Hydrometallurgy 142 (2014) 131–136

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Hydrometallurgy

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Eh–pH diagrams from 333.15 to 453.15 K for lithium–titaniumcomposite oxides and their synthesis in aqueous solution

Lin Li a, Yunjiao Li a,c,⁎, Cang Xu a, Vladimiros G. Papangelakis b,a, Guang Chu a, Guiliang Li c,Xuanyu Wang a, Long Kong a,c

a School of Metallurgy and Environment, Central South University, Changsha, Hunan 410083, PR Chinab Department of Chemical Engineering and Applied Chemistry, University of Toronto, 200 College Street, Toronto, Ontario M5S 3E5, Canadac Citic Dameng Mining Industries Limited, 18 Zhujin Road, Nanning, Guangxi 530028, PR China

⁎ Corresponding author at: Central South Univeand Environment, 932# South Lushan Road, ChangsTel.: +86 731 88830476; fax: +86 731 88710171.

E-mail address: yunjiaoli6601@hotmail.com (Y. Li).

0304-386X/$ – see front matter © 2013 Elsevier B.V. All rhttp://dx.doi.org/10.1016/j.hydromet.2013.11.010

a b s t r a c t

a r t i c l e i n f o

Article history:Received 28 April 2013Received in revised form 31 October 2013Accepted 6 November 2013Available online 4 December 2013

Keywords:Lithium–titanium composite oxidesLi–Ti–H2O systemPotential–pH diagramsSynthesis

The potential-pH diagrams for lithium–titanium–water system at temperatures of 333.15, 363.15, 393.15 and453.15 K and ion activities of 0.01, 0.1 and 1 of the dissolved specieswere constructed to predict the predominantareas of lithium–titanium composite oxides in the Li–Ti–H2O system. Empirical functions were applied toestimate the thermodynamic data, which are unavailable in the literature. The presented Eh–pH diagramsshow that temperature and ion activity have significant effects upon the stable regions of various speciesconsidered in Li–Ti–H2O system. With an increase in temperature, or ion activity, the dominant regions ofLi2TiO3(hc), Li4Ti5O12(hc) and Li4TiO4(hc) shift towards lower pH zones. Also, the predominant areas of Li4Ti5O12(hc)

and Li2TiO3(hc) shrink, while that of Li4TiO4(hc) enlarges. This demonstrates that the production of Li4Ti5O12 andLi2TiO3 is thermodynamically feasible by awet process. Experimentswere subsequently performed in light of theconstructed potential-pH diagrams. The results indicated that the potential-pH diagrams are consistent withexperiment. Therefore, the preparation of Li4Ti5O12 and Li2TiO3 in aqueous solution is feasible and the processcan be controlled in practice.

© 2013 Elsevier B.V. All rights reserved.

1. Introduction

A predominance area diagram depicts regions of stable speciesexisting in a given system at different equilibrium conditions. It is veryuseful in predicting reactions that produce the dominant species andmay occur under a particular set of conditions. Eh–pH diagrams areamong the most well-known types of stability diagrams (Osseo-Asareand Brown, 1979).

A large volume of published Eh–pH diagrams is restricted to thetemperature of 298.15 K because of lack of thermodynamic data athigher temperatures. However, many significant processes are carriedout at elevated temperatures. Because chemical equilibria in aqueoussolutions are affected with temperature significantly, it is alwaysinsightful to construct and use Eh–pH diagrams at higher temperatures(Townsend and Jr., 1970).

Various lithium–titanium composite oxides have been widelystudied and applied in different fields recently. For example, with its

rsity, School of Metallurgyha, Hunan 410083, PR China.

ights reserved.

excellent cycling performance and working life, the spinel Li4Ti5O12,which is known as a zero-strain insertion material, is identified as oneof the most promising alternative anode materials for lithium ionbatteries (Allen et al., 2006; Gao et al., 2007; Ohzuku et al., 1995). Onthe other hand, Li2TiO3 is one of the most promising candidatesamong tritium breeding materials, and has several useful propertiessuch as high lithium density, good chemical stability, excellent thermalconductivity, low activation characteristic and good tritium releasebehavior (Davis and Haasz, 1996; Roux et al., 1996; Casadio et al.,2004; Hoshino and Oikawa, 2011). Li2TiO3 is used as the precursor forlithium-ion sieve materials (Chen et al., 2009; Shen et al., 2012) forthe extraction of lithium from brines by adsorption due to its reactionefficiency and stable structure (Jiang, 2012; Liu et al., 2010; Zhanget al., 2010).

