Experimental and DFT mechanistic study of dehydrohalogenation of 1

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Kinetics, Catalysis, and Reaction Engineering

Experimental and DFT mechanistic study of dehydrohalogenation of 1-chloro-1,1-difluoroethane over metal fluorides Wenfeng Han, Bing Liu, Yikun Kang, Zhikun Wang, Wei Yu, Hong Yang, Yongnan Liu, Jiaqin Lu, Haodong Tang, Ying Li, and Weiyu Song Ind. Eng. Chem. Res., Just Accepted Manuscript • Publication Date (Web): 05 Sep 2019 Downloaded from pubs.acs.org on September 5, 2019

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Experimental and DFT mechanistic study of dehydrohalogenation of 1chloro-1,1-difluoroethane over metal fluorides Wenfeng Han,*a Bing Liu,a Yikun Kang,b Zhikun Wang,a Wei Yu,a Hong Yang,a Yongnan Liu,a Jiaqin Lu,a Haodong Tang,a Ying Li,a Weiyu Song*b a Institute of Industrial Catalysis, Zhejiang University of Technology, Chaowang Road 18, Hangzhou 310014, Zhejiang, P. R. China. b State Key Laboratory of Heavy Oil Processing, College of Science, China University of Petroleum, Beijing, 18 Fuxue Road, Beijing, 102249, China * E-mail: [email protected] (W. H.) [email protected] (W. S.) Abstract VDF (vinylidene fluoride) is one of the major fluorinated monomers. Currently, it is produced via the pyrolysis of 1-chloro-1,1-difluoroethane at above 650 oC without any catalyst. Herein, we propose that metal fluorides are promising catalysts which selectively promote the pyrolysis at 300-450 oC. With various metal fluorides as the catalysts, the conversion rate increases with the amount of acidic sites which is also reinforced by the bader charges q. The affinity to Cl of the metal fluorides is responsible for the selectivity. However, the high affinity to Cl also leads to the chlorination of metal fluorides forming metal chlorides followed by the deactivation of catalyst. Different from other metal fluorides, F defects play a major role on the performance of AlF3. With increase in F defects, the selectivity changes from vinylidene chlorofluoride (dehydrofluorination) to VDF (dehydrochlorination), which further confirms the role of affinity to Cl on the selectivity. Keywords: F-defect; 1-Chloro-1,1-difluoroethane; vinylidene fluoride; metal fluoride; dehydrochlorination; dehydrofluorination 1. Introduction Metal fluorides are one kind of fundamental chemicals which finds wide application in metallurgy, electronics, optics, fluorine chemical industry and daily necessities

1-3.

In addition to the unique electronic and optical properties, metal

fluorides also function as catalysts with high stability in corrosive environment, 1


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especially in HCl and HF atmospheres

4-6.

As the fluorine atoms withdraw electrons,

the metals in metal fluorides usually exhibit strong Lewis acidic properties

7-8.

Consequently, they directly serve as the catalysts or as the catalyst supports. For instance, AlF3 is considered as a typical Lewis acid catalyst with strong and abundant Lewis acidic sites 9-11. Dehydrohalogenation of fluorinated and chlorinated carbons is one of the most efficient routes for the preparation of fluorine/chlorine containing olefins, such as VDF (vinylidene fluoride), VF (vinyl fluoride), TrFE (trifluoroethylene) and HFO-1234ze (1, 1, 1, 3, 3-pentafluoropropane) 5, 12-17. Take VDF as the example, generally it is produced via the dehydrochlorination (DeHCl) of HCFC-142b (CH3CClF2, 1-Chloro-1,1difluoroethane). Industrially, the reaction is carried out at temperatures above 650 oC in the absence of catalyst. Due to the high reaction temperatures, coke formation and low selectivity are resulted 18. Recently, we have discovered that nitrogen doped carbon and some of the metal fluorides are effective catalysts for the dehydrochlorination of HCFC-142b at reaction temperature of only 350 oC

14, 18-20.

In the presence of proper

catalyst, both high conversion of HCFC-142b and selectivity to VDF were achieved. However, the activity and stability vary with the kind of catalyst significantly. Hence, it necessitates the deep investigation of catalysts. As carbon-based catalysts are difficult to be regenerated following deactivation, the study is focused on the metal fluorides. It is well accepted that dehydrochlorination (DeHCl) and dehydrofluorination (DeHF) are catalyzed by Lewis acidic sites over the surface of catalysts 21-22. Due to the stability in HF and HCl corrosive atmosphere, fluorides such as AlF3, fluorinated Cr2O3 and MgF2 are usually adopted as the catalysts in the fluorine chemical industry. It was suggested that unsaturated coordination sites of Al and Mg are responsible for the strong Lewis acid 23-24. However, there are very few works concerning the relationship between the formation of acidic sites and coordination of metals. Actually, they function as the active sites for the reactions, and therefore coordination of metal plays a major role in the performance of catalyst. In addition, as reported by Li

25,

strong

Lewis acid is also the active site for coke deposition. Hence, it leads to the deactivation of catalyst rapidly. This conclusion has been confirmed by Fang 16, the stability of AlF3 2



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catalyst was improved significantly with the pre-deposition of carbon over strong acidic sites. Therefore, it necessitates the deep investigation of acidity of various metal fluorides which are potential catalysts for the reactions such as DeHCl and DeHF. DeHCl and DeHF are the competitive routes for the conversion of chlorine and fluorine containing hydrocarbons, typical known as hydrochlorofluorocarbons (HCFCs) to chlorinated or fluorinated monomers. As mentioned previously, DeHCl of HCFC142b to VDF is the major way industrially. Unfortunately, DeHCl, DeHF and Cl/F exchange reactions compete from each other, resulting in the spontaneous formation of VDF, Vinylidene chlorofluoride (VCF), CH3CF3 (HFC-143a) and CH3CCl2F (HCFC141b) as indicated in reactions (1), (2) and (3). Clearly, catalyst, especially the acidity plays a major role in the selectivity of DeHCl, DeHF and Cl/F exchange reactions 26. It is of significance to investigate the effects of unsaturated coordination, acidity, selectivity as well as the stability of various metal fluorides on the catalytic DeHCl and DeHF.

