When we talk about gases, we often start with something called the Ideal Gas Law. This law simplifies things a lot and is written like this: .
This formula assumes that gas particles don’t interact with each other and are really, really tiny. But, in the real world, gases don’t always act this way, especially under certain conditions. Let’s look at how real gases differ from what the Ideal Gas Law suggests, using a concept called Kinetic Molecular Theory.
The Ideal Gas Law assumes that gas particles don’t take up any space. But in reality, gas particles do have size. This size matters when there is high pressure.
When we push a gas into a smaller space, the room that the particles actually need becomes noticeable. This means that the volume we measure is smaller than what the Ideal Gas Law predicts.
Another idea in the Ideal Gas Law is that gas particles don’t affect each other; they just move around freely. However, in real life, especially when it’s cooler and the pressure is high, gas particles can attract or repel each other. These forces are called intermolecular forces, like van der Waals forces or dipole-dipole interactions.
Because of these forces, particles can stick together a little bit. This makes the pressure lower than what we expect based on the formula , since not all particles are hitting the walls of the container like they should.
Also, when temperatures drop, gas particles have less energy. This means they move around less and start to interact with each other more. The closer these particles get, the stronger those intermolecular forces become. This causes even more differences from ideal behavior.
In short, real gases don’t behave like the Ideal Gas Law says for two main reasons:
When we have conditions that are very different from normal, like really high pressure or very low temperatures, these effects become even clearer. Understanding how real gases behave is important because it helps us use the Ideal Gas Law correctly and know when we need to make changes for real-life situations.
When we talk about gases, we often start with something called the Ideal Gas Law. This law simplifies things a lot and is written like this: .
This formula assumes that gas particles don’t interact with each other and are really, really tiny. But, in the real world, gases don’t always act this way, especially under certain conditions. Let’s look at how real gases differ from what the Ideal Gas Law suggests, using a concept called Kinetic Molecular Theory.
The Ideal Gas Law assumes that gas particles don’t take up any space. But in reality, gas particles do have size. This size matters when there is high pressure.
When we push a gas into a smaller space, the room that the particles actually need becomes noticeable. This means that the volume we measure is smaller than what the Ideal Gas Law predicts.
Another idea in the Ideal Gas Law is that gas particles don’t affect each other; they just move around freely. However, in real life, especially when it’s cooler and the pressure is high, gas particles can attract or repel each other. These forces are called intermolecular forces, like van der Waals forces or dipole-dipole interactions.
Because of these forces, particles can stick together a little bit. This makes the pressure lower than what we expect based on the formula , since not all particles are hitting the walls of the container like they should.
Also, when temperatures drop, gas particles have less energy. This means they move around less and start to interact with each other more. The closer these particles get, the stronger those intermolecular forces become. This causes even more differences from ideal behavior.
In short, real gases don’t behave like the Ideal Gas Law says for two main reasons:
When we have conditions that are very different from normal, like really high pressure or very low temperatures, these effects become even clearer. Understanding how real gases behave is important because it helps us use the Ideal Gas Law correctly and know when we need to make changes for real-life situations.