Effect of CO2 Bubbling into Aqueous Solutions Used for

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Article 2

Effect of CO Bubbling into Aqueous Solutions Used for Electrochemical Reduction of CO for Energy Conversion and Storage 2

Heng Zhong, Katsushi Fujii, Yoshiaki Nakano, and Fangming Jin J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/jp509043h • Publication Date (Web): 04 Dec 2014 Downloaded from http://pubs.acs.org on December 7, 2014

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The Journal of Physical Chemistry

Effect of CO2 Bubbling into Aqueous Solutions Used for Electrochemical Reduction of CO2 for Energy Conversion and Storage

Heng Zhong,† Katsushi Fujii,*,‡ Yoshiaki Nakano,§ and Fangming Jin

†

‖

Research Center for Advanced Science and Technology, The University of Tokyo, Tokyo 153-8904,

Japan. ‡

Global Solar plus Initiative, The University of Tokyo, Tokyo 153-8904, Japan.

§

Department of Electrical Engineering and Information Systems, School of Engineering, The University

of Tokyo, Tokyo 113-8656, Japan. ‖

School of Environmental Science and Engineering, Shanghai Jiao Tong University, 800 Dongchuan RD,

Shanghai 200240, China

*

Corresponding Author: Katsushi Fujii

Tel/Fax: +81-3-5452-5428 E-mail: [email protected]

1

ACS Paragon Plus Environment


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ABSTRACT CO2 bubbling into electrolytes is widely used for the electrochemical and photoelectrochemical reduction of CO2 as a carbon source. However, the effects of the CO2 bubbling into the electrolytes have rarely been studied and are not clear. The CO2 bubbling into different electrolytes such as KHCO3, K2CO3, KCl, and KOH was investigated in this research. The concentrations of total dissolved carbon (TC), H2CO3* (sum of dissolved CO2 and H2CO3), HCO3-, and CO32- in different solutions before and after bubbling with CO2 until saturation were evaluated. The CO2 bubbling caused the pH of the electrolytes to decrease and affected the existence forms of the dissolved carbonaceous species. After bubbling with CO2, the major form of dissolved CO2 in KHCO3, K2CO3, and KOH was HCO3-, while H2CO3* was the dominant species in KCl, HCl, and H2O. The ratio of H2CO3*/TC reduces as the concentration of KHCO3 increases, which is possibly the reason for the decrease of CO2 reduction by the solution of high KHCO3 concentration. Furthermore, CO2 partial pressure increasing caused by the CO2 bubbling strongly enhanced the concentration of CO2 in H2O, which is in accordance with the Henry’s law.

Keywords: CO2 bubbling, carbonate equilibrium, total carbon concentration, CO2 reduction, electrolyte.

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1. INTRODUCTION Electrical energy generated from the renewable energies such as solar, wind, and tide is regarded as one of the most promising ways to replace fossil fuels and to combat climate change.1-3 However, electrical energy generated from these renewable energies is not stable because these energy sources fluctuate with time and seasons. Meanwhile, renewable energies are not equally distributed around the world. Solar energy is mainly concentrated in the equatorial area, while tide energy is mainly found in offshore places not suitable for human habitation. Therefore, electrical energy needs to be stored and transported to better utilize the renewable energies. Battery technologies used to store electricity have been researched for over one hundred years and are regarded as one of the most promising ways to store the electrical energy.4,5 However, the energy densities of the batteries are still not high enough for energy transportation. Hydrogen evolution from water splitting by using electrical energy is one of the most useful methods to transform electrical energy into chemical energy and store it.6-8 Hydrogen can be used as a clean energy because only water is generated after its combustion. However, hydrogen is not easy to store and transport because it is hard to liquefy. Electrochemical reduction of CO2 into hydrocarbons or other organic chemical fuels such as CH4 and CH3OH is another promising way to transform electrical energy into chemical energy.9-11 Compared with hydrogen, organic chemical fuels contain higher energy density. In addition, some liquid organic chemicals such as CH3OH are much easier to store and transport than hydrogen gas. Furthermore, reducing CO2 is an effective way to solve the problem of the increasing concentration of atmospheric CO2 that cause climate change.12,13 3



