Solvothermal synthesis of pyrite FeS₂ nanocubes and their superior high rate lithium storage properties

Abstract

Iron pyrite nanocubes with particle sizes of around 80–120 nm have been synthesized via a facile solvothermal method. Time-dependent characterization found that the iron(II) chloride and sulfur precursors first react to form sheet-like amorphous Fe_1−xS, which transforms into pyrite FeS₂ nanocubes upon further heating. As an anode material for lithium ion batteries, the as-synthesized pyrite FeS₂ nanocubes were found to deliver a reversible discharge capacity of 540 mA h g⁻¹ after cycling for 150 cycles at a current density of 1 A g⁻¹. Even at higher current density of 5 A g⁻¹, the pyrite FeS₂ electrode still managed to give a stable discharge capacity of about 220 mA h g⁻¹. The enhanced lithium storage properties are attributed to its higher specific surface area that can provide more lithium ion (Li⁺) reaction sites, leading to less polarization and better cycling performance.

---

Introduction

Electricity is one of the prevailing forms of energy used in the world today. It plays a crucial part in many aspects of our lives, ranging from lighting to communication and even entertainment. With the increase in global population and development of modern technologies, the applications for electrical energy have become more diversified and the demand for electricity is rising. Presently, most of the electricity in use is generated through the burning of non-renewable fossil fuels in stand-alone power plants. However, the rapid depletion of fossil fuels and the environmental pollution caused by the combustion of these energy resources encouraged a shift of focus towards the use of renewable and cleaner sources like solar and wind power. As these renewable resources are intermittent in nature, efficient storage systems such as lithium ion batteries (LIBs) are greatly needed to store and deliver these energies when and where required.¹ In fact, since the commercialization of LIB in 1991, it has found its way into a vast variety of applications. LIB is currently one of the most widely used forms of rechargeable energy storage systems for portable electronics and is also viewed as a potential candidate in emerging applications such as electric vehicles, complementary energy storage for renewable resources and grid load levelling.^2,3 The main reasons for such widespread usage of LIBs include their high volumetric and gravimetric energy densities, excellent cycling performances and design versatility,⁴ allowing the fabrication of batteries of different shape and sizes, depending on the application requirements.

Conventional commercial LIB uses graphite as anode material because of its good cycling stability, Earth-abundance and low cost.^5,6 However, graphite also suffers from several problems such as its low theoretical capacity of 372 mA h g⁻¹, poor rate capability and low lithiation voltage that can result in the formation of lithium dendrites and short-circuiting after repeated cycling.^7,8 Hence, a lot of research efforts have been put into the search for alternative electrode materials that are cheap, safe, environmentally friendly and, yet, can provide higher energy density and power to meet the requirements for emerging applications.^3,9 Recently, materials such as metal oxides and sulfides that can react with lithium via a conversion reaction mechanism have gained tremendous research interest because of their higher theoretical lithium storage capacities than graphite, which stores lithium via an intercalation reaction.^10–15 Of these materials, iron sulfides, in particular iron pyrite (FeS₂), are appealing electrode choices as they have been extensively researched in the late 20th century in the lithium battery field. Iron sulfides are not only used in commercial primary lithium batteries and high temperature thermal batteries, they have also been studied as potential electrode choice in rechargeable energy storage systems at moderate temperatures (75–135 °C).¹⁶ Pyrite FeS₂ possesses various advantages such as a high theoretical capacity of 890 mA h g⁻¹, low environmental impact and, most importantly, affordable cost.^16–18 However, the poor cycling performance and capacity retention of FeS₂ at room temperature limited its use in room temperature LIB.^19,20 Nanostructuring is viewed as an effective approach to alleviate these problems. Nanomaterials have higher specific surface area than their bulk counterparts and, hence, can provide more contact area between the electrode material and electrolyte, increasing the number of lithium ion (Li⁺) reaction sites. This, together with the shorter Li⁺ diffusion length in nanomaterials, may help to enhance the lithium storage kinetics and thus improve the cycling properties and rate capability of the electrode.^3,21 Previous works done by other research groups have demonstrated the improvement in cycling performance of some LIB anode materials achieved through nanostructuring.^22,23 Besides the size of the electrode materials, its morphology is also widely believed to have great influence on its electrochemical properties. For example, Xu et al.¹³ reported that carbon-coated iron monosulfide (C@FeS) nanosheets display better cycling stability than C@FeS nanoparticles and C@FeS nanoplates. Rui et al.¹¹ also reported that hierarchical urchin-like Co₃O₄ hollow spheres exhibit higher lithium storage capacity and better rate capability than Co₃O₄ hexapods. Hence, in the past years, much effort have been put into the synthesis of iron pyrite nanomaterials with different morphologies to improve the cycling performance and capacity retention of FeS₂ at room temperature.

