2.1 The Urban Ozone Problem: An Overview

The classic examples of urban photochemical oxidant pollution are found in cities like Los Angeles and Mexico City. This phenomenon, often called photochemical smog, is characterized by high concentrations of oxidants, the most prominent of which is ozone (O₃). Ozone is not emitted directly but is a secondary pollutant formed in the atmosphere through a complex series of chemical reactions involving organic compounds (hydrocarbons), nitrogen oxides (NOx), and sunlight. Because sunlight is a critical ingredient, high ozone levels are typically a summertime problem.

The daily rhythm of urban life is clearly imprinted on the concentrations of these pollutants. Figure 2 shows a typical diurnal (daily) variation of NO, NO₂, and total oxidants in Pasadena, California.

[Insert Figure 2: Diurnal variation of pollutants in Pasadena, CA]

The sequence is revealing:

1. Early Morning: Concentrations of nitric oxide (NO), a primary pollutant from vehicle exhaust, rise sharply, peaking during the morning commute.
2. Mid-Morning: As NO is oxidized in the atmosphere, its concentration falls while the concentration of nitrogen dioxide (NO₂) rises, reaching its peak after the NO peak.
3. Afternoon: After NO has been significantly depleted and NO₂ has peaked, the concentration of ozone (and other oxidants) begins to grow, typically reaching its maximum in the afternoon.

2.2 The Core Chemistry of Tropospheric Ozone Formation

At its most basic level, the formation of ozone in the troposphere is governed by a simple three-reaction cycle involving nitrogen oxides.

1. Photolysis of NO₂: Nitrogen dioxide absorbs sunlight and breaks apart, producing nitric oxide and a ground-state oxygen atom.
2. Ozone Formation: The highly reactive oxygen atom immediately combines with an oxygen molecule to form ozone. (M is any third molecule, like N₂ or O₂, that absorbs excess energy).
3. Ozone Destruction: The newly formed ozone can then be destroyed by reacting with the nitric oxide produced in the first step, reforming NO₂.

Under these conditions, a rapid equilibrium is established. Assuming a steady state where the rate of ozone formation (Reaction 2) equals its rate of destruction (Reaction 3), we can derive the photostationary state or the Leighton relationship:

[O₃] = k₁[NO₂] / k₃[NO]

This equation reveals that the ozone concentration is directly proportional to the rate of NO₂ photolysis (which depends on sunlight intensity, k₁) and the ratio of NO₂ to NO concentration. However, this simple model presents a major contradiction: if NO destroys ozone as quickly as NO₂ photolysis creates it, high concentrations of ozone should not accumulate. The high ozone levels observed in polluted urban air clearly indicate that another chemical pathway must be at play—one that converts NO to NO₂ without consuming an ozone molecule. Deviations from this photostationary state relationship exist, because as we will see shortly, peroxy radicals can also react with NO to make NO₂.

2.3 The Role of Hydrocarbons and the Hydroxyl Radical (OH)

The missing piece of the puzzle lies in the atmospheric oxidation of Volatile Organic Compounds (VOCs), or hydrocarbons. Smog chamber experiments provide a crucial clue. As seen in Figure 3, the observed loss rate of propene (a common VOC) in a simulated urban atmosphere is far greater than can be explained by its known reactions with ozone and oxygen atoms. This points to the existence of a much more powerful and reactive oxidant.

[Insert Figure 3: Propene loss rate in a smog chamber]

This powerful oxidant is the hydroxyl radical (OH), a highly reactive species with an unpaired electron. The OH radical initiates a chain reaction that rapidly oxidizes hydrocarbons and, critically, converts NO to NO₂. Let’s examine this process using propene as an example.

1. Chain Initiation: The OH radical attacks the propene molecule.
2. Peroxy Radical Formation: The resulting alkyl radicals instantly react with molecular oxygen to form alkylperoxy radicals (RO₂).
3. NO to NO₂ Conversion: This is the key step. The peroxy radicals (both RO₂ and the simpler hydroperoxy radical, HO₂) are powerful enough to oxidize NO to NO₂, fueling the production of ozone.
4. Radical Regeneration: The subsequent reactions of the products not only form stable molecules like acetaldehyde and formaldehyde but also regenerate the OH radical.

The OH radical, therefore, acts as a catalyst. As shown in the schematic below, it drives a reaction cycle that consumes hydrocarbons and NO while producing NO₂ and other products, ultimately leading to the buildup of ozone.

[Insert Figure 4: Schematic of OH-initiated hydrocarbon oxidation]

2.4 Peroxyacetyl Nitrate (PAN): A Key Photochemical Oxidant

The oxidation of hydrocarbons produces a variety of secondary pollutants. One of the most significant is peroxyacetyl nitrate (PAN). Its formation begins with acetaldehyde (CH₃CHO), itself a product of hydrocarbon oxidation. Acetaldehyde reacts with OH radicals to form a peroxyacetyl radical.

