The Chemistry of Oxygen Sensors
Direct-Reading Technologies for Protecting Workers in Confined Spaces
Like other real-time detection instruments, oxygen sensors play a key role in protecting worker health and safety. Understanding how these instruments work is integral for achieving optimum performance—and for protecting workers in confined spaces against the potentially fatal effects of oxygen deficiency or enrichment.
An oxygen sensor measures the atmospheric oxygen concentration (or, in some cases, the oxygen partial pressure) to warn of oxygen deficiency or enrichment conditions. Both conditions may pose dangers to health or safety. Either excessively low or high O2 concentrations are problematic. Dry air is composed of about 21 percent O2 by volume, but numerous situations may result in O2 deficiency, which OSHA’s confined spaces standard for general industry defines as O2 content < 19.5 percent by volume. The standard requires the testing of confined space atmospheres to (among other things) verify that the atmospheric O2 content is acceptable. The OSHA standard for confined spaces in shipyards stipulates an acceptable upper concentration of 22 percent O2 by volume. Atmospheres with dangerously low O2 may occur as O2 is consumed without a contaminant being added to the atmosphere. Chemical oxidation (that is, rusting) is possible in enclosed spaces where steel or iron and water are present. The ferrous metals will oxidize in the presence of water and O2 until the available O2 is totally consumed, at which point the system becomes stable and rusting ceases. Combustion is a faster form of oxidation, usually accompanied by flame and smoke. In addition to consuming O2, combustion produces many byproducts, including some that may be toxic, so the drop in O2 percent by volume may also be accompanied by an increase in toxicity. Also, some chemicals can absorb or adsorb O2. For example, newly cast concrete will absorb O2 from the air as it cures. In closed systems that contain living organisms (whether aerobic bacteria or people), O2 will be consumed. Some of the many byproducts of metabolism may be toxic, so this also presents a case where a drop in O2 may be accompanied by an increase in toxicity. Carbon dioxide (CO2) is the chief toxic byproduct of human respiration and may also be produced by aerobic microorganisms. Atmospheric dilution with virtually any gas can also lead to low O2 conditions. Figure 1 depicts fluctuating O2 and CO2 concentrations measured in a new-construction manhole (not yet connected to an active sewer system) following the death of a worker there. This figure shows that atmospheric conditions in the confined space appeared to be safe on some days and were deadly on others; the death could have been prevented with proper testing of the atmosphere within the manhole.
Figure 1.
Direct-reading field measurements of O2 and CO2 concentrations over time in a new-construction manhole where a confined space fatality had recently occurred. Field measurements for CO2 were not completed for days 0, 1, and 11. Originally published in “Exposure of Unsuspecting Workers to Deadly Atmospheres in Below-ground Confined Spaces and Investigation of Related Whole-Air Sample Composition Using Adsorption Gas Chromatography,”
Journal of Occupational and Environmental Hygiene
(Dec. 2014). Reprinted with permission.
Atmospheres with dangerously high O2 may also occur, most commonly due to the release of relatively pure O2 gas from certain types of welding or torch-cutting equipment, and from medical use of O2. However, some chemical reactions can also produce oxygen-enriched environments. For example, hydrogen peroxide (H2O2) may decompose into water (H2O) and O2 gas, especially in the presence of a catalyst such as manganese dioxide or transition metals, as shown in this reaction: 2H2O2 ? 2H2O + O2(g) Some organic peroxides and inorganic oxidizers might also decompose to produce O2 gas, as with the heating of potassium chlorate (KClO3). Also, when H2O2 and sodium hypochlorite (NaOCl) are mixed, the two species will react vigorously to produce O2 gas. A low O2 reading obtained with a correctly calibrated and used O2 sensor indicates that either O2 has been consumed from the atmosphere tested or another gas has been added to dilute the O2 normally present. Since the approximate atmospheric nitrogen/O2 volumetric ratio is 4:1, the O2 concentration will fall 1 percent by volume for each 5 percent total dilution. High-toxicity gases such as hydrogen sulfide (H2S) can create an extremely dangerous atmosphere while only minimally diluting O2 content. For example, the addition of sufficient H2S to create a deadly 5,000 parts-per-million atmospheric concentration (0.5 percent H2S by volume) would result in an O2 decrease of only about 0.1 percent by volume. Such a minute decrease would not by itself be a cause for alarm. For this reason, an O2 sensor is almost always used in a multi-gas instrument to allow detection of low or high O2 concentrations as well as other important air contaminants at toxicologically relevant concentrations.