Iron pyrite nanoparticles were first synthesized more than ten years ago.²⁴ Since then, a few different approaches such as microwave irradiation,^25,26 hot injection method,^27–33 solid-state synthesis,^18,34 hydrothermal^35–37 and solvothermal^38,39 reactions have been used for the synthesis of nanostructured pyrite FeS₂. In this work, pyrite FeS₂ nanocubes are synthesized using a simple one-step solvothermal method without the requirement of specialized single-source precursors and inert environment, and the need for post-annealing process. The synthesized nanoparticles are characterized for its phase, size and morphology. Investigation into the formation mechanism of the nanoparticles in the solvothermal reaction has also been carried out, and studies on the lithium storage properties and cycling performance of the synthesized pyrite FeS₂ nanoparticles are also performed.

Experimental

Preparation of FeS₂ nanocubes

The pyrite FeS₂ nanocubes were synthesized via a facile solvothermal process. In a typical reaction, 15 mL of oleylamine was added into a 50 mL teflon-lined stainless steel autoclave containing 0.5 mmol of iron(II) chloride (FeCl₂, anhydrous, 99.5%, Alfa Aesar), 0.5 mmol of 1,2-dodecanediol (90%, Aldrich) and 3 mmol of sulfur powder (S, -325 mesh, 99.5%). The reaction mixture was sonicated for an hour to ensure homogenous mixing before the autoclave was sealed and maintained at 180 °C in a convection oven for 18 h. The autoclave was then cooled to room temperature naturally and the precipitate collected via centrifugation. The collected powder was washed several times with hexane to remove impurities and then dried in air at 70 °C overnight before collection for further characterization. To study the formation mechanism of the pyrite FeS₂ nanocubes, several pots of reactions with the same amount of reactants and fill ratio were subjected to the same solvothermal treatment at 180 °C for 3 h, 12 h, 18 h and 24 h.

Materials characterization

The phase structure of the synthesized powder was characterized by powder X-ray diffraction (XRD) (Bruker AXS, D8 Advance) using a Cu Kα radiation (λ = 0.15418 nm) and scanned in the 2θ range of 10–70° at a step width of 0.05°. Raman spectroscopy was performed using a confocal Raman microscopy system (WITec alpha 300 SR) with a 633 nm diode laser to further study the phase and phase purity of the synthesized powder. The size and morphology of the as-obtained nanoparticles were examined with a field emission scanning electron microscope (FESEM) (JEOL JSM-7600F). For both the Raman and FESEM characterization, the samples were prepared by dispersing the dried as-synthesized powder in hexane via sonication and drop-cast a few drops of the dispersion onto a small piece (≈5 mm × 5 mm) of silicon wafer. High resolution transmission electron microscopy (HRTEM) (JEOL JEM-2010) was performed at an accelerating voltage of 200 kV to study its size, morphology and crystallographic structure. The TEM sample was prepared via a similar method as the Raman and FESEM samples but the dispersion was drop-casted onto a carbon-laced copper grid instead of silicon wafer. Nitrogen adsorption/desorption isotherm was measured on a Micromeritics ASAP 2020 at −196 °C. The sample was degassed at 120 °C for 4 h under vacuum prior to measurement and the specific surface area was calculated using the Brunauer–Emmet–Teller (BET) method.

