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Self-Organization of Layered Inorganic Membranes in Microfluidic Devices Qingpu Wang, Megan Rae Bentley, and Oliver Steinbock J. Phys. Chem. C, Just Accepted Manuscript • Publication Date (Web): 09 Jun 2017 Downloaded from http://pubs.acs.org on June 14, 2017

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The Journal of Physical Chemistry

Self-Organization of Layered Inorganic Membranes in Microfluidic Devices

Qingpu Wang, Megan R. Bentley, and Oliver Steinbock* Department of Chemistry and Biochemistry, Florida State University, Tallahassee, Florida 32306-4390, USA.

ABSTRACT

Inorganic precipitate membranes play an important role in chemobrionics and origin of life research. They can involve a range of catalytic materials, affect crystal habits, and show complex permeabilities. We produce such membranes in a microfluidic device at the reactive interface between laminar streams of hydroxide and Co(II) solutions. The resulting linear membranes show striking color bands that over time expand in the direction of the Co(II) solution. The cumulative layer thicknesses (here up to 600 µm) obey square root laws indicating diffusion control. The effective diffusion coefficients are proportional to the hydroxide concentration but the membrane growth slows down with increasing concentrations of Co(II). Based on spatially resolved Raman spectra and other techniques, we present chemical assignments of the involved materials. Electron microscopy reveals that the important constituent β-Co(OH)2 crystallizes as thin hexagonal microplatelets. Under drying, the membrane curls into spirals revealing mechanical differences between the layers.

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INTRODUCTION

Chemobrionics is the study of self-organizing processes that create life-like morphologies from seemingly simple inorganic reactions. These macroscopic structures are typically the result of diffusion and fluid motion controlling precipitation reactions.1-3 The most iconic examples are so-called “chemical gardens” which are hollow tubes with radii between 1 µm and 1 cm that form when salt crystals are placed into an alkaline solution containing silicates, carbonates, phosphates, or other anions.4-9 For the classical case of sodium silicate, the tube wall consists of an outer layer of amorphous silica and a thicker interior layer of metal hydroxides or oxides.10 Their growth is triggered by the spontaneous formation of a colloidal, semipermeable membrane around the seed that compartmentalizes the system and controls an osmotically driven inflow of water. Once the fragile membrane ruptures, a buoyant jet of salt solution templates the either steady or oscillatory growth of the tube.11 Over the past decade, the study of chemical gardens has seen a renaissance due to possible applications in materials science and origin of life research. The tubes are microporous12 and often catalytically active as exemplified by aluminosilicate13 and silica-ZnO systems14. They can be based on polyoxometalates15,16, super-paramagnetic magnetite17, or include trapped fluorescent quantum dots18 bestowing sensing capabilities on the system. Furthermore, applications as three-dimensional microfluidic networks have been suggested.19 Their relevance to prebiotic chemistry and the origin of life lies in similarities to the chimneys at hydrothermal vents on the ocean floor.20,21 Russell et al.22 suggested that life on Earth began in the precipitate materials of off-axis alkaline vents that provided spatial confinement in the form of


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interconnected pores with catalytic surfaces, geochemically driven pH gradients, and reactants such as CO, CH4, and H2. Also the growth dynamics of the tubular structures have attracted considerable attention with several groups formulating quantitative theories for the tube radius23, the growth speed24, and observed pressure oscillations25. In addition, Roszol and Steinbock26 found, indirectly via weight measurements, that the wall thickness in silica-Cu(OH)2 systems increases over time and that this thickening occurs strictly in the direction of the interior copper solution. Many of these quantitative experiments replaced the seed crystal with a corresponding solution that is injected into the silicate solution at constant flow rates. Further efforts to control and simplify the growth phenomenon include a study of tube growth by injection into a sodium hydroxide solution demonstrating that no additional ions such as silicate are needed.27 In 2015, Haudin et al.28 investigated chemical gardens that were produced by delivering cobalt(II) chloride into thin layers of silicate solution. These experiments yielded quasi-two dimensional analogs of chemical gardens as well as several unexpected patterns.29,30 All of these structures develop a striking palette of colors similar to the one shown in Figure 1a. To date, the corresponding compounds have not been identified systematically and little is known about the chemical composition and structure of the membranes forming across the steep concentration gradients between the solutions. Here we continue very recent work on precipitation reactions in microfluidic devices that simplify the geometry of the membrane to a thin, linear wall.31,32 Specific goals were the in situ and ex situ characterization of the complex membrane constituents and the measurement of their individual thickening dynamics. Our study accomplishes these goals for a simplified version of the two-dimensional experiments by Haudin et al.28, which excludes silicate and exposes the