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Electrochemical reduction of CO2 in aqueous solutions has been studied for many decades because the electrochemical method is simple and can be operated under ambient conditions.14-16 There have been recent reports on photocatalytic and photoelectrochemical reductions of CO2 into CO, hydrocarbons and other organics in aqueous solutions by directly utilizing solar energy.17-19 In these researches CO2 bubbling into the aqueous electrolytes is generally used as the carbon source. Although detailed mechanism of CO2 reduction is still unclear, dissolved CO2, rather than HCO3- or CO32-, is generally accepted as the active species.10,14,17 However, the absorption of CO2 in the electrolyte under the CO2 bubbling and its effect have rarely been studied and are not yet clear. The CO2 gas can dissolve into and react with H2O to form H2CO3, which further decomposes into HCO3- and CO32- as shown in reactions (1) to (4).20-22 CO2 (g) ⇄ CO2 (aq)

(1)

CO2 (aq) + H2O (l) ⇄ H2CO3 (aq)

(2)

H2CO3 (aq) ⇄ H+ (aq) + HCO3- (aq)

(3)

HCO3- (aq) ⇄ H+ (aq) + CO32- (aq)

(4)

The concentrations of different carbonaceous species (CO2, H2CO3, HCO3-, and CO32-) in aqueous solutions without CO2 bubbling have been widely studied in sea water research.20-22 However, the concentrations of these carbonaceous species under CO2 bubbling used in electrochemical reactions have not yet been reported. Meanwhile, the pH, ion composition and concentration, and CO2 absorption into the electrolyte are affected by the CO2 bubbling according to reactions (1) to (4), which probably influences the CO2 reduction. Therefore, the effects of the CO2 bubbling into typical electrolytes used for the CO2 reduction are discussed here to help further understanding the mechanism of CO2 reduction. 4


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2. EXPERIMENTAL SECTION KHCO3 (99.5%), K2CO3 (99.5%), KCl (99.8%), and KOH (flakes, 86.0%) (Kanto chemical Co., Inc, Japan) were used directly without any purification. All the solutions were made by mixing the chemicals with deionized water (18.2 MΩ, Organo Corporation, PRA-0015). CO2 gas (99.995%, Taiyo Nippon Sanso, Japan) was used for the bubbling when needed. The CO2 gas was bubbled into 30 mL of different solutions contained in a 40-mL TOC sample vial with flow rates ranging from 50 to 100 mL min-1. The sample vial was not sealed and opened to the atmosphere. A pH meter was always placed in the sample solutions during the whole process of CO2 bubbling to record the pH changes. The total carbon concentrations (TC) of these solutions before and after bubbling with CO2 were analyzed by a Shimadzu TOC analyzer (TOC-V CPH). All the experiments and TOC analysis were conducted at room temperature (23oC) unless specified.

3. RESULTS AND DISCUSSION 3.1. Carbonate equilibrium in water solution. As shown in reactions (1) to (4), the CO2 can dissolve into and react with H2O to form H2CO3, HCO3-, and CO32-. The concentration of the H2CO3 is, however, much smaller than that of the dissolved CO2 (CO2 (aq)) (≤ 0.3 %).20-22 The sum of the two electronic neutral forms is chemically inseparable.20-22 Therefore, [H2CO3*] is generally used to denote the sum of [CO2 (aq)] and [H2CO3] in the aqueous solution. Note that the carbonate equilibrium in water solution is not denoted by the reaction path ways. It is generally denoted as follows:20-22

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(5) Then the equilibrium constants of K1 and K2 are given as: −

K1 =

−

[ H + ][ HCO 3 ] [ H + ][ HCO 3 ] ≈ * [CO 2 ( aq )] [ H 2 CO 3 ]

(6)

2−

K2 =

[ H + ][CO 3 ] − [ HCO 3 ]

(7)

where pK1= 6.35 and pK2= 10.33 at 25oC.23 The total concentration of dissolved carbonate forms (TC) is expressed as follows: TC = [H2CO3*] + [HCO3-] + [CO32-]

(8)

Then the [CO32-], [HCO3-], and [H2CO3*] can be derived from equations (6) to (8) as follows: 2−

[CO 3 ] =

−

K 1 K 2 TC [ H ] + K 1[ H + ] + K 1 K 2

(9)

K 1 [ H + ]TC [ H ] + K 1[ H + ] + K 1 K 2

(10)

[ HCO 3 ] =

+ 2

+ 2

[H2CO3*] = TC - [HCO3-] - [CO32-]

(11)

Since the K1 and K2 are constants, the dependence of [HCO3-], [CO32-], and [H2CO3*] on the [H+] at given TC in aqueous solution can be plotted in accordance with equations (9) to (11) as shown in Figure 1. The distribution of carbonaceous species is dependent on [H+] and independent of TC. The TC decides the maximum concentration of each species. The H2CO3* is the dominant species at pH below 5. The HCO3- becomes the major anion when the pH ranges from 7.5 to 9. There is only CO32- when the pH is above 12. When [HCO3-] is the same as [H2CO3*], it can be derived immediately from equation (6) that the pH is equal to pK1, which is 6.35. This is indicated by the intersection A in Figure 1. A similar

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reasoning holds for intersection B in Figure 1, where the pH is equal to pK2 (10.33) and [HCO3-] is equal to [CO32-].

