In the Laboratory
Laboratory Application and Demonstration of Automotive Oxygen Sensors T. Schober* and J. Friedrich Institut für Festkörperforschung, Forschungszentrum Jülich, 52425 Jülich, Germany; *
[email protected] Control and determination of oxygen in gas streams is important in modern technology, including automotive technology and semiconductor manufacture. As a consequence, chemistry and chemical engineering students need knowledge and experience with oxygen sensors and control systems. To achieve this, one can utilize automotive exhaust gas sensors. These are inexpensive (at about $150 U.S.) and readily available. They may be found in almost every modern car, measuring the oxygen partial pressure in the exhaust line. Thus, any auto new parts supplier stocks a variety of these sensors. They are usually manufactured from Y stabilized zirconia (YSZ) electrolytes. The partial pressure of oxygen gas, p′O2, is determined by measuring a difference in the oxygen activity across this YSZ electrolyte. To do this, the electrolyte needs to be at an elevated temperature, usually between 600 and 700 °C. An internal rodlike heater brings the closed-end zirconia tube to the desired working temperature, which is in the neighborhood of 630 °C. The zirconia electrolyte has Pt coatings on both sides, which serve as cathode and anode for the electrochemical cell and for electrical contact. These commercial sensors are referred to as lambda sensors. Taking into account the cathode reaction 1 ⁄2 O2 + 2e{ O2{ and the anode reaction O2{ 2e{ + 1⁄2 O2 leads to the basic Nernst equation, which applies to these sensors, and one which the student can confirm in some of the examples: p′O U = RT ln 2 (1) 4F p′′O 2
Given the high temperature of the Pt-coated zirconia tube, we may tacitly assume that the equilibrium of the following reaction describing the formation and decomposition of water is maintained: H2(g) + 1⁄2 O2(g)
H2O(g)
(2a)
The equilibrium constant KH2O for eq 2a is given by KH O = 2
p H O/p Ho O 2
p H /p Ho 2 2
2
p O /p Oo 2
1/ 2
= exp ∆S°/R exp {∆H°/RT (2b)
2
where ∆S ° and ∆H ° are the standard entropy and enthalpy changes of reaction 2a. The reference pressures are 1 bar. Numerical values are ∆S ° = {55.22 J mol{1 K{1 and ∆H ° = {247.8 kJ mol{1. Experimental Details A schematic outline of the experimental setup is depicted in Figure 1. A central water-cooled SS316 vessel contains two tapped holes into which oxygen sensors #1 and #2 of type Bosch 0258104004 are inserted. Sealing of the sensors is achieved by the use the commercial sealing rings provided with them and Teflon tape around the threads. To allow unlimited access of the gas to the active sensor head, the cap of the sensor shield is cut off. This exposes the whole diameter of the Zr oxide thimble to the gas stream. The heating leads are connected to a suitable 12-V dc power supply or to a car battery. The approximate operating temperature of the sensor head may be determined by the following technique. In a
where R the gas constant, T the absolute temperature, F Faraday’s constant, p′O2 the oxygen pressure in the gas stream and p′′O2 a constant reference pressure (e.g., 0.21 bar in air.) Commercial lambda sensors have a variety of possible applications in the laboratory and may also be used for purposes of demonstration. The following possibilities are discussed here: Measurement of oxygen impurities in gas streams Operation of a hydrogen–oxygen fuel cell Measurement of water vapor pressure in Ar–H2 mixtures Removal of trace oxygen from gas streams Adding controlled amounts of oxygen to gas streams Generation of low pO2 (≈10{22 bar) Ar–H2O–H2 gas mixtures by electrochemical decomposition of water vapor Generation of low pO2 (≈10{19 bar) in CO2–CO gas mixtures by electrochemical decomposition of CO2
These demonstrations are relatively simple and inexpensive to perform, and most undergraduate laboratories are already equipped with the necessary equipment. Along with the Nernst equation, the student needs some other basic theoretical information about the chemistry involved.
Figure 1. Experimental setup with 2 automotive oxygen sensors in a test vessel. A portion of the shielding cap of the sensors was removed. The gas stream is controlled by a flow controller. The humidity is controlled by thermostated adjustment of the temperature.
