Combining Intermediate Steps How To Handle Oxygen Molecules In Chemical Reactions

Introduction

In the realm of chemistry, understanding reaction mechanisms is crucial for comprehending how chemical reactions occur. Often, complex reactions proceed through a series of elementary steps, known as intermediate steps, before reaching the final products. These intermediate steps involve the formation of short-lived species called intermediates, which are subsequently consumed in the following steps. Combining these intermediate steps allows us to visualize the overall reaction process and stoichiometry. This article delves into the concept of combining intermediate steps, using a specific example involving nitrogen and oxygen molecules to illustrate the process. Our focus will be on how to handle molecules, particularly oxygen, when combining these steps to derive the overall balanced chemical equation.

Understanding Intermediate Steps in Chemical Reactions

Chemical reactions rarely occur in a single, concerted step. Instead, they often proceed through a series of elementary steps that collectively describe the reaction mechanism. Each step involves the collision and rearrangement of atoms and molecules, leading to the formation of intermediates. Intermediates are transient species that are formed in one step and consumed in a subsequent step. They do not appear in the overall balanced chemical equation, as they are not reactants or products of the overall reaction. Recognizing and understanding intermediate steps are pivotal in elucidating the reaction pathway and kinetics.

In the given scenario, we have two intermediate steps:

1. $N_2(g) + O_2(g) ightarrow 2NO(g)$
2. $2NO(g) + O_2(g) ightarrow 2NO_2(g)$

In the first step, nitrogen gas ($N_2$) reacts with oxygen gas ($O_2$) to produce nitrogen monoxide ($NO$). This reaction is endothermic and requires high temperatures to proceed effectively. In the second step, the nitrogen monoxide formed in the first step reacts with additional oxygen gas to form nitrogen dioxide ($NO_2$). Nitrogen monoxide ($NO$) acts as an intermediate in this reaction mechanism. It is produced in the first step and consumed in the second step. Therefore, it will not appear in the overall balanced equation.

Combining Intermediate Steps: A Step-by-Step Approach

To obtain the overall balanced chemical equation, we need to combine the intermediate steps. This involves adding the reactions together and canceling out any species that appear on both sides of the equation (i.e., intermediates). The procedure is as follows:

1. Write down all the intermediate steps:

   • $N_2(g) + O_2(g) ightarrow 2NO(g)$
   • $2NO(g) + O_2(g) ightarrow 2NO_2(g)$

2. Identify any species that appear on both the reactant and product sides of the equations. In this case, $2NO(g)$ appears on the product side of the first equation and the reactant side of the second equation. Since it's an intermediate, it will be canceled out.

3. Add the equations together, canceling out the intermediate species:

   $N_2(g) + O_2(g) + 2NO(g) + O_2(g) ightarrow 2NO(g) + 2NO_2(g)$

4. Simplify the equation by canceling out the common species ($2NO(g)$) and combining like terms:

   $N_2(g) + O_2(g) + O_2(g) ightarrow 2NO_2(g)$

   Which simplifies to:

   $N_2(g) + 2O_2(g) ightarrow 2NO_2(g)$

The overall balanced chemical equation is: $N_2(g) + 2O_2(g) ightarrow 2NO_2(g)$.

This equation represents the overall reaction, showing that nitrogen gas reacts with two molecules of oxygen gas to produce two molecules of nitrogen dioxide. The intermediate, nitrogen monoxide, does not appear in the overall equation, as it is both produced and consumed during the reaction.

The Role of Oxygen Molecules in the Reaction

In this reaction, oxygen molecules ($O_2$) play a crucial role as a reactant. In the first step, one molecule of oxygen reacts with one molecule of nitrogen to form two molecules of nitrogen monoxide. In the second step, another molecule of oxygen reacts with two molecules of nitrogen monoxide to form two molecules of nitrogen dioxide. Therefore, in the overall reaction, two molecules of oxygen are required for every molecule of nitrogen.

When combining the intermediate steps, it is essential to account for the stoichiometry of the oxygen molecules. In this case, we have one molecule of oxygen in the first step and one molecule of oxygen in the second step. When we add the equations together, we combine these oxygen molecules to obtain two molecules of oxygen in the overall balanced equation.

To best describe what Jason should do with the oxygen molecules, he should combine them on the reactant side of the overall equation. This means adding the oxygen molecules from each intermediate step together. In this specific example, one $O_2$ molecule from the first step combines with one $O_2$ molecule from the second step, resulting in a total of two $O_2$ molecules in the balanced overall equation.

Conclusion

Combining intermediate steps is a fundamental technique in chemistry for understanding complex reaction mechanisms. By identifying intermediates and canceling them out, we can derive the overall balanced chemical equation, which represents the stoichiometry of the reaction. In the example discussed, the reaction between nitrogen and oxygen proceeds through two intermediate steps, involving the formation of nitrogen monoxide as an intermediate. By correctly combining these steps and accounting for the oxygen molecules, we arrive at the overall balanced equation: $N_2(g) + 2O_2(g) ightarrow 2NO_2(g)$. This process highlights the importance of understanding reaction mechanisms and the role of intermediates in chemical reactions.

In summary, when Jason combines the two intermediate steps, he must carefully consider the oxygen molecules involved in each step. The correct approach is to combine the oxygen molecules from each step on the reactant side to derive the overall balanced equation. This method ensures that the stoichiometry of the reaction is accurately represented, providing a comprehensive understanding of the chemical transformation.