According to literature reports (Chen et al., 2008; Li et al. 2009, 2013;Sugita et al., 1990), Li4Ti5O12 was prepared bywet processes in aqueoussolution, along with a subsequent heat treatment. Metastable cubicα-Li2TiO3 and monoclinic β-Li2TiO3 were obtained by a hydrothermalmethod and a wet-chemistry route with a thermal treatment,respectively, in the publications (Laumann et al., 2010, 2011; Li et al.,2012; Tomiha et al., 2002). Although these materials are synthesizedin aqueous solution, the thermodynamics of the synthesis of Li–Ti–Ocomposites is rarely reported. For example, no Eh–pH diagrams of theLi–Ti–H2O system at elevated temperatures have been reportedpreviously. Qiu et al. (2010) used a two-parameter model and the

Table 1Various species considered in the Li–Ti–H2O system and their thermodynamic data at298.15 K.

Component ΔfG298.15Θ (Jmol−1) S298.15

Θ (Jmol−1)

Ti3+ −349.78a −238.78①

TiO2+ −577.39a −152.70②

TiO22+ −467.23a −98.20②

Li+ −293.30b 13.39b

H+ 0b 0b

H2O −237.14b 69.95b

TiH2(c) −80.30a 29.70b

TiO3(c) N/Ac 53.53③

TiO3·2H2O(c) −1173.00a 113.58⑤

TiO2(c) −888.40a 49.94b

Ti2O3(c) −1434.20a 78.78b

TiO2(hc) −821.30a 109.94⑤

Ti2O3(hc) −1388.00a 118.78⑤

Li4Ti5O12(c) −5852.48d 356.62④

Li4TiO4(c) −2209.72d 133.71④

Li2TiO3(c) −1579.80d 91.79b

Li4Ti5O12(hc) −5683.73d 506.62⑤

Li4TiO4(hc) −2209.72d 133.71⑤

Li2TiO3(hc) −1544.05d 121.79⑤

① to ⑤: from empirical relations as in main text.a Kelsall and Robbins (1990).b Dean (1991).c Not available and not used.d Qiu et al. (2010).

132 L. Li et al. / Hydrometallurgy 142 (2014) 131–136

homologous linear rule (Jiang, 1977) to estimate the standard Gibbsfree energy of formation of Li4Ti5O12 and Li4TiO4 and constructedpotential-pH diagrams of the Li–Ti–H2O system at 298.15 K. However,this study is restricted to room temperature only.

In this paper, a set of unknown thermodynamic properties ofLi–Ti–O compounds were estimated using empirical correlations.Then, the potential-pH diagrams for Li–Ti–H2O system at temperaturesof 333.15, 363.15, 393.15 and 453.15 K and ion activities of 0.01, 0.1 and1 for the dissolved species were constructed for predicting theformation of various Li–Ti–O composites. The effects of temperatureand ion activity on the synthesis of Li–Ti–O composites were furtherinvestigated by experiments.

2. Thermodynamic data and equations

2.1. Theoretical considerations

An electrochemical half-reaction could be expressed in the followinggeneral form:

aAþ nHþ þ ze ¼ bBþ cH2O:

Because ΔrGT = −zFET, the Nernst relationship results as shown inEq. (2):

ΔrGT ¼ ΔrGΘT þ RTln

acH2OabB

aaAanHþ

: ð1Þ

ET ¼ EΘT−RTzF

lnacH2Oa

bB

aaAanHþ

ð2Þ

where, z denotes the number of electron participating in the reaction;F refers to the Faraday constant, C·mol−1; R represents the gas constant,J·mol−1·K−1; T is the absolute temperature, K; aH2O; aA; aB; aHþ are theactivities of H2O, A, B and H+, respectively.

The standard Gibbs free energy of reaction is computed from theGibbs free energy of formation of species involved in a given reactionas follows:

ΔrGΘT ¼

XΔ fG

ΘT productsð Þ−

XΔ fG

ΘT reactantsð Þ ð3Þ

where ΔrGTΘ is the standard Gibbs free energy of reaction at T; ΔfGT

Θ

refers to the standard Gibbs free energy of formation of speciesparticipating in a reaction.

All the Li–Ti–O composites in the Li–Ti–H2O system are solidcompounds. The values of the average heat capacity of most speciesconsidered in the Li–Ti–H2O system are not available in the publishedliterature or databases. If we accepted that the reaction heat capacitychange (calculated by an equation analogous to Eq. (3)) were zero,the standard free energy of formation at an elevated temperatureT would then be given by (Townsend and Jr., 1973; Stephen, 1980):

Δ fGΘT ¼ Δ fG

Θ298:15− T−298:15ð ÞΔSΘ298:15 ð4Þ

where, ΔS298.15Θ is the absolute entropy value of species.This assumption accepts that the overall contributions of the heat

capacity of the species participating in the reaction cancel out. However,to apply Eq. (4), one needs to know the absolute entropy values of allelements participating in the formation reaction.

2.2. Thermodynamic data

Species considered in the Li–Ti–H2O system are listed in Table 1,together with their corresponding thermodynamic data. Since hydratesare expected to be the dominant solid compounds in aqueous solution(Kelsall and Robbins, 1990), only the hydrates of oxides and complex

oxides were taken into account. The standard Gibbs free energy offormation for these species at 298.15 K was taken from literature(Dean, 1991; Kelsall and Robbins, 1990; Qiu et al., 2010) and the valuesof the standard entropy of certain elements were estimated from theempirical relations as follows:

①. Eq. (5) proposed by Powell and Latimer (1951) was used tocalculate the entropy value of a monatomic ion.

