Question

Nitrogen dioxide \(\left(\mathrm{NO}_{2}\right)\) is the only important gaseous species in the lower atmosphere that absorbs visible light. (a) Write the Lewis structure(s) for \(\mathrm{NO}_{2}\). (b) How does this structure account for the fact that \(\mathrm{NO}_{2}\) dimerizes to form \(\mathrm{N}_{2} \mathrm{O}_{4}\) ? Based on what you can find about this dimerization reaction in the text, would you expect to find the \(\mathrm{NO}_{2}\) that forms in an urban environment to be in the form of dimer? Explain. (c) What would you expect as products, if any, for the reaction of \(\mathrm{NO}_{2}\) with CO? (d) Would you expect \(\mathrm{NO}_{2}\) generated in an urban environment to migrate to the stratosphere? Explain.

Short answer

In summary, NO2 has a Lewis structure with one double-bonded O, one single-bonded O, and an unpaired electron on N. Dimerization occurs through the reaction 2NO2 ⟶ N2O4, but it may not be found primarily as a dimer in urban environments due to various factors. NO2 reacts with CO to form NO and CO2. Migration of NO2 generated in urban environments to the stratosphere is less likely due to its reactivity, low concentration, and heavier mass compared to the major constituents of the atmosphere.

Step-by-step solution

(Step 1: Writing the Lewis structure of NO2)

To create the Lewis structure for NO2: 1. Count the total number of valence electrons: N has 5, and each O has 6, thus we have 5 + 2(6) = 17 electrons. 2. Identify the central atom, which is Nitrogen in this case. 3. Place two electrons between the central atom and each outer atom, resulting in two electron pairs acting as a bond between N and each O atom. 4. Complete the octets for the outer atoms (O atoms). 5. Distribute the remaining electrons in order to give the central atom an octet if possible. After setting up the Lewis structure, we have one double-bonded O, one single-bonded O, and one unpaired electron on N.

(Step 2: Accounting for Dimerization)

The presence of an unpaired electron on nitrogen in NO2 makes it a free radical, which is highly reactive. Thus, NO2 can dimerize through the reaction: 2NO2 ⟶ N2O4 This dimerization reaction balances the unpaired electrons and forms a more stable molecule - the dimer N2O4. In an urban environment, NO2 is in equilibrium with N2O4, and various factors (temperature, pressure, concentration) could affect the amount of dimer. However, due to the high reactivity of NO2 and the presence of other reactive species in the urban atmosphere, it is less likely to find the majority of NO2 in the form of N2O4.

(Step 3: Predicting the Reaction with CO)

The reaction between NO2 and CO can be represented as follows: NO2 + CO ⟶ Products Considering the fact that NO2 is an oxidizing agent and CO is a reducing agent, we can expect the following redox reaction to occur: NO2 + CO ⟶ NO + CO2 In this reaction, CO is oxidized to CO2, and NO2 is reduced to NO.

(Step 4: Discussing the Migration of NO2 to the Stratosphere)

Nitrogen dioxide (NO2) is mainly produced in the lower atmosphere due to human activities such as combustion of fossil fuels in automobiles and industries. As a reactive species, NO2 takes part in various chemical reactions with other species present in the urban environment. Hence, its removal from the lower atmosphere occurs relatively quickly. Considering its reactivity, low concentration, and the fact that NO2 is heavier than the major constituents of the atmosphere (oxygen and nitrogen), it is less likely for the NO2 generated in an urban environment to migrate to the stratosphere in substantial amounts.

Key concepts

Lewis structure of NO2

Understanding the Lewis structure of nitrogen dioxide (NO2) forms the basis for comprehending its chemistry. The process begins by counting the valence electrons which, for NO2, totals 17. Nitrogen being the central atom shapes the connection with two oxygen atoms. We arrive at a structure where nitrogen has one single bond, one double bond with oxygen atoms, and a single unpaired electron that contributes to NO2's reactivity as a free radical.

The presence of this unpaired electron is vital; it indicates that NO2 can engage in additional reactions – a point often missing in step-by-step solutions but crucial for appreciating the molecule's behavior in various environments.

Dimerization of NO2

The tendency of NO2 to dimerize into dinitrogen tetroxide (N2O4) is a fascinating aspect of its chemistry. The dimerization is a response to the unstable free radical nature of NO2, where two NO2 molecules react to pair their unpaired electrons, resulting in a more stable N2O4 molecule. This equilibrium, however, is affected by temperature and other atmospheric conditions.

Urban Atmosphere and Dimerization

In an urban setting, where the concentration of NO2 is high, dimerization competes with various other reactions. The exercise often suggests a steady equilibrium, yet in reality, the presence of sunlight, pollutants, and temperature variations disrupt this, leading to less N2O4 than one might expect.

NO2 reaction with CO

Reactions involving NO2, such as its interaction with carbon monoxide (CO), are essential for understanding air quality control measures. NO2 is an oxidizing agent, while CO acts as a reducing agent. When they react, one typically observes a redox process resulting in nitrogen oxide (NO) and carbon dioxide (CO2).

This chemical reaction is crucial in urban pollution chemistry. It demonstrates how NO2 plays a role in the oxidization of harmful pollutants like CO, emphasizing why mitigating NO2 emissions is a key environmental goal. Teaching content often underestimates such practical implications, but they are central for students understanding the relevance of these chemical reactions.

NO2 in the Stratosphere

The movement of NO2 into the stratosphere involves complex atmospheric dynamics not fully explored in many classroom materials. Generated predominantly in the lower atmosphere through human activities, NO2 is unlikely to rise into stratosphere layers due to its reactivity and heaviness compared to the main components of air.

Implications for the Ozone Layer

While NO2 is a minor player in stratospheric chemistry, it's important for students to recognize its potential impact if it does reach those heights. It can participate in ozone depletion processes, which is a crucial reason why scientists monitor its concentration and behavior in the atmosphere.