How a catalytic converter works. Before catalytic converters were developed, waste gases made by a car engine blew straight down the exhaust tailpipe and into the. How do catalytic converters work? Chris Woodford. Last updated: June 1. Blackened buildings and choking. Los Angeles, London, Paris, or Beijing. Cars, buses, and. But their. engine pollution spoils the places where we live and. Fortunately, most vehicles are now fitted with. Let's take a closer look at these brilliant gadgets and how they. Photo: An experimental new catalytic converter. Picture courtesy of Southwest Research Institute and. US Department of Energy/National Renewable Energy Laboratory (Do.
E/NREL). Why engines make pollution. Photo: The columns of the Parthenon in Athens, Greece have been blackened by vehicle pollution. Athens is one of the world's most traffic- polluted cities. Photo by courtesy of U. S. Geological Survey. Car engines run on gasoline or diesel. Most of our petroleum is formed when the. Petroleum is made up of hydrocarbons. In theory, if you burn any kind of hydrocarbon fuel with oxygen from the air, you release a lot. In practice, though, there may be too. That means you generally get some. The pollutant gases made by car engines include a. VOCs (volatile organic. What is a catalytic converter? Pollutant gases are made of harmful molecules, but those molecules. So if we could find a way of. That's the job that a catalytic converter does. Photo: This scientist is working to develop. Photo by Warren Gretz courtesy of Do. E National Renewable Energy Laboratory. These gadgets are much simpler than they sound. A catalyst. is simply a chemical that makes a chemical reaction go faster without itself. It's a bit like an athletics coach who stands. The. coach doesn't run anywhere; he just stands there, waves his arms about. In a catalytic converter, the. The catalyst is made from platinum or a similar, platinum- like metal. A catalytic converter is a large metal box, bolted to the underside of your car, that has two pipes coming out of it. One of them (the converter's "input") is connected to the engine and brings in hot, polluted fumes from the engine's cylinders (where the fuel burns and produces power). The second pipe (the converter's "output") is connected to the tailpipe (exhaust). As the gases from the engine fumes blow over the catalyst, chemical reactions take place on its surface, breaking apart the pollutant gases and converting them into other gases that are safe enough to blow harmlessly out into the air. One very important thing to note about catalytic converters is that they require you to. What happens inside the converter? Photo: Engineers are constantly trying to improve the performance of. This is an example of a low- temperature oxidation catalyst made from tin oxide and platinum. Photo by CPL Bryant V courtesy of NASA Langley Research Center (NASA- La. RC). Inside the converter, the gases flow through a dense honeycomb. The honeycomb structure means the gases touch a. Typically, there are two different catalysts in a. One of them tackles nitrogen oxide pollution using. This breaks up nitrogen oxides into nitrogen and. The other catalyst works by an opposite chemical process called oxidation (adding. Another oxidation reaction turns unburned hydrocarbons in the exhaust into carbon dioxide and water. In effect, three different chemical reactions are going on at the same time. That's why we talk about three- way catalytic converters. Some, less- effective converters carry out. After the catalyst has done its job, what emerges from the exhaust is. How effective are catalytic converters? Chart: Effectiveness of catalytic converters. Cats make a big difference to emissions, with three- way converters giving a significant extra benefit over two- way converters. Figures show pollutants in grams per kilometer at 8. Chart drawn by Explain that Stuff. US EPA (1. 99. 0), quoted in table 3. Air Pollution from Motor Vehicles: Standards and Technologies for Controlling Emissions, Faiz et al, World Bank, 1. Catalytic converters are mainly designed to reduce immediate, local air pollution—dirty air where you're driving—and this chart certainly seems to suggest that they're effective. Even so, people sometimes question whether they're really as green as they seem. One problem is that they only really work at high temperatures (over 3. C/6. 00°F or so), when the engine has had chance to warm things up, which might take 1. So they are ineffective for the first part of a journey (or any part of a very short journey). Another issue is whether they increase greenhouse gas emissions. We think of carbon dioxide as a safe gas, because it's not toxic in everyday concentrations. Nevertheless, it isn't entirely harmless, because we now know it's the major cause of global warming and climate change. Some people believe catalytic converters make climate change worse because they turn carbon monoxide into carbon dioxide. In fact, the carbon monoxide your car produces would eventually turn into carbon dioxide in the atmosphere all by itself, so a catalytic converter makes no difference on that score: it simply reduces the carbon monoxide a car pumps into the street as it drives along, improving the local air quality. But when it comes to climate change, auto engineers and environmentalists have long pointed out another serious issue. Although cats turn most nitrogen oxides into nitrogen and oxygen, they also produce small amounts of nitrous oxide (N2. O) in the process, a greenhouse gas that's over 3. The trouble is that with so many vehicles on the road, even small amounts of nitrous oxide add up to a major problem. Back in 2. 00. 