SΘ298:15 ¼ 3.

2R lnMþ 37−270Z

.r2e

ð5Þ

Where, M is the atomic weight; Z is the absolute value of ioniccharge; re is the effective radius of an ion.

②. Eq. (6) given by Connick and Powell (1953)was used to estimatethe entropies of aqueous oxy-anions.

SΘ298:15 ¼ 43:5−46:5 Z−0:28nð Þ ð6Þ

Where, n is the number of O atoms excluding those included inthe hydroxyl groups.Since there is no reliable function for the entropy estimationof oxy-cations, the entropies of TiO2+ and TiO2

2+ were alsoestimated from Eq. (6).

③. Latimer (1952) proposed an additive method for estimating thestandard entropy of a solid compound, in which the elementvalue was assigned to the cation whereas the value of the anionwas variable and dependent on the cation charge. Latimer'sapproach was adopted in the present work using the correctvalence state for cations. Specifically, the standard entropy ofTiO3 was calculated from this method by assuming that theentropy contribution of O2− on Ti6+ is equal with that on Ti4+,i.e. 4.1858 J·mol−1 of O2−.

④. Guo and Zhao (2006) proposed a two-parameter model tocalculate the standard entropy of binary complex oxides. Itis supposed that the standard entropy of a complex oxideaMmOx·bNnOy (MmOx and NnOy are the simple oxides, and a, bare coefficients) consists of the entropy contributions of the

Fig. 2. Potential-pH diagram for Li–Ti–H2O system at 363.15 K. Activities of dissolvedspecies: 1 , 0.1 , 0.01 .

133L. Li et al. / Hydrometallurgy 142 (2014) 131–136

corresponding simple oxides plus the entropy of interaction, asshown in Eq. (7).

S ¼ aAþ bBþ abaþ b

A′ þ B′� �

þ D ð7Þ

In Eq. (7), S is the standard entropy of the complex oxideaMmOx·bNnOy. A and B are the contribution entropies ofcorresponding simple oxides, whereas A′andB′ are the entropiesof interaction. D is a constant equal to −8.932 J·mol−1. Eq. (7)was used to estimate the standard entropies of Li4Ti5O12 andLi4TiO4.

⑤. Qiu et al.(2010) accounted for the effect of water of hydrationfor TiO2(hc), Ti2O3(hc), Li4Ti5O12(hc), Li2TiO3(hc) and Li4TiO4(hc),where subscript (hc) refers to hydrated compound (Kelsall andRobbins, 1990), assuming hydration numbers 2, 4/3, 5, 1 and 0,respectively. Latimer (1952) reported that the entropies ofhydrates may be estimated by adding an extra 39.34 J for thecontribution of 1 mol of hydratedwater. The entropy differencesbetween solid compounds and their corresponding hydrateswere then calculated. It was found that the contribution of39.34 J per hydratedwaterwas applicable for saline compounds.However, the differences between metal oxides and theircorresponding hydrates were smaller with the contribution permole of water of hydration being closer to 30 J. This value wasalso accepted in the present work.

All the reactions considered in the Li–Ti–H2O system are provided inAppendix A, along with the standard Gibbs free energy of reaction at298.15 K (Qiu et al., 2010). The relationships between the potentialand pH as well as the standard Gibbs free energy of reaction at highertemperature are shown in Appendix B.

3. Results and discussion

Based on all reactions listed in Appendix B, the Eh–pH diagramsat temperatures 333.15, 363.15, 393.15 and 453.15 K were plottedwith the help of the software Origin 8.0 and are shown in Figs. 1–4.The activities of dissolved species were set at 0.01, 0.1 and 1 at alltemperatures.

Fig. 1. Potential-pH diagram for Li–Ti–H2O system at 333.15 K. Activities of dissolvedspecies: 1 , 0.1 , 0.01 .

As shown in Figs. 1–4, titanium occurs in diverse forms underdifferent conditions. Hexavalent titanium exists as TiO2

2+ at low pHand high potential, and is converted into TiO3·2H2O with an increasein pH. Tetravalent titanium ion occurs in the form of TiO2+, whichchanges to hydrous titanium dioxide as pH increases. Trivalent titaniumis in the form of Ti3+ at low pH and Eh, and transforms into hydratedTi2O3(hc) at high pH. Divalent titanium is not predominant (Kelsall andRobbins, 1990), but TiH2 can be produced at low pH and low potentialconditions. As pH is further increased in the presence of Li+, thepredominant regions of Li4Ti5O12(hc), Li2TiO3(hc) and Li4TiO4(hc) areseen in sequence and adjacent to each other. The diagrams demonstratethat Li4Ti5O12(hc) and Li2TiO3(hc) are thermodynamically stable inaqueous solutions. A wet chemistry technique may be employed tosynthesize these types of lithium–titanium composite oxides.