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advecting cobalt solution to an adjacent laminar stream of sodium hydroxide. An example of the quasi-two-dimensional patterns for the latter reactants is shown in Figure 1b. We also report qualitative findings on the mechanical behavior of the membrane.

EXPERIMENTAL SECTION

We perform reference experiments on precipitate patterns forming in quasi-two-dimensional systems (see Figure 1 and Figure S1 in the Supporting Information) following the methods reported for confined chemical gardens by Haudin et al.28 The setup used for these measurements is a horizontal Hele-Shaw cell consisting of two transparent Plexiglas plates (20 cm × 20 cm) separated by a gap of 0.5 mm. The cell is initially filled with 0.5 M NaOH (Sigma-Aldrich) solution or 0.5 M Na2SiO3 (Fisher) solution. Then 3.2 mL of 0.5 M CoCl2 (Sigma-Aldrich) solution are injected into the cell using a programmable syringe pump (New Era Pump Systems, NE 4000). The flow rate is 0.66 mL/min and identical to the one in ref. 28. We do not degas the reactant solutions which hence contain dissolved oxygen. Photos are acquired using a digital camera (Nikon D3300) equipped with a Tamron 1:1 macro lens (90 mm, f/2.8). In the main part of our study, reactions are carried out in a microfluidic device that is similar but larger than the one described in ref. 31. As shown in Figure 2a, the device consists of a cut Parafilm membrane placed between two Plexiglas plates (5 cm × 8 cm). Several holes are drilled into the plates to create the inlets and to secure the device using screws and nuts. Two barb fittings are glued onto the upper Plexiglas plate using an epoxy adhesive. A Y-shaped pattern is cut out from the Parafilm as the channel for the reaction using an inexpensive electronic cutting tool (Silhouette Portrait) connected to a PC. The straight part of the channel is 6 cm long, 3 mm wide, and approximate 0.13 mm high. Once the construction of the microfluidic device is


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completed, two syringes containing NaOH solution (left) and CoCl2 solution (right) are connected to the barb fittings with plastic tubing (Tygon, inner diameter 1/16’’). The syringes are discharged using a programmable syringe pump (New Era Pump Systems, NE 1600). Prior to the experiment, we fill the channel with water to prevent undesired (primarily surface-tension induced) effects that are observed if we inject solutions directly into the air-filled device. To avoid pressure variations caused by the discontinuous detachment of droplets from the outflow port, we guide the solution to a waste container via a small membrane bridge made of filter paper. All microfluidic experiments are performed at room temperature (22 °C) and at a constant pump rate of 1 mL/h per syringe. The reaction time varies from about 1 h to 10 h, depending on the concentrations of the solutions used. The progress of the precipitation processes within the microfluidic device is observed using an inverted microscope (Leica DM IRB). The microscope is connected to a camera (Nikon D3300) and photos are collected using a PC and digiCamControl software. During microscopy, the sample is illuminated using white LEDs and all optical data acquired is based on the reflection of this light. For further characterization, the precipitate is carefully extracted from the device, rinsed in ultrapure water, and dried under ambient conditions. Micro-Raman spectroscopy (JY Horiba LabRam HR800) is carried out at an excitation wavelength of 633 nm; scanning electron microscopy (JEOL 7401 FE-SEM) is performed at 10 kV on gold-sputtered samples.