Electrochemical measurements

The working electrode was fabricated by mixing the as-obtained powder, carbon nanotubes (CNT, P3-SWNT, carbon solutions) and poly(vinylidene fluoride) (PVDF) at a weight ratio of 8 : 1 : 1 and dispersing them in N-methylpyrrolidone (NMP, 99+%, Alfa Aesar) to form a slurry. The resultant slurry was then pasted onto a circular copper foil (diameter = 14 mm) and dried in an oven under vacuum at 50 °C overnight. After drying, CR2032 coin cells were assembled using the prepared slurry on copper foil as the working electrode, lithium foil as the counter and reference electrode and Celgard 2400 membrane as the separator. The electrolyte used in these battery cells was made of 1 M lithium fluorophosphate (LiPF₆) in ethylene carbonate (EC)–dimethyl carbonate (DMC) (1/1, w/w). The whole battery-assembling process was carried out in an argon-filled glovebox (MBraun, Germany) where both the moisture and oxygen content were less than 1 ppm. The assembled cells were cycled galvanostatically in the voltage range of 0.001–3 V at room temperature under a current density of 1 A g⁻¹ using a NEWARE multi-channel battery test system. The cyclic voltammetry (CV) measurements were conducted using a Solartron 1470E analytical equipment in the voltage window of 0.001–3 V at room temperature under a constant scanning rate of 0.2 mV s⁻¹. As a comparative study, natural-occurring iron pyrite mineral was purchased. The purchased mineral was grinded using an agate mortar into fine powder and its electrochemical properties was tested using the same procedures as the synthesized pyrite powder.

Results and discussion

Phase and morphology

X-ray diffraction (Fig. 1) shows that the powder obtained after solvothermal reaction at 180 °C for 18 h has a cubic pyrite FeS₂ (JCPDS 042-1340) phase with a lattice parameter of 5.4179 Å. The sharp and well-defined diffraction peaks in Fig. 1 suggests that the as-synthesized powder has good crystallinity even without annealing. No additional peak was observed in the XRD pattern, confirming the good purity of the sample. However, as iron and sulfur are able to combine in different stoichiometric ratio to form various stable iron sulfide compounds like FeS, Fe₃S₄, Fe₇S₈, Fe_1−xS and FeS₂, and these phases have quite a number of overlapping diffraction peaks, it is very challenging to distinguish the diffraction peaks of one phase from another, especially when the phase is present in minute amount with very few and low intensity diffraction peaks in the XRD pattern. Hence, Raman spectroscopy, with its lower impurity detection limit compared to XRD and its ability to differentiate between the different phases of iron sulfide like FeS, Fe₃S₄ and FeS₂, was used to further verify the phase purity of the synthesized sample.⁴⁰ It can be seen from Fig. 2 that the synthesized pyrite FeS₂ powder after solvothermal reaction at 180 °C for 18 h only has two obvious Raman peaks present in the frequency range of 300–400 cm⁻¹. This distinctive two-peak spectrum is indicative of the presence of S–S bond in the sample.^40,41 The position of the Raman peaks in the obtained spectrum match well with those of pyrite iron sulfide that were reported in the literature (339 cm⁻¹ and 378 cm⁻¹),^28,42,43 with the Raman peak near 339 cm⁻¹ caused by the displacements of the S atoms perpendicular to the S–S bond axis (E_g mode) and that near 378 cm⁻¹ resulting from the in-phase stretching vibrations of the S₂ dumbbells (A_g) in the pyrite FeS₂ crystal. From the Raman spectrum in Fig. 2, it can be seen that there is also a small peak at around 424 cm⁻¹ which can be matched to that caused by the third Raman active mode (T_g mode) in a pyrite crystal.^37,44 The absence of peaks at around 210 cm⁻¹ and 280 cm⁻¹ indicates the nonexistence of FeS in the synthesized sample, further confirming its phase purity.⁴⁵

Fig. 1 XRD pattern of the synthesized pyrite FeS₂ after solvothermal reaction at 180 °C for 18 h.

Fig. 2 Raman spectrum of the synthesized pyrite FeS₂ powder after solvothermal reaction at 180 °C for 18 h.

