Article pubs.acs.org/cm
In Situ Infrared Spectroscopic Study of Atomic Layer-Deposited TiO2 Thin Films by Nonaqueous Routes Karla Bernal Ramos,*,† Guylhaine Clavel,‡ Catherine Marichy,‡ Wilfredo Cabrera,† Nicola Pinna,§ and Yves J. Chabal† †
Department of Materials Science and Engineering, University of Texas at Dallas, Richardson, Texas 75080, United States Department of Chemistry, CICECO, University of Aveiro, 3810-193 Aveiro, Portugal § Humboldt-Universität zu Berlin, Institut für Chemie, Brook-Taylor-Strasse 2, 12489 Berlin, Germany ‡
ABSTRACT: The mechanisms of growth of TiO2 thin films by atomic layer deposition (ALD) using either acetic acid or ozone as the oxygen source and titanium isopropoxide as the metal source are investigated by in situ Fourier transform infrared spectroscopy (FTIR) and ex situ X-ray photoelectron spectroscopy. The FTIR study of the acetic acid-based process clearly shows a ligand exchange leading to the formation of surface acetate species (vibrational bands at 1527 and 1440 cm−1) during the acetic acid pulse. Their removal during the metal alkoxide pulse takes place via the elimination of an ester and the formation of Ti−O−Ti bonds. These findings confirm the expected ester elimination condensation mechanism and demonstrate that the reaction proceeds without intermediate surface hydroxyl species. The in situ FTIR study of the O3-based ALD process demonstrates similarities with the process described above, with formation of surface formate and/or carbonate species upon exposure of the surface titanium alkoxide species to ozone. These surface species are removed by the subsequent titanium isopropoxide pulse, leading to the formation of Ti−O−Ti bonds. KEYWORDS: atomic layer deposition, titanium dioxide, nonaqueous sol−gel routes, FTIR
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INTRODUCTION Deposition of metal oxide thin films with precise control of the thickness and composition at the atomic scale is critical to achieving the current demands of the manufacturing of the next-generation integrated circuits.1 In this respect, atomic layer deposition (ALD) is a promising technique that provides a well-controlled, uniform, and conformal film growth with a high degree of both reproducibility and controllability.2 ALD of metal oxides is generally achieved by the reaction between a metal precursor and an oxygen source (e.g., water, oxygen plasma, and ozone). Depending on the oxygen source used in an ALD process, different reactions that account for metal oxide formation have been identified. However, when water or hydrogen peroxide is used (and to some extent also with ozone), the main reaction pathway observed consists of a succession of hydrolysis and condensation steps with the formation of an OH-terminated surface after the oxygen source pulse. The film grows then via a reaction between the supplied metal precursor and the hydroxyl surface groups. A few ALD processes have been developed in which the metal−oxygen− metal (M−O−M) bond is supposed to be obtained without an intermediate hydrolysis step.3−7 Some of these deposition processes were inspired by liquid phase syntheses of metal oxide and hybrid materials, and they will be identified, in this paper as for their solution synthesis analogue, by the term nonaqueous sol−gel routes. For instance, the first nonaqueous © XXXX American Chemical Society
sol−gel route transposed to ALD, reported by Brei et al. and later by Ritala et al., made use of a metal alkoxide as the oxygen source.3,8 Even though metal oxides could be deposited by this approach, the depositions were performed at higher temperatures than similar approaches in solution, and thus, the decomposition of the metal alkoxide leading to the formation of OH- surface species could not be safely excluded. This different behavior can be explained by the different nature of the reactions in solution and ALD in terms of the time scale, precursor concentrations, the formation of intermediate species, and catalytic effects.4 Another approach that does not suffer from the same drawback is the formation of metal oxides using metal alkoxides and carboxylic acids as the oxygen source, which proceeds via a ligand exchange followed by an ester elimination condensation step.9 Besides these nonaqueous sol− gel routes, the formation of surfaces species other than hydroxyl groups has also been observed in the case of ozone-based processes. It is assumed that they are the active sites for the next precursor pulse reaction and that the expected mechanisms are different from those for hydrolyzed surfaces.5,7,10−13 Unfortunately, the reactions responsible for oxide growth, i.e., the condensation steps of those surface species Received: January 15, 2013
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Figure 1. Difference infrared absorbance spectra of the four first (a) and four last (b) ALD half-cycles referenced to their respective preceding halfcycle. nth TIP and nth Ac mean that the spectra were recorded after the titanium isopropoxide/purge and acetic acid/purge of cycle n, respectively.
