Energy Fuels 2010, 24, 3917–3927 Published on Web 06/08/2010
: DOI:10.1021/ef1002167
Kinetics of the Oxidation of a Co-precipitated Mixture of Cu and Al2O3 by O2 for Chemical-Looping Combustion S. Y. Chuang,† J. S. Dennis,*,† A. N. Hayhurst,† and S. A. Scott‡ †
Department of Chemical Engineering and Biotechnology, University of Cambridge, Pembroke Street, Cambridge CB2 3RA, United Kingdom, and ‡Department of Engineering, University of Cambridge, Trumpington Street, Cambridge CB2 1PZ, United Kingdom Received February 25, 2010. Revised Manuscript Received May 5, 2010
A co-precipitated mixture of CuO and Al2O3 is a good oxygen carrier for chemical-looping combustion in, e.g., CuO þ CO f Cu þ CO2. The kinetics of regeneration of this reduced oxygen carrier (particle size of 355-500 μm) by O2 were measured from 300 to 750 C. Care was taken to ensure that these measurements were not affected by interphase mass transfer or mixing of gases in the bed and sampling system. The order of reaction with respect to O2 was found to be approximately unity, and the rates of regeneration in subsequent cycles could be approximated by those in the first. Below 600 C, oxidation of the reduced oxygen carrier was incomplete. Above 600 C, oxidation was complete to CuO and was controlled to a considerable extent by external mass transfer. At these higher temperatures, regeneration involved a shrinking core mechanism and the two consecutive steps: 2Cu þ 1=2O2 f Cu2 O, Cu2 O þ 1=2O2 f 2CuO,
0 ΔH1073 K ¼ - 167:3 kJ=mol 0 ΔH1073 K ¼ - 131:5 kJ=mol
ð1Þ ð2Þ
At 900 C, the fully oxidized carrier decomposed to Cu2O and O2. object of the technique. To overcome this, Scott et al.1 proposed a modified version of CLC, in which oxidation and reduction occur successively in one fluidized bed. In this scheme, a solid fuel, such as coal, is fed to the fuel reactor and gasified by the fluidizing agent, e.g., steam or a combination of steam and CO2, to form synthesis gas. The synthesis gas is then oxidized in the bed of CuO to produce H2O and CO2. When the inventory of CuO becomes depleted, the feed of solid fuel is stopped and the remaining carbon in the bed is allowed to gasify and combust further until all of it has been used up. The fluidizing agent is then switched to atmospheric air, which regenerates the depleted CuO. This scheme depends upon the reduction of the oxygen carrier being sufficiently exothermic to balance the endothermic gasification reactions; this has been achieved in our work using a Cu-based oxygen carrier, which has exothermic reduction reactions (to either Cu2O or Cu) of sufficient magnitude.2 Other than lignite, this scheme has also been used for bituminous coal.3 There is little literature on the kinetics of the regeneration reaction
1. Introduction Chemical-looping combustion (CLC) conventionally involves two interconnected fluidized beds, referred to as the fuel and air reactors, respectively. The granular material fluidized in each is a solid metal oxide, which can be circulated continuously between the reactors, acting as an oxygen carrier. A stream of fuel gas, e.g., a hydrocarbon CnH2m, enters the fuel reactor and reacts with the oxygen carrier, MexOy, to form a reduced oxide, MexOy-1 (or possibly the metal, Me), steam, and CO2 according to ð2n þ mÞMex Oy þ Cn H2m f ð2n þ mÞMex Oy - 1 þ mH2 O ð3Þ
þ nCO2
The steam can be condensed and removed, leaving behind pure CO2 for storage in a suitable geological structure, such as a saline aquifer or a depleted oil and gas field. The reduced carrier is transported to the air reactor, where it is oxidized and regenerated by the incoming air Mex Oy - 1 þ 1=2O2 f Mex Oy
ð4Þ
Cu þ 1=2O2 f CuO
Subsequently, the regenerated oxygen carrier is transported back to the fuel reactor to begin a new cycle of operation. The use of CLC with solid fuels presents problems mainly because the particles of fuel and oxygen carrier cannot be easily separated. Without separation, the solid fuel would enter the oxidation reactor along with the recycling oxygen carrier and give CO2 in the off-gas, thereby defeating the
ð5Þ
Garcı´ a-Labiano et al.4 found that, for a wet-impregnated oxygen carrier containing 10 wt % CuO supported on alumina, (1) Scott, S. A.; Dennis, J. S.; Hayhurst, A. N.; Brown, T. A. AIChE J. 2006, 52, 3325–3328. (2) Dennis, J. S.; Scott, S. A. Fuel 2010, 89, 1623–1640. (3) Dennis, J. S.; M€ uller, C. R.; Scott, S. A. Fuel 2010, DOI: 10.1016/j. fuel.2010.01.037. (4) Garcı´ a-Labiano, F.; de Diego, L. F.; Adanes, J.; Abad, A.; Gayan, P. Ind. Eng. Chem. Res. 2004, 43, 8168–8177.