The pyrite FeS₂ particles obtained after solvothermal reaction at 180 °C for 18 h have a homogenous cubic morphology with edge lengths in the range of 80–120 nm (Fig. 3(A)–(C)). Well-aligned lattice fringes can be observed on the face of the pyrite FeS₂ nanocube, as shown in Fig. 3(D), demonstrating the good crystallinity of the sample. A spacing of 2.70 Å was measured from the HRTEM image, which is consistent with the lattice spacing in the (2 0 0) plane of pyrite FeS₂, further confirming the cubic iron pyrite phase of the synthesized particles. Nitrogen adsorption/desorption measurement on the as-synthesized nanocubes determined a Brunauer–Emmett–Teller (BET) surface area of 10.0 m²g⁻¹ (ESI, Fig. S1†).

Fig. 3 (A) FESEM, (B and C) TEM, and (D) HRTEM images of the as-synthesized pyrite FeS₂ after solvothermal reaction at 180 °C for 18 h.

Formation mechanism of as-prepared FeS₂ nanocubes

To understand the formation mechanism of the pyrite FeS₂ nanocubes, four different pots of reaction were carried out with all parameters kept constant except for the reaction durations (3 h, 12 h, 18 h and 24 h). Fig. 4 shows the XRD patterns of the powders synthesized for various durations, illustrating the change in phase as the reaction proceeds.

Fig. 4 XRD patterns of the powders synthesized for various durations.

It can be seen from Fig. 4 that the powder synthesized for 3 hours has an amorphous nature, with only a very broad hump at around 2θ ≈ 30° which coincides with the hump caused by the XRD holder and no obvious diffraction peaks. Due to the presence of only one broad hump in the XRD pattern for the powder synthesized for 3 h, no conclusive phase matching can be done for this sample. Obvious and sharp diffraction peaks that can be matched to cubic iron pyrite (FeS₂) phase (JCPDS 042-1340) appeared after 12 hours of reaction. Similar XRD patterns were obtained for the samples synthesized for 18 h and 24 h with no observable impurity peaks for all three samples, indicating that all of them are pure iron pyrite phase. These three samples also gave similar Raman spectrum (Fig. 5) with peak positions (340 cm⁻¹ and 379 cm⁻¹) matching well to those reported in the literature for iron pyrite (339 cm⁻¹ and 378 cm⁻¹).^28,42,43 No other obvious Raman peaks were observed for all three spectra, which are consistent with the phase purity indicated by their XRD patterns. A very small peak was also observed for the powder synthesized for 18 and 24 h, matching well with one of the T_g mode peak. The powder synthesized for 3 h displays a different Raman spectrum (Fig. 5) from the other three samples, with a Raman peak at 295 cm⁻¹ and another at 359 cm⁻¹, matching well to the Fe_1−xS phase (292 cm⁻¹ and 354 cm⁻¹).²⁸ From the XRD and Raman spectroscopy results, it can be determined that the anhydrous FeCl₂ and S precursors react for 3 hours to form amorphous Fe_1−xS and this Fe_1−xS phase slowly reacts with the remaining S in the reaction mixture to crystallize into pure pyrite FeS₂ phase. The Raman spectra in Fig. 5 do not have any other detectable and observable peaks besides those from the amorphous Fe_1−xS and pyrite FeS₂. Hence, other iron sulfide phases like FeS (Raman peaks at 210 cm⁻¹ and 280 cm⁻¹) are most likely not formed nor present during this reaction.

Fig. 5 Raman spectra of the powders synthesized for various durations.

The FESEM image (Fig. 6(A)) shows that the amorphous Fe_1−xS formed after 3 h of synthesis has a sheet-like morphology that are highly aggregated. After heating for another 9 hours (12 h sample) in an autoclave at 180 °C, the sheet-like amorphous Fe_1−xS transforms into pyrite FeS₂. The FESEM image for this sample (Fig. 6(B)) shows that nanocubes start to appear in-between the sheets. When the reaction duration is increased to 18 h, well-defined pyrite FeS₂ nanocubes with sizes in the range of 80–120 nm are formed (Fig. 3(A) and 6(C)). Subjecting the mixture to further reaction duration of 24 h yields larger cubic particles with a wider size distribution ranging from about 100–250 nm, as shown in Fig. 6(D).

Fig. 6 FESEM images of the powders synthesized for (A) 3 h, (B) 12 h, (C) 18 h and (D) 24 h.