CH3CClF2 → CH2=CF2 + HCl

(1)

CH3CClF2 → CH2=CClF + HF

(2)

2CH3CClF2 → CH3CF3 + CH3CCl2F

(3)

In the present study, various metal fluorides, including KF, MgF2, CaF2, SrF2, BaF2, LaF3, CrF3, ZnF2 and AlF3 were adopted as the catalysts for the dehydrohalogenation of HCFC-142b. The conversion, reaction rate, selectivity and stability were correlated with the acidity, surface area, and chlorination of metal fluorides. Density functional theory (DFT) calculation was carried out to further explore the factors affecting selectivity and conversion rate and elucidate the formation and role of F-defects (unsaturated metal sites). 2. Experimental 2.1 Preparation of catalysts Metal fluorides were prepared via the precipitation of metal nitrates with NH4F in the solution. The metal precursors, including KNO3, Mg(NO3)2•6H2O, Ca(NO3)2•4H2O, 3


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Sr(NO3)2,

Ba(NO3)2,

La(NO3)3•6H2O,

Cr(NO3)3•9H2O,

Zn(NO3)2•6H2O

and

Al(NO3)3•9H2O were purchased from Aladdin Company (Shanghai, China) with analytical purity (>99.0%) without further purification. During the preparation of metal fluorides, 0.1 mol nitrate was dissolved in 200 mL deionized water. With vigorous stirring, NH4F was slowly added to the solution. The molar ratio of NH4F to metal was 1:2.2 for divalent metals and 1:3.3 for trivalent metals. Following stirring for 2 h, the precipitates were obtained via filtration. All the precipitates were dried at 110 oC for 10 h and calcined at 500 oC for 4 h at air atmosphere with a ramp rate of 10 oC/min. 2.2 Characterization of catalysts X-ray diffraction (XRD) experiments were carried out for the identification of crystal structures. XRD patterns were obtained over a Kratos AXIS Ultra DLD analytical instrument. A monochromatic Al K radiation source (1486.6 eV) with an analyzer pass energy of 80 eV was operated at 3 mA and 15 kV. BET specific surface areas and N2 adsorption-desorption isotherms of catalysts were determined at -196 oC on a Micromeretics ASAP 2020 instrument in static measurement mode. Before the measurement, the samples were degassed at 250 oC for 10 h. Transmission electron microscopy (TEM) characterization was adopted for the observation of microstructures of various metal fluorides with a JEOL 2100F transmission electron microscope at an accelerating voltage of 200 kV. The surface compositions of catalysts were determined by X-ray energy spectrometer (EDS). The acidity was measured by the technique of ammonia temperature programmed desorption (NH3-TPD). The experiments were conducted over a self-made instrument with thermal conductivity detector (TCD) and mass spectrometer for detection of the desorption of NH3. During the experiments, the sample was heated in a flow of He to 500 oC at a rate of 10 oC /min, and kept at 500 oC for 30 min. After cooling down to 100 oC, the sample was heated from 100 oC to 600 oC with a heating rate of 10 oC /min in He atmosphere. 2.3 Catalytic activity evaluation The catalytic activities of all the metal fluorides were evaluated for the pyrolysis of HCFC-142b at temperatures between 300 oC and 450 oC. The gas hourly space velocity (GHSV) of HCFC-142b was kept at 600 h-1. All the reactions were performed 4



with a fixed bed reactor (pure nickel tube with an i.d. of 22 mm). Feeding flowrates of HCFC-142b (>99.0%, Juhua Group Co. Ltd., Quzhou, China) balanced by equivalent amounts of N2 (>99.9, Minxing Gas Co. Ltd., Hangzhou, China) were controlled by mass flowrate controllers, respectively. Prior to the reactions, 2 mL catalysts were loaded to the isotherm zone of the reactor. A thermocouple was placed in the middle of the catalyst bed for the detection of reaction temperature. The reaction system was purged with N2 at reaction temperatures before the introduction of reactant (for the removal of air and steam in the system). Following reaction, the effluent flow from the reactor was first washed by KOH solution in a scrubber for the removal of HCl and HF. Then, the stream was further dried by the NaOH pellets. The composition was analyzed by a Jie Dao GC-1690 gas chromatograph equipped with a thermal conductivity detector (TCD). 2.4 Density functional theory (DFT) calculation The process of dehydrochlorination and dehydrofluorination is further reinforced by the DFT calculations. DFT calculations were conducted in the Vienna Ab Initio Simulation Package (VASP) 27-28. It adopts projector augmented wave (PAW) method to describe the interaction between the ions and the electrons with frozen-core approximation 29-31. The Perdew-Burke-Ernzerhof (PBE) electron exchange-correlation functional was used 32. Γ-centered k point meshes of 1×2×1 were used for the Brillouim zone integration. The stable point was identified by the conjugate gradient method until the forces acting on each ion were smaller than 0.05 eV/Å. The energy criterion for convergence of the electron density was set at 10−4 eV. Transition state was calculated by climbing image nudged elastic band method (CINEB) and verified by vibrational analyses with only one imaginary frequency. Stability of different terminations of metal fluorides (111) has been explored by many investigations 33-35. The bulk equilibrium lattice constant of AlF3 was a= 6.93, b= 12.00, c= 7.13.The CaF2 was a=b=c=5.47. The SrF2 was a=b=c=5.79. The SrF2 was a=b=c=6.20. The KF was a=b=c=5.34. For KF we choose the exposed (110) side for research. For other metal fluorides, the most stable (111) surfaces were applied for the investigation of adsorption and reaction mechanism for dehydrohalogenation. (3 × 3) 5


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terminations of (111) surfaces of metal fluorides were selected. The surfaces were cleaved in excess of 6 atomic layers, which was sufficient to converge the surface structure

36-37.