*To whom correspondence should be addressed. Telephone: þ44-(0)1223-334787. Fax: þ44-(0)1223-334796. E-mail:
[email protected]. r 2010 American Chemical Society
0 ΔH1073 K ¼ - 149:4 kJ=mol
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: DOI:10.1021/ef1002167
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Figure 1. Schematic diagram of the experimental apparatus.
the oxidation of the reduced carrier was controlled by chemical kinetics, with the order of the reaction with respect to O2 being unity and the activation energy being 15 kJ mol-1. Our previous work5 has concentrated on particles, typically containing 82.5 wt % CuO and 17.5 wt % Al2O3, made by co-precipitating the nitrates of copper and aluminum with sodium carbonate, followed by washing to remove all traces of Naþ and calcining to decompose the resulting basic carbonates. These particles maintained quantitative conversion of the content of copper for up to 18 consecutive cycles, consisting of reduction by ∼2.5 mol % CO, with a balance of N2, and reoxidation with ∼21 mol % O2, with a balance of N2, at temperatures between 800 and 900 C in a laboratory-scale fluidized bed. In particular, no evidence of attrition or agglomeration of the particles was detected. The kinetics of the reduction of such a carrier by CO and H2 have also been studied;6,7 the reduction appears to take place indirectly via the intermediate, Cu2O, at high temperatures (∼800 C) and directly to Cu at low temperatures (∼300 C). This paper is concerned with extending the research on the copper-based carrier to measure the kinetics of the oxidation of the reduced carrier by O2. This is important because the rate at which the particles oxidize influences the design of the chemical-looping reactor.
the bed was fluidized with 10.4 vol % O2 in N2 (BOC Ltd.) at a constant flow rate of 7.35 10-5 m3/s (as measured at 25 C and 100 kPa). Experiments were conducted from 300 to 900 C, corresponding to U/Umf being 2.5-7.9, where U is the superficial velocity of gas through the bed and Umf is its minimum value for fluidization. When measuring the order of reaction, the concentration of oxygen was varied by diluting with pure N2 while still keeping the total volumetric flow rate constant at 73.5 cm3/s (at 293 K and 1 bar). When more than 10.4% O2 in N2 was needed, a cylinder containing 21.0 ( 0.5 vol % O2 in N2 (BOC Ltd.) was used. In an experiment, a batch of ∼0.01 g of the reduced carrier (details of its production will be given below) was tipped into the bed via a metal funnel. The off-gas was sampled continuously via a quartz probe to a quadrupole mass spectrometer (Hiden HPR20) and a paramagnetic analyzer, both of which measured [O2] in the off-gases. The paramagnetic analyzer was located in a unit (ABB EL3020), which also contained non-dispersive infrared analysers measuring [CO], [CO2], and [SO2]. Upon entering the unit, the gas flow divided, so that it passed through the individual analysers in parallel at equal rates of flow to each. The sampling system was modeled as a plug flow reactor in series with a continuous stirred tank. The sudden addition or removal of O2 from the fluidizing gas confirmed that the sampling system followed a first-order response; i.e., the true concentration of oxygen, [O2]t, can be approximated with [O2]t = [O2]m þ τd[O2]m/dt, where [O2]m is the measured concentration of oxygen and τ is a time constant for the sampling system. These measurements were made by fluidizing the bed with pure N2 and rapidly withdrawing the probe from the bed, so that it switched from sampling N2 to atmospheric air. The value of τ was found to be 0.6 s with the mass spectrometer and 3.0 s for the paramagnetic analyzer. The measured values of [O2]m were corrected using the equation to derive [O2]t. The rate of consumption of O2 was measured as the product of the flow rate through the bed and the drop in [O2], after the reduced oxygen carrier had been added to the bed. Measurements from the mass spectrometer were calibrated in such a way that the total amounts of O2 consumed, as measured by the mass spectrometer and paramagnetic analyzer, were identical. Measurements from the mass spectrometer were used whenever possible, because it had a lower value of τ; i.e., the measurements of [O2] required a smaller correction. More on this will be given in section 3.1.1. However, the mass spectrometric measurements had the problems of (i) a baseline drifting with time and (ii) its signal-to-noise ratio being low. Therefore, the paramagnetic
2. Experimental Section The experimental setup used in this work is shown in Figure 1. The laboratory-scale fluidized bed consisted of a vertical quartz tube (inner diameter, 29.5 mm; length, 460 mm), with a sintered quartz plate as a distributor (110 mm from the base of the tube) and a plenum chamber below the distributor. Silica sand (2 10-5 m3; 355-425 μm; unfluidized bed depth, 30 mm) was the material fluidized. The tube was housed in an electric furnace, and the temperature of the bed was measured with a K-type thermocouple. In the majority of the experiments to study the kinetics, (5) Chuang, S. Y.; Dennis, J. S.; Hayhurst, A. N.; Scott, S. A. Combust. Flame 2008, 154, 109–121. (6) Chuang, S. Y.; Dennis, J. S.; Hayhurst, A. N.; Scott, S. A. Proc. Combust. Inst. 2009, 32, 2633–2640. (7) Chuang, S. Y.; Dennis, J. S.; Hayhurst, A. N.; Scott, S. A. Chem. Eng. Res. Des., manuscript submitted.
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Figure 2. Rates of oxidation of 0.0122 g of co-precipitated mixture of Cu and Al2O3 sized to 355-500 μm, at 700 C and 1.22% O2 in N2, measured by (a) paramagnetic analyzer and (b) mass spectrometer. The measurements are shown in the raw, smoothed, and deconvoluted forms. The corresponding deconvoluted rates are shown in panels c and d, plotted against time and conversion.