Figure 2. Difference infrared absorbance spectra during initial 1−2 (a) and 12−13 (b) ALD cycles after titanium isopropoxide (nth TIP) and ozone (nth O3) pulses of 20s (200 sccm of O3) referenced to each preceding treatment. reservoir at room temperature. When ozone was used as the oxygen source, O3 was introduced by pneumatic valves in which the exposure time and flow rate ranged from 5 to 20 s pulse and 100 to 500 sccm, respectively, while the TIP parameters were kept identical to those of the TIP/acetic acid process. Ozone was produced using an In-USA Ozone generator with a constant concentration of 200 g/nm3. After each half-ALD cycle, in situ infrared absorbance spectroscopy was conducted using a Thermo Nicolet 6700 interferometer with a 400−4000 cm−1 frequency range. During the IR measurement, the sample was kept at 60 °C and oriented so that the IR beam incident angle was close to the Brewster angle. The substrate was kept in the same exact position during the whole ALD and IR characterization; only its temperature was adjusted. X-ray photoelectron spectroscopy (XPS) was used to investigate the composition and chemical state of the film. XPS measurements were conducted by using an Al Kα (1486.6 eV) X-ray source at a chamber base pressure of 10−10 Torr, and spectra were recorded using a 16-channel detector with a hemispherical analyzer. Sputtering was performed directly on the sample using an argon ion gun.
with the incoming metal precursor, are still not clearly identified. In this work, in situ Fourier transform infrared spectroscopy (FTIR) is used to investigate the reaction mechanism responsible for the deposition of TiO2 during two different processes that are supposed to be hydroxyl-free. The study focuses on the initial surface reaction mechanisms of TiO2 thin film growth by reaction of a titanium alkoxide with either a carboxylic acid or ozone as the oxygen source. Both reaction pathways will be discussed and compared.
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EXPERIMENTAL SECTION
Double-side-polished Si(100) float-zone grown substrates (lightly Pdoped, ρ ∼ 10 Ω cm) were used. The substrates were treated in a piranha solution (1:3 H2O2/H2SO4 mixture) for 30 min to produce OH-terminated oxide surfaces, typically 1−2 nm thick.14 After being thoroughly rinsed with deionized water and blown dry with nitrogen (N2), the sample was loaded into the ALD reactor, with a base pressure of 10−4 Torr. The deposition took place in a homemade ALD tool in which the reaction chamber was connected to a spectrophotometer, allowing in situ FTIR study. The ALD reactor was described previously.15 The Si substrate was kept at 200 °C during ALD growth. The substrate was exposed to alternate pulses of titanium isopropoxide (TIP) and the oxygen source, each followed by a 5 min purging with ultrapure N2 gas to prevent cross gas phase reactions between precursors. The TIP precursor was exposed for 1 s through an ampule heated to 80 °C. Deuterated acetic acid-d4 was used to differentiate between the species coming from the alkoxide and those coming from the acetic acid. A 0.03 s pulse was introduced, drawn directly from the
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RESULTS AND DISCUSSION
In situ Fourier transform infrared (FTIR) measurements were performed to study the mechanism of growth of TiO2 deposited by both a titanium isopropoxide/acetic acid and a titanium isopropoxide/ozone precursor combination. The IR absorbance spectra of the film were recorded after each precursor/purge sequence, i.e., after each half-cycle. Each spectrum was referenced to the previous one; i.e., the spectrum of a half-cycle became the background of the subsequent halfB
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well as the CH3 bending modes at 1332, 1366, and 1380 cm−1 associated with the TIP absorption, are clearly observed.12,16,17 Moreover, the symmetric and asymmetric stretching modes of the Ti−O−C bond, located at 1020 and 1130 cm−1, are also observed.18−20 The subsequent acetic acid (“1st Ac”) exposure led to the loss of the characteristic TIP-related modes and the appearance of symmetric and asymmetric stretching modes of the O−C−O bond of acetate species, at 1438 and 1527 cm−1, respectively.21−23 In fact, even though the formation of such species is not involved in TIP exposure, one can note that they were already present in a very small amount in the first spectrum. This is attributed to a small contamination of the ALD chamber from previous deposition by the acetic acid. Acetates may have different coordination modes, η1, η2, and