The three top atomic layers of the (111) surfaces was allowed to relax

and the rest of fifteen atomic layers were fixed. The vacuum gap thickness was set to 12 Å. The adsorption energy was calculated by Eads,X=Eslab-X−(Eslab+EX)

(4)

Eslab,X, Eslab and EX are the energy of surface covered by X, the energy of Surface without X adsorption and the energy of X, respectively 38. Based on the approach developed by Reuter and Scheffler

39,

for the surface

system the surface free energy γ(T,P) in an fluorine pressure p and temperature T is defined as γ =

1 slab(T,P) 2𝐴[G

– NMμM(T,P) – NFμF(T,P)]

(5)

where Gslab is the free energy of each species slab system; 2A is the total surface area unit cell; μM(T,P) and μF(T,P) are the chemical potentials of metal and F atoms; NM and NF are the numbers of Co and O atoms in the slab. For the metal fluorides MFx (M = K, Ca, Sr, Ba and Al) crystal, thermal equilibrium between surface and bulk requires that μM(T,P) + xμF(T,P) = gbulk(T,P)

(6)

where gbulk(T,P) is the Gibbs free energy of per MFx formula unit. Combine this constraint with Eq.(5) to get γ =

1 slab(T,P) 2𝐴[G

– NMgbulk(T,P) – (xNM-NF)μF(T,P)],

(7)

then the surface free energy depends only on the fluorine chemical potential. The boundary of the calculated fluorine chemical potential is 1 𝑥ΔGf(0,0)

≤ μF(0,0) – μFgas ≤ 0,

(8)

the left and right boundaries represent fluorine-poor and fluorine-rich limit respectively.The formation heat of the oxide, ΔGf(T,P), is calculated by ΔGf(T,P) = gbulk(T,P) - gM(T,P) - xgFgas,

(9)

where gbulk(T,P) is the Gibbs free energy of metallic solid, and μFgas is calculated by 6



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1

the Gibbs free energy of F2 molecule, gFgas = 2gF2gas(T,P). In this study, the total energy calculated by DFT (in zero temperature and zero pressure environment) approximates Gibbs free energy, and to simplify calculations, the zero-point vibrations, vibrational entropy contributions, and enthalpy changes have been neglected. 3. Results 3.1 Phase structures of metal fluorides To investigate the crystal structures of prepared metal fluorides, including KF, MgF2, CaF2, SrF2, BaF2, LaF3, CrF3, ZnF2 and AlF3, XRD experiments were carried out and the results are presented in Fig. 1. Except for ZnF2 and CrF3, all the XRD patterns agree well with standard diffraction peaks, respectively. No any impurity is identified, indicating the successful preparation of pure metal fluorides. In addition, sharp diffraction peaks for theses fluorides suggesting that the crystallites of fluorides are well developed. For AlF3, both α-AlF3 and β-AlF3 are identified. Hence, the mixture of α-AlF3 and β-AlF3 were prepared via the precipitation of Al(NO3)3 with NH4F. With regard to Zn and Cr nitrates, probably fluorinated ZnO and Cr2O3 are obtained based on the patterns. We suggest that ZnF2 and CrF3 were oxidized during the calcination at 500 oC in air atmosphere. AlF3 ZnF2 CrF3 LaF3 BaF2 SrF2 CaF2 MgF2 KF

10

20

30

40 50 60 2 Theta， °

70

80

Figure 1. XRD patterns of all samples. The standard profiles (presented as dots) of AlF3 (PDF #80-1007/ PDF #84-1642), LaF3 (PDF #76-0510), BaF2 (PDF #85-1341), SrF2 7


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(PDF #88-2294), CaF2 (PDF #99-0051), MgF2 (PDF #72-1150) and KF (PDF #361458) were included as a reference. 3.2 Surface area, porosity and morphology The nitrogen adsorption-desorption isotherms of all the metal fluorides were displayed in Fig. S1. All the metal fluorides exhibit typical type IV isotherms with a well-defined capillary condensation step and H3 hysteresis loops (according to classification of IUPAC) 40. H3 hysteresis loops are usually observed with aggregates of particles giving rise to slit-shape pores. The small hysteresis loops of AlF3, BaF2, SrF2 and others indicating that these samples prepared by precipitation method possess low surface area and small pore volume. The specific surface area and pore volume are listed in Table S1. Except for the CrF3, all the values of specific surface area are lower than 2 m2/g and the pore volumes are less than 0.11 cm3/g. Clearly, during the evaluation of as catalysts, the significant difference of catalytic performance can not be attributed to the difference in surface area and porosity. Typical fluorides were further investigated with TEM for the observation of micro structures. Fig. 2 demonstrated the TEM images of KF, SrF2, BaF2 and AlF3 at different magnitudes. Selected area electron diffraction (SAED) analysis is also included for the identification of lattice fringes (exposed facets).

8



Figure 2. Different magnification TEM images of KF (a1, a2, a3), SrF2(b1, b2, b3), BaF2(c1, c2, c3) and AlF3(d1, d2, d3); Insets of images show the corresponding SAED patterns. As shown in Fig. 2, KF presents flake shapes and while irregular particles are observed for SrF2, BaF2 and AlF3. The particle sizes are larger than 100 nm (Fig. 2a1, b1, c1 and d1). Clear lattice fringes are identified for the fluorides investigated. From the corresponding SAED patterns, we can see that all samples have obvious diffraction 9


