Question

The combustion of methane, \(\mathrm{CH}_{4}\), is an exothermic process. Therefore, the products of this reaction must possess (higher/lower) total potential energy than do the reactants.

Short answer

The products of the methane combustion reaction possess LOWER total potential energy than the reactants, as energy is released in the exothermic process, transferring potential energy from the reactants to the surroundings in the form of heat.

Step-by-step solution

Understand exothermic processes

Exothermic processes are chemical reactions that release energy in the form of heat. This means that during the reaction, energy is transferred from the reactants (higher potential energy) to the surroundings (lower potential energy) as heat.

Analyze the combustion of methane

The combustion of methane (\(\mathrm{CH}_{4}\)) is a well-known exothermic reaction represented by the chemical equation: \[\mathrm{CH}_{4} + 2\mathrm{O}_{2} \rightarrow \mathrm{CO}_{2} + 2\mathrm{H}_{2}\mathrm{O} + \text{heat}\] Here, the reactants are methane and oxygen gas, while the products are carbon dioxide, water, and heat. Since heat is released during the reaction, it implies that potential energy is transferred from the reactants to the surroundings, resulting in products with lower potential energy.

Compare potential energy of reactants and products

In an exothermic process like the combustion of methane, energy is released as heat. Therefore, the reactants have higher potential energy as compared to the products, given that energy is transferred from the reactants to the surroundings.

Answer the question

Based on our understanding of exothermic processes and potential energy, the products of the methane combustion reaction must possess LOWER total potential energy than the reactants.

Key concepts

Combustion of Methane

Understanding the combustion of methane is key to grasping basic principles of chemistry and energy transfer. Methane \text{(\text{CH}_4)} is a common fuel source, and when it burns, it combines with oxygen (\text{O}_2) in the air to produce carbon dioxide (\text{CO}_2), water (\text{H}_2\text{O}), and energy, primarily in the form of heat. This reaction can be summarized as:
\[\begin{equation}\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O} + \text{heat}\text{\}\text{\}\text{\}\end{equation}\]
Students learning about this reaction are often struck by its immediate relevance, as it's similar to the burning of natural gas in stoves and heaters. The fascinating part is that despite the simplicity of the methane molecule, when it reacts with \text{O}_2, a significant amount of heat is released, making it an important source of energy in various applications. This makes the combustion of methane a cornerstone example of an exothermic reaction—a concept that will be explored further in the next section.

Chemical Reactions

Chemical reactions are processes where reactants transform into products through the breaking and forming of chemical bonds. These transformations can either absorb energy (endothermic) or release energy (exothermic). In the case of exothermic reactions, such as the combustion of methane, energy stored within chemical bonds is released into the surroundings, usually as heat or light.
Chemists use chemical equations to represent these reactions, where the left side shows the reactants and the right side shows the products. Various factors, including temperature, pressure, and the presence of catalysts, can influence the rate and direction of these reactions. It's essential for students to understand that energy changes are inherently tied to chemical reactions and that these changes can be detected and measured in a laboratory setting.

Potential Energy

Potential energy, in the context of chemical reactions, refers to the stored energy within the chemical bonds of molecules. The more energy that's stored, the higher the potential energy of the substance. During chemical reactions, the potential energy can change as bonds break and new ones form. An exothermic reaction, like the combustion of methane, results in products that have lower potential energy than the reactants.
This is because the energy that was held in the methane and oxygen molecules is released, mostly as heat, reducing the total potential energy remaining in the newly formed substances. It’s a fundamental principle that students must remember: the energy decrease in exothermic reactions doesn't vanish; instead, it's transferred to the surroundings, providing heat or light that can be harnessed for practical use. This interplay of potential energy and energy transfer is a central theme in understanding how energy moves and is transformed in physical and chemical processes.