The predominance area diagrams at different temperaturesdemonstrate that temperature and ion activity have significant effects

Fig. 3. Potential-pH diagram for Li–Ti–H2O system at 393.15 K. Activities of dissolvedspecies: 1 , 0.1 , 0.01 .

Fig. 5. The XRD patterns of the precursors obtained by reactions of hydrate TiO2 with LiOHin aqueous solutions under different conditions. a—precursor, b—precursor, c—precursorand d—precursor are the intermediate product obtained under the conditions given inTable 2, respectively.

Fig. 4. Potential-pH diagram for Li–Ti–H2O system at 453.15 K. Activities of dissolved spe-cies: 1 , 0.1 , 0.01 .

Table 2Conditions for the preparation of Li–Ti–O composite precursors.

Sample Li:Ti(molar ratio)

Initial concentrationof LiOH (mol·L−1)

Temperature(°C)

FinalpH

a 2.20 3.0 40 13.4b 2.20 3.0 70 13.4c 2.20 3.0 95 13.4d 2.00 1.0 200 7.0

134 L. Li et al. / Hydrometallurgy 142 (2014) 131–136

upon the stability of species considered in Li–Ti–H2O system. The mainobservations are summarized as follows:

(1) The equilibrium between TiO2(hc) and Li4Ti5O12(hc) is hardlyaffected by temperature. However, the equilibrium line betweenTiO2+ and TiO2(hc) moves to lower pH with an increase intemperature, resulting in an expansion of the stability field ofTiO2(hc).

(2) For Li2TiO3(hc), the dominant area varies slightly withtemperature, because the equilibrium lines of Li4Ti5O12(hc)/Li2TiO3(hc) and Li2TiO3(hc)/Li4TiO4(hc) move in the same directionand the space between them changes tinily. The stability regionof Li2TiO3(hc) is between pH 12.5 to 13.6 at 333.15 K. It movesto pH ranging from 9.9 to 10.9 at 453.15 K and at 0.1 activity ofall dissolved species. This indicates that the synthesis of Li2TiO3

requires less alkaline conditions at elevated temperature.(3) With an increase in temperature, the stable field of Li4TiO4(hc)

increases markedly and moves to lower pH values zone.

Although Li4TiO4(hc) was found to be thermodynamically stable overa wide pH range, especially at higher temperatures, few reports on thesynthesis of Li4TiO4 from aqueous solution are found in the literature.Nevertheless, experiments were performed to prove the feasibility ofsynthesizing lithium–titanium composite oxides in aqueous solutions.

4. Experimental confirmation

Analytically pure ammonia solution (25–28 vol.%, Xilong ChemicalCo., Ltd.), chemically pure titanium tetrachloride (purity ≥ 98%,Sinopharm Chemical Reagent Co., Ltd.) and lithium hydroxidemonohydrate (purity ≥ 98.9%, Sichuan Tianqi Lithium Industries, Inc.)were used as raw materials. A 2 mol·L−1 TiCl4 aqueous solution wasprepared from chemically pure TiCl4 in an ice-water bath undermagnetic stirring. A dilute TiCl4 solution of 0.5 mol·L−1 was contactedwith 2 mol·L−1 ammonia. Amorphous hydrous TiO2 was obtainedafter filtering and water-washing thoroughly to remove Cl− (Li et al.,2013). Then, 1 or 3 mol·L−1 lithium hydroxide solution was mixedwith hydrous TiO2 at different Li:Ti molar ratios. The mixture washeated up to the target temperature ranging from 40 °C to 200 °C andmaintained for 6 h. The reactions between hydrous TiO2 and lithiumhydroxide at temperatures below 100 °C were performed in a waterbath (DF-101S, Gongyi Yuhua Instrument Co., Ltd.) with magneticstirring, and the reaction at 200 °C was conducted in an 1Cr18Ni9Ti

stainless steel autoclave (GS-0.25, Weihai Jingda Chemical MachineryCo., Ltd.) equipped with a mechanical stirrer. The resulting slurry wasfiltered, and the obtained precipitate was dried to the dehydratedform at 120 °C for 20 h. Finally, the dried precursors were heat-treated at 800 °C for 6 h to obtain the final Li–Ti–O composite products.

The powders produced under the conditions given in Table 2 werecharacterized by powder X-ray diffraction (XRD, Rigaku D/max-2500),using a beam of Cu Kα-radiation (λ = 1.5418 Å), with scanning anglefrom 10° to 70° and step size of 0.02° at room temperature(CuKα1-radiation λ = 1.544056 Å).

