Question

A gas deviates from ideal behaviour at a high pressure because its molecules (a) have kinetic energy (b) are bound by covalent bonds (c) attract one another (d) show the Tyndall effect.

Short answer

A gas deviates from ideal behavior at high pressure because its molecules attract one another.

Step-by-step solution

Understand Ideal Gas Behavior

An ideal gas is a hypothetical gas whose molecules occupy negligible space and have no interactions, and which consequently obeys the gas laws exactly. Real gases deviate from ideal behavior due to factors that are not present in an ideal gas.

Analyze the Options

Let's examine each answer choice: (a) All molecules have kinetic energy, this does not cause a deviation. (b) The existence of covalent bonds is not related to the behavior of gases themselves. (c) Real gas molecules attract one another, leading to deviations from ideal gas behavior at high pressure. (d) The Tyndall effect is about the scattering of light in colloids or in fine suspensions, and is unrelated to gas behavior.

Identify the Correct Option

Since real gas molecules do indeed attract each other, this intermolecular force causes a deviation from ideal gas behavior when the gas is at high pressure. Therefore, the correct answer is (c) attract one another.

Key concepts

Real Gas Deviations

When studying gases, it's important to recognize that real gases often don't conform to the ideal gas law, especially under conditions of high pressure. Real gases exhibit deviations due to the fact that their molecules occupy space and experience forces of attraction and repulsion among each other. These factors are ignored in the model of an ideal gas, which assumes no volume and no interactive forces between molecules.

In the context of real gases, the volume occupied by the gas molecules becomes significant at high pressures, as the molecules are forced closer together. This decreases the actual volume available for the gas's molecules to move in, which is not accounted for in the ideal gas law. Additionally, the intermolecular forces, which are more pronounced at close range, can cause attraction between molecules, leading to a further deviation from ideal behavior.

To accurately describe the behavior of real gases, modifications to the ideal gas law, like the Van der Waals equation, incorporate terms to correct for these two significant factors. This allows the equation to more closely predict the behavior of real gases under various conditions.

Kinetic Theory of Gases

The kinetic theory of gases provides a molecular-level model for understanding gas behavior. It is based on certain assumptions that simplify the complex motion of gas molecules. According to the kinetic theory, gas molecules are in constant, random motion and this motion is the cause of gas pressure – which is essentially the force exerted by colliding gas molecules against the walls of their container.

The assumptions of the kinetic theory include:

• Molecules are in constant linear motion.
• Collisions between molecules, and with the walls of the container, are perfectly elastic, meaning there is no net loss of energy.
• The size of the gas molecules is negligible compared to the volume of their container.
• There are no intermolecular forces acting between the molecules.

While these assumptions work well at predicting the behavior of an ideal gas, real gases can behave differently due to intermolecular forces and the finite size of molecules, especially at high pressures and low temperatures.

Intermolecular Forces

Intermolecular forces are the attractive and repulsive forces between molecules, and they play a critical role in determining the physical properties and behavior of a substance. In gases, these forces are typically weak compared to those in liquids and solids, but they become more significant in real gases at high pressures and low temperatures.

Types of Intermolecular Forces

The main types of intermolecular forces are:

• Dispersion forces: Also known as London forces, these are weak forces that result from temporary fluctuations in the electron distribution within molecules.
• Dipole-dipole interactions: These occur between molecules that have permanent dipoles, meaning they have a positive end and a negative end.
• Hydrogen bonding: A strong type of dipole-dipole interaction that occurs between a hydrogen atom bonded to a highly electronegative atom, like oxygen or nitrogen, and another electronegative atom.

When gases are compressed, the distance between molecules decreases, causing these forces to become more noticeable. This attraction can lead to a decrease in pressure and volume than what would be predicted by the ideal gas law, symbolizing a deviation from ideal gas behavior.