(XRD) patterns of most known crystalline compounds was used to identify the species present in the oxygen carriers. The particles were also mounted on epoxy (Agar Scientific) and polished with silicon carbide paper (RS; P1200) until the cross-sections were exposed, which were then viewed under the microscope.
analyzer was used when the rate of reaction was slow and the experiment was of long duration (>45 s). The reduced oxygen carrier was prepared by reducing the coprecipitated mixture of CuO and Al2O3, prepared as described elsewhere5 and containing 82.5 wt % CuO, with 2.5 vol % CO in N2 in the fluidized bed at 800 C. Upon completion of the reaction, the furnace was turned off and the bed was allowed to cool to room temperature with N2 flowing through it. The particles were then recovered from the sand and sieved to 355-500 and 850-1000 μm. The porosity of the reduced carrier, ε, was determined from Hg porosimetry to be 0.77 ( 0.02. Assuming that all of the CuO was converted to Cu, the ratio of the mass fractions of Cu to Al2O3 was ∼79:21. The theoretical skeletal density, Fs, of the resulting reduced carrier is then given by the sum of the product of the skeletal density and the relative fraction of the components, i.e., (0.79 8940) þ (0.21 4000) = 7900 kg/m3. Accordingly, the bulk density, Fb, of the resulting carrier was ca. 7900 (1 - 0.77) = 1820 kg/m3. From N2 absorption analysis, the BET surface area of the particles, SBET, was ∼6.3 m2/g. The mean radius of the pores, rp = 2ε/SBETFb, was therefore ∼130 nm. The corresponding value from mercury porosimetry is ∼200 nm. A co-precipitated mixture of Cu2O and Al2O3 was also used at some stages of this work. It was made by heating the co-precipitated mixture of CuO and Al2O3 in the fluidized bed at 950 C in pure N2, so that the CuO decomposed to Cu2O and O2, as shown previously.5 When the evolution of O2 had become negligible, the furnace was turned off and the bed was cooled to room temperature with N2 flowing through it. Assuming that all of the CuO was converted to Cu2O, the ratio of the mass fractions of Cu2O/Al2O3 was ∼81:19. The theoretical Fs of the resulting carrier was then (0.81 6000) þ (0.19 4000) = 5620 kg/m3. Mercury porosimetry showed that the porosity of such particles was ∼0.66. Accordingly, Fb for the resulting carrier was 5620 (1 - 0.66) = 1910 kg/m3, with a BET surface area of ∼11 m2/g. Particles were also characterized using an X-ray diffractometer (Philips PW1830) and optical microscope (Digital Blue QX5), to gain a better understanding of the mechanism of the reaction. The diffractometer operated at 40 kV and 40 mA, using Ni-filtered Cu KR radiation. A database containing the X-ray diffraction
3. Results 3.1. Effect of the Sampling System on the Interpretation of the Experimental Measurements. The factors possibly resulting in a loss of accuracy when measuring the kinetics are now discussed. 3.1.1. Effect of the Analyzer. Panels a and b of Figure 2 show the change in [O2] when a batch of 0.0122 g of reduced oxygen carrier, sized to 355-500 μm, was suddenly added to a bed of sand at 700 C fluidized by 1.22 vol % O2 in N2. Measurements were made with both the mass spectrometer and paramagnetic analyzer. The raw measurements in each case were smoothed by taking a moving average of every seven consecutive data points, measured every 0.5 s. The smoothed measurements were then deconvoluted using the technique described in section 2. The effect of τ is more obvious for the case of the paramagnetic analyzer than the mass spectrometer, because the sampling system associated with the former possessed a bigger τ. It can be seen that, even after smoothing, the mass spectrometric measurements are still relatively noisy, even though features such as the minimum, start, and end points are fairly obvious. The smoothness of the mass spectrometric measurements could be further increased by taking the average of a larger number of data points. However, this was not performed to avoid the accompanying loss in trend of the results, e.g., the mole fraction at which the maximum drop occurred. It should be noted that, at the same [O2] in the fluidizing gases at lower temperatures (e500 C), the fall in [O2] was very small and the signal-to-noise ratio of the mass spectrometric measurements became lower, rendering the resulting measurements 3919
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Figure 3. Effects of yO2, temperature, and dp on the rates of reaction of the base case, which was performed with reduced carrier sized to 355-500 μm and at 700 C and yO2 = 0.0122.
Figure 4. Specific rates of oxidation of reduced carrier (355-500 μm) by 1.23 vol % O2 in N2 at 750 C for three different masses of the carrier. Measurements were made with the mass spectrometer.
