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Chapter 9: Problem 33

Why is the sign of \(\Delta H\) negative for an exothermic process?

Short Answer

Expert verified
Answer: The sign of 螖H for an exothermic process is negative. This is because heat energy is released to the surroundings during the process, causing the energy of the products to be lower than the energy of the reactants. As a result, when calculating 螖H (H_products - H_reactants), the value is negative.

Step by step solution

01

Definition of Exothermic Process

An exothermic process is a chemical reaction or physical process that releases heat to the surroundings. In an exothermic process, the energy stored in chemical bonds is transformed into heat as a result of the reaction, and the system loses energy (in the form of heat) to the surroundings.
02

Enthalpy Change 螖H

The enthalpy change, represented as 螖H, is a measure of the energy change within a system during a chemical reaction or a physical process. It represents the difference between the products' enthalpy and the reactants' enthalpy, expressed mathematically as: 螖H = H_products - H_reactants
03

Energy Flow in Exothermic Processes

In an exothermic process, heat energy is released from the system to the surroundings, meaning the energy of the products is lower than the energy of the reactants. This occurs because energy stored in chemical bonds gets transformed into heat as the reaction progresses, leading to a decrease in the total energy within the system.
04

Sign of 螖H for Exothermic Processes

Since the total energy of the products is lower than the total energy of the reactants in an exothermic process, the difference in enthalpy (螖H) between them will be negative. By calculating the difference in enthalpy, 螖H = H_products - H_reactants, we find that 螖H is negative for an exothermic process because H_products < H_reactants. As a result, we can conclude that the sign of 螖H is negative for an exothermic process because heat energy is released to the surroundings, leading to a decrease in the system's energy.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Enthalpy Change
Understanding enthalpy change is crucial to grasping exothermic processes. Enthalpy, denoted as \( H \), is a measure of the total heat content in a thermodynamic system. It includes both internal energy and the energy required to make room for the system by displacing its surroundings. When a chemical reaction or physical process occurs, the change in enthalpy \( \Delta H \) signifies the heat absorbed or released under constant pressure. It's computed as the difference between the enthalpy of the products and the reactants:\[\Delta H = H_{\text{products}} - H_{\text{reactants}}\]In an exothermic process, \( \Delta H \) is negative because more energy is released in forming the products than is consumed in breaking the reactants' bonds. This release of energy leads to a decrease in the system's enthalpy.
Energy Flow
Energy flow in chemical reactions clarifies why some reactions feel hot to the touch. In an exothermic reaction, energy moves from the system to its surroundings, primarily in the form of heat. Imagine a campfire; as the wood burns, it releases heat energy, raising the temperature of the surrounding air. Likewise, in an exothermic chemical reaction, the system's energy decreases because it's transferring heat to its surroundings.This energy flow is crucial for understanding why exothermic processes have a negative \( \Delta H \). The energy stored within chemical bonds in the reactants is greater than that in the products. Thus, as bonds form, excess energy is released, causing the system's total energy to drop.
Chemical Bonds
Chemical bonds are the essential links between atoms in molecules, storing potential energy. When bonds break and form during a reaction, energy is either absorbed or released. In an exothermic reaction, the energy needed to break the bonds in the reactants is less than the energy released when new bonds form in the products. This results in a net release of energy鈥攁s chemical bonds are rearranged, heat gets liberated, warming the surroundings.Key points to remember are:
  • Breaking bonds always requires energy.
  • Forming bonds releases energy.
  • In exothermic reactions, new bonds release more energy than the old bonds consumed.
This dynamic balance is why exothermic reactions are typically spontaneous, proceed faster, and result in a negative \( \Delta H \).

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Most popular questions from this chapter

When measuring the heat of combustion of a very small amount of material, would you prefer to use a calorimeter having a heat capacity that is small or large?

In a high-temperature gas-phase reaction, methanol \(\left(\mathrm{CH}_{3} \mathrm{OH}\right)\) reacts with \(\mathrm{N}_{2}\) to produce \(\mathrm{HCN}\) and \(\mathrm{NH}_{3}\). The reaction is endothermic, requiring \(164 \mathrm{kJ}\) of thermal energy per mole of methanol under standard conditions. a. Write a balanced chemical equation for this reaction. b. Is energy a reactant or a product? c. What is the change in enthalpy under standard conditions if \(60.0 \mathrm{g}\) of \(\mathrm{CH}_{3} \mathrm{OH}(g)\) reacts with excess \(\mathrm{N}_{2}(g),\) forming \(\mathrm{HCN}(g),\) and \(\mathrm{NH}_{3}(g) ?\)

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