RESULTS AND DISCUSSION

In our experiments, a linear precipitate wall forms at the fluid interface when NaOH solution and CoCl2 solution are injected into the microfluidic device. The wall extends from the corner of the mixing point down to the outflow port. As the wall thickens, we find that the membrane growth



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occurs almost exclusively in the direction of the CoCl2 solution. Figure 3a-d show the same area at 20 min intervals. Notice that the hydroxide side of the wall is further to the left and outside of the field of view. In most of our measurements, the combined flow rate of the two reactant streams corresponds to an average fluid velocity v = 1.42 mm/s and a Reynolds number of Re ≈ 1. Considering the length of the reaction channel, we find an average transit time tT of 42 s, which is very short compared to the duration of the overall growth experiment (1-10 h). Accordingly, the membrane growth starts at the mixing point and very quickly (on the time scale of tT) spreads along the channel. Subsequently, we do not observe marked differences in the width and overall appearance of the membrane along the channel. However, there are small variations at the extreme beginning and end of the channel, which we interpret as effects of the flow field. Also notice that over time, v increases and tT decreases slightly due to the channel-width reducing growth of the membrane. Figure 3a-d also show that some (green) dendrite-like structures emerge gradually on the surface of the precipitate membrane that is in contact with the cobalt solution. This phenomenon is observed in most of our experiments except those with the largest ratio of concentrations studied ([OH-]/[Co2+] = 5). A magnified view of this type of structure is shown in Figure 3e. As confirmed by variations of the microscope’s focal plane, these dendrite-like structures extend with only small changes along the vertical direction of the channel. In some experiments, we noticed that the base of the structures nucleated at defects in bright, pink area of the membrane. These defects are discernible in our optical micrographs (e.g. black arrow in Figure 3d). We reemphasize that the cobalt-based membrane growth occurs exclusively in the direction of the cobalt solution. This finding is in agreement with an earlier study on manganese-based precipitation that found membrane growth to proceed only in the direction of the Mn2+


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solution31. The latter study concluded that the precipitate has very different permeabilities for hydroxide and metal ions favoring the transport of the negatively charged ion while effectively blocking the metal cation. Our observations also support this conclusion. The latter study noted no spatial variations in the chemical composition of the membrane. In the cobalt system, however, we notice striking color differences across the product structure, which strongly suggest layer-like changes in compositions. To study the growth of each layer, we record their dynamics based on image sequences obtained by digital microscopy. The colors in our figures represent the unprocessed video signal but are biased by the white LED illumination, the CMOS sensor in the camera, and the digital reproduction of the figures. Image processing was strictly limited to the adjustment of image brightness and contrast. Figure 4a shows a representative time-space plot of the growing membrane. In this experiment, 0.5 M NaOH solution and 0.5 M CoCl2 solution are injected into the microfluidic device, while we collect images at 10 s intervals for 8 h. We then select a vertical position in these images, at which no unusual features (such as membrane curvature or defects) are observed. At this position, we extract the spatial color profiles. Each profile spans 1406 pixels and corresponds to a 0.6 µm × 900 µm band across the membrane. These profiles are then stacked into the timespace plot which reveals three distinct bands of increasing thickness: green (in contact with cobalt solution), blue, and pink. We also observe a thin black (or brown) region which is in contact with the hydroxide solution. Notice that the value w = 0 corresponds to the initial position of the newly formed membrane. A schematic illustration of the time-space plot construction is given in Figure S2 in the Supporting Information.