Electrochemical properties

The Li-ion storage properties of the pyrite FeS₂ nanocubes synthesized after solvothermal reaction at 180 °C for 18 h was examined by conducting a series of electrochemical measurements based on the half-cell configuration. This sample was chosen for the electrochemical test because of its nano-size and smaller particle size distribution, providing better understanding of the effect of nanostructuring where the dimensions of the particles is less than 100 nm. The assembled cell was cycled galvanostatically between 0.001 and 3 V (vs. Li/Li⁺) at a current density of 1 A g⁻¹ to investigate its cycling performance. A plot of the discharge capacity of the working electrode against the cycle number for the cell is shown in Fig. 7, where the synthesized pyrite FeS₂ exhibits a high initial discharge capacity of 980 mA h g⁻¹, exceeding the theoretical capacity of FeS₂ (890 mA h g⁻¹). This phenomenon has been observed in quite a few material systems from insertion-typed anodes such as carbon, to alloying-typed anodes such as tin antimony, and conversion-typed anodes such as transition metal oxides, especially during the first discharge cycle.^46,47 The reason for this excess lithium storage capacity is attributed to the formation of a veil-like layer, more widely known as the solid/electrolyte interphase (SEI) layer, on the electrode surface.⁴⁷ After the first cycle, the discharge capacity slowly decreases from 980 mA h g⁻¹ (1st cycle) to 649 mA h g⁻¹ (2nd cycle), reaching a capacity of about 620 mA h g⁻¹ at its 3rd cycle, which is about 60% of the initial discharge capacity. From the 3rd to 70th cycle, the discharge capacity of the as-synthesized pyrite FeS₂ nanoparticles remains fairly constant, with a difference of not more than 5% between the two capacity values. After the first 100 charge–discharge cycles, the working electrode has a discharge capacity of 534 mA h g⁻¹, retaining 54% of its initial discharge capacity. At the end of the 150th charge–discharge cycle, the working electrode still manage to maintain a discharge capacity of 540 mA h g⁻¹, retaining about 55% of its initial discharge capacity.

Fig. 7 Cycling performance of the synthesized and commercial pyrite FeS₂ electrode at a current density of 1 A g⁻¹ between 0.001 and 3 V (vs. Li/Li⁺).

For comparison with the cycling performance of the as-synthesized pyrite nanocubes, the electrochemical properties of a natural-occurring iron pyrite mineral sample (FeS₂, naturally occurring mineral, grains, ≈1.5–4.8 mm, Alfa Aesar) were tested and the result is shown in Fig. 7. The as-purchased iron pyrite mineral was grinded using a mortar and pestle into smaller particles so that it can be used to prepare battery half-cells for electrochemical testing. FESEM image of the grinded iron pyrite mineral (ESI, Fig. S2†) shows that the sample consists mainly of micron-sized granular particles with some particles in the nanometre range. The phase purity of the commercial mineral sample was confirmed using XRD (ESI, Fig. S3†) which shows only iron pyrite diffraction peaks and no observable impurity peak. When subjected to galvanostatic charge–discharge cycling, the commercial pyrite FeS₂ exhibits an initial discharge capacity of 639 mA h g⁻¹ (Fig. 7) which is much lower than that of the as-synthesized iron pyrite nanocubes. After the first cycle, the discharge capacity of the commercial pyrite FeS₂ electrode dropped to 459 mA h g⁻¹ (2nd cycle) and this capacity slowly decreases to 181 mA h g⁻¹ (100th cycle), retaining only about 28% of its initial discharge capacity. The improved cycling performance and capacity retention of the pyrite nanocubes compared to the commercial pyrite sample are attributed to its smaller particle size and particle size distribution. These result in a larger BET surface area and, hence, an improvement in the lithium storage properties.