-1

Concentration (mol L )

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


-

[HCO3 ]

0.5

2-

[H2CO3*]

[CO3 ]

0.4 0.3

B

A 0.2 0.1

-

[H2CO3*]

[HCO3 ]

2-

[CO3 ]

0.0 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

pH Fig. 1 Distribution of carbonaceous species in water solution with pH at different total dissolved carbon concentrations (TC) (TC: solid line: 0.5 mol L-1; dashed line: 0.1 mol L-1, temperature: 25oC).

3.2. Effect of CO2 bubbling on pH changes in different solutions. As shown in reactions (1) to (4), the dissolved CO2 in the aqueous solutions decomposes into HCO3- and CO32- and produces H+, which probably leads to a pH decrease in the solutions. Time dependences of the pH changes in different solutions were investigated by measuring the pH under a constant CO2 bubbling rate of 50 mL min-1 in order to study the effect of CO2 bubbling on the pH. The results are shown in Figure 2. The pHs of the solutions obviously decrease after bubbling with CO2 in all cases except for 0.1 mol L-1 HCl, in which the pH remains almost the same during the experiment. As shown in reaction (5), the pK1 and pK2 of the H+ generation from the reaction of CO2 with H2O are 6.35 and 10.33, respectively. This suggests that the

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H+ generated from the decomposition of the absorbed CO2 is extremely lower than the concentration of H+ contained in 0.1 mol L-1 HCl. Therefore, no significant change is observed in the pH of 0.1 mol L-1 HCl.

pH

14 12 10

pH

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12 11 10 9 8 7

8

0 0.5 mol/L KHCO3

6

0.1 mol/L KHCO3

0.5 mol/L K2CO3

30

60 90 120 150 Time (min) 0.1 mol/L KOH

4

0.5 mol/L KCl

2 0

0.1 mol/L HCl

0

5

10

15

20

25

30

35

40

Time (min)

Fig. 2 pHs as functions of time in different solutions under the bubbling of CO2 (CO2 flow rate: 50 mL min-1, temperature: room temperature (23oC)).

In solutions of KOH, KHCO3, and KCl, the pH dropped quickly in the first 10 min but kept stable after that. This means the CO2 absorptions in these electrolytes reached a steady state after bubbling of CO2 for 10 min. In 0.1 mol L-1 KOH, the pH dropped from 13.0 to 7.2 after CO2 bubbling. This can be attributed to the reaction of H2CO3* (dissolved CO2) with the OH- contained in the KOH solution as shown in reaction (12). H2CO3* (aq) + OH- (aq) ⇄ HCO3- (aq) + H2O (l)

8


(12)

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This reaction can be derived from reactions (1) to (4). Therefore, the KOH solution changed into KHCO3 solution after bubbling with CO2 until saturation. The pH of 0.1 mol L-1 KOH after the CO2 bubbling (pH 7.2) was very similar to that of 0.1 mol L-1 KHCO3 after the CO2 bubbling (pH 7.1). This suggests that the final solutions of 0.1 mol L-1 KOH and 0.1 mol L-1 KHCO3 are similar after the CO2 bubbling. In the KHCO3 solution, both pHs were ca. 8.3 in 0.1 and 0.5 mol L-1 KHCO3 before the CO2 bubbling. The basicity before bubbling was caused by the OH- generated from the reaction of HCO3with H2O as shown in reaction (13). HCO3- (aq) + H2O (l) ⇄ H2CO3* (aq) + OH- (aq)

pK3 = 7.65 (13)

The equilibrium constant K3 of this equation can be derived as: =

[ ∗ ] ×[ ] [ ]

[ ∗ ]×

= [ ]×[ ] =

(14)

Therefore, pK3 = pKw – pK1 = 7.65

(15)