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In order to characterize the fairly abrupt color transitions between the green, blue, and pink bands, we select three transition colors in terms of specific RGB values (RGB refers to the red, green, and blue components of the digital image). Image noise makes it necessary to consider small intervals around these values, which are specified in the Figure S3 in the Supporting Information. The resulting sets of (w, t) coordinates that obey our RGB conditions characterizes the front-like motion of the color transitions. In Figure 4b, these point sets are superposed onto the original time-space plot as red, yellow, and blue dots. For further analysis, we average the three data sets individually and for each time to yield functions ‫ݓ‬ ഥ(t) describing the evolution of the color transitions. These curves are shown in Figure 4c and are fitted individually to the square root function ‫ݓ‬ ഥ(t) = (Deff t)1/2 + w0 which assumes diffusion-controlled growth kinetics. For the example in Figure 4c, there is very good agreement between the measurements and the fits. We hence conclude that—at least for the reaction conditions of Figure 4—the growth of the materials corresponding to the three different colors, is in each case governed by diffusion while the transforming reactions can be assumed to occur instantaneously. Notice that a diffusioncontrolled thickening of the membrane combined with kinetically controlled transitions at time τ do not describe our data. The latter mechanism yields transition dynamics of the form ‫ݓ‬ ഥ(t) = [Deff (t – τ)]1/2 that in the time-space plot, would generate constant time delays between the transition curves. This clearly not the case. Systematic variations of the reactant concentrations also revealed conditions where the system dynamics are more complex. Especially at low cobalt concentrations, we observe a thick gel-like layer that forms within the first few seconds of the reaction. Possibly for related reasons, efforts to measure the three Deff values as a function of [Co2+] yielded unsatisfactory results. Systematic variations of the initial hydroxide concentration at a constant cobalt concentration of 0.5 M,


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however, did reveal simple dependencies (Figure 4d). The values of all three effective diffusion coefficients Deff increase with increasing hydroxide concentrations and obey proportional dependencies Deff ∝[OH-] with proportionality constants of 2.77 × 10-7 cm2/(s M), 3.50 × 10-7 cm2/(s M), and 3.90 × 10-7 cm2/(s M) for the pink to blue, blue to green, and green to solution transitions, respectively. In Figure 4, these transitions are color-coded as blue, yellow/green, and red. The finding of proportional dependencies is consistent with a previous theoretical study on the membrane growth rate of manganese-based membranes31. In our study, however, we do not only consider the overall thickness but also internal compositional changes. We also employed a simpler analysis for the characterization of the membrane growth rate with respect to different reactant concentrations. This method does not assume a particular function but rather measures the time tw needed for a 300 µm thick membrane to grow to a width of 500 µm. The latter widths are chosen to minimize the effects of possible, initial transients while avoiding extremely long growth times. Results of such measurements are shown in Figure 5. We find that an increase in the initial hydroxide concentration causes faster growth and hence decreasing values of tw. An increase in the initial cobalt concentration, however, causes slower thickening and increasing values of tw. Notice that the former result is in agreement with the observed dependence of Deff on [OH-] (Figure 4d). The latter result surprisingly shows the opposite trend that one would expect based on the earlier study of manganese-based membranes which found increasing values of Deff with increasing metal ion concentrations31. The reasons for this difference are unknown but one can speculate that the internal compositional changes in the cobalt system and the existence of a very thick gel layer are contributing factors. In the following, we analyze the microscopic morphology and the layered composition of the precipitate membranes. For experiments with 0.5 M NaOH and 0.5 M CoCl2, the precipitate is



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removed from the device after 1 h of growth, rinsed with water, and dried under ambient conditions for SEM imaging. Figure 6a shows a representative micrograph of the cobalt-based precipitate illustrating that the membrane’s striking rectangular shape is preserved during drying. The exposed, bright surface of the structure in (a) was, during the experiment, in contact with the Plexiglas of the microfluidic device. Closer inspection of this surface reveals systematic changes in roughness and texture (b). The following, higher-magnification images (c-h) are arranged from the hydroxide to the cobalt side. Near the original hydroxide solution, we find irregular, disordered nanoflakes. Their large dimension spans 100-500 nm but their thickness is much smaller. The edges of the nanoflakes are disordered and uneven. As we move closer to the central area of the membrane, we find hexagonal plate-like crystals with a diameter of about 2 µm (Figure 6f). Most of these platelets are closely stacked and coplanar to the membrane surface, whereas some stand upright revealing a very thin width of about 70 nm (Figure 6d). In certain areas, the hexagonal crystals do not fully cover the membrane surface and expose smaller, irregular objects that might be of an amorphous nature as well as smaller crystals with an amorphous overgrowth (Figure 6e, f). Adjacent to the area of hexagonal crystals and farther towards the cobalt-solution side, we find another type of thin nanostructure (Figure 6g) and finally at the cobalt-solution interface, the flake-like structures shown in Figure 6h. Notice that the border colors of four images in Figure 6 correspond to the colors (black, brown, pink, blue/green) of the different macroscopic regions of the membrane. The compositional changes across the membrane are quantified by micro-Raman spectroscopy. Figure 7a-d show Raman spectra obtained from four different areas. These areas are crossreferenced with Roman numerals according to Figure 7e which shows an optical micrograph of the dried membrane. Notice that each measurement region spans about 1 µm of the 280 µm wide


