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Chapter 7: Problem 14

What is the difference between \(\Delta H\) and \(\Delta E ?\)

Short Answer

Expert verified
The primary difference between 螖H and 螖E lies in their representation of different energy aspects of a system. 螖H, the change in enthalpy, measures heat absorbed or released at constant pressure during a chemical reaction. 螖E, the change in internal energy, focuses on the total internal energy change of the system. The relationship between them is given by the equation: \(\Delta H = \Delta E + P\Delta V\). The difference between these two quantities depends on the process and conditions under which it occurs.

Step by step solution

01

Definition of 螖H

螖H, also known as the change in enthalpy, represents the heat absorbed or released during a chemical reaction at constant pressure. It is an extensive property and depends on the number of substances involved in the reaction. Enthalpy is the measure of the total energy of a thermodynamic system, including internal energy, pressure, and volume. When a reaction takes place at constant pressure, the difference in enthalpy of the products and reactants signifies the amount of heat absorbed or released by the reaction.
02

Definition of 螖E

螖E, also known as the change in internal energy, represents the difference in the total internal energy of the system before and after the reaction. Internal energy is a sum of the potential and kinetic energies of the molecules in the system. The change in internal energy of the system (螖E) is determined by the heat added to or removed from the system, and the work done on or by the system.
03

Comparison between 螖H and 螖E

Both 螖H and 螖E are thermodynamic quantities used to describe the energy changes in a system, but they represent different aspects of the system's energy. The primary difference between 螖H and 螖E is that 螖H is specifically concerned with the heat changes at constant pressure, whereas 螖E is concerned with the total change in internal energy. The relationship between 螖H and 螖E can be described by the equation: \[\Delta H = \Delta E + P\Delta V\] Where P is the constant external pressure, and 螖V is the change in volume of the system. When a reaction occurs at constant pressure, if there is no significant change in volume, 螖H and 螖E values are almost the same. However, if there is a noticeable change in volume during the reaction, then the 螖H and 螖E values can be significantly different due to the work done on or by the system. In conclusion, 螖H represents the change in enthalpy (heat content) of a system at constant pressure, while 螖E represents the change in the total internal energy of the system. The difference between the two scales depends on the particular process taking place and the conditions under which it occurs.

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Key Concepts

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

Enthalpy Change (螖H)
Enthalpy change, denoted as 螖H, is a central concept in thermodynamics that helps us understand energy changes in chemical reactions. It measures the heat absorbed or released when a process occurs at constant pressure. This is important because many reactions we observe in real life take place in open systems, such as in labs or in nature, where pressure remains constant.

Key characteristics of 螖H include:
  • It is an extensive property, meaning it depends on the amount of substance involved.
  • 螖H accounts for heat changes but does not include work done due to pressure and volume changes.
  • When 螖H is positive, heat is absorbed (endothermic reaction); when negative, heat is released (exothermic reaction).
Understanding enthalpy change is crucial for predicting whether a reaction will give off energy as heat or require energy input. This concept is widely used in fields like chemistry, physics, and engineering to analyze and optimize chemical processes.
Internal Energy Change (螖E)
The internal energy change, represented by 螖E, relates to the total energy changes within a system during a chemical reaction. It includes both kinetic and potential energy of the molecules, influenced by heat exchange and work done on or by the system.

Important points about 螖E:
  • 螖E is comprehensive, including both heat (q) and work (w), formatted as 螖E = q + w.
  • It illustrates the conservation of energy principle, encapsulating all energy changes within the system.
  • The value of 螖E can change even if no heat is transferred, if work is done on or by the system.
Considering 螖E helps us understand energetic transactions from a broader perspective, beyond mere heat exchange. It is instrumental in system analysis, especially in closed systems where volume may change, affecting internal energy.
Energy Changes in Chemical Reactions
Energy changes in chemical reactions are fundamental to understanding how substances interact and transform. These changes are often assessed through 螖H and 螖E, which offer different perspectives of the process.

Let's break it down:
  • When analyzing a reaction at constant pressure with minimal volume change, 螖H provides a good measure of the heat flow.
  • In cases where work done involves significant volume changes, 螖E becomes vital to account for both heat and work contributions.
  • The relationship between 螖H and 螖E is expressed as: \[螖H = 螖E + P螖V\]which adjusts for work done against external pressure.
Understanding these energy metrics illuminates why some reactions release energy, making them feel warm, while others absorb energy, causing them to feel cold. This knowledge is indispensable in fields like thermodynamics, material science, and practical applications such as energy production and chemical engineering.

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Most popular questions from this chapter

What is incomplete combustion of fossil fuels? Why can this be a problem?

The enthalpy change for a reaction is a state function and it is an extensive property. Explain.

Standard enthalpies of formation are relative values. What are \(\Delta H_{\mathrm{f}}^{\circ}\) values relative to?

Consider the following reaction: $$2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l) \quad \Delta H=-572 \mathrm{kJ}$$ a. How much heat is evolved for the production of 1.00 mole of \(\mathrm{H}_{2} \mathrm{O}(l) ?\) b. How much heat is evolved when 4.03 g hydrogen are reacted with excess oxygen? c. How much heat is evolved when \(186 \mathrm{g}\) oxygen are reacted with excess hydrogen?

Combustion of table sugar produces \(\mathrm{CO}_{2}(g)\) and \(\mathrm{H}_{2} \mathrm{O}(l) .\) When \(1.46 \mathrm{g}\) table sugar is combusted in a constant-volume (bomb) calorimeter, \(24.00 \mathrm{kJ}\) of heat is liberated. a. Assuming that table sugar is pure sucrose, \(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}(s)\) write the balanced equation for the combustion reaction. b. Calculate \(\Delta E\) in \(\mathrm{kJ} / \mathrm{mol} \mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}\) for the combustion reaction of sucrose.
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