too noisy and unsuitable for proper analysis. Although this problem can be overcome by adding a bigger batch of particles, the effect of interphase mass transfer can become significant, as will be discussed later. Using the dotted horizontal line in panels a and b of Figure 2 joining the initial and final mole fractions of [O2], the change in mole fraction from its initial value was multiplied by the molar flow rate of fluidizing gas to yield the actual rate of the reaction. The derived rates of the reaction are shown in Figure 2c. It can be seen that the maximum rates were reached in times ∼2 and 7 s using the mass spectrometer and paramagnetic analyzer, respectively. This delay in reaching the maximum rates is partly attributable to the time taken to mix the particles of the carrier in the sand, to heat them up to the temperature of the bed, and for air to be expelled from each particle. However, the delay is mainly caused by the failure of the procedure of deconvolution to fully offset the effect of mixing in the bed and sampling system. The difference in times at which the maximum rate was reached is probably owing to the effect of a bigger τ for the sampling system with the paramagnetic analyzer, coupled with the fact that the response of neither system is entirely first-order. Thus, the effect of improperly correcting for mixing is more significant for the slower instrument, with its larger τ, compared to the faster mass spectrometer. The conversion of the reduced carrier to CuO, X, was determined by integrating the area under the curve in Figure 2c to obtain the actual number of moles of O2 consumed, which was then divided by the theoretical number of moles of O2 consumed, assuming that all Cu in the sample was fully oxidized. Corresponding plots of the rate against X for these two cases are shown in Figure 2d. The higher accuracy of the mass spectrometer over the paramagnetic analyzer can be confirmed from the maximum rate occurring closer to X = 0; i.e., the maximum rate from the mass spectrometric measurements is a better estimate of the initial rate. Therefore, whenever possible, the mass spectrometric measurements were used, except in cases when the noise-to-signal ratio was too high or when the drift in baseline was significant. 3.1.2. Effects of the Temperature, [O2], and Size of the Particle. Using the oxidation of the reduced carrier (sized to dp = 355-500 μm) at 700 C and yO2, the mole fraction of O2 in the fluidizing gas, of 0.0122, as the base case, the effects of a variation in, respectively, temperature, yO2, and diameter of the particle, dp, are examined in Figure 3, in which changes made to the base case are labeled on the corresponding graphs. It should, however, be noted that the measurements for the experiments in which dp and temperature were
changed were made with the paramagnetic analyzer. This was because the corresponding mass spectrometric measurements of the off-gas concentrations at low rates of reaction were too noisy to be used. The measurements for the base case and the experiment in which yO2 was changed were made using the mass spectrometer. It can be seen that, when yO2 was increased from 0.0122 to 0.075, the maximum in the curve was shifted from X ∼ 0.06 to 0.4, making it a poorer estimate of the initial rate of reaction. This was probably because the time of the reaction was short (∼10 s) and the effect of mixing in the bed and sampling system could not be sufficiently minimized with the procedure of deconvolution, even when measured by the faster mass spectrometer. When dp was increased to 850-1000 μm, the time of the reaction increased from ∼40 to 150 s and the X at which the maximum rate occurred remained at ∼0.06. This demonstrates that the longer the time of the reaction, the greater the likelihood of obtaining more realistic measurements, even if the slower paramagnetic analyzer is used. When the temperature was decreased to 300 C, the time of the reaction increased to ∼600 s and the X at which the maximum rate was reached dropped to ∼0.05. This again confirms the higher level of accuracy from long experiments. The lower than unity conversion (∼0.3) at 300 C will be discussed later. In summary, a longer time of reaction was necessary if accurate measurements were to be made. This could be accomplished using larger particles or reducing the operating temperature and yO2. 3.1.3. Effect of Interphase Mass Transfer. Interphase mass transfer refers to the transfer of reactants from the bubble to the particulate phase in the fluidized bed. If the rate of interphase mass transfer is slower than the rate of reaction in the particulate phase, the former will become the controlling mechanism. To check the effect of interphase mass transfer, batches of reduced carrier with different masses were oxidized with 1.23 vol % O2 in N2 at 750 C. The rate of interphase mass transfer with Geldart Group B solids is not significantly affected by the temperature.8 The resulting rates of reaction are shown in Figure 4. There is no appreciable change in the specific rate when the mass of carrier used was increased from 0.0121 to 0.0219 g; however, the maximum rate dropped by ∼50% and the time of reaction increased by ∼50%, when the mass was increased to 0.087 g. This indicates the onset of interphase mass transfer. Because the mass of the carrier used in this study was always (8) Davidson, J. F.; Harrison, D. Fluidization; Academic Press: New York, 1963.
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Figure 5. Variation of the final conversion with the temperature for the oxidation of the carrier containing Cu to CuO at 1.3 vol % O2 in N2. The range of X measured is shown by the vertical bars.
∼0.01 g or below, the effect of interphase mass transfer was negligible. 3.2. Variation of X with the Temperature. Preliminary experiments were performed at 1.3 vol % O2 in N2, using the oxygen carrier fully reduced to Cu and sized to 355500 μm, at 300-900 C, to obtain the relationship between X and temperature. At least four experiments were performed at each temperature, and the results are shown in Figure 5. This shows that X < 1 below 600 C. It should be noted that X was calculated by integrating the area under curves, such as those in Figure 2c. For the reactions at low temperatures (500 C. The detection of both Cu2O and CuO could mean that the oxidation of the reduced carrier is a two-stage process via the intermediate, Cu2O, at these low temperatures. Alternatively, the oxidation reaction might proceed via both the direct (Cu f CuO) and indirect (Cu f Cu2O f CuO) routes. The mechanisms will be discussed further, later. From Figure 5, it can also be seen that X dropped at 900 C, because Cu2O is the thermodynamically stable oxide at these conditions5 (see Figure 6). The presence of Cu2O at 900 C was confirmed by XRD analysis of the oxidized particles, which were retrieved from the bed by cooling it to room temperature with N2 flowing through it. This was to prevent further oxidation to CuO, should the partially oxidized sample come into contact with atmospheric air. To ensure that the oxidation of the reduced oxygen carrier was not affected by such decomposition at 900 C, experiments to measure kinetics were only performed at 750 C and below. 3.3. Order of the Reaction. As noted in section 3.1.1, the maximum of the deconvoluted curve of the rate of reaction against conversion, X, was taken as the initial rate of reaction. The reaction was assumed to follow the rate equation: r = k[O2]n. To derive the order of reaction, n, and rate constant, k, [O2] was varied. The variation of the initial rate with [O2] is shown in Figure 7, for carriers containing, initially, either Cu or Cu2O, at selected, fixed temperatures. It is important to note that there is always a proportional relationship between the initial rate and [O2] at low [O2], but this rate levels off at higher [O2]. This might suggest that the
Figure 6. Phase diagram of the oxidation of Cu to CuO.