The galvanostatic charge (Li ion extraction) and discharge (Li ion insertion) voltage profiles of the as-synthesized and commercial iron pyrite electrode when cycled at a current density of 1 Ag⁻¹ between 0.001 and 3 V (vs. Li/Li⁺) is depicted in Fig. 8. During the first cycle, both samples display a discharge profile with one plateau that is associated with formation of Fe, Li₂S and Li-rich phases. The second cycle discharge curves exhibit two plateaus instead of one, indicating a change in the reduction mechanism of the electrode material. It can be seen from Fig. 8 that both the as-synthesized and commercial iron pyrite electrode displays similar voltage profiles, implying that both materials undergo similar electrode reactions during charging and discharging delay in transfer of electrons and Li⁺ at the active material/electrolyte interface is known to result in the polarization in LIB³⁶ and, from the difference between the voltage of the charge and discharge plateau, it can be seen that the as-synthesized pyrite has better electrons and Li⁺ than the commercial sample.

Fig. 8 Galvanostatic charge–discharge voltage profiles of the as-synthesized and commercial iron pyrite electrode when cycled at a current density of 1 A g⁻¹ between 0.001 and 3 V (vs. Li/Li⁺).

To better understand the reaction between the electrode materials and Li⁺ during cycling, cyclic voltammetry (CV) of the as-synthesized iron pyrite electrode recorded at a constant scan rate of 0.2 mV s⁻¹ between 0.001 and 3 V (vs. Li/Li⁺) under ambient temperature for the first, second and fifth cycle is shown in Fig. 9. In the first cycle, two peaks (1.05 V and 0.77 V) were observed during the discharge process of the as-synthesized iron pyrite electrode, associating to the reaction of Li with FeS₂ to form Li₂S and Fe.^16,36 Two peaks were also observed in the first charge process of the as-synthesized iron pyrite electrode where the peak centred at 1.94 V is attributed to the oxidation of Fe to Li₂FeS₂ while that at 2.54 V is associated to the formation of Li_2−xFeS₂ (0 < x < 0.8). Peaks at 1.38 V and 1.98 V during the discharge of the second and fifth cycle are related to the formation of Li_2−xFeS₂ while peaks at around 1.94 V and 2.50 V during the charge step are attributed to the delithiation process of Li₂FeS₂ to form Li_2−xFeS₂. Conversion between Li_2−xFeS₂ and Li₂FeS₂ is reversible, which enables the application of iron sulfides in rechargeable LIBs.¹³

Fig. 9 Cyclic voltammetry of the as-synthesized iron pyrite electrode recorded at a constant scan rate of 0.2 mV s⁻¹ between 0.001 and 3 V (vs. Li/Li⁺) under ambient temperature.

Good rate capability is one of the essential criteria in making high power and fast charging lithium ion batteries. With this idea in mind, the cycling performance of the pyrite FeS₂ electrode at various current densities was tested and the results are presented in Fig. 10. The iron pyrite electrode has a high reversible discharge capacity of around 850 mA h g⁻¹ when cycled galvanostatically between 0.001 and 3 V (vs. Li/Li⁺) at a current density of 100 mA g⁻¹. This discharge capacity value is very near its theoretical capacity of 890 mA h g⁻¹. Subsequent cycling yields a discharge capacity of about 780 mA h g⁻¹, 650 mA h g⁻¹, 530 mA h g⁻¹ and 220 mA h g⁻¹ at a current density of 200 mA g⁻¹, 500 mA g⁻¹, 1 A g⁻¹ and 5 A g⁻¹ respectively. It is worth mentioning that even at a high current density of 1 A g⁻¹ and 5 A g⁻¹, the as-synthesized FeS₂ electrode is still able to exhibit a stable cycling performance, giving an average discharge capacity of about 530 mA h g⁻¹ and 220 mA h g⁻¹ respectively.

Fig. 10 Cycling performance of the pyrite FeS₂ electrode at various current densities.

Conclusions

In conclusion, pyrite FeS₂ with good purity has been successfully synthesized using a facile solvothermal method. The as-synthesized pyrite FeS₂ particles have a homogenous cubic morphology with particle sizes around 80–120 nm. Obvious and well-aligned lattice fringes over the whole cubic face shows that the iron pyrite nanocube has the characteristic of a single crystal. Time-dependent characterization found that the precursors first react to form sheet-like amorphous Fe_1−xS, and with further heating transforms into pyrite FeS₂ nanocubes of about 100 nm. The electrochemical properties of the as-synthesized pyrite FeS₂ nanocubes have been characterized. It was found that the iron pyrite electrode exhibits a stable cycling performance at high current density of 1 A g⁻¹, with a reversible discharge capacity of 540 mA h g⁻¹ after 150 charge–discharge cycles. Even at higher current density of 5 A g⁻¹, the pyrite FeS₂ electrode still managed to give a stable discharge capacity of about 220 mA h g⁻¹. The enhanced lithium storage properties are attributed to its higher specific surface area that can provide more lithium ion (Li⁺) reaction sites, leading to lesser polarization and better cycling performance.