JChemEd.chem.wisc.edu • Vol. 76 No. 12 December 1999 • Journal of Chemical Education
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dismantled sensor, a thermocouple is inserted together with the heater into the Zr oxide thimble. Upon heating, a saturation temperature of ≈900 K is established in the thimble. This temperature was found to vary by about ± 15 K from one sensor to the next. For our calculations here we use T ≈ 900 K. The upstream sensor #1 may either be connected to a potentiostat or to a high-impedance volt meter. The downstream sensor #2 is also connected to a high-impedance volt meter. The gas flow (Ar, nonexplosive Ar–4% H2 or CO2) goes from the steel bottles to a flow controller (1–10 mL/ min) and then to the sensor vessel. Alternatively, it is bubbled through a humidifying flask where the water vapor pressure of the gas is regulated in the range 20 to 100 mbar by adjustment of the water temperature. The downstream outlet of the sensor vessel is connected with a hose to a bubbletype flow controller to avoid back penetration of oxygen into the vessel. The reference side of the oxygen sensors is always open to air (pO2 ≈ 0.21 bar).
when a load is applied to the output. It is seen that for I = 0, essentially the open cell voltage (OCV) is obtained. From the water equilibrium we obtain p′O2 ≈ 3.1 × 10{24 bar while p′′O2 = 0.21 bar . Using eq 1, these values lead to an OCV of 1.018 V. Increasing the load leads to a decrease of the voltage. The initial steep decrease probably is due to an activation overpotential where the kinetics of the electrode reactions are determined by charge transfer. Let us assume that the decrease in Figure 2 is linear starting with the OCV Urev. Then the output voltage U is given by U = Urev – R tot I
(3)
where R tot is the total polarization loss, which is essentially ohmic. Here, R tot ≈ 40 Ω. For the power Pw we obtain Pw = IUrev – I 2R tot. For our data, this almost quadratic function is depicted in Figure 3. Differentiating the last expression with respect to I leads to the maximum power, Pmax: Pmax = Urev2 /4R tot
(4)
Results
which is roughly 5.5 mW in this case.
Measurement of Oxygen Impurity Level in Flowing Ar If Ar gas of 99.999% purity is used with a flow rate of ≈5 mL/min, a Nernstian voltage of about 100 mV is observed using both sensors after flushing for about 1 h. Using eq 1, this voltage translates to a p′O2 of about 1.15 mbar. This elevated value does not reflect the purity of the bottled gas but is a consequence of outgassing from the walls of the system. Flushing for many hours may eventually raise this voltage to a level of ≈150 mV. Operation of a Hydrogen–Oxygen Fuel Cell A fully operational O2–H2 fuel cell is obtained when Ar–4% H2 is first suitably humidified (e.g., pH2O = 24 mbar) and then admitted to the test chamber. Through the combined use of hydrogen and water vapor a very low oxygen partial pressure is established at the sensor head in the test cell, as given by the water equilibrium (eq 2b). In effect, an oxygen concentration cell is established under these circumstances. Figure 2 shows the voltage output as a function of the current I
Measurement of Water Vapor Pressure in Ar–H2 Mixtures Oxygen sensors may also be conveniently used to measure water vapor pressures in certain gases. Consider the case of water vapor in Ar–4% H2 gas. Owing to the water equilibrium in eq 2b, such a mixture results at the hot sensor surface in a partial pressure of oxygen that then can be determined by measuring the Nernstian potential of the sensor. Here, a constant reference pressure of pO2 = 0.21 bar is assumed. In this experiment, Ar–4% H2 gas was humidified at known temperatures T resulting in fixed values for the water vapor pressures. This mixture was then admitted to the sensor chamber and the resulting voltages at sensors #1 and #2 were recorded. Figure 4 shows the results obtained from sensor #1 (n) and sensor #2 (,) and the theoretically expected curve based on eq 2b. Good agreement between the three curves is noted. A small offset of a few mV is noted between the theoretical and experimental curves. Such an offset seems to be normal for automotive oxygen sensors and has been discussed elsewhere
Figure 2. EMF of the test cell when operated as a fuel cell.
Figure 3. Power output of the fuel cell as a function of current I.