Figure 7. Plots of the rate of consumption of O2 (by unit mass of the carrier) against [O2] for carriers containing Cu or Cu2O at selected temperatures. The range of rates measured is shown by the vertical bars. The dotted straight lines are tangents of the early parts of the curves.
reaction was following, e.g., a Langmuir-Hinshelwood mechanism, in which case n would change from 1 to 0, when [O2] is increased. However, it is more likely that the rates of reaction at high [O2] were too fast for reliable measurements to be made. This is illustrated in panels a and b of Figure 8, which give plots of the rate against time and rate against conversion for the oxidation of the carrier containing Cu at 7.3 vol % O2 in N2 at 400 and 650 C, at which [O2] = 1.31 and 0.96 mol/m3, respectively. In Figure 7, the measurement at 650 C was made with the mass spectrometer and the measurement at 400 C was made with the paramagnetic analyzer, because the mass spectrometric measurements at 400 C were noisy and the baseline drifted. As seen in Figure 8a, at 650 C, the reaction took ∼20 s to complete, and the resulting plot of the rate against conversion in Figure 8b shows a maximum at X ∼ 0.4, which made it a poor estimate of the initial rate. The reaction at 400 C took ∼100 s to finish, but most of the O2 was consumed in the first 15 s, when X ∼ 0.25. The corresponding maximum rate was reached at X ∼ 0.15. Because the reaction by and large stopped after 15 s, when X ∼ 0.25, it is likely that the maximum rate is a poor estimate of the initial rate. Hence, it was probably the inability of the paramagnetic analyzer to measure an accurate initial rate that gave rise to the leveling off of the rate at lower [O2] at 400 C, as seen in Figure 7, relative to the mass spectrometric measurements at higher 3921
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Figure 9. Plots of the rate of oxidation of the reduced oxygen carrier (containing Cu and Al2O3) by 1.3 vol % O2 in N2. ([ and ]) Carrier sized to 355-500 μm and (9 and 0) carrier sized to 850-1000 μm. ([ and 9) theoretical external mass-transfer rates and (] and 0) measured rates.
where εmf is the voidage of the particulate phase of the bed (taken to be 0.42) and Rep is the Reynolds number for the particulate phase, calculated as Umfdp/ν, where ν is the kinematic viscosity of the fluidizing gas. The Schmidt number is Sc = ν/DG. For the reduced oxygen carrier sized to 850-1000 μm, the mass of a spherical particle is mp = 7.5 10-4 g/particle for dp = 9.25 10-4 m and S = 2.7 10-6 m2. At 700 C and 1.22 vol % O2 in N2, CB ∼ 0.15 mol/m3, U = 0.343 m/s, Umf = 0.058 m/s, εs = 0.42, and DG = 1.5 10-4 m2/s, giving Sh = 1.9 and kg = 0.31 m/s, by applying eq 6. Hence, the maximum rate of external mass transfer per unit mass of carrier was ∼1.7 10-4 mol g-1 s-1, for the extreme case where CS = 0. This is fairly close to the measured maximum rate at the same temperature and particle size and 1.22 vol % O2 in N2, as seen in Figure 3. Figure 9 shows a plot of the logarithm of the initial specific rate of reaction against 1/T for the oxidation of the reduced carrier at 1.3% O2 in N2, for both of the size ranges of reduced oxygen carrier used. It should be noted that the maximum rates of these reactions were reached at X below 0.1; therefore, they were reasonable estimates of the initial rates. The rate of oxidation of the larger carrier particles exhibits a weak dependence upon the temperature, suggesting that the reaction was controlled mainly by external mass transfer. The rate of reaction of the smaller particles, on the other hand, more strongly depends upon the temperature and is therefore less likely to be controlled solely by external mass transfer. The drop in rate at 900 C (i.e., 103/T = 0.85 K-1) for the small particles can be attributed to the increasing tendency to form Cu2O rather than CuO as the temperature increases beyond 900 C, with the simple kinetic expression involving simple first-order dependence upon [O2], with the single rate constant unlikely to be applicable. However, large particles did not manifest this fall in rate beyond 900 C, essentially because any changes in intrinsic chemical kinetics were masked by the effects of external mass transfer. This provides additional evidence that the measurements from the smaller particles are more representative of the intrinsic kinetics. Hence, kinetic information was only obtained from experiments on the smaller particles (355-500 μm) and at temperatures e750 C. 3.5. Effect of Cycling. The rates of oxidation in the first and seventh cycles were compared at 300 and 700 C, as shown in Figure 10. The reduction of the carrier was carried out with 2.5 vol % CO in N2, and the oxidation of the carrier was carried out with 2.5 vol % O2 in N2 at 700 C and 15.0 vol % O2 in N2 at 300 C. The oxidation and reduction periods were
Figure 8. Plots of the specific rate of consumption of O2 by the carrier containing Cu (355-500 μm) at 7.3 vol % O2 in N2 at 400 and 650 C against (a) time and (b) conversion. The plot of the rate against conversion for the consumption of O2 by the carrier containing Cu (355-500 μm) at 1.2 vol % O2 in N2 at 400 C ([O2] = 0.22 mol/m3) is also shown in panel b.
