A system absorbs 196 kJ of heat and the surroundings do 117 kJ of work on the system. What is the change in internal energy of the system?

Verified step by step guidance

1. Identify the first law of thermodynamics, which states that the change in internal energy (\( \Delta U \)) of a system is equal to the heat added to the system (\( q \)) plus the work done on the system (\( w \)). This can be expressed as: \( \Delta U = q + w \).
2. Recognize that the problem provides the values for heat absorbed by the system and work done on the system. Here, \( q = 196 \text{ kJ} \) and \( w = 117 \text{ kJ} \).
3. Substitute the given values into the equation: \( \Delta U = 196 \text{ kJ} + 117 \text{ kJ} \).
4. Perform the addition to find the change in internal energy: \( \Delta U = 196 \text{ kJ} + 117 \text{ kJ} \).
5. Conclude that the result from the addition gives the change in internal energy of the system.

Key Concepts

First Law of Thermodynamics

The First Law of Thermodynamics states that energy cannot be created or destroyed, only transformed from one form to another. In a closed system, the change in internal energy is equal to the heat added to the system minus the work done by the system. This principle is fundamental for understanding energy conservation in thermodynamic processes.

Internal Energy

Internal energy is the total energy contained within a system, encompassing both kinetic and potential energy at the molecular level. It is a state function, meaning it depends only on the current state of the system, not on how it reached that state. Changes in internal energy are crucial for analyzing thermodynamic processes.

Work and Heat Transfer

Work and heat are two primary ways energy is transferred into or out of a system. Work is done on the system when external forces act on it, while heat is energy transferred due to a temperature difference. Understanding how to quantify these transfers is essential for calculating changes in internal energy using the First Law of Thermodynamics.