On July 26, 1943, a black cloud enveloped much of Los Angeles. It was difficult to see more than a few blocks and people’s eyes and noses burned. In the midst of World War II, many residents feared they were under chemical attack.
Fortunately, the dark fog was not the result of chemical warfare, but residents couldn’t identify the culprit. At first, many blamed pollution from a manufacturing plant that made synthetic rubber. Public pressure to stop the toxic cloud forced the plant to close, but the problem persisted.
People began calling the periodic gaseous invasions “smog,” thinking (incorrectly) that it was a combination of smoke and fog. Scientists thought (correctly) that emissions from cars and manufacturing plants could cause smog, but they didn’t understand how.
What is is smog? Where does it come from? And how could the city of Los Angeles stop it?
Two scientists, Arie Jan Haagen-Smit and Eugene Houdry, were instrumental in beginning to answer these questions. To do so, they used their knowledge of chemical reaction kinetics, which is the focus of this learning module.
Energy and Reaction Rates
In a chemical reaction, the speed that reactant molecules convert into product molecules is called its reaction rate. The study of reaction rates and what variables can speed them up or slow them down is called chemical reaction kinetics.
Chemical reactions involve changes in energy (see “Energy Changes” in Chemical Reactions), and each reaction has an energy barrier, called activation energy, that reactants must overcome before they can form products.
Imagine riding a bike over a steep hill. Your legs push hard as the incline gets steeper and steeper. Finally, after spending much energy, you reach the top and—at last—coast down the other side. This is a lot like what happens with chemical reactions. Activation energy is the hill that reactants must climb over to form products.
Cars are culprits in the Los Angeles smog mystery because of the reactions that occur during and after combustion. To understand this, let’s consider how a car engine burns gasoline to produce power. Although gasoline contains many different hydrocarbon molecules—for now—assume it is made primarily of octane (C8H18) and follows the combustion reaction in Eq. 1 (see “Combustion reactions” in Chemical Reactions):
2 C8H18(l) + 25 O2(g) → 16 CO2(g) + 18 H2O(g)
At first glance, the reaction looks pretty clean, producing just CO2 (which is a problem for global warming but does not produce smog) and water. And an energy diagram for the reaction would look like Figure 1, where the height of the hill represents the activation energy.
Despite what Equation 1 shows, what would happen if you simply mixed air and gasoline together? Would it spontaneously combust? No. You must first provide enough energy (the activation energy) to get the reaction going. Cars have spark plugs that shoot tiny electrical sparks to add the energy to initiate the reaction. Then, because combustion is exothermic (see “Energy Changes” in Chemical Reactions), it generates heat that sustains the reaction.
Just like in our bicycle example, there are ways to speed up chemical reactions. What if a friend of yours was running behind you and pushing you up the hill? You would have more energy so it’s able to climb the hill much faster. The same thing occurs in chemical reactions. Increasing the energy of reactants speeds reaction up, and decreasing the energy of reactants slows reaction rates down. But why?
When reactant molecules collide with enough energy—energy greater than or equal to the activation energy—they form products. This is known as the collision theory of reactions. If we add energy to reactant molecules, more molecules collide with enough force to push them over the activation energy hill. In other words, more reactant molecules cross the activation energy threshold in a given time, and the reaction rate increases. If we decrease the energy of reactants, the opposite occurs.
Equation 1 suggests that gasoline combustion produces only carbon dioxide and water. But these products don’t contribute to Los Angeles’ smog problem. Arie Jan Haagen-Smit, a biochemist who had previously studied the compounds responsible for the flavor of pineapple and some wines, suspected there was more to the story.
Chemical Reactions in the Real World
Haagen-Smit became interested in the Los Angeles smog issue because of the damage that the pollution was doing to crops in the area. He knew that car engines did not only give off carbon dioxide and water, and that the reactions were more complicated.
It turns out that gasoline contains small amounts of sulfur compounds. When the sulfur compounds are provided with activation energy, as they are in an internal combustion engine, they react with oxygen to form sulfur oxides. Since many different compounds of sulfur and oxygen molecules are formed, including SO, SO2, etc., sulfur oxides are written as SOX. We also know that air contains nitrogen. Similar to sulfur, nitrogen reacts with oxygen during combustion to form nitrogen oxides (written as NOX for the same reason as SOX).