temperatures. If [O2] was further lowered at 400 C, a more realistic estimate of the initial rate could be obtained. This is illustrated by the rate of consumption of O2 at 1.2 vol % O2 in N2 ([O2] = 0.22 mol/m3), shown by the graph marked with crosses in Figure 8b. The overall conversion was the same as the case at the higher [O2] of 7.3 vol % O2 in N2 (0.96 mol/m3), but the X at which the maximum was reached was lower at ∼0.05. This confirms that the maximum rates at lower [O2] are better estimates of the initial rate. Therefore, to eliminate the inaccuracy caused by the slow response of the analyzer, only measurements at [O2] < 0.5 mol/m3 were used here to study intrinsic kinetics; in these cases, the maximum rates occur at low values of X. The proportionality between the rate and [O2] suggests that the order of the reaction is unity; alternatively, it might imply that the reaction is controlled by external mass transfer, for which n = 1. 3.4. Effect of External Mass Transfer. To establish the importance of external mass transfer, the theoretical rate of external mass transfer was calculated. The rate of mass transfer of O2 to a single particle of reduced oxygen carrier is kgS(CB - CS), where S is the external surface area of the particle, CB is the bulk concentration of O2, and CS is the concentration of O2 at the external surface of the particle. The external mass-transfer coefficient, kg, can be estimated from Sh = kgdp/DG, where DG is the molecular diffusivity of O2 in N2. The Sherwood number, Sh, was computed from La Nauze and Jung9 Sh ¼ 2εmf þ 0:69ðRep =εmf Þ1=2 Sc1=3
ð6Þ
(9) La Nauze, R. D.; Jung, K. Chem. Eng. Res. Des. 1985, 63, 3–33.
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Figure 11. Plots of kv against 1/T for carriers containing Cu and Cu2O (355-500 μm). The range of kv calculated is shown by the vertical bars.
only have given faster rates of reaction with short times of the reaction. Consequently, the measurements would then have been obscured to a larger extent by mixing and sampling effects, such as attaining the maximum rate at a high value of X. Hence, the measurement of the intrinsic kinetics involved a compromise; efforts were made to obtain realistic kinetic information from the carrier sized at 355-500 μm, using measurements between 300 and 750 C. Because the observed order of reaction was unity, the corresponding overall rate constant, koverall (m3 g-1 s-1), defined by the rate expression r = koverallC for unit mass of carrier particles, can be expressed as ð7Þ 1=koverall ¼ ð1=kg Sm þ 1=km ηÞ
Figure 10. Rates of oxidation of fully reduced carrier (850-1000 μm) by (a) 15 vol % O2 in N2 at 300 C and (b) 2.5 vol % O2 in N2 at 800 C.
8 min each. To make these measurements, the carrier was reacted for six cycles and, at the end of the seventh reduction, the bed was cooled to room temperature with pure N2 flowing through it at the same time. The content of the bed was then emptied into a crucible, and the reduced carrier was recovered by sieving. The size of the sand was 355-425 μm, and the oxygen carrier used was 850-1000 μm. This allowed for easy separation from the sand. These recovered particles were then reacted in a new bed of fresh sand at identical conditions in the oxidation stage of the cyclic operation. This procedure allowed for the rate of reaction per unit mass of particles to be established unequivocally, avoiding errors arising from the loss of mass of the carrier by attrition during the cycling. As can be seen in Figure 10, both the rates of reaction in the first and seventh cycles of operation at 300 and 700 C are identical within experimental error; i.e., the kinetics in later cycles of oxidation can be approximated by those in the first. Overall conversions of 0.5). This is because, as soon as the Cu2O was formed from Cu, it should theoretically be, in turn, capable of being oxidized to CuO. The presence of Cu2O in the partially oxidized sample in Figure 12c suggests that there might be a delay in the formation of CuO. There are two possibilities that could account for the delay in the formation of CuO. These are now investigated in detail below. 4.2.1. Equilibrium Concentration of O2 Necessary for the Oxidation of Cu2O to CuO. At 700 C, O2 molecules in the bulk gas initially diffuse through a boundary layer around the outside of a particle and react with the Cu on the external surface of the reduced carrier to form Cu2O. For Cu2O to be oxidized to CuO, the equilibrium partial pressure of O2 needs to be at least ∼1 10-4 bar at 700 C, using a value of Kp calculated from the NASA Glenn thermodynamic data,15 approximately equivalent to a mole fraction of 1 10-4 at atmospheric pressure. Oxidation of Cu2O to CuO, reaction 2, will not proceed until the mole fraction of O2 on the external surface of the reduced oxygen carrier, yext, exceeds 1 10-4. To estimate yext, a model based on the shrinking core mechanism16 was used CB - Ceqm # N ¼ " ð10Þ 1 1 1 1 1 þ þ 4πDeff r1 r0 4πr1 2 ks 4πr0 2 kg
Figure 13. Comparison of yext and yO2,eqm at 700 C and yO2 = 0.013 at different extents of reaction (3.16), for r0 = 4.63 10-4 m.