μ coordination, namely, the monodentate, also called ester, chelating bidentate, and bridge binding modes, respectively. By considering the shift in frequency between the observed asymmetric and symmetric stretching modes, the monodentate coordination mode can be excluded.24,25 Therefore, only the two different acetate bidentate geometries need to be considered: a chelating structure, in which each acetate binds to a single Ti center, and a bridging structure, in which the acetate bridges between two Ti centers. The frequency separation (Δν ∼ 85 cm−1) between νasym(OCO) and νsym(OCO) suggests that CH3COO− acts as a chelating bidentate ligand.21,23,26,27 The ligand exchange, i.e., the replacement of the surface alkoxide by acetate species, was observed in the initial ALD cycles. Figure 1b shows the steadiness of this ligand exchange during the complete process of TiO2 growth. However, there is a difference in the intensity of the acetate band after the acetic pulses (positive band) and the alkoxide pulses (negative band) within one cycle. This indicates an incomplete removal of the
cycle spectrum. Thus, those differential IR spectra enhance the change in the surface species composition after each step.
Figure 3. Infrared absorbance spectra of TiO2 after 13 ALD cycles for (a) acetic acid and (b) ozone processes referenced to the initial SiO2 surface.
TIP/Acetic Acid Process. Differential IR spectra of the initial and last ALD cycles of the TIP/acetic acid process are presented in Figure 1. The spectrum of the first TIP pulse (“1st TIP”) on the hydroxylated SiO2 surface (Figure 1a) shows the loss of the characteristic signature of -OH groups at 3740 cm−1 (negative peak) and the appearance of the alkoxide-related vibrational modes (positive peaks), proving the reaction of titanium isopropoxide with the surface hydroxyl group. The symmetric CH3 mode at 2870 cm−1 and the two nondegenerate asymmetric CH3 stretching modes at 2973 and 2930 cm−1, as
Figure 4. X-ray photoemission spectra (black curves) of Ti 2p (a), O 1s (b), C 1s (c), and Si 2p (d) orbitals of the TiO2 film after 13 ALD cycles deposited from TIP and acetic acid. The colored curves represent the different contributions. C
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acetate species or acetic acid residues in the chamber because of the low temperature of its walls (100 °C) (see further discussion in the XPS part). The pulses of the titanium isopropoxide on the acetate surface remove the acetate species and restore an alkoxide-terminated surface (Ti−O−C band between 1020 and 1130 cm−1). From the observed Ti−O−C bands and their intensity, the reaction appears to be robust and steady after the first cycle. Furthermore, the IR spectra show no band in the 3600−3800 cm−1 region, i.e., no formation of intermediate OH groups at the surface of the film, consistent with the direct reaction of the alkoxide precursor with the acetate surface species. In solution, the reactions between metal alkoxides and carboxylic acids, as well as between metal acetates and metal alkoxides, are well-documented. For example, in traditional sol−gel chemistry, the reaction of titanium isopropoxide and acetic acid is used to produce the TiO2 gel, clusters, and particles. As the main reaction pathway, it has been demonstrated that the metal precursor is hydrolyzed by water that is generated in situ by the reaction of carboxylic acids and alcohol products. Additionally, a competitive mechanism has also been proposed, in which a direct esterification takes place between acetate and alkoxy ligands but could not be experimentally verified because the products of both mechanisms are identical.17,21,28 The reaction mechanism responsible for M−O−M bond formation by reaction of metal alkoxides with a carboxylic acid is somewhat different in the ALD process. Here, a direct condensation between metal alkoxides and carboxylates occurs without hydrolyzed species or H2O formation. A similar direct condensation reaction was already reported in solution between metal carboxylates and alkoxides, called ester elimination, which is similar to the acid-catalyzed transesterification reaction.29 This reaction was used for the formation of clusters as well as metal and mixed-metal oxides. Thus, from the in situ FTIR measurements, the general reaction pathway derived is in agreement with the mechanism previously suggested by Rauwel et al., based on GC−MS studies,9 and can be schematized as follows:
Figure 5. X-ray photoemission spectrum of the C 1s orbital before and the spectrum after Ar sputtering for 2 min.