As shown in Fig. 5, the detected intermediate products obtained byreactions of hydrate TiO2 with LiOH in aqueous solutions under definedconditions are cubic Li2TiO3 phase. No characteristic peaks of TiO2

or other elements were detected, even the product obtained byhydrothermal reaction at 200 °C. The products after heat-treatingthe precursors obtained at 40 °C (Fig. 5a-precursor) and 70 °C(Fig. 5b-precursor) at Li:Ti molar ratio of 2.2:1 and LiOH concentrationof 3.0 mol·L−1, are mainly Li2TiO3 with a trace amount of Li4Ti5O12

(Fig. 6a and b). Almost pure Li2TiO3was obtained from the intermediateproducts at either 95 °C (Fig. 5c-precursor) or 200 °C (Fig. 5d-precursor) albeit at a slightly lower Li:Ti molar ratio of 2.0:1 andlower LiOH concentration of 1.0 mol·L−1, as shown in Fig. 6c and d.Sugita et al. (1990) reported that a single phase of lithium titaniumwas formed in a relatively limited range of Li/Ti mole ratio (R) of0.75–0.85. The material became mostly amorphous after largedehydration at 230 °C, and showed a Li4/3Ti5/3O4 phase by heat-treating at 400 °C. Once the R is in larger than 0.9, namely, the lithium

Fig. 6.TheXRDpatterns of the products obtained by heat-treating the precursors at 800 °Cfor 6 h. a, b, c and d are the product after heat-treating the corresponding precursor,respectively.

No. Reactions ΔrG298.15Θ

(kJ mol−1)

A.1 Ti3+ + 2H+ + 5e = TiH2 269.48A.2 Ti2O3(hc) + 10H+ + 10e = 2TiH2 + 3H2O 515.98A.3 TiO2+ + 2H+ + e = Ti3+ + H2O −9.53A.4 TiO2(hc) + 4H+ + e = Ti3+ + 2H2O −2.76A.5 2TiO2(hc) + 2H+ + 2e = Ti2O3(hc) + H2O 17.46A.6 2Ti3+ + 3H2O = Ti2O3(hc) + 6H+ 22.98A.7 TiO2+ + H2O = TiO2(hc) + 2H+ −6.77A.8 4Li+ + 5TiO2(hc) + 2H2O = Li4Ti5O12(hc) + 4H+ 70.26A.9 Li4Ti5O12(hc) + 6Li+ + 3H2O = 5Li2TiO3(hc) + 6H+ 434.70A.10 Li2TiO3(hc) + 2Li+ + H2O = Li4TiO4(hc) + 2H+ 158.10A.11 TiO3·2H2O + 2H+ + 2e = TiO2(hc) + 3H2O −359.72A.12 5TiO3·2H2O + 4Li+ + 10e + 6H+ = Li4Ti5O12(hc) + 13H2O −1727.35A.13 TiO3·2H2O + 2Li+ + 2e = Li2TiO3(hc) + 2H2O −258.73A.14 TiO3·2H2O + 4Li+ + 2e = Li4TiO4(hc) + H2O + 2H+ −100.65A.15 TiO3·2H2O + 2H+ = TiO2

2+ + 3H2O −5.62A.16 2Li4Ti5O12(hc) + 18H+ + 10e = 5Ti2O3(hc) + 8Li+ + 9H2O −53.16A.17 2Li2TiO3(hc) + 6H+ + 2e = Ti2O3(hc) + 4Li+ + 3H2O −184.52A.18 TiO2

2+ + 2e = TiO2(hc) −354.08A.19 TiO2

2+ + 2H+ + 2e = TiO2+ + H2O −347.30A.20 2H+ + 2e = H2 0

+

Reaction Temp/K ΔrGTΘ

(kJ mol−1)Equation

A.1 333.15 260.08 E = −0.5391 + 0.0132lg(Ti3+) − 0.0264pH363.15 252.02 E = −0.5224 + 0.0144lg(Ti3+) − 0.0288pH393.15 243.97 E = −0.5057 + 0. 0156lg(Ti3+) − 0.0312pH

3+

135L. Li et al. / Hydrometallurgy 142 (2014) 131–136

is excessive, Li2TiO3 was formed. This demonstrated that the obtainedLi4Ti5O12 precursor during wet process is amorphous, which isconsistent with the results in the publication (Li et al., 2013), inwhich the intermediate product was obtained at 75 °C and at lowerpH in Li/Ti mole ratio of 2.0. The undetected trace Li4Ti5O12 precursorcrystallized during heat-treating and resulted in the appearance ofLi4Ti5O12 in the products of a and b. These experimental results are ingood agreement with the presented predominance area diagrams, asshown in Figs. 1–4.