(m/s) of the oxidation of Cu to Cu2O, i.e., reaction 1. The relation between kv and ks is given by 4 4πr1 2 ks ¼ πr1 3 kv 3 \ks ¼
r 1 kv 3
ð11Þ
The first term in the denominator on the right-hand side of eq 10 represents the resistance from diffusion through the product layer of Cu2O. The second term in the denominator on the right-hand side of eq 10 represents the resistance from the chemical reaction between O2 and Cu. The third term in the denominator on the right-hand side of eq 10 represents the resistance from the external mass transfer of O2 from the bulk to the external surface of the particle. The rate of external mass transfer is given by N = 4πr02kg(CB - Cext), where Cext is the concentration of O2 at the external surface of the reduced carrier. Rearranging N ð12Þ Cext ¼ CB 4πr0 2 kg Using the value of kv,2 determined from section 3.6 and the operating conditions used for the experiment to obtain the partially oxidized sample in Figure 12c (T = 700 C, yB = 0.013, and dp = 850-1000 μm), the variation of yext with r1 is shown in Figure 13. A simple heat balance showed that there was a rise in the temperature of the particle of ∼2 C. Therefore, it is reasonable to assume that the oxidation was isothermal. The equilibrium concentration of O2, yO2,eqm, necessary for reaction 2 to proceed has also been included in Figure 13. From the figure, it is clear that, at any extent of reaction 1, yext is sufficiently high for reaction 2 to proceed. Therefore, the delay in the formation of CuO was unlikely to be caused by yext. To ensure that the delay was also not caused by the rate of reaction 1 being slow, the plots of X against time were plotted for both reactions 1 and 2 occurring separately. To do so, the following was performed. The rate of consumption of Cu, -rCu, can be obtained by assuming that the reduced carrier particle was spherical, considering just reaction 1 d 4 dr1 3 πr1 FCu ¼ - 4πFCu r1 2 ð13Þ - rCu ¼ dt 3 dt
In this expression, N is the rate of consumption of O2 (mol/s), Ceqm is the equilibrium concentration of O2, r0 is the initial radius of a particle of reduced carrier, r1 is the radius of the unreacted shrinking core of Cu, and ks is the rate constant
Here, FCu is the number of moles of Cu per unit volume of the carrier particle. It was calculated by dividing the product of the mass fraction of Cu (∼0.79) and the bulk density of the reduced carrier by the relative molecular mass of Cu, giving FCu = 2.21 104 mol/m3. According to the stoichiometric
(15) McBride, B. J.; Zehe, M. J.; Gordon, S. NASA Glenn coefficients for calculating thermodynamic properties of individual species. National Aeronautics and Space Administration (NASA), Washington, D.C., 2002; Report TP-2002-21155. (16) Levenspiel, O. Chemical Reaction Engineering, 2nd ed.; John Wiley and Sons International: New York, 1972.
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coefficients of reaction 1, eq 13 was divided by 4 to obtain the rate of consumption of O2, -rO2, which was then equated to eq 10. - rCu dr1 ¼ - rO2 ¼ N ¼ - πFCu r1 2 ð14Þ 4 dt Integrating eq 14, the time, tCu, for the radius of the unreacted core to shrink to r1 is given by - FCu 1 tCu ¼ ðr1 2 - r0 2 Þ 4ðCB - Ceqm, Cu Þ 2Deff 1 r1 - r0 r1 3 - r0 3 3 3 ð15Þ ðr1 - r0 Þ þ þ 3r0 Deff ks 3r0 2 kg
Figure 14. Plots of X against time for the theoretical oxidation of Cu to Cu2O, Cu2O to CuO, and the actual oxidation of Cu to CuO at 700 C, yO2 = 0.013, and dp = 850-1000 μm.
where r0 is the initial radius of the particle and Ceqm,Cu is the equilibrium concentration of O2 needed for reaction 1 to proceed. Assuming that the overall size of the particle stayed constant, the conversion of Cu to Cu2O is related to the radius of the unreacted shrinking core of Cu by 3 r1 ð16Þ XCu ¼ 1 r0
4.2.2. Diffusion of Cu Atoms through Cu2O and CuO. Zhu et al.17 provided a possible mechanism for such a process, which might explain the failure to see CuO in the partially reacted particles in the current work. According to them, for the oxidation of copper rods, direct oxidation of Cu to CuO initially takes place at the interface between Cu and air. A thin layer of CuO is thus formed. Then, at the interface between Cu and CuO, the following reaction occurs:
The same procedure can be applied to the oxidation of Cu2O to CuO, assuming that we start with a particle completely converted to Cu2O, and the corresponding time and conversion are given by - FCu2 O 1 ðr2 2 - r0 2 Þ tCu2 O ¼ 2ðCB - Ceqm, Cu2 O Þ 2Deff 1 r2 - r0 r2 3 - r0 3 ð17Þ ðr2 3 - r0 3 Þ þ þ 3r0 Deff ks 3r0 2 kg
Cu þ CuO f Cu2 O