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Journal of Chemical Education • Vol. 76 No. 12 December 1999 • JChemEd.chem.wisc.edu
In the Laboratory
(2). Obviously, if Ar–4% H2 with an unknown water vapor pressure is introduced into the sensor chamber, then the observed potential at the sensors may be used to calculate the water vapor pressure.
Removing Trace Oxygen from Gas Streams (Electrochemical Pumping) If a potential is applied to sensor #1 to cause O2{ ions to be transported to the air reference side, oxygen may be removed from the gas stream. The number of moles of O 2 removed by the current I per unit time is (Faraday’s law): nO2 = I/4F
(5)
In a typical experiment a flow of 2 mL of Ar per minute was used. Using a potentiostat, a voltage of 1 V was applied to sensor #1 so that the outside electrode was positive. Typical currents were 0.2 mA. Sensor #2 was used to measure the pO2 of the purified gas using eq 1. Prior to pumping, sensor #2 would display a voltage of about 100 mV. After about 30 minutes a voltage of 0.945 V was obtained at sensor #2, signifying that pO2 ≈ 1.3 × 10{22 bar. It is seen that this method is highly effective in removing trace oxygen from inert gas streams.
Injecting O2 into Gas Streams Reversing the polarity of sensor #1 of the above experiment allows injecting oxygen into a gas stream with a rate given by I/4F. Sensor #2 is again used to measure the oxygen content. The ideal gas law allows calculation of pO2 at sensor head #2: pO2 = nO2RT /V
(6)
where V is the volume of gas flowing through the system per unit time. In a typical experiment, Ar would be flushed through the system at a rate of 5 mL/min until sensor #2 showed a reading of 130 mV, corresponding to a pO2 of 2.4 × 10{4 bar. Within a few minutes after the pumping current of 1 mA was switched on at sensor #1, the voltage at sensor #2 would
decrease to 98 mV. Calculating the theoretical sensor voltage using eqs 1 and 6 gives a value of 94 mV. The agreement between experiment and theory may be termed fair, in view of the possible experimental errors. This experiment shows that within certain experimental limits, electrochemical oxygen doping is also possible using the oxygen sensors.
Generating Low-pO2 (≈10{20 bar) Ar–H2O–H2 Gas Mixtures by Electrochemical Decomposition of Water Vapor Assume for the following that humidified Ar gas flows through the system. If a potential below the dissociation level of the ZrO2 electrolyte or water molecules is applied (i.e., U < 1 V) to sensor #1 (driving O2{ to the reference side) oxygen is constantly removed at a high rate. The mass action law now predicts that water molecules are dissociated, providing further oxygen and also hydrogen. If the flow rate is not too high, then the gas stream will eventually have a very low pO2 (≈10{20 bar) and also contain hydrogen. If voltages in excess of ~1 V are applied to the present sensors then 2 phenomena may occur: 1. Blackening of the zirconia electrolyte may occur provided U > ∆G/nF (neglecting polarization) where ∆G is the Gibbs free energy of formation of the zirconia compound and n is the number of electrons involved in the reaction (3). This blackening is a consequence of the oxygen deficiency at the surface, which locally leads to a high electron concentration and the formation of F-centers via the reaction (7) OOx → VO?? + 2e′ x and V ?? are regular oxygen ions in lattice powhere OO O sitions and oxygen vacancies, respectively. To a first approximation, this effect does not lead to new surface phases or to an impairment of the sensor, provided that the sensor is not operated too long in this mode and that it may recover in oxygen afterwards. 2. Molecular species may be dissociated provided U > ∆G ′/ nF, where ∆G ′ is the Gibbs free energy of formation of the molecule. For instance, water molecules require U > 998 mV at 1000 K, since ∆G = {192.6 kJ (4, 5). This process corresponds to steam electrolysis; for CO2, voltages in excess of about 1.02 V are required at 1000 K to achieve the reaction.
