In their function as converters, fuel cells produce electricity with the help of chemical reaction energy. To do this, the fuel cell requires a combination of a continuously supplied fuel — hydrogen or methanol — and oxygen.
Oxygen acts as an oxidant 1 in the fuel cell. But more on this later. The definition of an oxidizing agent is comparatively simple. An oxidizing agent is a chemical substance that is capable of oxidizing other substances, but is itself reduced during this process. Oxidizing agents can accept electrons, which are negatively charged elementary particles, while reducing agents release them.
One reactant transfers electrons to the other. Typically, redox reactions take the form of combustion for conversion to kinetic energy. For example, in cars, trucks, ships and airplanes when gasoline, diesel and kerosene are burned in the engine. Rockets and sparklers also work thanks to oxidizing agents and reducing agents. So does the previously lit lighter or match. Basically, the fuel cell uses the same principle. Just like a battery, it also has electrodes: an anode and a cathode.
Each of the electrodes is coated with a metal as a catalyst, for example nickel or platinum. The anode and cathode in the fuel cell are separated by an electrolyte. This gas-impermeable membrane ensures separate conduction of ions positive charge and electrons negative charge.
The continuously supplied fuel, for example hydrogen, flows around the anode. At the same time, oxygen reaches the cathode in its function as an oxidizing agent. The hydrogen oxidizes by releasing electrons to form protons.
The electrons subsequently flow from the anode to the cathode — in other words, current flows. The protons penetrate the electrolyte separating the anode and cathode and reach the cathode. In addition to the energy generated, the addition of oxygen as an oxidant in the fuel cell produces water and heat as a by-product.
In this way, the fuel cell produces environmentally friendly electricity overall. The reaction heat generated during the process can even be used for heating. One example is the combustion of octane, the principle component of gasoline:. Plants and animals use the oxygen from the air in respiration. The administration of oxygen-enriched air is an important medical practice when a patient is receiving an inadequate supply of oxygen because of shock, pneumonia, or some other illness.
The chemical industry employs oxygen for oxidizing many substances. A significant amount of oxygen produced commercially is important in the removal of carbon from iron during steel production. Large quantities of pure oxygen are also necessary in metal fabrication and in the cutting and welding of metals with oxyhydrogen and oxyacetylene torches. Liquid oxygen is important to the space industry. It is an oxidizing agent in rocket engines. It is also the source of gaseous oxygen for life support in space.
As we know, oxygen is very important to life. The energy required for the maintenance of normal body functions in human beings and in other organisms comes from the slow oxidation of chemical compounds. Oxygen is the final oxidizing agent in these reactions.
In humans, oxygen passes from the lungs into the blood, where it combines with hemoglobin, producing oxyhemoglobin. In this form, blood transports the oxygen to tissues, where it is transferred to the tissues. The ultimate products are carbon dioxide and water. The blood carries the carbon dioxide through the veins to the lungs, where the blood releases the carbon dioxide and collects another supply of oxygen.
Digestion and assimilation of food regenerate the materials consumed by oxidation in the body; the energy liberated is the same as if the food burned outside the body. Oxygen reacts directly at room temperature or at elevated temperatures with all other elements except the noble gases, the halogens, and few second- and third-row transition metals of low reactivity those with higher reduction potentials than copper.
Rust is an example of the reaction of oxygen with iron. The more active metals form peroxides or superoxides. Less active metals and the nonmetals give oxides. Two examples of these reactions are:. The oxides of halogens, at least one of the noble gases, and metals with higher reduction potentials than copper do not form by the direct action of the elements with oxygen.
Elemental oxygen also reacts with some compounds. If it is possible to oxidize any of the elements in a given compound, further oxidation by oxygen can occur. Because the sulfur does not exhibit its maximum oxidation state, we would expect H 2 S to react with oxygen. It does, yielding water and sulfur dioxide. The reaction is:. It is also possible to oxidize oxides such as CO and P 4 O 6 that contain an element with a lower oxidation state.
The ease with which elemental oxygen picks up electrons is mirrored by the difficulty of removing electrons from oxygen in most oxides. Of the elements, only the very reactive fluorine can oxidize oxides to form oxygen gas.
Most nonmetals react with oxygen to form nonmetal oxides. Depending on the available oxidation states for the element, a variety of oxides might form. The two common oxides of sulfur are sulfur dioxide, SO 2 , and sulfur trioxide, SO 3. The odor of burning sulfur comes from sulfur dioxide. The formation of ozone from oxygen is an endothermic reaction, in which the energy comes from an electrical discharge, heat, or ultraviolet light:.
The sharp odor associated with sparking electrical equipment is due, in part, to ozone. Ozone forms naturally in the upper atmosphere by the action of ultraviolet light from the sun on the oxygen there. Molecular oxygen O 2 is an oxidizing agent, but is a vital component of the air we breathe, not a dangerous toxin. The formula for ozone O 3 looks very similar to that of molecular oxygen. Let's examine why O 3 is a much stronger oxidizer than O 2. The Lewis structures in Figure 1 indicate that the ozone molecule has two equivalent resonance structures, which means the electrons are delocalized.
From the Lewis structure, we see that the bond order for O 2 is 2 a double bond , whereas the bond order for O 3 is 1.
Recall that a smaller bond order means a weaker longer bond. When the bond order is lower, electrons are held less tightly. Comparing the Lewis structures of molecular oxygen Figure 1a and ozone Figure 1b indicates that ozone has delocalized electrons.
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