0, the. Intergovernmental Panel on Climate Change noted: "The introduction of catalytic converters as a pollution control measure in the majority of industrialized countries is resulting in a substantial increase in. N2. O emissions from gasoline vehicles.". Fortunately, newer catalytic converters produce dramatically less nitrous oxide than older ones. Even so, while catalytic converters have certainly helped us to tackle short- term air pollution, there are. Who invented the catalytic converter? Whom do we thank for making streets and cities safer and cleaner? French chemical engineer. Eugene Houdry (1. United States, filing the invention on May 5, 1. US Patent 2,6. 74,5. Catalytic converter for exhaust gases). April 6, 1. 95. 4. Houdry had previously invented catalytic cracking, the industrial process by. After that, he experimented with making different kinds of vehicle fuels and making them cleaner. Although he recognized the growing problem of air pollution, his ideas were far ahead of their time. Fortunately, in the 1. In 1. 97. 3, the US Environmental Protection Agency (EPA) released a report demonstrating how lead harmed people's health, which began the slow process for removing lead from gasoline. The first practical catalytic converters appeared shortly afterward, in the mid- 1. Artwork: Eugene Houdry's original catalytic converter from his 1. It's essentially a set of concentric metal tubes (blue) through which the exhaust gases flow. Clean air is sucked in through ventilation holes (yellow) with the help of a venturi (orange). As in a modern cat, Houdry explains that "the deposited finely divided metal catalyst is preferably platinum," although other similar metals can be used; unlike a modern cat, the catalyst (green) isn't arranged in a honeycomb but mounted in sixteen separate rings (red) at intervals along the tube, with each one working in parallel. Artwork from US Patent 2,6. Catalytic converter for exhaust gases, courtesy of US Patent and Trademark Office. Houdry invented the basic oxidation catalyst for tackling carbon monoxide. Improved, three- way catalytic converters, which could also tackle nitrogen oxides, were designed in the early 1. Carl Keith (1. 92. John Mooney (1. 92. Engelhard Corporation. Artwork: In Carl Keith and John Mooney's improved design, there are two separate catalytic converters. Polluted gases flow from the engine (red, 1. Artwork from US Patent 3,8. Process and Apparatus, courtesy of US Patent and Trademark Office. Find out more. On this website. Books. News articles. Catalytic converter thefts double as metal prices rise by Nicola Beckford, BBC News, 6 November 2. Precious metals are making catalytic converters an attractive target for thieves. Inventor in cleaner engine claim: BBC News, 2. January 2. 01. 0. A Scottish inventor claims to have developed a cool- running engine that produces virtually no particulate (soot) emissions. As Platinum Soars, the Catalytic Converter Gets Hot by Matthew Phenix. Wired, 1. 7 February 2. Catalysis - Wikipedia. Catalysis () is the increase in the rate of a chemical reaction due to the participation of an additional substance called a catalyst[1] (), which is not consumed in the catalyzed reaction and can continue to act repeatedly. Often only tiny amounts of catalyst are required in principle.[2]In general, reactions occur faster with a catalyst because they require less activation energy. In catalyzed mechanisms, the catalyst usually reacts to form a temporary intermediate which then regenerates the original catalyst in a cyclic process. Catalysts may be classified as either homogeneous or heterogeneous. A homogeneous catalyst is one whose molecules are dispersed in the same phase (usually gaseous or liquid) as the reactant molecules. A heterogeneous catalyst is one whose molecules are not in the same phase as the reactants, which are typically gases or liquids that are adsorbed onto the surface of the solid catalyst. Enzymes and other biocatalysts are often considered as a third category. Technical perspective[edit]In the presence of a catalyst, less free energy is required to reach the transition state, but the total free energy from reactants to products does not change.[1] A catalyst may participate in multiple chemical transformations. The effect of a catalyst may vary due to the presence of other substances known as inhibitors or poisons (which reduce the catalytic activity) or promoters (which increase the activity and also affect the temperature of the reaction).[1]Catalyzed reactions have a lower activation energy (rate- limiting free energy of activation) than the corresponding uncatalyzed reaction, resulting in a higher reaction rate at the same temperature and for the same reactant concentrations. However, the detailed mechanics of catalysis is complex. Catalysts may affect the reaction environment favorably(like heat), or bind to the reagents to polarize bonds, e. Kinetically, catalytic reactions are typical chemical reactions; i. Usually, the catalyst participates in this slowest step, and rates are limited by amount of catalyst and its "activity". In heterogeneous catalysis, the diffusion of reagents to the surface and diffusion of products from the surface can be rate determining. A nanomaterial- based catalyst is an example of a heterogeneous catalyst. Analogous events associated with substrate binding and product dissociation apply to homogeneous catalysts. Although catalysts are not consumed by the reaction itself, they may be inhibited, deactivated, or destroyed by secondary processes. In heterogeneous catalysis, typical secondary processes include coking where the catalyst becomes covered by polymeric side products. Additionally, heterogeneous catalysts can dissolve into the