The experiment was carried out in the following way. A 2-mL/min stream of Ar was humidified at 35 °C, resulting in a pH2O of 56 mbar. The measured voltage at sensor #2 was 75 mV, corresponding to a high pO2 of 4.2 mbar. Applying a potential of 0.9 V below pertinent dissociation levels at sensor #1, with the polarity such that the outer reference electrode is positive, leads within 30 min to a Nernst potential of 830 mV at sensor #2. This value translates via eq 1 to a pO2 = 5.1 × 10{20 bar. This new gas composition is now highly reducing, since it contains hydrogen at a very low pO2.
Generating Low-pO2 (≈10{19 bar) CO2–CO Gas Mixtures by Electrochemical Decomposition of CO2 CO2–CO mixtures are commonly used in the laboratory when highly reducing conditions are called for (6 ). Consider the equilibrium constant of the reaction CO(g) + 1⁄2O2(g) CO2(g): Figure 4. EMF of the test cell when an Ar–4% H2 mixture with known humidity is used. Reference side: air, pO2 = 0.21 bar. Two experimental curves and one theoretical curve are depicted. Note the small offset between theoretical and experimental values.
K=
o pCO /pCO 2
2
o pCO /poCO pO /pO 2
1/2
= exp ∆S°/k exp {∆H°/kT (8)
2
JChemEd.chem.wisc.edu • Vol. 76 No. 12 December 1999 • Journal of Chemical Education
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where the symbols have their usual meaning. Solving for pO2 yields
pO
2
o pO 2
=
pCO
2 2
pCO
exp
{2∆S ° 2∆H ° exp R RT
(9)
Commonly used numerical values are ∆S° = {88.09 J mol{1 K{1 and ∆H ° = {283.3 kJ mol{1. As a numerical example, we assume that pCO2/pCO ≅ 103. Then the partial pressure of oxygen will be around 10{12 bar at 1100 K. In this experiment, CO2 gas alone is admitted to the sensor vessel at a rate of 3 mL/min. After a few hours of flushing, a purity analysis using sensor #1 or #2 yields a voltage of ~90 mV, corresponding to a pO2 of 1.9 mbar. As before, we used sensor #1 to electrochemically pump oxygen to the outside, thereby shifting the chemical equilibrium to the very low pO2 domain and effectively producing a CO–CO 2 mixture. Using an applied potential of 1.20 V (just slightly above the dissociation level of CO2) at sensor #1 resulted in a voltage of 0.753 V at sensor #2, corresponding to a pO2 of 2.7 × 10{18 bar. Increasing the potential at #1 to 1.7 V and reducing the flow rate to 2 mL/min yields a voltage of 0.8 V at sensor #2 (pO2 = 2.4 × 10{19 bar). These experiments show that strongly reducing gas mixtures can also be achieved by using CO2 alone plus electrochemical oxygen pumping. Finally, there is a further application for the present setup which was not tested in the limited scope of this work. Ex-
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periments on the effect of non-faradaic electrochemical modification of the catalytic activity (NEMCA, [7 ]) could easily be performed and demonstrated using the configuration described. The NEMCA effect is due to a spillover of ionic species (here O2{) from the solid electrolyte onto the gas-exposed catalyst electrode surface. Thereby, a double layer is established on the catalyst surface, which changes the work function and thus the chemisorptive bond strengths. Acknowledgments Helpful discussions with H. Wenzl and J. B. Condon are acknowledged. Literature Cited 1. Minh, N. Q.; Takahashi, T. Science and Technology of Ceramic Fuel Cells; Elsevier: Amsterdam, 1995; p 26. 2. Kumar, R. V. Ionics 1997, 3, 161–169. 3. Guo, X.; Yao-Qing, S.; Cui, K. Sensors Actuators 1996, B31, 139–145. 4. Logothetis, E. M.; Visser, J. H.; Soltis, R. E.; Rimai, L. Sensors Actuators 1992, B9, 183–189. 5. Carconi, P. L.; Casadio, S.; Moauro, A.; Petrucci, L.; Mari, C. M. Fusion Technol. 1995, 28, 556–560. 6. Mortimer, A. G.; Reed, G. P. Sensors Actuators 1995, B24–25, 328–335. 7. Pitselis, G. E.; Petrolekas, P. D.; Vayenas, C. G. Ionics 1997, 3, 110–116.
Journal of Chemical Education • Vol. 76 No. 12 December 1999 • JChemEd.chem.wisc.edu