In addition, inefficient combustion of the carbon in gasoline forms carbon monoxide (CO) along with carbon dioxide (CO2). Since engines are not 100% efficient, incomplete combustion will leave some organic compounds in the fuel unburned (for example methane, CH4). These are usually referred to as volatile organic carbons (VOCs).
So, if we revisit our combustion reaction, we see that it is not as simple as Equation 1. Multiple reactions are actually occurring at the same time:
Equations 2A through 2C
A. Gasoline(l) + O2(g) → CO(g) + CO2(g) + H20(g) + VOCs(g)
B. N2 + O2(g) → NOx(g)
C. S + O2(g) → SOx(g)
In the 1940s, the products on the right side of Equations 2A through 2C all exited car tailpipes. Haagen-Smit suspected emissions from cars were at least partially to blame for Los Angeles’s smog problem. To prove his theory, he needed to find out what compounds were causing damage to the area’s crops.
He began to investigate smog’s chemical composition, cooling samples of the gas until some of its ingredients condensed. Then he analyzed the resulting liquid’s chemical properties. Haagen-Smit found it contained the same kind of volatile organic compounds (VOC) that could result from incompletely burned gasoline.
Now he needed to find out if these VOCs resulted in smog’s toxic properties. Working with a colleague, he exposed plants like spinach, beets, and alfalfa to smoggy air. Haagen-Smit noticed that the toxic air turned some plant leaves silver, others bronze, and others a bleachy white. Then he exposed the plants to various VOCs, thinking that the culprit chemical would produce the same discoloration of plant leaves as the smog. But after testing 50 different VOCs, the experiments failed—none of the compounds damaged the plants in the same way, and he was no closer to the answer.
Finally, he tested the plants by exposing them to a mix of VOCs and the oxidant ozone (O3).
At last! When he exposed the plants to the mixture of ozone and VOCs, their discoloration matched that of the smog. He deduced that that difference was the ozone, and he was confident he’d found the culprit. However, he still didn’t know how ozone formed in air near the ground.
High up in the atmosphere, the ozone layer blocks the most dangerous form of UV radiation from reaching Earth’s surface (see “Temperature in the atmosphere” in Composition of Earth’s Atmosphere). Way up there, ozone is good. Near the ground, where humans and other animals can breathe it, ozone is bad because it causes respiratory health issues like asthma and lung infections. An example of how smog can change the air quality of an area is shown in Figure 2.
Haagen-Smit remembered reading that compounds could oxidize in air in the presence of sunlight, and he developed an experiment that synthesized ozone using VOCs, nitrogen oxides (NOx), and sunlight. Making smog in the lab proved it was possible for harmful ozone to form at ground level from the chemicals in automobile exhaust under the right conditions. Today, we understand ground-level ozone pollution forms according to the following reaction:
NOx(g) + VOC(g) + Heat + Sunlight → O3(g)
At cool temperatures and low levels of sunlight (like in winter or northern latitudes), this reaction would be so slow that we wouldn’t need to worry about ground-level ozone pollution forming. So how do heat and sunlight speed this reaction up? They add energy! Not only do they add energy to push a few reactants over the activation energy hill, they add so much energy that reaction speeds increase enough to cause an ozone pollution problem.
Besides heat and sunlight, can you think of any other ways to increase or decrease the energy of reactants (and therefore speed up or slow down reaction rates)? Well, read on.
Variables That Affect Reaction Rates
Collision theory tells us that heat gives reactants more kinetic energy, which causes them to move around faster (like how boiling water moves around the pot). When reactant molecules move faster, more collide with enough energy to react, and the reaction rate speeds up. Colder temperatures do the opposite.
All of the car emission pollutant reactions shown in Equations 2A through 2C occur faster because gasoline combustion is exothermic and gives off heat. If the temperature of combustion could be reduced in a way that still allowed the reaction to power a car, these pollution reactions might not be a problem. So far, no one has discovered a way to do this.
Concentration is the amount of a substance in a given volume of solution. The higher the concentration of a substance, the more molecules there are within the same volume. A lower concentration has fewer molecules. For example, adding sugar to a hot cup of coffee or tea would increase the concentration of sugar with every spoonful.