Electrochemical methods are typically used to measure atmospheric O2 concentrations. The most common methods rely on balanced oxidation-reduction reactions in which the total flow of electrical current produced is proportional to O2 concentration. Electrical current may be expressed numerically as amperes; thus, these approaches are amperometric detection methods. As explained in a paper published in the April 2014 issue of the journal
, two fundamental amperometric O2 sensor approaches are possible: one uses a consumable lead anode in a galvanic fuel cell design, and the other involves an oxygen pump electrolytic mechanism.
Lead Anode (Fuel Cell) Oxygen Sensor
In a lead anode sensor, O2 is reduced to hydroxyl ions at a cathode, in the presence of aqueous base: O2 + 2H2O + 4e- ? 4OH- In the half-reaction that accompanies this reduction, a metallic lead anode is oxidized to lead oxide (PbO): 2Pb + 4OH- ?2PbO + 2H2O +4e- The overall balanced reaction may be written as: 2Pb + O2 ? 2PbO In the lead anode O2 sensor, four electrons are generated for each O2 molecule reduced at the cathode. This stoichiometric relationship allows the amount of O2 present at the cathode to be quantified as electrical current proportional to O2 concentration. PbO is created while metallic lead is consumed, so the anode has a finite lifetime determined by how much O2 enters the sensor. The reaction is spontaneous, requiring only the presence of O2 at the cathode to generate current, and is typically catalyzed by platinum in the sensing electrode.
Oxygen Pump O2 Sensor
The “oxygen pump” O2 sensor avoids the use of lead, which is restricted in European markets due to hazardous substance regulations. In this type of sensor, four electrons are needed for each O2 molecule that is reduced to water at a cathode by the following reaction: 4H+ + O2 + 4e- ? 2H2O In the accompanying half-reaction, which occurs at the counter electrode (an anode), water is oxidized to produce O2: 2H2O ? 4H+ + O2 + 4e- The amount of current necessary to reduce the incoming O2 at the sensing electrode is proportional to the O2 concentration in the atmosphere being sampled. The consumption and production of O2, protons, water, and electrons is balanced. If only O2 is considered, the net reaction is: O2 ? O2 For each O2 molecule that enters the sensor and is reduced at the cathode, another O2 molecule is produced at the anode and then expelled. The external current applied to the sensor to accomplish this is proportional to the number of O2 molecules reacted. This mechanism underlies the term “oxygen pump.”
When comparing percent-by-volume to part-per-million concentrations, a 1 percent concentration is equal to 10,000 ppm. An electrochemical sensor designed to detect most toxic gases is typically only challenged with concentrations up to hundreds of ppm, while O2 measurements must be made for hundreds of thousands of ppm. Thus, a means is needed to limit the ingress of O2 into an amperometric sensor. Regardless of whether a sensor is a lead anode or oxygen pump type, two possible approaches may be used to limit O2. The April 2014
paper describes these approaches: one involves the use of a thin capillary, while the other depends on a polymeric membrane to limit O2 entry into a sensor. A capillary restrictor about the diameter of a human hair limits the flow of O2 into a sensor, allowing for true readings of percent by volume with very little effect related to atmospheric pressure. In the membrane permeation-limiting approach, the O2 reading is based on the difference between the partial pressure of O2 in the air and within the sensor. This differential pressure varies according to total atmospheric pressure, even when O2 content in air remains constant at 20.9 percent by volume. While the sensor output actually measures the partial pressure of O2, the detector readout is typically converted into percent by volume. Most O2 sensors produced for the industrial hygiene and safety market employ a capillary-limiting design.
Capillary-Limited O2 Sensor
The capillary-limited O2 sensor is most commonly used in handheld multi-gas direct-reading instruments. This design provides for a true percent-by-volume measurement as long as nitrogen (N2) is the major component of both the atmosphere the instument is being used to analyze and of the instrument’s calibration gas. This requirement is based on the interaction between O2 molecules and those of the gas into which the O2 molecules are mixed as they diffuse through a capillary. Matrix gases with lower molecular weight than N2 will increase the diffusion rate of O2 molecules through a capillary inlet and will lead to inaccurate (high) sensor readings. The presence of high concentrations of matrix gases with a higher molecular weight than N2, such as carbon dioxide (CO2), will lead to decreased O2 readings. Also, high levels of CO2 can lower electrolyte pH, which may reduce the concentration of hydroxyl ions that are needed to support the electrochemistry occurring at the sensor electrodes. Minor fluctuations in temperature and barometric pressure will not change the capillary-limited O2 sensor reading. However, in general, large pressure changes can cause erroneous readings as the sensor pressure slowly equalizes with the atmosphere through the capillary.
