First law of thermodynamics

The first law of thermodynamics is a fundamental concept in the study of chemistry and physics. It is also known as the law of conservation of energy. This law states that energy cannot be created or destroyed in an isolated system. This means that the total energy of an isolated system remains constant. Instead, energy can only be converted from one form to another or transferred from one part of the system to another. Let's understand the first law of thermodynamics in depth.

Understanding energy

Before we dive into the first law of thermodynamics, it is important to understand the concept of energy. Energy is the capacity to do work or produce heat. It exists in various forms, including kinetic energy, potential energy, thermal energy, chemical energy and electrical energy, etc. The unit of measurement of energy in the International System of Units (SI) is the joule (J).

Kinetic energy is the energy an object has because of its motion. For example, a moving car or a flowing river has kinetic energy. The formula to calculate kinetic energy is:

KE = 0.5 * m * v²

Where m is the mass of the object and v is its velocity.

Potential energy is the energy stored in an object because of its position or state. For example, a drawn bow or a book on a shelf has potential energy. One form of potential energy is gravitational potential energy, which is calculated as follows:

PE = m * g * h

where m is the mass, g is the acceleration due to gravity, and h is the height above the reference point.

Stated in the law

Now that we have a basic understanding of energy, let's explain the first law of thermodynamics in a simple way: The change in the internal energy of a system is equal to the heat added to the system minus the work done on the system.

Mathematically, the first law of thermodynamics is expressed as follows:

ΔU = Q - W

Where:

• ΔU is the change in the internal energy of the system.
• Q is the heat added to the system.
• W is the work done by the system.

This equation implies that when a system absorbs heat (Q is positive), its internal energy increases. Conversely, if the system does work (W is positive), its internal energy decreases.

Examples of the first law

Let us consider some scenarios to better understand how the first law of thermodynamics works in different systems.

Example 1: Gas in a cylinder

Imagine a gas enclosed in a cylinder containing a piston. If we heat the gas, energy is transferred to the gas in the form of heat (Q). This causes the gas to expand and does work (W) on the piston by pushing it upward. According to the first law, the change in internal energy (ΔU) of the gas depends on the heat added and the work done.

ΔU = Q - W

If 500 joules of heat energy is added to the gas, and the gas does 200 joules of work on the piston, the change in internal energy will be:

ΔU = 500 J - 200 J = 300 J

Thus, the internal energy of the gas increased by 300 joules.

Example 2: Boiling water in a closed vessel

Consider boiling water in a closed, rigid container. As heat is added to the system, the temperature of the water increases. However, since the container is rigid and cannot expand, no work is done (W = 0). Therefore, any heat added to the system goes directly into increasing the internal energy of the system.

ΔU = Q - 0 = Q

If 1000 joules of heat is added then the change in internal energy will also be 1000 joules.

Visualizing the first law

Let's see how these concepts can be visualized. Consider the simple example of a gas in a cylinder with a moving piston, which helps to visualize the transformation and transfer of energy:

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