In areas with more manufacturing plants and traffic, concentrations of NOx and VOCs in the atmosphere are higher. These elevated concentrations mean more ozone-producing reactions, which explains why highly populated urban areas suffer more from ground-level ozone pollution than rural locations.
A cube of sugar in coffee or tea might take a few minutes to completely dissolve, whereas, a spoonful of granulated sugar would dissolve almost instantaneously. The difference comes down to surface area. In the cube, most of the sugar molecules aren’t exposed to the liquid until the outer layers dissolve. With the granulated sugar, however, the liquid tea or coffee immediately surrounds each individual grain.
It works similarly in chemical reactions that involve solids. Increasing the surface area of the solid reactants by grinding them into a fine powder exposes more molecules, increasing the likelihood that they’ll collide.
If you wanted a cube of sugar to dissolve faster in a cup of coffee or tea, what would you do? Grab a spoon and stir it! Mixing reactants together is another way to add kinetic energy: Increase the number of reactant molecules that collide with one another and you will increase reaction rates.
Sunlight, temperature, concentration, surface area, and mixing aren’t the only way to affect reaction rates. Now that Haagen-Smit had found that car and manufacturing emissions indirectly produced smog via reactions that occurred in the atmosphere, a French mechanical engineer named Eugene Houdry would use another method of controlling reaction rates to begin solving the problem of ozone pollution.
As smog continued to be covered in the news, Eugene Houdry also became concerned about smog pollution in Los Angeles and other large cities. Houdry was especially worried because the number of automobiles and industrial plants was rapidly growing at the time. He knew the problem would only get worse.
Haagen-Smit’s work had identified the culprits in the formation of smog. Houdry began building on this work by developing a device to reduce the VOC emissions in automobile tailpipes. Houdry studied catalysts that could speed up the rates of chemical reactions that would reduce VOC pollution from forming. He used a material that lowers activation energy, called a catalyst.
How does a catalyst work?
Catalysts speed up reactions by lowering the activation energy required (Figure 3). Unlike the variables listed above, catalysts do not increase the energy of reactants or the number of collisions, and catalysts are not chemically changed as part of the reaction. They just make the energy barrier smaller. In the analogy of the bike riding over a hill, a catalyst shrinks the hill, making it easier to pedal to the other side (Figure 3).
Ironically, Houdry first began his work on catalysts in the 1920s on a chemical process used to make gasoline. The process he developed increased the amount of gasoline that could be extracted from oil and other petroleum products and helped make the growth of automobiles possible. However, Houdry had no idea then that gasoline emissions would create ozone pollution. Decades later, in the 1950s, Houdry used his knowledge of kinetics and catalysts to invent the “catalytic converter” to reduce smog emissions. Houdry’s first catalytic converter helped break down the VOCs and carbon monoxide that could be released by cars. Since VOCs were a part of the reaction that formed ozone, they also reduced smog in the process.
Have catalytic converters solved the ozone pollution problem?
Over time, scientists and engineers have improved on Houdry’s design. Today, converters are more efficient at stopping pollution by reducing more pollutants cars emit. However, they haven’t completely solved the problem of ozone pollution because catalytic converters are not yet perfectly efficient, and cars are not the only source of these pollutants. It’s true that cars today emit significantly fewer pollutants than in the 1950s, but there are also many more cars on the road, and all that pollution adds up, especially in larger cities with high traffic.
What other problems might reaction kinetics help us solve?
Ground-level ozone is just one of many problems that reaction kinetics and catalysts can help solve. What about all the CO2 released from car tailpipes and other sources like power plants that burn coal and natural gas? Well, scientists are studying catalysts that can break down CO2 into carbon and oxygen gas.
Earth’s climate is changing. Hopefully advances in science and alternative fuels will help us tackle climate change as well as ozone pollution. Some scientists have developed new catalysts that make jet fuel from carbon dioxide. Imagine a world where instead of drilling for more oil and gas, we simply made fuel from CO2 harvested from the atmosphere.
Table of Contents
- Energy and Reaction Rates
- Chemical Reactions in the Real World
- Variables That Affect Reaction Rates
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