Problem 41 - Chemistry The Molecular Science

Question

Is this reaction predicted to favor the products at low temperatures, at high temperatures, or both? Explain your answer briefly. \(\mathrm{Mg}(\mathrm{s})+\frac{1}{2} \mathrm{O}_{2}(\mathrm{~g}) \longrightarrow \mathrm{MgO}(\mathrm{s}) \quad \Delta_{\mathrm{r}} H^{\circ}=-601.70 \mathrm{~kJ} / \mathrm{mol}\)

Step-by-Step Solution

Answer

The reaction favors products at low temperatures due to its exothermic nature.

Step 1: Analyze the Reaction Type

The given reaction shows magnesium (\(\mathrm{Mg}(\mathrm{s})\) reacting with oxygen (\(\mathrm{O}_2(\mathrm{g})\) to form magnesium oxide (\(\mathrm{MgO}(\mathrm{s})\). This is a synthesis reaction where a metal reacts with oxygen to form a metal oxide. The process is typically exothermic as indicated by the negative enthalpy change (\(\Delta_{\mathrm{r}}\H^{\circ} = -601.70 \mathrm{~kJ} / \mathrm{mol}\)).

Step 2: Apply the Concept of Exothermic Reactions

In exothermic reactions, energy is released as heat. A large negative enthalpy, such as \(-601.70 \mathrm{kJ/mol}\), suggests that the reaction releases a significant amount of energy, making it exothermic. Exothermic reactions are typically favored at low temperatures because at lower temperatures, the energy released helps the system reach a more stable state.

Step 3: Consider Temperature Effects

According to Le Chatelier's principle, decreasing the temperature will shift the equilibrium towards the products in an exothermic reaction. Since the reaction releases heat, lowering the temperature emphasizes the production of more heat (i.e., more product formation) to counterbalance the temperature change.

Step 4: Conclusion on Temperature Favorability

The strong exothermic nature of the reaction, combined with the principles of equilibrium, suggest that the reaction favors the formation of products at low temperatures. At high temperatures, the reverse reaction might become more favorable due to the application of Le Chatelier's principle, as the system works to absorb excess heat.

Key Concepts

Synthesis Reaction

A synthesis reaction is a fundamental type of chemical reaction where two or more simple substances combine to form a more complex product. In the case of our example, magnesium (\(\mathrm{Mg}(\mathrm{s})\)) reacts with oxygen (\(\mathrm{O}_2(\mathrm{g})\)) to create magnesium oxide (\(\mathrm{MgO}(\mathrm{s})\)). This type of reaction is characterized by the combination of its reactants into one chemical compound.
These reactions are often depicted by the general formula \(A + B \rightarrow AB\), where \(A\) and \(B\) are reactants and \(AB\) is the product.

• Synthesis reactions are common in the formation of metal oxides from metals and oxygen.
• They are usually accompanied by energy changes, either releasing or absorbing heat.
• Synthesis reactions are instrumental in industries for forming necessary compounds.

In the day-to-day, these reactions are evident when rust (iron oxide) forms and when plants photosynthesize, showcasing their importance across various contexts.

Le Chatelier's Principle

Le Chatelier's Principle is a fascinating concept used to predict the effect of a change in conditions on a chemical equilibrium. When a system at equilibrium is disturbed by changing the conditions (such as temperature, pressure, or concentration of reactants or products), the equilibrium will shift to counteract the change and restore a new balance.
Applying this principle to our synthesis reaction, we have an exothermic process, which means it releases heat.

• If the temperature is reduced, the system combats this change by producing more products to release more heat, enhancing product formation.
• Conversely, increasing the temperature pushes the equilibrium towards reactants, as the system absorbs extra heat favoring the reverse reaction.

Utilizing Le Chatelier's Principle in exothermic reactions provides predictions about how external temperature changes can affect the yield of the reaction. Understanding this principle helps us optimize conditions for industrial chemical production and manage reactions effectively.

Enthalpy Change

Enthalpy change, represented as \(\Delta H\), reflects the heat change in a reaction at constant pressure. It is a critical factor in determining whether a reaction is exothermic or endothermic. The magnesium oxide formation reaction exhibits an enthalpy change (\(\Delta_r H^{\circ} = -601.70 \mathrm{~kJ/mol}\)), signifying it is exothermic.
Exothermic reactions have a negative \(\Delta H\), demonstrating that they release energy to their surroundings. This release of energy makes the products more stable compared to the reactants.

• Reactions with large negative \(\Delta H\) values signify substantial energy release.
• Enthalpy change helps in understanding not only the energy dynamics but also the temperature suitability of the reaction.
• Analyzing \(\Delta H\) is crucial to industrial applications, as managing energy changes effectively can enhance reaction efficiency.

Knowing the enthalpy change of reactions assists scientists and engineers in designing reactors and processes that maximize yield, while minimizing environmental impact.