The reaction (13) is actually the reverse reaction of reaction (12). In the KHCO3 solutions before CO2 bubbling, reaction (13) was the dominant reaction. During the bubbling of CO2, the dissolved CO2 increased the concentration of H2CO3* and then made reaction (13) shift to the left, which means reaction (12) became the dominant. As a result, the OH- was consumed by H2CO3* and leaded to the pH decrease and [HCO3-] increase. This suggested that reaction (12) was the dominant reaction under CO2 bubbling. The pHs in 0.1 and 0.5 mol L-1 KHCO3 after bubbling with CO2 were 7.0 and 7.6, respectively. Less pH drop was observed in higher concentration KHCO3 solution. This is because increasing the

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HCO3- concentration can prevent the H2CO3* from reacting with OH- as shown in reaction (12), and then prevent the pH drop. In 0.5 mol L-1 K2CO3, it took over 2 h for the pH drop to reach the steady state than other solutions under CO2 bubbling as shown in the inset of Figure 2. The process of the pH decrease can be roughly separated into three parts: a quick drop in the first 30 min, a slow drop between 30 to 90 min, and a relatively quick drop from 90 to 120 min. In the K2CO3 solution before the CO2 bubbling, OH- were generated from the reaction of CO32- with H2O as shown in reaction (16). CO32- (aq) + H2O (l) ⇄ HCO3- (aq) + OH- (aq)

pK4 = 3.67

(16)

The equilibrium constant K4 of this equation can be derived as: =

[ ] ×[ ] [ ]

[ ]×

= [ ]×[ ] =

(17)

Therefore, pK4 = pKw – pK2 = 3.67

(18)

This leads to the basicity of the K2CO3 solution before bubbling. Although the generated HCO3- from reaction (16) can further react with H2O to form OH- as shown in reaction (13), pK3 of reaction (13) is much bigger than pK4 of reaction (16). The effect of reaction (13) in K2CO3 solution can be omitted. Therefore, reaction (16) was the dominant reaction in K2CO3 solution without CO2 bubbling. In the beginning of CO2 bubbling into K2CO3, the absorbed CO2 reacted with the OH- and produced HCO3- in accordance with reaction (12). This leaded to the quick pH drop in the first 30 min. The decrease of OHcaused the reaction (16) shifted to the right side and then leaded to the CO32- transforming into HCO3-. Although both reaction (12) and reaction (16) produced HCO3-, the concentration of HCO3- was very low in the starting state, which was not able to affect reaction (12) and (16) obviously. However, as the 10


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reactions processed, the increased concentration of HCO3- prevented the reaction (12) and (16). Since the CO2 was keep bubbling into the solution, reaction (12) should still dominate in the system. As a result, the pH drop slow down from 30 min. After all the CO32- transformed into HCO3-, situations became the same as bubbling of CO2 into KHCO3 solutions. The pH drop speeded up again from 90 min. Therefore, the pH drop pattern of CO2 bubbling into K2CO3 showed three different parts. The detailed mechanism should be controlled by the reaction kinetics which will be study in the future. By combining reactions (12) and (16), the total reaction of CO2 bubbling into K2CO3 solution can be derived as follows: CO32- (aq) + H2CO3* (aq) ⇄ 2HCO3- (aq)

(19)

Finally, the 0.5 mol L-1 K2CO3 became 1.0 mol L-1 KHCO3 after bubbling with CO2 until saturation. Therefore, the reaction (12) should be the dominant reaction to drive the pH decrease in all the solutions under CO2 bubbling. 3.3. Total carbon concentration of different solutions before and after CO2 bubbling. The concentrations of different carbonaceous species (H2CO3*, HCO3-, and CO32-) can be calculated through equations (9) to (11) if the TC is known. The TCs of different solutions before and after bubbling with CO2 were measured by TOC analyzer and the results are shown in Figure 3. The TCs in 0.1, 0.5 and 1.5 mol L-1 KHCO3 were 0.105, 0.523 and 1.575 mol L-1 before bubbling and 0.141, 0.549 and 1.599 mol L-1 after bubbling, respectively. Therefore, the TCs increased 0.036, 0.026, and 0.024 mol L-1 in 0.1, 0.5, and 1.5 mol L-1 KHCO3 after bubbling with CO2, respectively. These TC increases were caused by the absorption of CO2 into the solutions. However, the absolute increases in TC reduced from 0.036 to 0.024 mol L-1 as the KHCO3 concentration increased from 0.1 to 1.5 mol L-1. Hence, more CO2 was absorbed 11