Permeation-Limited O2 Sensor
In an O2 sensor where diffusion of atmospheric gases is limited by a membrane, entry of O2 into the sensor is driven by the O2 partial pressure differential across the membrane. This results in a partial pressure reading, which is usually corrected to percent O2 by volume in the detector’s electronics. When a sensor is calibrated with gas containing a given percent O2 by volume at a particular atmospheric pressure, subsequent changes in both atmospheric O2 concentration or in the barometric pressure can change the sensor readings. At lower barometric pressure (that is, at a higher altitude), any gas is less dense, and for a fixed volume of gas mixture at higher pressure, each gas present will have a higher partial pressure than the mixture with identical percent-by-volume composition at a lower total pressure. This is the reason that the partial pressure of O2 in a “normal” atmosphere (20.9 percent O2) is lower when atmospheric pressure is lower, and higher when atmospheric pressure increases. A capillary-limited O2 sensor will always tend to read 20.9 percent O2 in an atmosphere with normal O2 content despite gradual changes in altitude or barometric pressure. As altitude increases (and total pressure decreases), a permeation-limited O2 sensor reading will decrease even with constant exposure to 20.9 percent O2 by volume. A permeation-limited O2 sensor is desirable for monitoring O2 content in gases used in underwater diving, where partial pressure and matrix effects of diluent gases are of concern. For example, helium is often used to replace N2 in breathing air for deep diving where N2 narcosis is possible. A capillary-limited O2 sensor will not be accurate when changing from an N2/O2 mixture to one composed of helium/O2 due to the matrix effect noted previously. A permeation-limited O2 sensor would also be desirable in any other situation where light or heavy gases may be encountered—for example, where liquid helium (lower molecular weight compared to air) or CO2 (greater molecular weight compared to air) are used.
Sensors can fail for a variety of reasons. Once all the available surface area of lead anode O2 sensor is converted to PbO, the sensor will no longer measure O2. Electrolyte leakage due to physical damage can also cause a sensor to fail, as can desiccation resulting from storage in a hot, dry environment. Blockage of a capillary pore or a permeation membrane by liquids or dirt will also cause failure. Freezing an oxygen sensor can cause the sensor to burst, which may occur with sensors stored at extremely low temperatures. The stoichiometric conversion of metallic lead to PbO results in expansion of the anode material as an O2 sensor ages; lead anode sensor designs provide space within the sensor to account for this. An older lead anode that has undergone substantial conversion from metallic lead to PbO will be more susceptible to rupture from frozen electrolyte than a new sensor exposed to similar temperatures. When a sensor is ruptured due to freezing, very high O2 readings are likely as the diffusion-limiting capillary or permeation membrane has been circumvented and O2 is flowing into the sensor unimpeded. The concern in this situation is that the caustic electrolyte could contact the sensor circuit board and other internal circuitry. Also, electrolyte neutralized by exposure to high levels of CO2 can be a source of O2 sensor damage.
The technology used in real-time detection instruments has shown remarkable progress in recent years. Oxygen sensors are an essential tool for protecting workers in confined spaces. Just as knowledge of hazards is vital for industrial hygienists, knowledge of how these instruments work is equally important.
an industrial hygienist and AIHA member, recently retired from the U.S. government. Currently, he is a compliance officer for the State of Hawaii. He can be reached at
Acknowledgment: The author thanks Chris Wrenn and Lee Monteith for their contributions to developing the material discussed in this article.
Code of Federal Regulations:
Title 29, Part 1910.146, Subpart J
; and
Title 29, Part 1915.12, Subpart B
Powell Fabrication and Manufacturing, Inc.: “The Bleach Strength Test—A Chemical Test Method to Determine the Strength of Sodium Hypochlorite” (
: “Oxygen Sensing for Industrial Safety—Evolution and New Approaches” (April 2014).