O3 pulse could be similar to those reported for ALD processes using alternative oxygen sources, such as carboxylic acid. The differential IR spectra of the TIP/ozone process, when a 20 s pulse and 200 sccm of O3 are used, are presented in Figure 2. The loss of the -OH group at 3740 cm−1 and the symmetric and asymmetric Ti−O−C bands at 1018 and 1116 cm−1, respectively, together with CHx stretching vibrations (2800− 3000 cm−1) after the first TIP (1st TIP) pulse can be observed in Figure 2a. After the ozone exposure, the TIP-related modes are removed and two small bands at 1605 and 1580 cm−1 are observed. Those bands are assigned to monodentate formate and to chelate bidentate carbonate species, respectively.30−32 The peak intensities appear to be weaker for this process than for the acetic acid one. However, the reactions are reaching a steady regime after a few cycles, as seen in Figure 2b. Similar results were observed when the O3 parameters were varied to determine the influence of pulse length (5, 10, and 20 s) and flow (100, 200, and 500 sccm) during the ALD process. The relative intensities of the carbonate and formate peaks are relatively stable when these parameters are varied. Previous studies established that different species are formed when ozone is used as the oxygen source in an ALD process. For example, Kwon et al.7,10 reported that the main growth mechanism for Al2O3 deposition from trimethylaluminum involves the formation of formate species and oxidation of the substrate after ozone exposure at a high flow rate (200 sccm). Formate formation was also reported by Goldstein et al. for the same process.13 Later, Kwon et al. reported the preferential formation of alkoxide species and the reduction of substrate oxidation when the ozone flow rate is reduced.7 In the case of titanium tetraisopropoxide, Rai et al. observed bands attributed to CO32− at 1620 cm−1 and around 1200−1250 cm−1 when reacting with plasma O2, as well as CO and CO2 formation.33 In the case of the O3 process, they observed weak bands in the 1300−1700 cm−1 region, also attributed to carbonate species.6 However, the reactions responsible for the growth of oxide from those species with metal complexes (e.g., metal alkoxides) are still not clearly identified.10,12 Moreover, in some cases, it was reported that ozone can produce water as a byproduct leading to a hydroxyl-terminated surface.34,35 Thus, in those cases, the metal precursor half-reactions are similar to processes using H2O or H2O2 sources, consisting of a succession of hydrolysis and condensation steps similar to the aqueous sol−gel chemistry in solution. In this study, no
≡M−OR′ + RCOOH → ≡M−OOCR + R′OH ≡M−OOCR + M−OR′ → ≡M−O−M≡+RCOOR′
During the TIP pulse, the metal alkoxide reacts with either the -OH surface groups in the case of the first TIP exposure or the acetate species formed during the previous cycle, leading to the release of an ester and the formation of Ti−O−Ti bonds. During the acetic acid exposure, there is a ligand exchange between the evaporated carboxylic acid and the surface alkoxy species, leading to the formation of acetate surface groups, upon elimination of the corresponding alcohol. This observation proves that surface carboxylates can be removed by reaction with metal alkoxides, even at relatively low temperatures. TIP/Ozone Process. From the mechanistic studies found in the literature on ozone-based ALD processes, the reaction between ozone and metal alkoxide leads to the formation of alkoxide, carbonate, and formate species and/or hydroxyl groups at the surface depending on the metal precursor and the O3 flow rate. Those surface species are the reactive sites for the next precursor pulse. Consequently, the reaction mechanisms leading to the removal of those surface groups generated by the D
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Figure 6. X-ray photoemission spectra (black curves) of Ti 2p (a), O 1s (b), C 1s (c), and Si 2p (d) orbitals of the TiO2 film after 13 ALD cycles of TIP and ozone. The colored curves represent the different peak contributions.
