Question

A real gas behaves like an ideal gas if its A) pressure and temperature are both high B) pressure and temperature are both low C) pressure is high and temperature is low D) pressure is low and temperature is high

Short answer

A real gas behaves most like an ideal gas under conditions of low pressure and high temperature, which corresponds to option D).

Step-by-step solution

Understanding Ideal Gas Behavior

An ideal gas is a theoretical gas composed of a set of randomly-moving, non-interacting point particles. The behavior of an ideal gas is described by the Ideal Gas Law, which relates the pressure (P), volume (V), temperature (T), and number of moles (n) in the following equation: PV = nRT, where R is the universal gas constant. For a real gas to behave like an ideal gas, the interactions between the particles and the volume occupied by the gas particles must be negligible.

Conditions for Ideal Gas Behavior

In order for the particles' interactions to be insignificant, the particles must be far apart relative to their size, which happens at low pressures. At high temperatures, the kinetic energy of the particles is so high that the effect of intermolecular forces becomes less significant. Therefore, for a real gas to behave like an ideal gas, the pressure should be low, and the temperature should be high.

Choosing the Correct Option

According to the Ideal Gas Law and the behavior of real gases, the correct conditions for real gas to behave like an ideal gas are low pressure and high temperature. Thus, the correct answer is D) pressure is low and temperature is high.

Key concepts

Ideal Gas Law

The Ideal Gas Law is a cornerstone of chemical and physical understanding, providing a clear formula that connects the state variables of a gas under theoretical ideal conditions. It is written as the equation: \( PV = nRT \), where \(P\) stands for pressure, \(V\) denotes volume, \(T\) is temperature, and \(n\) is the number of moles of the gas, with \(R\) being the universal gas constant. The beauty of this law is that it presupposes a gas where the molecules do not interact and occupy no space.

For students improving their grasp of this law, it's crucial to internalize that it describes an ideal scenario which rarely exists in the real world. However, under certain conditions, real gases can exhibit behavior closely approximating this ideal state. This is particularly true when the gas particles are spaced out enough that their individual volumes and the forces between them are negligible in comparison to the entire system. Such conditions are typically found at low pressures, which enhance the gas particles' separation, and high temperatures, which provide the particles with enough energy to overcome any possible attractions between them.

Real Gases

While the Ideal Gas Law offers a simplified model, real gases display behaviors that deviate from the ideal due to factors like intermolecular forces and the actual volume of the gas particles. The behavior of real gases can be approximated by the Van der Waals equation, which provides corrections for the volume occupied by gas molecules and the attractive forces between them.

Understanding the differences between ideal and real gases is a stepping stone to mastering topics in both chemistry and physics. To approach this, students should consider how variables such as pressure and temperature affect a real gas. At high pressures, gas particles are forced closer together, which makes the volume they occupy and the interactions between them more significant. Conversely, at low temperatures, particles move slower, and thus, intermolecular forces have a more pronounced effect. It's under the opposite conditions—namely low pressure and high temperature—that real gases can mimic the ideal behavior, as particles are both separated and energetically able to defy attraction.

Intermolecular Forces

When studying gases, one cannot overlook intermolecular forces—the forces of attraction or repulsion which act between neighboring particles, such as atoms, molecules, or ions. These forces are varied, including London dispersion forces, dipole-dipole interactions, and hydrogen bonding. They are pivotal in determining the physical properties of substances and play a significant role in how real gases behave.

For a more intuitive understanding, imagining the particles as individual people in a large room can be helpful. If the room is cold or crowded (high pressure, low temperature), people are more likely to huddle together, similar to how particles experience stronger intermolecular forces. But if the room is warm and spacious (low pressure, high temperature), people tend to spread out, akin to how gas particles behave more ideally. Thus, by relating these concepts to everyday situations, students can better visualize and comprehend the intricacies of intermolecular forces and their impact on gas behavior.