Further oxidation takes place via the diffusion of Cu atoms through the Cu2O and CuO layers. The buildup of the layer of Cu2O will be faster than that of CuO, which will remain thin. This is because the diffusion coefficient18 of Cu in Cu2O is of the order of 10-13 m2/s, while the diffusion coefficient19 of Cu in CuO is of the order of 10-19 m2/s, so that it is harder for Cu to diffuse through the CuO layer. In the current work, when the oxidation was stopped at X ∼ 0.3, as in Figure 12c, the partially oxidized sample should theoretically have a black outer layer surrounding the dark red unreacted Cu metal core. The black outer layer would have corresponded to the grains of unreacted Cu and Cu2O covered by a thin layer of CuO. During the period when the partially oxidized sample was cooled to room temperature in pure N2, no O2 was available for the direct oxidation of Cu to CuO. Therefore, it was likely that, at the granular level, the Cu atoms diffused through the Cu2O layer and reacted with the thin layer of CuO to produce Cu2O. By the time the bed had been cooled to room temperature, all CuO would have been converted to Cu2O, giving a brown outer layer of Cu2O and an inner core of dark red Cu, as seen in Figure 12c. Had the sample not been removed from the bed at X ∼ 0.3 and oxidation had been continued, there would have been a point at which all Cu had become exhausted. Then, the Cu atoms from Cu2O would have diffused outward through the thin CuO layer and reacted with O2 at the exterior to form CuO, thereby leading to the progressive buildup of a layer of CuO, as seen in Figure 12d. Therefore, this possibility is more likely to be the explanation for the failure to see CuO at X = 0.3. It should be noted here that the overall rate of oxidation of the reduced carrier cannot be controlled by the diffusion of Cu atoms through layers of product (Cu2O or CuO). If it were,
and XCu2 O ¼ 1 -
3 r2 r0
ð19Þ
ð18Þ
where r2 is the radius of the unreacted shrinking core of Cu2O, FCu2O is the number of moles of Cu2O per unit volume of carrier particle (calculated to be 1.08 104 mol/m3), and Ceqm,Cu2O is the equilibrium concentration of O2 needed for reaction 2 to proceed. The plots of XCu and XCu2O against time, for dp = 850-1000 μm, T = 700 C, and yO2 = 0.013, are shown in Figure 14. The critical point to make here is that reactions 1 and 2 are being considered independently of each other; i.e., the flux of O2 to r1 is not affected by the consumption of O2 at r2. The plot of experimental X against time has also been included. At these conditions, Ceqm,Cu ∼ 6.24 10-14 mol/m3, Ceqm,Cu2O ∼ 1.25 10-2 mol/m3, DK ∼ 4.4 10-5 m2/s, DG ∼ 1.6 10-4 m2/s, and Deff ∼ 2.0 10-5 m2/s (calculated using Deff = ε2/(1/DK þ 1/DG)). In Figure 12c, a significant proportion of Cu2O was detected, without any CuO having apparently been formed at that stage. In Figure 12c, the experimental X is ∼0.3, corresponding to XCu = 0.6, if only reaction 1 has, indeed, occurred; experimentally, reaction 2 has not occurred, so that the experimental XCu2O = 0. However, from the theoretical plots in Figure 14, in the time taken for XCu to fall to 0.6, XCu2O would, in fact, have been as much as ∼0.55. Substituting this value of 0.55 into eq 18, the ratio of r2/r0 would be ∼0.76, suggesting that, at XCu = 0.6, ∼24% of the radius of the particle should have been CuO, which it was not, as seen in Figure 12c. On the basis of these approximate calculations, it is concluded that equilibrium considerations cannot account for the delay in the formation of CuO.
(17) Zhu, Y.; Kimura, K.; Isshiki, M. Corros. Sci. 2004, 46, 2445– 2454. (18) Grzesik, Z.; Migdalska, M.; Mrowec, S. J. Phys. Chem. Solids 2008, 69, 928–933. (19) Rebane, J. A.; Nikolay, V. Y.; Chicherin, D. S.; Tretyakov, Y. D.; Leonyuk, L. I.; Yakunin, V. G. J. Mater. Chem. 1997, 7, 2085– 2089.
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the rate of reaction would be independent of the gaseous [O2]; in practice, the rate is strongly dependent upon the gaseous concentration of O2.
approximated with those in the first, at both low and high temperatures. (4) At low temperatures (∼300 C), the reduced oxygen carrier appeared to be oxidized via both the direct and indirect pathways, with the indirect pathway being more dominant. The reaction took place evenly throughout the particles of the reduced oxygen carrier. (5) At high temperatures (∼700 C), the reduced oxygen carrier appeared to be oxidized entirely via the indirect pathway. Diffusion of gases through the pores of the particles was slow with respect to the chemical reaction. The reaction proceeded via a shrinking core mechanism. A likely explanation has been advanced to account for the failure to see CuO in the partially reacted particles based on the different mobilities of Cu atoms in Cu2O and CuO.
5. Conclusions The major conclusions to be drawn from the work are as follows: (1) The reduced oxygen carrier was incompletely oxidized to CuO, Cu2O, and CuAl2O4 at T < 600 C. At 600 C and above, the oxidation was complete, but when the temperature was increased to 900 C and beyond, the oxidized carrier decomposed to Cu2O and O2. (2) The order of the intrinsic chemical reaction with respect to [O2] was ∼1 from 300 to 500 C, at low [O2]. However, the reaction was controlled to a considerable extent by external mass transfer at temperatures in excess of 600 C, giving an “apparent” order of unity. (3) The rates of reaction in subsequent cycles could be
Acknowledgment. The authors are grateful to the Engineering and Physical Sciences and Research Council (EPSRC) for financial support.
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