009);
c) H2O2 → 2OH, k = 2.77 × 10–4 s−1;
d) HO2 + O3 → OH + 2O2, k = 1.2 × 106 L mol−1 s−1 (Nizkorodov et al. 2000);
l) OH + O3 → HO2 + O2, k = 1.1 × 108 L mol−1 s−1 (Sehested et al. 1984);
m) OH + OH → H2O2, k = 4.64 × 109 (Elliot and Bartels 2009);
n) OH + HO2 → O2 + H2O, k = 8.59 × 109 (Elliot and Bartels 2009);
plus a process of OH generation, … → OH.
Rate of OH generation, mol L–1s–1 / (dnCl2/dt)/Vliq / (dnClO3−/dt)/Vliq(Experiment) / 47.5 / 3.55
0 / 13.5 / 2.31
1 × 10–9 / 13.5 / 2.34
1 × 10–8 / 13.4 / 2.60
1 × 10–7 / 13.0 / 5.21
1 × 10–6 / 8.5 / 31.4
1 × 10–5 / 2.68 × 10–4 / 78.8
1 × 10–4 / 8.71 × 10–3 / 78.8
1 × 10–3 / The steady state is not established in the calculations.
0.1 / 8.09 × 10–3 / 3.70 × 10–2
1 / 2.34 × 10–4 / 3.63 × 10–3
10 / 7.40 × 10–5 / 3.41 × 10–4
100 / 2.34 × 10–5 / 2.84 × 10–5
1000 / 7.39 × 10–6 / 1.75 × 10–6
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