CH₃CHO + OH → CH₃CO + H₂O (Reaction 10) CH₃CO + O₂ → CH₃C(O)O₂ (Reaction 11)

This peroxyacetyl radical now stands at a chemical crossroads. Its fate depends entirely on what it collides with next in the atmospheric soup: if it meets an NO molecule, it fuels further ozone production by oxidizing NO to NO₂ (Reactions 12-15); if it meets an NO₂ molecule, it forms the stable reservoir species PAN.

CH₃C(O)O₂ + NO₂ ↔ CH₃C(O)O₂NO₂ (PAN) (Reaction 16)

PAN is a significant photochemical oxidant in its own right, known as a potent eye irritant and a cause of damage to plants. Its atmospheric lifetime is highly dependent on temperature. In the warm lower troposphere, it is relatively short-lived (about 30 minutes at 298 K), thermally decomposing back into its precursor radicals. However, in the cold upper troposphere, it can persist for months, acting as an important reservoir for NOx. It can transport nitrogen oxides over long distances before decomposing and releasing them in remote regions.

2.5 Radical Sources and Chain Termination

For the entire photooxidation process to begin, there must be an initial source of hydroxyl (OH) radicals. There are several key pathways for radical production in the atmosphere:

• Photolysis of Ozone: In the presence of water vapor, the photolysis of ozone can produce two OH radicals.
• Photolysis of Nitrous Acid (HONO): HONO can build up overnight and is rapidly photolyzed by morning sunlight, providing a crucial “kick-start” of OH radicals for the day’s chemistry.
• Photolysis of Peroxides: Both hydrogen peroxide (H₂O₂) and organic peroxides (ROOH) can photolyze to produce OH radicals.
• Photolysis of Aldehydes: Aldehydes, such as formaldehyde (HCHO), are major radical sources.

The hydrocarbon oxidation cycle is a chain reaction, but it does not continue indefinitely. It is eventually stopped by chain-terminating reactions that remove the chain-carrying radicals by forming relatively stable products. The most important termination reactions are:

HO₂ + HO₂ → H₂O₂ + O₂ (Reaction 19) RO₂ + HO₂ → ROOH + O₂ (Reaction 20) OH + NO₂ + M → HNO₃ + M (Reaction 21)

2.6 Organic Reactivity and Ozone Control Strategies

Not all VOCs are created equal; they have vastly different abilities to promote ozone formation. The atmospheric lifetime of a VOC, largely determined by its reaction rate with the OH radical, is a key indicator of its reactivity.

The balance between VOCs and NOx is critical. In a typical urban mix of pollutants, the chain-terminating reaction of OH + NO₂ (Reaction 21) becomes dominant when the VOC-to-NO₂ ratio is low (e.g., less than 5.5:1). Under these conditions, OH radicals are removed before they can oxidize VOCs, thereby inhibiting ozone formation. Conversely, when the ratio is high, OH preferentially reacts with VOCs, accelerating the production of radicals and ozone.

To quantify this, scientists use a metric called Maximum Incremental Reactivity (MIR), which measures the grams of ozone formed per gram of a specific VOC added to a test atmosphere.

This table clearly shows that some compounds, like ethene and m-xylene, are far more effective at producing ozone than others, like methane and methanol. This concept is the basis for control strategies like reformulating gasoline to change the composition of VOC emissions, not just the total amount.

The complex, non-linear relationship between initial VOCs, initial NOx, and maximum ozone produced can be visualized using an ozone isopleth diagram. These contour plots are essential tools for developing effective air pollution control strategies.

[Insert Figure 5: Ozone isopleth diagram for Atlanta, GA]

The diagram is divided by a “ridge line.”

• VOC-Limited Region (above the ridge): Here, there is an abundance of NOx relative to VOCs. Ozone formation is limited by the availability of VOCs. As shown for the base case in Atlanta, reducing VOCs is an effective strategy to lower ozone.
• NOx-Limited Region (below the ridge): Here, there is an abundance of VOCs relative to NOx. Ozone formation is limited by the availability of NOx.

Crucially, in a VOC-limited regime like the Atlanta base case, reducing NOx can have the counterintuitive effect of increasing peak ozone levels. This occurs because reducing NOx lessens the rate of the chain-terminating OH + NO₂ reaction, allowing more OH radicals to participate in the ozone-forming oxidation of VOCs.

Nighttime Chemistry

After sunset, the chemistry shifts dramatically. With no sunlight, photolysis stops, and ozone is no longer produced. Instead, the ozone remaining from the daytime can react with NO₂ to form the nitrate radical (NO₃).

O₃ + NO₂ → O₂ + NO₃ (Reaction 28)

The NO₃ radical is rapidly destroyed by sunlight, so it is only significant at night. It can react with other species or with NO₂ to form dinitrogen pentoxide (N₂O₅).

NO₃ + NO₂ + M ↔ N₂O₅ + M (Reactions 29, 30)

This nighttime formation of N₂O₅ and its subsequent reaction with water to form nitric acid (Reaction 32) is a crucial pathway that links the urban smog cycle to the regional problem of acid deposition, which we will explore next.