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rings, which indicates that all samples are in the form of polycrystalline, which agree well with the observation XRD experiments. For KF, the lattice fringes are determined to be 0.26 nm, well corresponding to the KF (200) crystal planes. By contrast, lattice fringes of 0.33 nm, 0.35 nm and 0.45 nm are detected for SrF2, BaF2 and AlF3 respectively. Clearly, all these fringes are well assigned to (111) planes. It is consistent with the results of XRD that the planes of strongest diffraction in the corresponding XRD patterns are overserved by the lattice fringes. According to the TEM results, majorly (200) facets are exposed for KF. However, for SrF2, BaF2 and AlF3, the main exposed facets are (111) planes. 3.3 Catalytic activity Table 1 lists the catalytic activities of all the metal fluoride catalysts for the catalytic pyrolysis of HCFC-142b to vinylidene fluoride as a function of reaction temperatures. The reactions were carried out at 300 oC -450 oC, pressure of 1 bar and GHSV (HCFC142b) of 600 h−1 respectively. In addition to the target product of CH2=CF2 (VDF, dehydrochlorination product of 1,1-chlorofluoroethane), the by-products include CH2=CClF (VCF, dehydrofluorination product of 1,1-chlorofluoroethane), 1,1,1trifluoroethane (HFC-143a, CH3CF3, the product of Cl/F exchange reaction) and other products during the catalyst activity evaluation. As a typical solid base catalyst 41, the activity of KF is negligible at all temperatures investigated. In the presence of strong Lewis acid catalyst, AlF3 42, the highest conversion levels of HCFC-142b are achieved. Clearly, Lewis acid facilitates the pyrolysis of HCFC-142b rather than Lewis base. This conclusion is reinforced by the activity of CaF2 catalyst. HCFC-142b shows rather low conversion over CaF2. With neutral ZnF2 as the catalyst, it exhibits improved catalytic performance at 300 oC ,350 oC and 400 oC. High conversion levels are achieved at elevated reaction temperatures, such as 450 oC. For the catalysts with moderate Lewis acidity, such as MgF2, SrF2, BaF2 CrF3 and LaF3, although lower than that of AlF3, high conversion levels of HCFC-142b are obtained. It is well accepted that fluorinated Cr2O3 possesses high affinity to fluorocarbons and excellent dehydrofluorination activity to hydrofluorocarbons 12-13, 17. Consequently, as indicated by XRD pattern, although CrF3 are partially oxidized to Cr2O3, high conversion of HCFC-142b is still approached. 10



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Table 1. The conversion of HCFC-142b and selectivity over various fluorides as a function of reaction temperature at pressure of 1 bar and GHSV of 600 h−1. Catalysts

KF

MgF2

CaF2

SrF2

BaF2

LaF3

CrF3

ZnF2

AlF3

T,

Selectivity, %

oC

Reaction rate （mmol/h/g）

Conversion, %

CH2=CF2

CH3CF3

CH2=CClF

Others

300

1.35

6.2

31.9

32.7

35.4

0.0

350

1.41

6.5

29.2

30.6

40.2

0.0

400

0.67

3.1

29.8

35.1

35.1

0.0

450

0.91

4.2

84.4

9.9

5.7

0.0

300

13.95

53.6

7.6

28.9

56.9

6.6

350

23.95

92.0

3.2

28.5

61.4

6.9

400

25.56

98.2

6.2

24.9

62.4

6.5

450

25.85

99.3

13.6

19.3

61.9

5.2

300

1.11

5.1

56.8

14.6

28.6

0.0

350

4.87

22.4

51.9

11.1

37.0

0.0

400

3.41

15.7

72.3

3.9

23.8

0.0

450

5.37

24.7

79.1

2.8

18.1

0.0

300

10.06

38.3

53.3

1.6

45.1

0.0

350

19.58

74.5

52.5

1.2

46.3

0.0

400

23.44

89.2

59.8

1.9

38.3

0.0

450

24.67

93.9

50.5

6.5

42.8

0.2

300

6.71

40.1

98.2

2.1

5.0

0.0

350

10.08

60.3

96.5

1.1

6.0

0.0

400

13.03

77.9

94.8

0.7

4.6

0.0

450

13.43

80.3

90.1

1.1

8.8

0.0

300

8.41

62.8

25.3

36.7

25.3

12.7

350

11.59

86.5

27.0

38.8

22.8

11.4

400

12.56

93.8

33.3

39.4

18.3

9.0

450

12.58

93.9

44.1

38.2

12.1

5.6

300

15.19

73.5

3.6

4.0

92.4

0.0

350

14.94

72.3

5.3

1.0

93.7

0.0

400

20.52

99.3

11.3

36.9

34.2

17.6

450

14.47

70.0

29.5

54.0

12.9

3.6

300

0.96

4.0

80.9

0.0

19.1

0.0

350

0.65

2.7

89.2

0.0

10.8

0.0

400

5.25

21.9

89.1

5.2

5.7

0.0

450

17.78

74.1

87.4

9.3

2.0

1.3

300

29.54

80.3

0.2

53.4

33.6

12.8

350

36.13

98.2

2.5

2.1

73.9

21.5

400

36.64

99.6

7.9

0.0

64.8

27.3

450

36.68

99.7

16.9

32.0

30.5

20.6

11


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In addition to the conversion of HCFC-142b, the selectivity also differs from catalysts significantly. Although the highest reaction rate of conversion are obtained over AlF3, the major products are CH3CF3 at 300 oC and VCF at elevated reaction temperatures. The selectivity to target product VDF is rather low (less than 17%). As indicated in Reactions (2) and (3), CH3CF3 and VCF are formed via Cl/F exchange and dehydrofluorination reactions, respectively. Clearly, at relatively low temperatures (300 oC), AlF3 catalyst favors the Cl/F exchange for HCFC-142b and while at temperatures above (>300 oC), dehydrofluorination dominates the catalytic reaction on AlF3. Similarly, MgF2 and CrF3 (or fluorinated Cr2O3) exhibit high selectivity to VCF (dehydrofluorination reaction) which agree well with the literature that MgF2 and CrF3 (or fluorinated Cr2O3) are competitive catalysts for dehydrofluorination

5, 43-44.

Although LaF3 although presents high activity for the conversion of HCFC-142b, it promotes the formation of CH3CF3 (Cl/F exchange reaction) significantly. Meanwhile, large amounts of VCF and VDF are also produced. Clearly, both DeHF and DeHCl reactions together with Cl/F exchange reaction take place spontaneously. With less strong or neutral acidity, CaF2, SrF2, BaF2 and ZnF2 display much higher selectivity to VDF than other metal fluorides. In the presence of CaF2, a selectivity between 50-80% to VDF is achieved at reaction temperatures investigated. ZnF2 shows even higher selectivity (80-90%) although the reaction rate and conversion of HCFC142b are low. Both SrF2 and BaF2 possess high reaction rate of HCFC-142b conversion and selectivity to VDF. It is worth noting that BaF2 as the catalyst presents higher activity as the conversion of HCFC-142b is about 78% and possesses highest selectivity to VDF (close to 95%) among all catalysts at 400 oC. It is consistent with our previous reports that SrF2 and BaF2 are potential catalysts for the DeHCl of HCFC-142b to VDF 14, 20.