oxygen bonded to two Ti atoms (Ti−O−Ti), and its shoulder (blue curve) at 531.9 eV is assigned to Si from the SiOx interface, existing between the Si substrate and titania thin film, and to carboxylate species. Examination of the Si 2p spectrum in Figure 4d shows four peaks, corresponding to two contributions each split into their respective Si 2p and Si 2p3/2 contributions. The lower-intensity contribution with its stronger contribution centered at 99.3 eV is assigned to bulk Si, while the higher-intensity contribution with its stronger contribution centered at 102 eV is assigned to the suboxide, SiOx, at the interface. Figure 4c shows the C 1s spectrum. The peak located at 285 eV (red curve) is due to hydrocarbon contamination resulting from the exposure of the film to air during the ex situ transfer and loading into the XPS chamber. The peak at 289 eV (blue curve) is assigned to the carbon of the acetate species left in the film, with some contribution from C−O−Ti bonds characterized by the peak at 286.8 eV (green curve). This finding is consistent with the IR spectra in Figure 3, showing the presence of acetate species in the film after 13 ALD cycles. Figure 5 presents the XPS C 1s spectra of a TiO2 film before and after Ar sputtering for 2 min. This particular TiO2 film was growth in a hot wall reactor (150 °C) to clarify the issue of the acetic acid residues on the chamber. The spectra show an important amount of carbon before sputtering due to air exposure during transport to the XPS chamber. The C concentration decreases with sputtering time, proving that there is not a considerable amount of C left in the bulk of the film and that the accumulation of carbon (observed in Figure 3) is mainly due to the desorption of acetic acid residues accumulated on the chamber walls because of their relatively low temperature. To avoid the accumulation of carbon when acetic acid is used as an oxygen source, the walls of the ALD reactor must therefore be kept at a temperature of ≥150 °C.
hydroxyl species are observed in the FTIR spectra using alternative oxygen sources such as acetic acid and ozone. This finding underscores the importance of studies of the initial ALD cycles for the fundamental understanding of the chemical reactions. In general, a systematic study of the surface chemistry controlling the film properties needs to be performed to achieve the development of high-quality thin film materials. The absorption spectra, recorded after the last exposure of the acetic acid and ozone after 13 cycles and referenced to the respective initial SiO2 surfaces, are shown in Figure 3. The vibrational modes at 430 and 850 cm−1 are assigned to the longitudinal and transverse optical (LO and TO, respectively) phonon modes of titanium dioxide, respectively.36 The absorption spectrum of the 13th acetic acid exposure clearly shows the TiO2 bands. Along with these modes, an increase in the level of acetate species at 1440 and 1527 cm−1 is consistent with an incomplete reaction during the alkoxide pulse or the contamination of acetic acid gas residues on the chamber walls, leading to the remaining acetate groups. For the ozone process, the TiO2 vibrational bands are weaker than those obtained with the acetic acid process, meaning that TiO2 growth is faster when a carboxylic acid is used as the oxygen source in this case. To verify a potential organic contamination and to identify the chemical states of the species present in the films, we conducted ex situ XPS studies with the TiO2 films after 13 ALD cycles. XPS Characterization. The Ti 2p XPS spectrum (Figure 4a) of the TiO2 film grown by the acetic acid/TIP process shows the presence of the two well-formed peaks at 459 and 464.7 eV corresponding to the 2p3/2 (red curve) and 2p1/2 (blue curve) orbitals, respectively, of Ti atoms in the Ti4+ state,37 thus confirming the growth of TiO2 on the surface. The O 1s spectrum of the TiO2 film in Figure 4b shows a peak and a shoulder. The peak (red curve) at 530.5 eV is assigned to E
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Figure 6 shows the XPS spectra of the TiO2 film grown by TIP/ozone with a 20 s pulse and 200 sccm of O3. The Ti 2p spectra are shown in Figure 6a, where peaks at 458.5 (red curve) and 464.2 eV (blue curve) are observed, corresponding to the Ti 2p3/2 and 2p1/2 orbitals, respectively. In Figure 6b, the O 1s spectrum of the TiO2 film is presented. The peak at 530.5 eV (blue curve), assigned to the Ti−O−Ti bond, has an intensity much lower than that of the peak when acetic acid is used as the oxygen source, whereas the peak at 532.5 eV (red curve), corresponding to SiO2, has a higher intensity. The Si 2p spectra in Figure 6d show a broad peak at 100.4 eV indicating that more suboxide is formed at the interface when ozone is used. Figure 6c shows the C 1s spectrum, once again with a contribution of hydrocarbon contamination (at 284.9 eV, red curve) in the spectrum. The peak at 288.3 eV (blue curve) is assigned to the C associated with carbonates and/or formate groups formed after ozone exposure. The peak at 286.4 eV (green curve) corresponds to the C−O−Ti bonds. A comparison of the Si 2p XPS spectra between both ALD processes (using acetic acid and O3) shows that the initial surface reactions on the silicon oxide surface when acetic acid is used do not involve additional oxidation of the silicon substrate, allowing the detection of the Si bulk at 99.3 eV; in contrast, when ozone is used, no Si bulk signal is observed, suggesting the formation of a thicker interfacial layer.
