The reaction \(\mathrm{N}_{2}+3 \mathrm{H}_{2} \rightarrow 2 \mathrm{NH}_{3}\) is exothermic. Draw a reaction-energy profile for the reaction. Label the gap that represents \(\Delta E_{\mathrm{rxn}} .\)

An exothermic reaction is a chemical process where energy is released into the surrounding environment, usually in the form of heat. This type of reaction is characterized by the products having lower energy than the reactants.

To visualize exothermic reactions, we think of them occurring within an energy landscape. Here's how it works: During the reaction, the chemical bonds of the reactants are broken, and new bonds are formed to create the products. Because the new bonds are more stable (releasing energy), the overall energy of the system decreases.

Let's consider the reaction where nitrogen (N_{2}) combines with hydrogen (H_{2}) to form ammonia (NH_{3}). This reaction releases energy, so if we were to draw an energy profile for it, we'd observe that the line representing the energy of the products is lower than that of the reactants—indicating that energy was given off during the reaction. The difference in height between the reactants and products in this energy profile corresponds to Delta E_{rxn}, which is negative for exothermic reactions.

Chemical kinetics is the study of the speeds—or rates—at which chemical reactions occur. It's about understanding how different variables like concentration, temperature, and the presence of catalysts affect the frequency and speed of collisions between reactant molecules.

The speed of a reaction is quantified by the reaction rate, which measures how quickly reactants are converted into products. But not all collisions between molecules result in a reaction; they only do so if the molecules collide with sufficient energy and proper orientation.

In the context of our example, nitrogen and hydrogen molecules collide countless times before they react to form ammonia. The rate at which this happens can be influenced by the conditions of the reaction. For instance, increasing the temperature grants more energy to the molecules, which can lead to more effective collisions and a faster rate of reaction. Kinetics provides a framework for analyzing these factors and predicting how a reaction will progress over time.

Activation energy, denoted as E_{a}, is the minimum amount of energy required to initiate a chemical reaction. Think of it as the energy barrier that reactants must overcome to be converted into products.

In an energy profile, activation energy is visualized as the peak of the curve that represents the transition state—the highest point along the reaction pathway. This peak corresponds to an unstable, intermediate condition known as the activated complex.

Returning to our exothermic reaction, when nitrogen (N_{2}) and hydrogen (H_{2}) molecules collide with sufficient energy to overcome this barrier, they rearrange into ammonia (NH_{3}). The larger the activation energy, the fewer molecules that will have sufficient energy to react at a given temperature. Conversely, a lower activation energy means that more molecules can overcome the barrier, leading to a faster reaction rate.

Many enzymes and catalysts work by lowering the activation energy of a reaction, making the process faster even at lower temperatures. Hence, understanding activation energy is crucial for manipulating reaction rates in both natural and industrial processes.