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in lower concentration of KHCO3, which suggested that the lower concentration of KHCO3 promoted the absorption of CO2 into the solution. This can be also explained from reaction (12). Decreasing the HCO3- concentration caused the reaction (12) shifted to the right, which leaded to a decrease of H2CO3* concentration. As the H2CO3* concentration decrease, more CO2 was absorbed into the solution to maintain the H2CO3* concentration which was controlled by the Henry’s law. Hori et al. reported that the faradaic efficiency of H2 generation increased from ca. 11% to ca. 45% while the faradaic efficiency of CO2 reduction dropped from ca. 81% to ca. 45% as the concentration of KHCO3 increased from 0.1 to 2.0 mol L-1 in the electrochemical reduction of CO2 at a Cu electrode.10 We also observed that the CO2 reduction decreased with the increase of KHCO3 concentration under electrochemical reduction of CO2 in 0.1, 0.5 and 1.5 mol/L KHCO3 on a Cu electrode at various potentials using a potentiostatic method. This phenomenon was probably because more CO2 was absorbed in the lower concentration of KHCO3 electrolytes and hence improved the CO2 reduction according to our results. In the case of K2CO3, the TC doubled after the CO2 bubbling. The TCs changed from 0.55 to 1.17 and 1.61 to 3.16 mol L-1 in 0.5 and 1.5 mol L-1 K2CO3 after CO2 bubbling, respectively. This proved that all the original CO32- was transformed into HCO3- after bubbling with CO2, and the same amount of HCO3- was generated from the absorbed CO2 depending on the electroneutrality of the solution as discussed in the former section. As a result, the 0.5 and 1.5 mol L-1 K2CO3 solutions change into 1.0 and 3.0 mol L-1 KHCO3 solutions after the CO2 bubbling, respectively, which can be explained by reaction (19). In the solutions of 0.1 mol L-1 HCl, 0.5 mol L-1 KCl, and pure H2O, the TCs were almost zero before CO2 bubbling. The TCs were very similar in these solutions after bubbling with CO2, ranging 12


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from 0.03 to 0.04 mol L-1. This suggests that the major ions such as K+ and Cl- in these solutions have no obvious effect on the absorption of the CO2.

-1

Before bubbling

3.5

After bubbling

16 14

3.0 12 2.5

10

2.0

8

1.5

pH

Total carbon concentration (mol L )

6 4

0.5

2

0.0

0

M

0.5

M

KH

CO KH 3 1. 5 CO M KH 3 CO 0. 5 M K 3 2 CO 1. 5 3 M K 2 CO 0.1 M 3 HC 0.5 l M K Cl 0.1 M KO H 1. 0 M KO H H 2O

1.0

0. 1

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Fig. 3 Total dissolved carbon concentration (TC) and pH changes in different electrolytes before and after bubbling with CO2 until saturated (open column: TC before bubbling, shaded column: TC after bubbling, open circle: pH before bubbling, solid circle: pH after bubbling, M: mol L-1, CO2 flow rate: 100 mL min-1, temperature: room temperature (23oC)).

The TC increased from zero to 1.1 mol L-1 after the CO2 bubbling in the OH- abundant solution of 1.0 mol L-1 KOH. Considering that the pH of the KOH solution after bubbling with CO2 was ca. 7.8, the dominant anion changed into HCO3- in accordance with the reaction of OH- with H2CO3* in reaction (12). As a result, the solutions of 1.0 mol L-1 KOH, 1.0 mol L-1 KHCO3, and 0.5 mol L-1 K2CO3 probably change into 1.0 mol L-1 CO2 saturated KHCO3 after the CO2 bubbling.

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If we assume the CO2 is a simple gas and can be applied to Henry’s law, which describes the equilibrium between vapor and liquid as below:

c H 2 CO 3 * = κ p CO

2

(20)