Acknowledgements

The authors would like to thank Mr Sim Daohao for his help with the Raman Spectroscopy measurement.

References

1. J. Cabana, L. Monconduit, D. Larcher and M. R. Palacín, Adv. Mater., 2010, 22, E170–E192.
2. Z. G. Yang, J. L. Zhang, M. C. W. Kintner-Meyer, X. C. Lu, D. W. Choi, J. P. Lemmon and J. Liu, Chem. Rev., 2011, 111, 3577–3613.
3. M. Wagemaker and F. M. Mulder, Acc. Chem. Res., 2013, 46, 1206–1215.
4. J. M. Tarascon and M. Armand, Nature, 2001, 414, 359–367.
5. S. Flandrois and B. Simon, Carbon, 1999, 37, 165–180.
6. C. Liu, F. Li, L.-P. Ma and H.-M. Cheng, Adv. Mater., 2010, 22, E28–E62.
7. J. R. Szczech and S. Jin, Energy Environ. Sci., 2011, 4, 56–72.
8. C. Kang, I. Lahiri, R. Baskaran, W.-G. Kim, Y.-K. Sun and W. Choi, J. Power Sources, 2012, 219, 364–370.
9. M. Armand and J. M. Tarascon, Nature, 2008, 451, 652–657.
10. G. M. Zhou, D. W. Wang, F. Li, L. L. Zhang, N. Li, Z. S. Wu, L. Wen, G. Q. Lu and H. M. Cheng, Chem. Mater., 2010, 22, 5306–5313.
11. X. Rui, H. Tan, D. Sim, W. Liu, C. Xu, H. H. Hng, R. Yazami, T. M. Lim and Q. Yan, J. Power Sources, 2013, 222, 97–102.
12. Q. H. Wang, L. F. Jiao, Y. Han, H. M. Du, W. X. Peng, Q. N. Huan, D. W. Song, Y. C. Si, Y. J. Wang and H. T. Yuan, J. Phys. Chem. C, 2011, 115, 8300–8304.
13. C. Xu, Y. Zeng, X. Rui, N. Xiao, J. Zhu, W. Zhang, J. Chen, W. Liu, H. Tan, H. H. Hng and Q. Yan, ACS Nano, 2012, 6, 4713–4721.
14. F. Liao, J. Światowska, V. Maurice, A. Seyeux, L. H. Klein, S. Zanna and P. Marcus, Electrochim. Acta, 2014, 120, 359–368.
15. D. Zhang, G. Wu, J. Xiang, J. Jin, Y. Cai and G. Li, Mater. Sci. Eng., B, 2013, 178, 483–488.
16. H. Siyu, L. Xinyu, L. QingYu and C. Jun, J. Alloys Compd., 2009, 472, L9–L12.
17. E. Strauss, D. Golodnitsky and E. Peled, Electrochim. Acta, 2000, 45, 1519–1525.
18. D. Zhang, Y. J. Mai, J. Y. Xiang, X. H. Xia, Y. Q. Qiao and J. P. Tu, J. Power Sources, 2012, 217, 229–235.
19. J.-W. Choi, G. Cheruvally, H.-J. Ahn, K.-W. Kim and J.-H. Ahn, J. Power Sources, 2006, 163, 158–165.
20. X. Feng, X. He, W. Pu, C. Jiang and C. Wan, Ionics, 2007, 13, 375–377.
21. D. Bresser, E. Paillard, M. Copley, P. Bishop, M. Winter and S. Passerini, J. Power Sources, 2012, 219, 217–222.
22. N. Li, C. R. Martin and B. Scrosati, J. Power Sources, 2001, 97–98, 240–243.
23. B. C. Kim, K. Takada, N. Ohta, Y. Seino, L. Zhang, H. Wada and T. Sasaki, Solid State Ionics, 2005, 176, 2383–2387.
24. J. P. Wilcoxon, P. P. Newcomer and G. A. Samara, Solid State Commun., 1996, 98, 581–585.