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membrane. The spectrum in (a) displays intense bands at 190, 473, and 680 cm-1 and bands of lower intensity at 516 and 611 cm-1. These bands are essentially identical to those reported in the literature for Co3O4.33 The Raman spectrum in (b) has a strong band at 499 cm-1 and broader features at higher wavenumbers up to 800 cm-1. Both characteristics are present in a published spectrum of CoO(OH).33 Furthermore, these compounds are known to be black and brown, respectively and hence, match the observed colors in our experiments. The spectrum in Figure 7c matches the spectrum of commercially obtained, pink β-Co(OH)2 (see Figure S4 in the Supporting Information for the reference data) with bands at 149, 434, 566, and 712 cm-1. Lastly, the spectrum in (d) corresponds to the green region of the membrane, which we tentatively assign to the compound α-Co(OH)2 because the synthesized pure phase α-Co(OH)2 was described as greenish.34,35 Additional characterizations using x-ray diffraction and infrared spectroscopy are presented in Figure S5 and S6 in the Supporting Information. Notice that earlier studies reported that the polymorph α-Co(OH)2 can continuously transform into β-Co(OH)2,36 which is in agreement with our observations of the membrane growth. Moreover, the green color is likely caused by the presence of Co(II) ions that are tetrahedrally coordinated to hydroxide ions and interstitial anions. β-Co(OH)2 has a brucite-like structure and Co(II) is coordinated to hydroxide ions forming edge-shared octahedrons, which assemble into charge-neutral layers with an inter-layer spacing of 4.6 Å.37 The octahedral symmetry of βCo(OH)2 appears to be the main reason of its pink appearance. These and the other color variations and assignments are summarized in Figure 8. Notice that the blue layer only appears in the hydrated precipitate, but quickly disappears to become part of the green layer if the precipitate is dried or the flow is stopped (see Figure S7 in the Supporting Information). The



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unstable blue color could be caused by the presence of an ion such as [Co(OH)4]2- (known to be dark blue38) in the hydrated, gel-like membrane. Lastly, we discuss physical changes of the precipitate membrane during drying. Our experimental approach of forming the membrane in a microfluidic channel gives us the unique opportunity to quantify the changes in membrane size and density due to water loss and other processes. This question is of importance to the characterization of all chemical-garden like structures because their analyses nearly exclusively rely on dried samples.1 For this purpose, we extract the membrane from the device, rinse it with water, and dry it in a Petri dish under ambient conditions for 24 h. The comparison of the original, hydrated sample and the dried sample does not reveal any volume change that could be discerned from microscope images (see Figure S7 and movie S3 in the Supporting Information). Consequently, measurements of the thickness and volume of the dried membrane (by SEM or other methods) are sufficiently accurate descriptions of the original, hydrated structure. Nevertheless, there is indirect evidence of small volume changes. Figure 9 shows a 4.5 cm long segment of membrane that within about one minute of drying curled up into a spiral (see movie S4 in the Supporting Information). The spiral is curved in the plane of the microfluidic device and accordingly, the curling is caused by a difference in the volume rate ∆V/V between the different layers, similar to the curling of a bimetallic strip. The main participants in this process are the blue/green layers whereas the wide pink area breaks into brick-like segments that can detach from or remain on the blue/green strip. These layers clearly have sufficient elasticity to endure the build-up of curvature.


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Also notice that the surface that originally was in contact with the cobalt solution, is the highest curved, inner part of the spiral-shaped strip. Therefore, this layer must experience a stronger (or faster) volume decrease than the central parts of the membrane. The colors in Figure 9 do not readily match our earlier discussion because the curling membrane was synthesized at a small concentration ratio of [Co2+]/[OH-] = 0.2 for which the pink-to-blue transition region appears slightly greenish and the blue color is much darker. Moreover, if the membrane is left in the microfluidic device and the solutions are replaced by air, curling of the membrane does not occur. Accordingly, there is an adhesive contact between the membrane and the Plexiglas surfaces.