3.4 DFT calculation In order to perform subsequent calculations on relatively reasonable models, the surface energies of KF (110), CaF2 (111), SrF2 (111), BaF2 (111) and AlF3 (111) ideal surfaces and their defect surfaces as a function of fluorine chemical potential are 12



presented in Fig. 3. For the defect models of KF (110), CaF2 (111), SrF2 (111) and BaF2 (111), we simulate the formation of one fluorine vacancy (i.e. 1/9 ML F atoms leave the surface) on each of their surface respectively. For the defect models of AlF3 (111), we simulated the surface energy of surface with 1, 2, 3, 4, 5, and 6 fluorine vacancies (i.e. 1/14, 1/7, 3/14, 2/7, 5/14 and 3/7 ML F atoms leave the surface) separately (Fig .3e).

Figure 3. The surface energies of (a) KF (110), (b) CaF2 (111), (c) SrF2 (111), (d) BaF2 (111) and (e) AlF3 (111) ideal surfaces and their defect surfaces as a function of fluorine chemical potential. It can be seen from Fig. 3 that with the increase of the fluorine chemical potential, the surface energy of KF (110), CaF2 (111), SrF2 (111) and BaF2 (111) with fluorine defects are always higher than their ideal surface, so they do not tend to form fluorine defects. While for AlF3 (111), the surface energy of the F-defect surface is lower than the ideal surface in most of the fluorine chemical potential change (except for a small portion at the F-rich region), and the more the defects, the more surface energy decreases. Therefore, we choose the ideal surface for KF (110), CaF2 (111), SrF2 (111) and BaF2 (111) (Fig. 4a-d), while for AlF3 (111), we choose 3/7 ML F-defect surface (Fig. 4f) for subsequent calculations.

13


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Figure 4.The schematic diagram of the ideal surface structure of CaF2 (111) (a), SrF2 (111) (b), BaF2 (111) (c), KF (110) (d), AlF3 (111) (e) and AlF3 (111) (f) with 3/7 ML F defects. The large balls represent the metal elements and small balls represent the fluorine atoms.

We believe that the first step of the reaction is the dissociation of C-F (DeF) or CCl (DeCl) bonds, the expression is CH3CF2Cl + * → CH3CFCl· + F*

(10)

CH3CF2Cl + * → CH3CF2· + Cl*,

(11)

and we calculated the reaction enthalpy (ER-DeCl and ER-DeCl) of the first step to describe the catalytic performance (Fig. 5 and Table S2). The smaller ER indicate the higher catalytic activity. As shown in Fig. 5, the general trend of ER is AlF3 < BaF2 < SrF2 < CaF2 < KF, which indicates that the conversion rate of metal fluoride is AlF3 > BaF2 > SrF2 > CaF2 > KF. In addition, ER-DeF and ER-DeCl describe the selectivity on different mental fluorides for DeHF and DeHCl, respectively. For KF, CaF2, SrF2 and BaF2, ER-DeCl < ER-DeF, which means they have selectivity for VCF, while for AlF3, it is more inclined to generate VDF (ER-DeF < ER-DeCl). This is consistent with the experimentally measured conversion and selectivity.

14



Figure 5. Reaction enthalpy of DeF (ER-DeF) and DeCl (ER-DeCl) on different metal fluorides.

4. Discussion 4.1 Effect of surface acidity on the conversion and selectivity In order to confirm the effect of the surface acidity on the catalytic performance, all the catalysts were further characterized by temperature-programmed desorption of ammonia (NH3-TPD). The profiles are displayed in Fig. 6. The peak areas of NH3-TPD obtained by integral are listed in Table S3. With the peak area of KF as reference, the relative acidic site amounts (peak area ratio with that of KF as 1) are also included in Table S2. NH3-TPD is usually adopted for the investigation of the acid strength of sites and the amounts of acidic sites on the surface of catalysts. As demonstrated in Fig. 6, basically there is no NH3 desorption peak for KF indicating the absence of acidic sites. As it is known that dehydrochlorination and dehydrofluorination of HCFCs need Lewis acid sites, the absence of acidic sites results in the poor catalytic performance of KF catalyst for the pyrolysis of HCFC-142b. There is a broad desorption peak for CaF2 in the range 350-700 oC. No strong peaks are observed in the case of CaF2 catalyst and the peak area is relatively low. That's why the catalytic performance of CaF2 catalyst is just a little bit higher than that of KF catalyst.

15


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o

667 C

AlF3 ZnF2 CrF3

Intensity， a.u.

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


LaF3 BaF2 o

520 C

SrF2 CaF2 MgF2

o

o

432 C

634 C

x10

KF

0 100 200 300 400 500 600 700 800 o

T, C Figure 6. Temperature-programmed desorption profiles of ammonia on metal fluoride

catalysts.

As for traditional strong Lewis acid metal fluorides of AlF3, MgF2 and CrF3, it is observed that these catalysts have strong NH3 desorption peaks. As listed in Table S3, AlF3, MgF2 and CrF3 catalyst possess much larger amounts of acidic sites than the rest of catalysts. There is a broad NH3 desorption peak for AlF3 at temperatures between 320 oC and 700 oC with the peak temperature at 667 °C. As argued previously, highest conversion rate of HCFC-142b was obtained with AlF3 compared to other catalysts (Table 1) at all reaction temperatures investigated. The same discussion can be applied to MgF2 and CrF3 catalysts. Similarly, SrF2 also exhibits large desorption peak at 520°C and agrees well with its high activity. It is notable that MgF2 catalysts possess large amounts of medium to strong acidic sites, well consistent with the high reaction rate. It is observed that a weak desorption peak for ZnF2 in the range 550-650 oC centered at 600 °C with the peak area close to KF. It explains low reaction rates at 300oC and 350 oC. We suggest that the formation of ZnOCl and ZnOF is responsible for the relatively high reaction rates at reaction temperatures of 400 oC and 450 oC

45-46.