25. E. J. Kim and B. Batchelor, Mater. Res. Bull., 2009, 44, 1553–1558.
26. M.-L. Li, Q.-Z. Yao, G.-T. Zhou, X.-F. Qu, C.-F. Mu and S.-Q. Fu, CrystEngComm, 2011, 13, 5936–5942.
27. J. Puthussery, S. Seefeld, N. Berry, M. Gibbs and M. Law, J. Am. Chem. Soc., 2010, 133, 716–719.
28. Y. Bi, Y. Yuan, C. L. Exstrom, S. A. Darveau and J. Huang, Nano Lett., 2011, 11, 4953–4957.
29. W. Li, M. Doblinger, A. Vaneski, A. L. Rogach, F. Jackel and J. Feldmann, J. Mater. Chem., 2011, 21, 17946–17952.
30. A. Kirkeminde, B. A. Ruzicka, R. Wang, S. Puna, H. Zhao and S. Ren, ACS Appl. Mater. Interfaces, 2012, 4, 1174–1177.
31. M. Gong, A. Kirkeminde and S. Ren, Sci. Rep., 2013, 3, 2092.
32. L. Zhu, B. J. Richardson and Q. Yu, Nanoscale, 2014, 6, 1029–1037.
33. M. Alam Khan, M. O. Manasreh and Y. M. Kang, Mater. Lett., 2014, 126, 181–184.
34. T. A. Yersak, H. A. Macpherson, S. C. Kim, V. D. Le, C. S. Kang, S. B. Son, Y. H. Kim, J. E. Trevey, K. H. Oh, C. Stol dt and S. H. Lee, Adv. Energy Mater., 2013, 3, 120–127.
35. C. Wadia, Y. Wu, S. Gul, S. K. Volkman, J. Guo and A. P. Alivisatos, Chem. Mater., 2009, 21, 2568–2570.
36. D. Zhang, X. L. Wang, Y. J. Mai, X. H. Xia, C. D. Gu and J. P. Tu, J. Appl. Electrochem., 2012, 42, 263–269.
37. L. Zhu, B. Richardson, J. Tanumihardja and Q. Yu, CrystEngComm, 2012, 14, 4188–4195.
38. Q. Xuefeng, X. Yi and Q. Yitai, Mater. Lett., 2001, 48, 109–111.
39. S. Kar and S. Chaudhuri, Chem. Phys. Lett., 2004, 398, 22–26.
40. J. A. Breier, C. R. German and S. N. White, Geochem., Geophys., Geosyst., 2009, 10, Q05T05.
41. H. D. Lutz and B. Müller, Phys. Chem. Miner., 1991, 18, 265–268.
42. S. B. Turcotte, R. E. Benner, A. M. Riley, J. Li, M. E. Wadsworth and D. Bodily, Appl. Opt., 1993, 32, 935–938.
43. A. K. Kleppe and A. P. Jephcoat, Mineral. Mag., 2004, 68, 433–441.
44. H. Vogt, T. Chattopadhyay and H. J. Stolz, J. Phys. Chem. Solids, 1983, 44, 869–873.
45. A. Boughriet, R. S. Figueiredo, J. Laureyns and P. Recourt, J. Chem. Soc., Faraday Trans., 1997, 93, 3209–3215.
46. H. Li, G. Richter and J. Maier, Adv. Mater., 2003, 15, 736–739.
47. S. Laruelle, S. Grugeon, P. Poizot, M. Dolle, L. Dupont and J. M. Tarascon, J. Electrochem. Soc., 2002, 149, A627–A634.

---

Footnote

† Electronic supplementary information (ESI) available: BET isotherm profile of as-synthesized iron pyrite nanocubes, FESEM and XRD pattern of the grinded commercial iron pyrite mineral. See DOI: 10.1039/c4ra08527b

---