CONCLUSIONS

Motivated by recent studies of two-dimensional chemical gardens28 -30, we have demonstrated the production of cobalt-based inorganic membranes in a microfluidic device. This device creates a sharp interface between laminar flows of Co2+ and OH- solutions, which then develops into a membrane with distinct color bands. The membrane growth occurs in the presence of very steep concentration gradients measuring about 1-100 mol/(L mm). These gradients are externally sustained for up to 10 h and hence create an excellent setting for the formation of materials under far-from-equilibrium conditions. Micro-Raman spectroscopy allows the analysis of this macroscopic layers and reveals compounds including α- and β-Co(OH)2 as well as mixed valence oxides. The temporal evolution of the membrane thickness and the different band transitions obey simple square-root laws which show that the growth and the compositional transitions are



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diffusion controlled. The individual effective diffusion coefficients are proportional to the employed hydroxide concentration and seem to decrease with increasing cobalt concentration, although the latter dependence had to be inferred from a simpler analysis. These findings can be compared to the work by Batista et al.31 who studied primarily manganese-based membrane growth in a microfluidic device and did not analyze transformations within the membranes. The effective diffusion coefficients, however, increased with both increasing Mn2+ and OHconcentrations. The reasons for this difference between the two systems are unclear but might be related to the internal reactions within the membrane. The compounds and their crystal habits detected in our membranes are of interest to materials science and for industrial applications. The membrane constituent Co(OH)2, for example, is a much-studied additive in rechargeable Ni/Co batteries39-41 and has also seen uses as a precursor for heterogeneous catalysts42,43. The detected oxides of cobalt are also a crucial precursor of ceramics with very high specific capacitances44,45. The presence of Co3O4 and the non-stoichiometric CoO(OH) require redox reactions which seem to involve dissolved oxygen. This link should be further investigated in future experiments that exclude oxygen. In addition, we observed interesting plate-like, hexagonal crystals of the β-Co(OH)2. Such crystals have been synthesized earlier but their production required hexamethylenetetramine as a hydrolysis agent. Beyond these interesting features, the membrane shows macroscopic striation which suggest that our method could allow the engineering of various complex membranes. These membranes could have tailored macro-layers with different compositions and transport features. Furthermore, we envision the trapping of nanoscopic units such as quantum dots that could introduce sensor-like capabilities. The trapping CdSe/ZnS quantum dots has been recently reported for controlled chemical-garden tubes and it was also shown that the nanocrystals


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remained accessible to ions in the surrounding solution.18 This approach introduces a new design feature for microfluidic devices that is surprisingly simple and versatile. The identification of the different compounds across the membrane also provides a better understanding of the beautiful patterns in spatially confined, quasi-two-dimensional chemical gardens (Figure 1). The color sequence within the border of these patterns (Figure 1b) matches the layering in our linear membranes. They hence provide a geometrically simpler version of the curved interfaces in Hele-Shaw cells and most likely also of three-dimensional chemical garden tubes where a characterization of the layering would be extremely difficult. We hope that future studies will also apply our method to questions relating to prebiotic chemistry such as the possible formation of simple organic molecules and polymers (e.g. peptides46 and pyrophosphates47) in precipitate membranes. A microfluidic approach to these interesting questions offers not only greatly reduced reactant volumes but should also allow for a highly parallel exploration of the vast underlying parameter spaces.

ASSOCIATED CONTENT Supporting Information The supporting information is available free of charge on the ACS Publication website. Overhead photos of precipitation patterns in horizontal Hele-Shaw cells, schematics illustrating the time-space plot construction, histograms of the red, green, and blue image channels in Figure 4a, micro-Raman spectra of the pink layer of the cobalt-based membrane and commercially obtained Co(OH)2, X-ray diffraction pattern of the dried precipitate, IR spectra as



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well as micrographs of the dried and hydrated precipitate, and text describing additional details regarding the SI movies (PDF). Movie of a precipitate membrane growing in the microfluidic device (AVI). Movie of the precipitate growing in the microfluidic device as viewed though a microscope (AVI). Movie showing a precipitate membrane drying in the device for 24 h (AVI). Movie showing that the extracted precipitate membrane is winding up into a spiral when dried in air at ambient conditions (AVI).

