The ability of ozone to effectively treat wastewater is dependent on the nature of contaminant in question. For example, ozone will readily remove color from a dye solution, but has much more difficulty reducing the biochemical oxygen demand (BOD) of a sucrose stream. These differences in ozone effectiveness are directly due to the chemistry involved in the ozone induced degradation process. Other agents, such as peroxides, ultraviolet (UV) light, and/or high pH are often used in conjunction with ozone to enhance the destruction of pollutants. This communication serves to define how ozone works in oxidizing pollutants and how peroxides, UV, and high pH assist ozone in the oxidation process.
Ozone is a naturally occurring allotrope of oxygen. Its chemical structure is shown in Figure 1.
Note that each resonance is composed of one single bond and one double bond. The single bond is analogous to peroxide bonds, which are rather weak and lead to the formation of free radicals. The double bond is analogous to molecular oxygen (O2), which is strongly bound and rather unreactive.
Figure 1. Ozone Resonance StructuresThe interconversion between the two resonance structures above is so rapid that the observed ozone structure is a blend of the two resonance structures (below). Consequently, the strength of the two oxygen to oxygen bonds can be considered equal, each being 1.5 in order (Figure 2).
Figure 2. “Observed” Ozone StructureTo understand why ozone is an oxidant, we first need to understand the concept of oxidation states. The oxidation state refers to the formal net gain or loss of electron(s) from an atom relative to the number of electrons in its valence shell. For example, in water, the oxidation state of both hydrogen atoms is +1 (because hydrogen has formally yielded its electron to the oxygen atom), and the oxidation state of the oxygen atom is -2 (because oxygen has formally acquired an electron from each hydrogen atom). The oxidation state for oxygen atoms is usually -2. However, in both ozone and molecular oxygen, the oxygen atoms each have an oxidation state of 0. Therefore, ozone and oxygen are both oxidants because they are capable of drawing electrons from a source, decreasing the oxidation state of at least one of their oxygen atoms in the process. In practice, however, ozone is a more powerful oxidant than oxygen partly because ozone can readily react with a substrate on its own whereas oxygen usually requires a catalyst (such as a metal ion) to initiate a reaction.
The electrically promoted reduction of ozone (O3) results in the release of molecular oxygen (O2) and the formation of an oxygen atom with a -2 oxidation state (usually described in the form of water, Figure 3). This straightforward process carries a standard reduction potential (relative to the hydrogen electrode) of 2.07 V. This value is greater than the reduction potentials of most other materials, meaning that the ability of ozone to oxidize most species is thermodynamically allowed.
Figure 3. Reduction of OzoneThe chemically promoted reduction of ozone is not as simple because of the various reaction pathways (mechanisms) that can take place. One of the more simple examples of a chemical process utilizing ozone as an oxidant is the reaction of ozone with ethylene (Figure 4). In this process, ozone reacts with ethylene to form intermediate A that contains two carbon-oxygen linkages. Note that the carbon-carbon double bond no longer exists in A. The carbon atoms have been oxidized, and the oxygen atoms bound to carbon have been reduced from 0 to -1 oxidation state. This reaction will proceed further to yield O2 and product B. The oxygen atom in B is now in the -2 state, and the substrate (ethylene) has been oxidized and is one step closer to becoming completely oxidized (to carbon dioxide and water).
Figure 4. Reaction of Ethylene With Ozone (idealized, not balanced)Ozone readily reacts with most species containing multiple bonds (such as C=C, C=N, N=N, etc.) through mechanisms similar to the one shown in Figure 4. Ozone does not react with singly bonded functionality such as C-C, C-O, O-H at near the same rate. This is, in part, because there is no easy chemical pathway for the oxidation to take place. However, ozone does react with simple oxidizable ions such as S2-, to form oxyanions such as SO32- and SO42-. These oxidations are simple and the mechanisms only require contact of ozone with the ion. Consequently, the oxidation of these ions by ozone occurs rapidly.