Similarly, formation of intermediates such as LaOF is responsible for the high 16



conversion rate of HCFC-142b over LaF3 catalyst 6, 38.

Reaction rate, mmol/h/g

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

40 35 30 25 20 15 10 5 0 0 5 10 15 20 25 30 Relative amount of acidic sites

Figure 7. Conversion rate of HCFC-142b as a function of relative amount of acidic sites. Reactions were carried out at 400 oC, pressure of 1 bar and GHSV (HCFC-142b) of 600 h−1. Relative amount of acidic sites were determined by peak areas NH3-TPD with the peak area of KF as reference.

Clearly, the conversion rate of HCFC-142b correlates with the acidic sites of metal fluorides. Consequently, the relationship between amount of acidic sites and conversion rate was attempted and the results are disclosed in Fig. 7. Although different kinds of metal fluorides are adopted as the catalysts in the present study, a clear trend that HCFC-142b conversion rate increases with amount of acidic sites was identified. As mentioned previously, both DeHCl and DeHF as well as Cl/F exchange reactions contribute the conversion. Hence, we suggest that acidic sites of these metal fluorides are majorly responsible for DeHCl, DeHF and Cl/F exchange reactions. Acidic sites function as the active sites for the above reactions. In order to understand the influence of surface acidity from a theoretical perspective, we calculated the average bader charges q of the metal on different fluorides surface (Table 2), which describe surface acidity. The more positive charges on the metal surface (the larger the q value) indicates the stronger the surface acidity. 17


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Table 2. Calculated average bader charges q of metal atoms on different fluorides surface. surface KF (110) CaF2 (111) SrF2 (111) BaF2 (111) AlF3 (111)a q 0.859 1.669 1.703 1.710 2.289 aFor AlF (111), we choose the surface with 3/7 ML F defects as the research 3 object. In addition, other metal fluorides are studied with their ideal surface. As shown in the Table 2, the average bader charge q of KF is the smallest, which means that its surface acidity is the lowest relative to other fluorides. CaF2, BaF2 and SrF2 have similar crystal structure, so their average charge q has marginal difference, but they still have the same trend as surface acidity in experiment. This is because the electronegativity follows Ca > Sr > Ba, and it is well known that the smaller the electronegativity of metals atoms, the harder it is to bind the surrounding electrons. The bader q of AlF3 is the largest because of the absence of F on its surface. From the theoretical calculations above, we found the trend of q is consistent with the theoretical reaction enthalpy of the first step (ER) and experimental conversion rate, and the trend is consistent with the experimental results in a series of fluorides, that is, higher surface acidity leads to higher conversion. It is well accepted that C-F bonds usually have much higher bond dissociation energy than that of C-Cl bonds 47. Consequently, strong Lewis acidic sites facilitates the activation of C-F bonds and therefore the reactions of DeHF. Hence, it is reasonable that AlF3 catalyst with strong Lewis acidic sites exhibits very high selectivity to VCF (DeHF) instead of main product VDF (DeHCl). It is confirmed by fluorinated Cr2O3 catalyst. As indicated in Fig. 6, fluorinated Cr2O3 catalyst also possesses significant amounts of strong acidic sites (with NH3 desorption temperatures above 550 oC in Fig. 6). As a result, high selectivity to VCF is also achieved for fluorinated Cr2O3 catalyst in Table 1. However, although BaF2 and SrF2 catalysts also exhibit strong acidic sites, they exclusively favor the formation of VDF (DeHCl reaction). In addition, no strong acidic sites were detected for LaF3 catalyst, significant selectivity to VDF, VCF and CH3CF3 was observed. Clearly, the above results can not be solely explained by the difference in strength of acidic sites. 18



4.2 Effect of affinity of metal fluorides to F and Cl on selectivity For the catalytic conversion of HCFC-142b, the first step of the catalytic reaction is adsorption of HCFC-142b over the surface of catalyst followed by the activation of C-F bond (DeHF) or C-Cl bond (DeHCl). Clearly, the adsorption plays a major role in the reactions. Consequently, the affinity of catalyst surface to F and Cl is responsible for the selectivity of HCFC-142b conversion. During the adsorption, if the catalyst tends to interact with Cl in HCFC-142b (high affinity to Cl), C-Cl bond is first activated resulting in the reaction of DeHCl and formation of VDF. Otherwise, activation of CF is facilitated leading to the reaction of DeHF and formation of VCF.

100

Selectivity to VDF, %

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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80 60 40 20 0 -20 0

20 40 60 80 100 120 Δ G, kJ/mol r

Figure 8. Selectivity to VDF as a function of Gibbs free energy changes of reactions between metal fluorides and HCl. HCFC-142b pyrolysis reactions were carried out at 400 oC, pressure of 1 bar and GHSV (HCFC-142b) of 600 h−1. In the present study, the reactions between metal fluorides and HCl forming metal chlorides are adopted as the indicator of the affinity of the metal fluoride catalysts to Cl (see Reaction 4, where M denotes K, Mg, Ca, Sr, Ba, La, Cr, Zn and Al respectively). The Gibbs free energy changes of reactions between metal fluorides and HCl are listed in Table S4. The Gibbs free energy change varies with metal fluorides significantly indicating that the affinity of metal fluorides to Cl differs dramatically.