453.15 227.86 E = −0.4723 + 0.0180lg(Ti ) − 0.0360pHA.2 333.15 510.71 E = −0.5293 − 0.0661pH

363.15 506.20 E = −0.5246 − 0.0721pH393.15 501.68 E = −0.5200 − 0.0780pH453.15 492.66 E = −0.5106 − 0.0900pH

A.3 333.15 −8.96 E = 0.0929 + 0.0661lg{(Ti2+) / (Ti3+)} − 0.1322pH

363.15 −8.48 E = 0.0879 + 0.0721lg{(Ti2+] / (Ti3+)} − 0.1441pH

393.15 −8.00 E = 0.0829 + 0.0780lg{(Ti2+] / (Ti3+)} − 0.1556pH

453.15 −7.03 E = 0.0728 + 0.0900lg{(Ti2+] / (Ti3+)} − 0.1798pH

A.4 333.15 4.54 E = −0.0471 − 0.0661lg(Ti3+) − 0.2644pH363.15 10.81 E = −0.1120 − 0.0721lg(Ti3+) − 0.2883pH393.15 17.08 E = −0.1770 − 0.0780lg(Ti3+) − 0.3125pH453.15 29.60 E = −0.3068 − 0.0900lg(Ti3+) − 0.3598pH

A.5 333.15 18.55 E = −0.0961 − 0.0661pH363.15 19.48 E = −0.1010 − 0.0721pH393.15 20.42 E = −0.1058 − 0.0780pH453.15 22.28 E = −0.1155 − 0.0900pH

A.6 333.15 9.45 pH = 0.2470 − 0.3333lg(Ti3+)363.15 −2.14 pH = −0.0512 − 0.3333lg(Ti3+)393.15 −13.73 pH = −0.3040 − 0.3333lg(Ti3+)453.15 −36.92 pH = −0.7093 − 0.3333lg(Ti3+)

A.7 333.15 −13.51 pH = −1.0592 − 0.5lg(TiO2+)

5. Conclusions

Eh–pH diagrams for the Li–Ti–H2O system at temperatures 333.15,363.15, 393.15 and 453.15 K, as well as 0.01, 0.1 and 1 ion activities ofaqueous species were constructed. The predominance area diagramsat different temperatures show that temperature and activity of aque-ous species have strong effects on the stability regions. With increasingtemperature, the stability field of Li4Ti5O12(hc) and Li2TiO3(hc) decrease,while that of Li4TiO4(hc) increases. The existence of the predominantareas of lithium–titanium composite oxides indicates that it is possibleto synthesize Li4Ti5O12 and Li2TiO3 by wet routes. Pure Li2TiO3 and Li–Ti–O composite compounds with Li2TiO3 as the main phase were ob-tained from synthesis experiments conducted under the guidance ofthe potential-pH diagrams that were constructed. The experimentsproved that the preparation of Li–Ti–O composite oxides in aqueous so-lution, such as Li2TiO3 and Li4Ti5O12, is feasible and the phase structureof the product can be controlled by adjusting the process parameters.The experimental results are consistent with the information given inthe potential-pH diagrams.

363.15 −19.29 pH = −1.3874 − 0.5lg(TiO2+)393.15 −25.08 pH = −1.6655 − 0.5lg(TiO2+)453.15 −36.64 pH = −2.1112 − 0.5lg(TiO2+)

A.8 333.15 78.52 pH = 3.0777 − lg(Li+)363.15 85.62 pH = 3.0785 − lg(Li+)393.15 92.72 pH = 3.0793 − lg(Li+)453.15 106.91 pH = 3.0805 − lg(Li+)

A.9 333.15 441.27 pH = 11.5296 − lg(Li+)363.15 446.91 pH = 10.7122 − lg(Li+)393.15 452.54 pH = 10.0196 − lg(Li+)

Acknowledgments

The authors wish to thank the financial support from the NationalNatural Science Foundation of China (Program Nos. 51074194 and51010105015) and the kind support from the Government of GuangxiZhuang Autonomous Region (Glorious Laurel Scholar Program No.2011A025).

Appendix A. Reactions and their standard Gibbs free energy at298.15 K (Qiu et al., 2010)

A.21 4H + O2 + 4e = 2H2O −474.28

Appendix B. Relationships between potential and pH and thestandard Gibbs free energy of reactions at each temperature

(continued)

Reaction Temp/K ΔrGTΘ

(kJ mol−1)Equation

453.15 463.82 pH = 8.9094 − lg(Li+)A.10 333.15 161.04 pH = 12.6228 − lg(Li+)

363.15 163.58 pH = 11.7630 − lg(Li+)393.15 166.12 pH = 11.0343 − lg(Li+)453.15 171.22 pH = 9.8666 − lg(Li+)

A.11 333.15 −366.94 E = 1.9015 − 0.0661pH363.15 −373.12 E = 1.9336 − 0.0721pH393.15 −379.30 E = 1.9656 − 0.0780pH453.15 −391.68 E = 2.0298 − 0.0900pH

A.12 333.15 −1756.16 E = 1.8201 − 0.0396pH + 0.0264lg(Li+)363.15 −1779.80 E = 1.8446 − 0.0432pH + 0.0289lg(Li+)393.15 −1803.54 E = 1.8692 − 0.0468pH + 0.0312lg(Li+)453.15 −1851.03 E = 1.9184 − 0.0540pH + 0.0360lg(Li+)

A.13 333.15 −262.98 E = 1.3628 + 0.0661lg(Li+)363.15 −266.62 E = 1.3816 + 0.0721lg(Li+)393.15 −270.26 E = 1.4005 + 0.0780lg(Li+)453.15 −277.54 E = 1.4382 + 0.0900lg(Li+)