where the pCO2 is the partial pressure of the CO2 in the bulk atmosphere, κ is the Henry’s law constant for CO2 ( κ = 0.035 mol L-1 atm-1, at 25oC),24 and cH2CO3* is the concentration of dissolved CO2. Therefore, the concentration of H2CO3* is calculated to be 2.46×10-4 mmol/mol (1.38×10-5 mol L-1) and 0.622 mmol/mol (0.035 mol L-1) in water under CO2 partial pressure of ca. 3.95×10-4 atm (the real partial pressure of CO2 in the atmosphere) and 1 atm at 25oC, respectively. Our experiment results showed that the TC in pure H2O was observed to be ca. 0.036 mol L-1 after bubbling with CO2 until saturated in comparison that the TC was almost zero before CO2 bubbling. These results were very close to the theoretic results. According to Henry’s law, this suggests that the CO2 bubbling increased the CO2 partial pressure into about 1 atm and then lead to the concentration of CO2 in water increased over 2500 times. 3.4. [H2CO3*], [HCO3-], and [CO32-] in different solutions before and after CO2bubbling. The concentration of each carbonaceous species (H2CO3*, HCO3-, and CO32-) can be calculated by equation (9) to (11) from the measured TC and pH of different solutions before and after CO2 bubbling as discussed. Figure 4 shows the composition and concentrations of H2CO3*, HCO3-, and CO32- in different solutions before and after bubbling with CO2 until saturation. The major anion in the KHCO3 solutions is always HCO3- before and after CO2 bubbling. It is not affected by the KHCO3 concentration. However, the major anion changed from CO32- into HCO3- after bubbling in K2CO3 solutions, as already discussed. A small amount of H2CO3* was also found in both KHCO3 and K2CO3 solutions after CO2 bubbling. 14


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This was caused by the pH decrease that affected the distribution of carbonaceous species, which could be explained by Figure 1. In the case of 0.1 mol L-1 HCl, 0.5 mol L-1 KCl, and pure H2O, the dissolved carbonaceous forms are H2CO3* after CO2 bubbling. In the case of 1.0 mol L-1 KOH, the dominant anion is the HCO3- after bubbling, which holds ca. 96% of the TC. Therefore, the major anion in KHCO3, K2CO3, and KOH is HCO3-, while H2CO3* is the dominant species in HCl, KCl, and pure H2O after bubbling with CO2. It is widely accepted that the dissolved CO2 rather than HCO3- or CO32- is the real active species in the electrochemical and photoelectrochemical reduction of CO2.10,14,17 Therefore, the concentration of H2CO3* (represents the active species) and the ratio of H2CO3*/TC were calculated and are shown in Figure 5 to clarify the amount of the active species in different solutions. The [H2CO3*] after bubbling with CO2 increased from 33 to 45 mmol L-1 in the KHCO3 solutions with the concentration from 0.1 to 1.5 mol L-1. The ratio of H2CO3*/TC was 23.6% in 0.1 mol L-1 KHCO3, while it was only 2.8% in 1.5 mol L-1 KHCO3, after CO2 bubbling. Since a lower concentration of KHCO3 is better for the electrochemical CO2 reduction as reported by Hori et al.10 and from our other experiment results. The presented results suggest that a high H2CO3*/TC ratio in low concentration of KHCO3 probably increases the reduction of CO2. Although further evidences are needed, a possible explanation of why high H2CO3*/TC ratio in low concentration KHCO3 solutions promotes the CO2 reduction could be suggested as follow. Low concentration KHCO3 solution also contains low concentration of K+. Although K+ cannot be reduced at the operation potential (redox potential of K+/K is -2.925 V vs NHE @ 25oC), it should be still attracted around or even adsorbed on the Cu electrode which might prevent

15



the CO2 adsorption and/or reduction. As a result, the effect of K+ was smaller in low concentration KHCO3 solution with high H2CO3*/TC ratio. 16

-

[HCO3 ]

3.0

14

2-

[CO3 ]

2.5

12

[H2CO3*]

10

2.0

8 1.5

6

AF 3

BF

AF

BF

0. 1 M 0. KH 1 M CO 0. KH 3 5 M CO 0. KH 3 5 M CO 1. KH 3 5 M CO 1. KH 3 5 M CO 0. KH 3 5 C M O 3 K 0. 5 2 CO M 1. K 3 5 2 CO M 1. K 3 5 2 CO M K 3 2 C O

AF

0 BF

0.0 AF

2 BF

0.5 AF

4

BF

1.0

pH

-1

Concentration (mol L )

3.5

(a) Concentration (mol L )

1.2

14 -

12

2-

[CO3 ]

0.8

10

[H2CO3*]

8 0.6 6 0.4

4 2

0.0

0

M

1

M

H C lB 0. HC F 5 M lA 0. KC F 5 l 0. M K BF 1 M Cl 0. KO AF 1 M H 1. KO BF 0 M H 1. KO AF 0 M HB KO F H A H 2HO F 2 O B H 2HO F 2 O AF

0.2

0.

1

pH

-1

[HCO3 ]

1.0

0.