MFx + xHCl → MClx + xHF

19


(4)

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Then, the Gibbs free energy changes of these reactions are correlated with their selectivity to VDF as demonstrated in Fig. 8. Interestingly, clear relationship between Gibbs free energy change and selectivity to VDF is identified. As AlCl3 is in the form of gas-phase under reaction conditions, it evaporates easily resulting in the loss of Al species. Hence, AlF3 is not included in Fig. 8. It is confirmed that selectivity to VDF decreases with ΔrG of reactions between metal fluorides and HCl. Low ΔrG indicates the high probability of reaction trend and high affinity to Cl. It reinforces the conclusion that high affinity to Cl facilitates the activation of C-Cl bond in HCFC-142b and production of VDF. Clearly, affinity to Cl of the metal fluoride catalysts is responsible for the selectivity. In order to explore the essential factors that influence the selectivity by DFT calculation, we calculated the two parts of reaction enthalpy (ER) of the first step, namely C-F or C-Cl bond dissociation without catalyst (the reaction enthalpy is denote as EdeF· and EdeCl· respectively) and adsorption of F or Cl radicals on the surface (the reaction enthalpy is denote as EFads and EClads respectively), i.e. ER-DeCl = EdeCl· + EClads, ER-DeF = EdeF· + EFads. The reactions are as follow. CH3CF2Cl → CH3CFCl· + F·

(12)

CH3CF2Cl → CH3CF2· + Cl·,·

(13)

F· + * → F*

(14)

Cl· + * → Cl*

(15)

As calculated by DFT, EdeF (5.457 eV)· is higher than EdeCl·(3.823 eV), and the difference (EdeF· - EdeCl·) is 1.63 eV in the absence of catalyst, which means the dissociation of C-F bond is more difficult than C-Cl bond. The results of EFads and EClads in the presence of catalysts are shown in Fig. 9 and Table S5. Obviously, this series of catalysts all have better ability to adsorb F radicals. Therefore, the key to influence selectivity is the extent of the difference in the ability of metal fluoride to adsorb F and Cl radical (i.e. the difference in EFads and EClads) for each metal fluoride. That is, the metal fluoride with a difference of EFads and EClads less than 1.63 eV exhibits selectivity to VCF (DeHF), while those above 1.63 eV are tend to generate VDF (DeHCl). 20



Figure 9. The reaction enthalpy of adsorption of F (EFads) or Cl (EClads) radicals on metal fluorides. We calculated the difference between EFads and EClads on different fluorides to further describe the selectivity of different metal fluorides (Details are listed in Table S6). As shown in Fig. 10, the difference between EFads and EClads of KF, CaF2, SrF2 and BaF2 is lower than 1.63 eV, so all of them show selectivity to VCF (DeHCl). While AlF3 favors the formation of VDF (DeHF), which is due to the extent of difference between F and Cl radical adsorption capacity is large enough (i.e. the difference between EFads and EClads is higher than 1.63 eV).

Figure 10. The calculated difference between EFads and EClads for each metal fluoride. The dotted line represents a value of 1.63 eV.

21


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Page 23 of 31

Obviously, through the above DFT theoretical study, the metal fluorides in this study all tend to adsorb F radical, while the difference in selectivity is due to the degree of difference in adsorption capacity.

4.3 Effect of affinity of metal fluorides to F and Cl on stability As depicted in Table 1, both BaF2 and SrF2 achieve satisfactory conversion of HCFC-142b. In addition, both catalysts exhibit high selectivity to target product, VDF. Clearly, BaF2 and SrF2 are potential catalysts for the pyrolysis of HCFC-142b. However, for a proper catalyst, in addition to the conversion and selectivity, the stability is also the key issue. The stability of BaF2 and SrF2 catalysts are evaluated and the results are displayed in Fig. 11. 100

100

a

Selectivity of VDF /%

Conversion. of HCFC-142b, %

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


80 60 40 20 0

BaF2 0

1

2

SrF2

3 4 5 6 7 Time on stream, h

8

80 60 40 20 0

9

b

BaF2 0

1

2 3 4 5 6 Time on stream, h

SrF2 7

8

9

Figure 11. Conversion of HCFC-142b (a) and selectivity to VDF (b) over BaF2 and SrF2 catalysts as a function of time on stream at 350 oC. Reaction conditions: 1 bar, N2: HCFC-142b of 1:1, GHSV (HCFC-142b) of 600h-1. . As indicated in Fig. 11, although similar initial conversion levels of HCFC-142b are achieved over BaF2 and SrF2 catalysts, BaF2 shows much higher selectivity to VDF than that of SrF2. Hence, BaF2 possesses superior activity than SrF2. Unfortunately, conversion of HCFC-142b over BaF2 drops with time on stream dramatically. The conversion on BaF2 decreases from 70% to 46% in less than 10 h. By contrast, although

22



Page 24 of 31

with relatively low selectivity to VDF, the conversion over SrF2 catalyst keeps unchanged and no noticeable deactivation is observed. According to the thermodynamic data in Table S4, the Gibbs free energy of BaF2 catalyst for the exchange of Cl/F with HCl is -11.08 kJ/mol, while it is 23.42 kJ/mol for SrF2. We suggest that BaF2 catalyst tends to react with HCl and other Cl containing species transforming into BaClF and BaCl2 under reaction conditions 48. By contrast, SrF2 is much more stable than BaF2. Hence, we suggest that the transformation of BaF2 during reaction is responsible for the deactivation. To further verify this conclusion, the surface elemental content of BaF2 and SrF2 catalyst following different reaction times were determined by EDS and the results are listed in Table 3 and Table S7.

Table 3 The surface elemental content of BaF2 catalyst following different reaction time determined by EDS. Sample BaF2-0h BaF2-2h BaF2-4h BaF2-6h BaF2-8h

Elemental content, mol% O F Cl 5.2 58.08 0 2.41 25.39 24.64 2.32 19.01 25.26 2.11 20.96 25.86 1.98 19.38 26.67

C 0 19.26 26.29 23.38 24.28

Cl/F Ba 41.92 30.71 29.44 29.80 29.68

0 1.07 1.51 1.38 1.53

Prior to the reaction, fresh BaF2 catalyst is mainly composed of Ba and F, no Cl is detected. However, following reaction, significant amounts of Cl are found. Following reaction of 4 h, the content of Cl is higher than 25% which is even much higher than the content of F. The molar ratio of Cl/F is higher than 1.5. The results strongly suggests that significant amounts of F in BaF2 were replaced by Cl during pyrolysis of HCFC142b. Similar results have been identified during the dehydrohalogenation of 3-chloro1,1,1,3-tetrafluorobutane

26.