As can be seen, it is sometimes better to think of ozone as a highly reactive species capable of reacting with many species rather than just as a compound with a large reduction potential. This is because the practical action of ozone is often dependent on how it reacts with pollutants, and usually not in its ability to simply gain electrons (except in the case of simple ion oxidations). This thinking emphasizes that an appropriate reaction pathway must exit for ozone to react with a substrate. In other words, although the thermodynamics for ozone induced oxidation may be favorable (due to ozone’s high reduction potential), kinetic factors will most often dictate whether ozone will oxidize a pollutant in a reasonable time frame.
To assist ozone in reacting with difficult to oxidize pollutants, other agents such as hydrogen peroxide (H2O2), high energy light (ultraviolet range or UV), and/or high pH (hydroxyl ion or HO-) are often used. Processes utilizing one or more of these agents to assist ozone are called advanced oxidation processes (AOP). These chemical aids are each capable of breaking down, at least partially, pollutants (H2O2 is an oxidant, UV and HO- are not). Consequently, when used in conjunction with ozone, the breakdown of an organic pollutant to CO2 and H2O is facilitated. To better understand how each of these agents assist ozone, we need to look at how they react (or interact) with various organic compounds. It is not the purpose of this report to examine every possible reaction pathway, but to give the reader a feeling for the type of process these agents help promote. Therefore, we will look at only the most common examples.
Hydrogen peroxide can decompose into the hydroxyl radical (HO.), and H2O2 can be thought of as a source of this highly reactive free radical. This conversion occurs slowly in cold, dark conditions, but is facile in the presence of UV/visible light or at high temperatures. Free radicals, being oxidants, are capable of withdrawing atoms (often hydrogen) from a substrate, while turning the substrate into an unstable free radical. The resulting free radical based substrate is highly reactive, and may undergo internal reactions that form functionality on the substrate that is susceptible to attack by ozone. Consequently, the role of hydrogen peroxide in AOP is probably to make the pollutant more susceptible to ozone attack (major effect), and also to aid in the overall oxidation (minor effect).
UV light acts by providing energy to disrupt chemical bonds, since the energy of UV light is on the same order as that of covalent bonds. When a chemical bond is cleaved by UV light, the remaining fragments are often more susceptible to ozone attack. Also, these fragmented by-products themselves can be reactive and degrade further.
UV light can be used to cleave H2O2 , forming the highly reactive hydroxyl radical, HO.. UV light, in the presence of ozone, also promotes the formation of hydroxyl radicals and the generation of oxygen. UV is commonly used in this manner as an ozone “destruct,”, converting ozone to molecular oxygen. Since UV can cause the conversion of ozone to molecular oxygen, its location in an AOP must be thought out.
High pH, or more properly put, a high concentration of HO-, is also effective at breaking down many organic species. The mode of hydroxyl ion action is usually attack at polar bonds, resulting in bond breaking, hydrolysis, and/or elimination reactions. The hydroxyl ion is also thought by some to lead the direct formation of HO- under AOP conditions. However, it is difficult to see how that could occur since HO- is not known to easily give up an electron (the reduction potential of HO- is 2.80V, or about 0.73V greater than that of ozone)!
However, high pH may lead to hydroxyl radical formation through an indirect route. The hydroxyl ion reacts with ozone to form the hydroperoxide ion, HO2-. HO2- is the conjugate base of H2O2 and at a pH lower than 11.6 (which is the pKa of H2O2) will be predominantly converted to H2O2. Therefore, high pH catalyzes the formation of hydrogen peroxide, which is a source of hydroxyl radicals. This process is shown in Figure 5.
Net ReactionFigure 5. Formation of H2O2 Catalyzed by HO¯
Ozone is an effective agent for at least the partial oxidation of simple ions and species containing multiple bonds. But for more complete and efficient oxidation, other agents such as H2O2, UV, and/or high pH are often used in AOP processes. These components work by degrading the substrate to fragments that are more susceptible to ozone attack, oxidizing the substrate in reaction mechanisms not available to ozone, and by helping to create other reactive species such as HO- that have unique oxidation mechanisms.