A.14 333.15 −101.94 E = 0.5282 + 0.0661pH + 0.1322lg(Li+)363.15 −103.03 E = 0.5339 + 0.0721pH + 0.1441lg(Li+)393.15 −104.12 E = 0.5396 + 0.0780pH + 0.1560lg(Li+)453.15 −106.32 E = 0.5510 + 0.0900pH + 0.1798lg(Li+)

A.15 333.15 −5.58 pH = 0.4375 − 0.5lg(TiO22+)

363.15 −5.52 pH = 0.3972 − 0.5lg(TiO22+)

393.15 −5.46 pH = 0.3631 − 0.5lg(TiO22+)

453.15 −5.35 pH = 0.3084 − 0.5lg(TiO22+)

A.16 333.15 −64.30 E = 0.0666 − 0.1190pH − 0.0528lg(Li+)363.15 −73.82 E = 0.076 − 0.1297pH − 0.0576lg(Li+)393.15 −83.34 E = 0.0864 − 0.1403pH − 0.0624lg(Li+)453.15 −102.38 E = 0.1061 − 0.1618pH − 0.0718lg(Li+)

A.17 333.15 −189.37 E = 0.9814 − 0.1982pH − 0.1322lg(Li+)363.15 −193.52 E = 1.0029 − 0.2163pH − 0.1441lg(Li+)393.15 −197.68 E = 1.0244 − 0.2340pH − 0.1560lg(Li+)453.15 −206.00 E = 1.0675 − 0.2698pH − 0.1798lg(Li+)

A.18 333.15 −361.35 E = 1.8726 + 0.0330 g(TiO22+)

363.15 −367.60 E = 1.9050 + 0.0360lg(TiO22+)

393.15 −373.84 E = 1.9373 + 0.0390lg(TiO22+)

453.15 −386.33 E = 2.0020 + 0.0450lg(TiO22+)

A.19 333.15 −347.84 E = 1.8026 − 0.0661pH + 0.0330lg{(TiO2

2+) / (TiO2+)}363.15 −348.30 E = 1.8050 − 0.0721pH + 0.0360lg

{(TiO22+) / (TiO2+)}

393.15 −348.76 E = 1.8073 − 0.0780pH + 0.0390lg{(TiO2

2+) / (TiO2+)}453.15 −349.69 E = 1.8122 − 0.0900pH + 0.0450lg

{(TiO22+) / (TiO2+)}

A.20 333.15 −4.57 E = 0.0237 − 0.0660pH363.15 −8.49 E = 0.0440 − 0.0721pH393.15 −12.41 E = 0.0643 − 0.0780pH453.15 −20.24 E = 0.1049 − 0.0900pH

A.21 333.15 −472.00 E = 1.2230 − 0.0661pH363.15 −470.04 E = 1.2179 − 0.0721pH393.15 −468.08 E = 1.2128 − 0.0780pH453.15 −464.17 E = 1.2027 − 0.0900pH

Appendix B. (continued)

136 L. Li et al. / Hydrometallurgy 142 (2014) 131–136

References

Allen, J.L., Jow, T.R., et al., 2006. Low temperature performance of nanophase Li4Ti5O12.J. Power Sources 159 (2), 1340–1345.

Chen, B.Z., Ma, L.W., et al., 2009. Research progress on preparationmethods for precursorsfor lithium ion-sieve. Inorg. Chem. Ind. 41 (7), 1–4 (in Chinese with English abstract).

Chen, P.P., Li, Y.J., et al., 2008. In-situ synthesis of spinel Li4Ti5O12 by hydrolysis of TiCl4aqueous in an acidic aqueous solution. Metal Mater. Metall. Eng. 36 (6), 7–13(in Chinese with English abstract).

Casadio, S., van der Laan, J.G., et al., 2004. Tritium release kinetics from Li2TiO3 pebbles asprepared by soft-wet-chemistry. J. Nucl. Mater. 329 (333), 1252–1255.

Connick, R.E., Powell, R.E., 1953. The entropy of aqueous oxy-anions. J. Chem. Phys. 21(12), 2006–2007.

Davis, J.W., Haasz, A.A., 1996. Thermal diffusivity/conductivity of AECL Li2TiO3 ceramic.J. Nucl. Mater. 232 (1), 65–68.

Dean, J.A., 1991. (translated) In: Shang, J.F., Cao, S.J., et al. (Eds.), Lange's Handbook ofChemistry, thirteenth ed. Science Press, Beijing, pp. 1468–1533.

Gao, J., Ying, J.R., et al., 2007. High-density spherical Li4Ti5O12/C anode material withgood rate capability for lithium ion batteries. J. Power Sources 166 (1), 255–259(in Chinese with English abstract).

Guo, P.M., Zhao, P., 2006. Estimation of standard entropy of complex oxide usingtwo-parameter model. J. Iron Steel Res. 18 (9), 17–20 (in Chinese with Englishabstract).