As indicated in Table S7, large amounts of Cl are also

detected over SrF2 catalyst. However, the content of Cl for SrF2 is significantly lower than that of SrF2. The Cl/F ratio is around 0.9 following reaction of 4 h which is almost 50% lower than that of BaF2. The chlorination of SrF2 and BaF2 catalysts is partially 23


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attributed to the chlorination of SrO and BaO. As listed in Table 3 and Table S7, BaF2 and SrF2 contain small amounts of O. Therefore, BaO or SrO transform to BaCl2 and SrCl2 respectively. As argued previously, high affinity of BaF2 surface to Cl facilitates the activation of C-Cl bond and consequently promotes the production of VDF. However, it also leads to the chlorination of BaF2 forming BaCl2 which exhibits rather low affinity to Cl. As a result, the conversion of HCFC-142b decreases with time on stream following the transformation of BaF2 to BaCl2. By contrast, SrF2 has relatively low affinity to Cl (relatively high Gibbs free energy). Hence, high stability is achieved for SrF2 catalyst. This discussion is further supported by the results of HCFC-142b pyrolysis over BaF2 at elevated temperatures. As presented in Fig. S2, with the increase in reaction temperature, faster deactivation is resulted although no significant change in selectivity was detected. With the increase in reaction temperature, the Gibbs free energy of the reaction between BaF2 and HCl drops (Table S8). Therefore, at high temperatures, BaF2 has high affinity to Cl. However, SrF2 still shows positive Gibbs free energies with the increase in reaction temperature. 4.4 Effect of F-defects on the performance of metal fluoride catalysts As discussed above, AlF3 (111) surfaces tend to have a large number of F defects, while KF (110), CaF2 (111), SrF2 (111) and BaF2 (111) are difficult to form surface F defects. In view of the fact that only AlF3 tends to form defects in the above metal fluorides, we investigate the effect of defects on the catalytic activity on the surface of AlF3 (111). Similar to the above method, we simulated the reaction enthalpy of dissociation of C-F (DeF) or C-Cl (DeCl) bonds on the AlF3 (111) surface with 1/14, 1/7, 3/14, 2/7, 5/14 and 3/7 ML F defects respectively. As shown in Fig. 12a, AlF3 exhibits selectivity to VCF (DeHCl) on surfaces with less F defects (Details are listed in Table S9 and S10). With the F defects increase, the ability to dissociate C-Cl bond is weakened, and the ability to dissociate C-F bond is relatively improved on AlF3 (111). When the F defect ratio is large enough, AlF3 exhibits selectivity to VDF (DeHF). We can further infer the change in selectivity from 24



Fig. 12b. As the F defects increase, the difference between EFads and EClads gradually increases until it is finally higher than 1.63 eV, which shows a change from selectivity to VCF (DeHCl) to selectivity for VDF (DeHF). As discussed above, the surface F defects of AlF3(111) also have an effect on its selectivity.

Figure 12. The calculated (a) reaction enthalpy of DeF and DeCl and (b) difference between EFads and EClads on AlF3 (111) with different proportions of F defects. The dotted line represents a value of 1.63 eV.

5. Conclusions In this work, the performances of dehydrohalogenation over metal fluorides including KF, MgF2, CaF2, SrF2, BaF2, LaF3, CrF3, ZnF2 and AlF3 prepared via the precipitation synthesis were investigated. KF, ZnF2 and CaF2 are relatively inactive for the dehydrohalogenation of HCFC-142b. The traditional Lewis acid metal fluorides, such as AlF3, MgF2 and CrF3 exhibit higher activity with close to 100% conversion of HCFC-142b at 400 oC and 450 oC. Over catalysts, DeHCl, DeHF and Cl/F exchange reactions compete from each other, resulting in the spontaneous formation of VDF, Vinylidene chlorofluoride (VCF), CH3CF3 (HFC-143a) and CH3CCl2F (HCFC-141b). Acidic sites of these metal fluorides are majorly responsible for DeHCl, DeHF and Cl/F exchange reactions. We further confirmed the relationship between acidity and conversion by calculating the bader charges q, the higher the value of q means the stronger the surface acidity and thus the higher conversion rate. 25


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The affinity to Cl of the metal fluorides is responsible for the selectivity, i.e. metal fluorides with a value of EFads-EClads greater than 1.63 eV exhibits selectivity for VDF (DeHF), and vice versa to VCF (DeHCl). However, the high affinity to Cl also leads to the chlorination of metal fluorides forming metal chlorides. As a result, the conversion of HCFC-142b decreases with time on stream following the phase transformation. Finally, we investigated the effect of F defects on the conversion rate on AlF3 (111). As the F defects increase, the value of EFads-EClads increases, and the selectivity changes from VCF (DeHCl) to VDF (DeHF), which further confirming our study on the selectivity of different fluorides.

Associated content Supporting Information. N2 adsorption isotherms of all metal fluoride catalysts, Surface areas and pore volumes of fresh metal fluoride catalysts, reaction enthalpy of DeF (ER-DeF) and DeCl (ER-DeCl) on different metal fluorides, the peak area of NH3TPD obtained by integral and the peak area ratio with KF as reference, Gibbs free energy changes of reactions between metal fluorides and HCl, the reaction enthalpy of adsorption of F (EFads) or Cl (EClads) radicals, the difference between EFads and EClads for each metal fluoride, surface elemental content of SrF2 catalyst following different reaction time determined by EDS, effect of reaction temperature on the stability of BaF2 catalyst for the pyrolysis of HCFC-142b to vinylidene fluoride, Gibbs Free Energy of reaction between BaF2 or SrF2 and HCl at elevated temperatures, reaction enthalpy of DeF (ER-DeF) and DeCl (ER-DeCl) on AlF3 (111) with different proportions of F defects, difference between EFads and EClads on AlF3 (111) with different proportions of F defects. These materials are available free of charge via the Internet at http://pubs.acs.org.

Author information Corresponding authors *Tel: +86-15158074035. E-mail: [email protected] (W.F. Han) [email protected] (W.Y. Song) ORCID 26



Wenfeng Han: 0000-0003-2252-9311 Notes The authors declare no competing financial interest. Acknowledgements This research was supported by Zhejiang Provincial Natural Science Foundation of China under Grant No. LY19B060009.

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Experimental and DFT mechanistic study of dehydrohalogenation of 1

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