solution in a solid–liquid system or sublimate in a solid–gas system. Background[edit]The production of most industrially important chemicals involves catalysis. Similarly, most biochemically significant processes are catalysed. Research into catalysis is a major field in applied science and involves many areas of chemistry, notably organometallic chemistry and materials science. Catalysis is relevant to many aspects of environmental science, e. Catalytic reactions are preferred in environmentally friendly green chemistry due to the reduced amount of waste generated,[3] as opposed to stoichiometric reactions in which all reactants are consumed and more side products are formed. Many transition metals and transition metal complexes are used in catalysis as well. Catalysts called enzymes are important in biology. A catalyst works by providing an alternative reaction pathway to the reaction product. The rate of the reaction is increased as this alternative route has a lower activation energy than the reaction route not mediated by the catalyst. The disproportionation of hydrogen peroxide creates water and oxygen, as shown below. H2. O2 → 2 H2. O + O2. This reaction is preferable in the sense that the reaction products are more stable than the starting material, though the uncatalysed reaction is slow. In fact, the decomposition of hydrogen peroxide is so slow that hydrogen peroxide solutions are commercially available. This reaction is strongly affected by catalysts such as manganese dioxide, or the enzyme peroxidase in organisms. Upon the addition of a small amount of manganese dioxide, the hydrogen peroxide reacts rapidly. This effect is readily seen by the effervescence of oxygen.[4] The manganese dioxide is not consumed in the reaction, and thus may be recovered unchanged, and re- used indefinitely. Accordingly, manganese dioxide catalyses this reaction.[5]General principles[edit]Catalytic activity is usually denoted by the symbol z[6] and measured in mol/s, a unit which was called katal and defined the SI unit for catalytic activity since 1. Catalytic activity is not a kind of reaction rate, but a property of the catalyst under certain conditions, in relation to a specific chemical reaction. Catalytic activity of one katal (Symbol 1 kat = 1 mol/s) of a catalyst means an amount of that catalyst (substance, in Mol) that leads to a net reaction of one Mol per second of the reactants to the resulting reagents or other outcome which was intended for this chemical reaction. A catalyst may and usually will have different catalytic activity for distinct reactions. See katal for an example. There are further derived SI units related to catalytic activity, see the above reference for details. Typical mechanism[edit]Catalysts generally react with one or more reactants to form intermediates that subsequently give the final reaction product, in the process regenerating the catalyst. The following is a typical reaction scheme, where C represents the catalyst, X and Y are reactants, and Z is the product of the reaction of X and Y: Although the catalyst is consumed by reaction 1, it is subsequently produced by reaction 4, so it does not occur in the overall reaction equation: X + Y → ZAs a catalyst is regenerated in a reaction, often only small amounts are needed to increase the rate of the reaction. In practice, however, catalysts are sometimes consumed in secondary processes. The catalyst does usually appear in the rate equation. For example, if the rate- determining step in the above reaction scheme is the first step. X + C → XC, the catalyzed reaction will be second order with rate equation v = kcat[X][C], which is proportional to the catalyst concentration [C]. However [C] remains constant during the reaction so that the catalyzed reaction is pseudo- first order: v = kobs[X], where kobs = kcat[C]. As an example of a detailed mechanism at the microscopic level, in 2. Danish researchers first revealed the sequence of events when oxygen and hydrogen combine on the surface of titanium dioxide (Ti. O2, or titania) to produce water. With a time- lapse series of scanning tunneling microscopy images, they determined the molecules undergo adsorption, dissociation and diffusion before reacting. The intermediate reaction states were: HO2, H2. O2, then H3. O2 and the final reaction product (water molecule dimers), after which the water molecule desorbs from the catalyst surface.[7][8]Reaction energetics[edit]. Generic potential energy diagram showing the effect of a catalyst in a hypothetical exothermic chemical reaction X + Y to give Z. The presence of the catalyst opens a different reaction pathway (shown in red) with a lower activation energy. The final result and the overall thermodynamics are the same. Catalysts work by providing an (alternative) mechanism involving a different transition state and lower activation energy. Consequently, more molecular collisions have the energy needed to reach the transition state. Hence, catalysts can enable reactions that would otherwise be blocked or slowed by a kinetic barrier. The catalyst may increase reaction rate or selectivity, or enable the reaction at lower temperatures. This effect can be illustrated with an energy profile diagram. In the catalyzed elementary reaction, catalysts do not change the extent of a reaction: they have no effect on the chemical equilibrium of a reaction because the rate of both the forward and the reverse reaction are both affected (see also thermodynamics). The second law of thermodynamics describes why a catalyst does not change the chemical equilibrium of a reaction. Suppose there was such a catalyst that shifted an equilibrium.
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