ACKNOWLEDGEMENTS

This material is based upon work supported by the National Science Foundation under Grant No. 1609495. We thank Elias Nakouzi, Dayton Syme, and Bruno C. Batista for technical help and discussions.

AUTHOR INFORMATION

Corresponding Author: * E-mail: [email protected]. Phone: +1 850-644-4824. Notes: The authors declare no competing financial interest.


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Figure Captions

Figure 1. Precipitation patterns in a thin, horizontal solution layer confined to a Hele-Shaw cell. (a) 0.5 M cobalt chloride (pH 6.2) is injected into the cell containing 0.5 M sodium silicate (pH 13.3). Field of view: 1.7 cm × 1.7 cm. (b) An identical cobalt solution is injected into 0.5 M sodium hydroxide (pH 13.4). Field of view: 1.2 cm × 1.2 cm. In both cases, 3.2 mL of cobalt solution are delivered at a rate of 0.66 mL/min. Photos were taken about 2 h after injection.

Figure 2. (a) A cut Parafilm membrane is placed between two Plexiglas plates. NaOH solution (left) and CoCl2 solution (right) are injected into the device using a syringe pump. (b) Photograph of the resulting membrane (thin vertical line) in our microfluidic device. See also movie S1 in the Supporting Information. (c) Micrograph of the precipitate in the device. Field of view: 1.3 × 1.9 mm2.

Figure 3. (a–d) Image sequence of the thickening membrane at 20 min intervals, showing only the side in contact with the CoCl2 solution. The initial concentrations of the reactants are [NaOH] = 1.75 M and [CoCl2] = 0.5 M. Field of view: 0.32 × 1.3 mm2. Frame (a) was recorded 72 min after start of the injection. See also movie S2 in the Supporting Information. (e) Dendrite-like growth on the side in contact with CoCl2 solution. Field of view: 0.50 × 0.50 mm2.

Figure 4. (a) Time-space plot of the growing membrane compiled from time-lapse videos. Reactant concentrations: [NaOH] = 0.5 M and [CoCl2] = 0.5 M. (b) The same data as in (a) with superposed color borders defined in terms of specific RGB values. (c) Color borders from (b) with superposed square root fits. (d) Effective diffusion coefficients Deff as a function of the concentration of OH- for [Co2+] = 0.5 M = const. The continuous lines are best fits assuming proportional dependencies. The small error bars in (d) are standard deviations from three independent experiments.



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Figure 5. Reaction time tw required for the width of the precipitate to change from 300 µm to 500 µm. The plots show tw as a function of the concentration of (a) OH- for [Co2+] = 0.5 M = const. and (b) Co2+ for [OH-] = 1.0 M = const.

Figure 6. Scanning electron micrographs of a precipitate membrane grown for 1 h with [NaOH] = 0.5 M and [CoCl2] = 0.5 M. The scale bars correspond to (a) 200 µm, (b) 40 µm, (c-e) 1 µm, (f-g) 500 nm, and (h) 1 µm.

Figure 7. (a-d) Micro-Raman spectra of different regions on the dried precipitate membrane highlighted by the arrows in (e). These measurements were performed across the surface that was in contact with the Plexiglas plates.

Figure 8. Schematics of the compositional layers forming the precipitate membrane.

Figure 9. Drying of the extracted membrane in air under ambient conditions causes it to bend and wind up into a spiral. Reactant concentrations: [NaOH] = 1.0 M and [CoCl2] = 0.2 M. Reaction time in the microfluidic device: 1 h. Field of view: 1.3 × 1.9 mm2.


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Figure 1



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Figure 2


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Figure 3



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Figure 4


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Figure 6 (two column figure)


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Figure 8


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TOC GRAPHICS


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