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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(b) Fig. 4 Concentrations of H2CO3* (solid column), HCO3- (open column), and CO32- (shaded column) calculated from TC and pHs (solid circle) in different solutions before and after bubbling with CO2 until saturation ((a): in solutions of KHCO3 and K2CO3; (b) in solutions of HCl, KCl, KOH, and deionized H2O. BF: before bubbling, AF: after bubbling, M: mol L-1, temperature: room temperature (23oC)).

16


80 -1

[H2CO3*] (mmol L )

Before bubbling After bubbling

100

70 80

60 50

60

40 40

30 20

20

10 0

0

0. 1 M 0. KH 5 M C 1. KH O3 5 M CO 0. KH 3 5 M CO 1. K 3 5 2C M O K 0. 2 C 3 1 O M 3 0. HC 5 l 0. M K 1 M Cl 1. K 0 O M H KO H H 2 O

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Ratio of H2CO3*/ TC (%)

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Fig. 5 H2CO3* concentrations and molar ratios of H2CO3* to TC in different solutions before and after bubbling with CO2 until saturation (open column: [H2CO3*] before bubbling, shaded column: [H2CO3*] after bubbling, solid circle: H2CO3*/TC ratio after bubbling, temperature: room temperature (23oC)).

The concentration of H2CO3* in the K2CO3 solution is generally higher than that in the KHCO3 solution. A 62 and 72 mmol L-1 of H2CO3* were obtained in 0.5 and 1.5 mol L-1 K2CO3, respectively, after the CO2 bubbling. However, the ratio of H2CO3*/TC is still very low: only 5% of the TC is in the form of H2CO3* in 1.5 mol L-1 K2CO3. In the case of 0.1 mol L-1 HCl, 0.5 mol L-1 KCl, and pure H2O, almost 100% of the TC is H2CO3*, and the concentrations of H2CO3* are around 30 to 40 mmol L-1, which is close to that in the 0.1 mol L-1 KHCO3 (33 mmol L-1). This suggests that the HCl, KCl, and H2O could be good electrolytes for the electrochemical reduction of CO2 from the viewpoint of a high active species ratio. However, considering the low pH of HCl with a large amount of H+ and the poor 17



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conductivity of pure H2O, they are not suitable for CO2 reduction. Only KCl is appropriate for the electrochemical reduction of CO2. Therefore, the KHCO3 and KCl solutions are two of the most widely used electrolytes in the electrochemical and photoelectrochemical reduction of CO2 in aqueous solutions probably because they have a high ratio of active species (H2CO3*) according to our results.10,17,25-27 Note that in the electrochemical reduction of CO2, H+ was consumed and the local pH around the electrode was increased, which was not good for the CO2 reduction because the standard redox potential of the CO2 reduction shifted negatively as the pH increased.11 Hence, the KHCO3 solution is probably better than KCl solution because the existence of HCO3- in KHCO3 solution has a buffer effect to maintain the pH of the solution. In the case of 0.1 mol L-1 KOH, the concentration and ratio of H2CO3* were both higher than those in 0.1 mol L-1 KHCO3. This suggests that KOH solution may possibly be a better electrolyte than KHCO3 solution in the electrochemical and photoelectrochemical reduction of CO2. 3.5. Thermodynamic calculation. Thermodynamic study suggested that the pH of H2O under CO2 bubbling was decided by carbonate equilibrium as shown in reaction (5). Since the equilibrium constants K2 is much smaller than K1, the pH was mainly decided by the decomposition of H2CO3 into H+ and HCO3- as shown in reaction (3). According to the definition of K1 as shown in equation (6), the pH can be obtained as: pH = pK1 + log[HCO3-] － log[H2CO3*]

(21)

Replacing the [H2CO3*] with P(CO2) using Henry’s law as shown in equation (20), and using pK1 = 6.35, the pH can be calculated as: pH = 7.81 + log[HCO3-] － log P(CO2) 18


(22)

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If started calculating the pH from reaction (12) or (13), which represented the equilibrium of KHCO3 solutions under or without CO2 bubbling, the same equation was derived. Assuming the pressure of CO2 was 3.95 × 10-4 and 1 atm before and under bubbling, respectively, and using the experimental data of [HCO3-] shown in Figure 3, the pHs of 0.1 and 0.5 mol L-1 KHCO3 under bubbling can be calculated to be 6.84 and 7.52, respectively. These were very close to our measured pHs (6.97 and 7.60), which suggested that a thermodynamic equilibrium was almost achieved under the CO2 bubbling until the pH remained unchanged. However, the calculated pHs of 0.1 and 0.5 mol L-1 KHCO3 before bubbling should be 10.24 and 10.92 respectively, which were totally disagree with the experiment results (both about 8.3). This suggested that the fresh made KHCO3 solution was far from equilibrium.