REFERENCES
(1) Mistry, K.; Allen, C.; Auth, C.; Beattie, B.; et al. Electron Devices Meeting, 2007. IEDM 2007. IEEE International 2007, 247−250. (2) Ritala, M.; Leskela, M. In Handbook of Thin Film Materials; Nalwa, H. S., Ed.; Academic Press: San Diego: 2002; Vol. 1. (3) Ritala, M.; Kukli, K.; Rahtu, A.; Raisanen, P. I.; Leskela, M.; Sajavaara, T.; Keinonen, J. Science 2000, 288, 319−321. (4) Clavel, G.; Rauwel, E.; Willinger, M. G.; Pinna, N. J. Mater. Chem. 2009, 19, 454−462. (5) Wang, Y.; Dai, M.; Ho, M.-T.; Wielunski, L. S.; Chabal, Y. J. Appl. Phys. Lett. 2007, 90, 022906-3. (6) Rai, V. R.; Agarwal, S. J. Phys. Chem. C 2008, 112, 9552−9554. (7) Kwon, J.; Dai, M.; Halls, M. D.; Chabal, Y. J. Appl. Phys. Lett. 2010, 97, 162903. (8) Brei, V. V.; Kaspersky, V. A.; Gulyanitskaya, N. U. React. Kinet. Catal. Lett. 1993, 50, 415−421. (9) Rauwel, E.; Clavel, G.; Willinger, M.-G.; Rauwel, P.; Pinna, N. Angew. Chem., Int. Ed. 2008, 47, 3592−3595. (10) Kwon, J.; Dai, M.; Halls, M. D.; Chabal, Y. J. Chem. Mater. 2008, 20, 3248−3250. (11) Lee, H. J.; Park, M. H.; Min, Y.-S.; Clavel, G.; Pinna, N.; Hwang, C. S. J. Phys. Chem. C 2010, 114, 12736−12741. (12) Rai, V. R.; Vandalon, V.; Agarwal, S. Langmuir 2010, 26, 13732−13735. (13) Goldstein, D. N.; McCormick, J. A.; George, S. M. J. Phys. Chem. C 2008, 112, 19530−19539. (14) Plummer, J. D.; Deal, M. D.; Griffin, P. B. Silicon VLSI Technology. (15) Kwon, J.; Dai, M.; Halls, M. D.; Langereis, E.; Chabal, Y. J.; Gordon, R. G. J. Phys. Chem. C 2009, 113, 654−660. (16) Moran, P. D.; Bowmaker, G. A.; Cooney, R. P. Inorg. Chem. 1998, 37, 2741−2748. (17) Birnie, D. P., III; Bendzko, N. J. Mater. Chem. Phys. 1999, 59, 26−35. (18) Jensen, H.; Soloviev, A.; Li, Z.; Søgaard, E. G. Appl. Surf. Sci. 2005, 246, 239−249. (19) Ivanova, T.; Harizanova, A.; Surtchev, M. Mater. Lett. 2002, 55, 327−333. (20) Harris, M. T.; Singhal, A.; Look, J. L.; Smith-Kristensen, J. R.; Lin, J. S.; Toth, L. M. J. Sol-Gel Sci. Technol. 1997, 8, 41−47. (21) Doeuff, S.; Henry, M.; Sanchez, C.; Livage, J. J. Non-Cryst. Solids 1987, 89, 206−216. (22) Garcia, A. R.; Silva, J. L. d.; Ilharco, L. M. Surf. Sci. 1998, 415, 183−193. (23) Pei, Z.-F.; Ponec, V. Appl. Surf. Sci. 1996, 103, 171−182. (24) Deacon, G. B.; Phillips, R. J. Coord. Chem. Rev. 1980, 33, 227− 250. (25) Deacon, G. B.; Huber, F.; Phillips, R. J. Inorg. Chim. Acta 1985, 104, 41−45. (26) Rotzinger, F. P.; Kesselman-Truttmann, J. M.; Hug, S. J.; Shklover, V.; Gratzel, M. J. Phys. Chem. B 2004, 108, 5004−5017. (27) Nolan, N. T.; Seery, M. K.; Pillai, S. C. J. Phys. Chem. C 2009, 113, 16151−16157. (28) Steunou, N.; Robert, F.; Boubekeur, K.; Ribot, F.; Sanchez, C. Inorg. Chim. Acta 1998, 279, 144−151. (29) Caruso, J.; Hampden-Smith, M. J. J. Sol-Gel Sci. Technol. 