Hoshino, T., Oikawa, F., 2011. Trial fabrication tests of advanced tritium breeder pebblesusing sol–gel method. Fusion Eng. Des. 86 (9–11), 2172–2175.

Jiang, M.Q., 1977. The quantitative relation between structure and property in organicmolecule—the homologous linear rule. Chemistry 4, 17–27 (in Chinese with Englishabstract).

Jiang, J.H., 2012. Study of an inorganic ion exchanger Li2TiO3. Adv. Mater. Res. 457–458,34–37.

Kelsall, G.H., Robbins, D.J., 1990. Thermodynamics of Ti–H2O–F(–Fe) systems at 298 K.J. Electroanal. Chem. 283, 135–157.

Laumann, A., Jensen, K.M.O., et al., 2011. In-situ synchrotron X-ray diffraction study of theformation of cubic Li2TiO3 under hydrothermal conditions. Eur. J. Inorg. Chem. 2011(14), 2221–2226.

Laumann, A., Fehr, K.T., et al., 2010. Metastable formation of low temperature cubicLi2TiO3 under hydrothermal conditions—its stability and structural properties. SolidState Ionics 181 (33–34), 1525–1529.

Latimer, W.M., 1952. The Oxidation States of the Elements and Their Potentials inAqueous Solutions, 2nd edition. Prentice-Hall, New York 359–365.

Liu, L.F., Zhang, Y.S., et al., 2010. Research progress of lithium extraction from seawater and brines by ion-sieve adorption method. J. Salt Chem. Ind. 39 (3), 40–44(in Chinese with English abstract).

Li, Y.J., Qiu, W.S., et al., 2009. Study on synthesis of Li4Ti5O12 by force hydrolysis of TiCl4.Min. Metall. Eng. 29 (3), 78–81 (in Chinese with English abstract).

Li, Y.J., Xu, C., et al., 2012. Synthesis of Li2TiO3 ceramic breeder powders by in-situhydrolysis and its characterization. Mater. Lett. 89 (15), 25–27 (in Chinese withEnglish abstract).

Li, Y.J., Wang, X.Y., Ming, X.Q., Xu, C., Kong, L., Li, L., Ye, W.Q., Ren, M.M., 2013.Preparation and electrochemical properties of Li4Ti5O12 synthesized by wetchemical method. J. Cent. South Univ. Sci. Technol. 44 (10) (in Chinese withEnglish abstract).

Osseo-Asare, K., Brown, T.H., 1979. A numerical method for computing hydrometallurgicalactivity–activity diagrams. Hydrometallurgy 4 (3), 217–232.

Ohzuku, T., Ueda, A., et al., 1995. Zero-strain insertion material of Li[Li1/3Ti5/3]O4 forrechargeable lithium cells. J. Electrochem. Soc. 142 (5), 1431–1435.

Powell, R.E., Latimer, W.M., 1951. The entropy of aqueous solutes. J. Chem. Phys. 19 (9),1139.

Qiu, W.S., Li, Y.J., et al., 2010. Eh–pH diagram of Li–Ti–H2O system at 298.15 K andsynthesis of Li4Ti5O12. Chin. J. Nonferrous Met. 20 (11), 2260–2268 (in Chinesewith English abstract).

Roux, N., Avon, J., et al., 1996. Low-temperature tritium releasing ceramics as potentialmaterials for the ITER breeding blanket. J. Nucl. Mater. 233 (237), 1431–1435.

Sugita, M., Tsuji, M., et al., 1990. Synthetic inorganic ion-exchange materials. LVIII.Hydrothermal synthesis of a new layered lithium titanate and its alkali ion exchange.Bull. Chem. Soc. Jpn. 63, 1978–1984.

Stephen, A.N., 1980. Thermodynamics of aqueous systemswith industrial applications. In:Herbert, E. Barner, Kust, Roger N. (Eds.), Application of Thermodynamics inHydrometallurgy. A State-of-the-Art Review. American Chemical Society,Washington, D.C., pp. 625–641.

Shen, H., Zhong, H., et al., 2012. Research of lithium ion sieve adsorbent. SiChuan Chem.Ind. 15 (3), 19–22 (in Chinese with English abstract).

Tomiha, M., Masaki, N., et al., 2002. Hydrothermal synthesis of alkali titanates from nanosize titania powder. J. Mater. Sci. 37 (11), 2341–2344.

Townsend, H.E., Jr., 1973. The importance of heat capacity in the construction ofpotential-pH diagrams at elevated temperature. Corros. Sci. 13 (4), 311–314.

Townsend, H.E., Jr., 1970. Potential-pH diagrams at elevated temperature for the systemFe–H2O. Corros. Sci. 10 (5), 343–358.

Zhang, L.F., Chen, B.Z., et al., 2010. Synthesis and adsorption property of H2TiO3 typeadsorbent. Chin. J. Nonferrous Met. 20 (9), 1849–1854 (in Chinese with Englishabstract).