4. CONCLUSIONS The CO2 bubbling has important effects on the electrochemical reduction of CO2. Firstly, the CO2 bubbling provides the carbon source. Secondly, the CO2 bubbling decreases the pH in the electrolytes, which affects the carbonate equilibrium in aqueous solutions, especially for K2CO3 and KOH solutions. The dominant dissolved carbonaceous species changes as CO32- → HCO3- → H2CO3* with decreasing pH from the equilibrium. A higher HCO3- concentration reduces the CO2 absorption into the KHCO3 solutions. The ratio of H2CO3*/TC also reduces as the concentration of KHCO3 increases. Thirdly, the CO2 bubbling changes the anion composition and concentration in the electrolytes. As a result of CO2 bubbling, KOH solution turned out to be KHCO3 solution saturated with CO2. The same is true for K2CO3 solution. The concentration of H2CO3* was higher in K2CO3 solution than those in other solutions after bubbling with CO2. However, the ratios of H2CO3*/TC were much higher in KHCO3 and 19



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KOH solutions with lower concentrations. The high H2CO3*/TC ratio is possibly the reason for the increase of CO2 reduction by the solution of low KHCO3 concentration from the comparison with the reported results.10 Lastly, the concentration of CO2 in water was improved tremendously under the CO2 bubbling due to the increased CO2 partial pressure. The KHCO3 and KCl solutions are suitable electrolytes for the electrochemical reduction of CO2 from the viewpoint of a high ratio of the active reactant of CO2.

Acknowledgement Prof. Sugiyama from the Department of Electrical Engineering and Information Systems in the University of Tokyo is acknowledged for help with the data analysis and discussion. We thank Dr. Yao and Dr. Yun from Professor Jin’s lab from The Shanghai Jiao Tong University for the experiment help.

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(11) Gattrell, M.; Gupta, N.; Co, A. A Review of the Aqueous Electrochemical Reduction of CO2 to Hydrocarbons at Copper. J Electroanal. Chem. 2006, 594, 1−19. (12) Bala, G.; Caldeira, K.; Mirin, A.; Wickett, M.; Delire, C. Multiceutury Changes to the Global Climate and Carbon Cycle: Results from a Coupled Climate and Carbon Cycle Model. J. Clim. 2005, 18, 4531−4544. (13) Solomon, S.; Plattner, G. K.; Knutti, R.; Friedlingstein, P. Irreversible Climate Change Due to Carbon Dioxide Emissions. Proc. Natl. Acad. Sci. U. S. A. 2009, 106, 1704−1709. (14) Hori, Y.; Suzuki, S. Electrolytic Reduction of Bicarbonate Ion at a Mercury-Electrode. J. Electrochem. Soc. 1983, 130, 2387−2390. (15) Hori, Y.; Suzuki, S. Electrolytic Reduction of Carbon Dioxide at Mercury Electrode in Aqueous Solution. B Chem. Soc. Jpn. 1982, 55, 660−665. (16) Frese, K. W.; Leach, S. Electrochemical Reduction of Carbon Dioxide to Methane, Methanol, and CO on Ru Electrodes. J. Electrochem. Soc. 1985, 132, 259−260. (17) Kumar, B.; Llorente, M.; Froehlich, J.; Dang, T.; Sathrum, A.; Kubiak, C. P. Photochemical and Photoelectrochemical Reduction of CO2. Annu. Rev. Phys. Chem. 2012, 63, 541−569. (18) Izumi, Y. Recent Advances in the Photocatalytic Conversion of Carbon Dioxide to Fuels with Water and/or Hydrogen Using Solar Energy and Beyond. Coord. Chem. Rev. 2013, 257, 171−186. (19) Kumar, B.; Smieja, J. M.; Kubiak, C. P. Photoreduction of CO2 on p-type Silicon Using Re(bipy-Bu')(CO)3Cl: Photovoltages Exceeding 600 mV for the Selective Reduction of CO2 to CO. J. Phys. Chem. C 2010, 114, 14220−14223. (20) Zeebe, R.; Wolf-Gladrow, D. CO2 in Seawater: Equilibrium, Kinetics, Isotopes; Elsevier Science: 22


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Table of contents (TOC) Image

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Effect of CO2 Bubbling into Aqueous Solutions Used for

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