1997, 8, 35−39. (30) Han, X.; Wang, X.; Xie, S.; Kuang, Q.; Ouyang, J.; Xie, Z.; Zheng, L. RSC Adv. 2012, 2, 3251−3253. (31) Brownson, J. R. S.; Tejedor-Tejedor, M. I.; Anderson, M. A. J. Phys. Chem. B 2006, 110, 12494−12499. (32) Liao, L.-F.; Lien, C.-F.; Shieh, D.-L.; Chen, M.-T.; Lin, J.-L. J. Phys. Chem. B 2002, 106, 11240−11245. (33) Rai, V. R.; Agarwal, S. J. Phys. Chem. C 2009, 113, 12962− 12965. (34) Knapas, K.; Ritala, M. Chem. Mater. 2008, 20, 5698−5705. (35) Rose, M.; Niinisto, J.; Endler, I.; Bartha, J. W.; Kucher, P.; Ritala, M. ACS Appl. Mater. Interfaces 2010, 2, 347−350.
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CONCLUSIONS These IR spectroscopic investigations of the ALD growth of TiO2 on SiO2-terminated Si surfaces have shown that the initial surface reaction takes place via the formation of carboxylate species under the elimination of an ester when carboxylic acid is used as the oxygen source. The growth rate is also higher than for the ozone-based process, as illustrated in the absorbance spectra in Figure 3. Carbonate and/or formate species are formed in the ozone-based TiO2 ALD, and less C accumulation is observed. The reaction between titanium isopropoxide and acetic acid has been clearly elucidated by in situ FTIR and the product species confirmed by FTIR and XPS characterization. However, the mechanism involved in TiO2 ALD using O3 is poorly understood, with disagreement among the few in situ studies previously performed on such systems. Nevertheless, combining the current in situ FTIR and the ex situ XPS characterization, we can conclude that monodentate formate and/or chelate bidentate carbonate species are formed. These species are not fully removed during the subsequent TIP exposure, leading to a low TiO2 growth rate.
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS The work at the University of Texas at Dallas was supported by the National Science Foundation (CHE-0911197) and partially by the NHARP of the Texas Higher Education Board. K.B.R. was partially supported by the Consejo Nacional de Ciencia y Tecnologia (CONACyT), Mexico. The work at the University of Aveiro was supported by the FCT project (PTDC/CTM/ 098361/2008) and FCT grant (SFRH/BD/71453/2010). F
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(36) Scarel, G.; Hirschmugl, C. J.; Yakovlev, V. V.; Sorbello, R. S.; Aita, C. R.; Tanaka, H.; Hisano, K. J. Appl. Phys. 2002, 91, 1118−1128. (37) Methaapanon, R.; Bent, S. F. J. Phys. Chem. C 2010, 114, 10498−10504.
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