Which Statement About Gasses Is True

Let's dive into the fascinating world of gases and explore some fundamental truths about their behavior and properties. Understanding these truths is crucial for anyone studying chemistry, physics, or even just trying to understand the world around them.

The Defining Characteristics of Gases

Gases are one of the three fundamental states of matter, the others being solids and liquids. What distinguishes a gas from the other two? Primarily, it's the freedom of movement and the weak intermolecular forces between the constituent particles (atoms or molecules).

Here are the defining characteristics of gases:

• Expandability: Gases expand to fill whatever container they occupy. Unlike solids or liquids, they don't have a fixed volume or shape.
• Compressibility: Gases are highly compressible. You can squeeze them into a smaller volume by increasing the pressure. This is the principle behind compressed air tanks and aerosol cans.
• Low Density: Compared to solids and liquids, gases have very low densities because the particles are widely spaced.
• Mixability: Gases mix readily and uniformly with each other, regardless of their chemical identities. This is why the air we breathe is a homogenous mixture of nitrogen, oxygen, and other gases.
• Fluidity: Gases, like liquids, are fluids. This means they can flow and conform to the shape of their container.

Key Statements About Gases: Unpacking the Truth

Now, let's examine some key statements about gases and determine which ones hold true based on scientific principles. We will delve into the kinetic molecular theory, gas laws, and real gas behavior to provide a comprehensive understanding.

1. Gases Have a Definite Shape and Volume

False. This statement is incorrect. One of the defining characteristics of gases, as mentioned earlier, is their lack of definite shape and volume. They take on the shape and volume of the container they occupy.

2. Gas Particles are in Constant, Random Motion

True. This statement aligns with the kinetic molecular theory of gases. This theory provides a microscopic explanation of gas behavior based on the following postulates:

• Gases consist of a large number of particles (atoms or molecules) that are in continuous, random motion.
• The volume of the individual gas particles is negligible compared to the volume of the container.
• Intermolecular forces (attractions and repulsions) between gas particles are negligible.
• Collisions between gas particles and the walls of the container are perfectly elastic (no energy is lost).
• The average kinetic energy of gas particles is directly proportional to the absolute temperature of the gas.

The constant, random motion of gas particles is responsible for their ability to expand, diffuse, and exert pressure.

3. Gas Pressure is Independent of Temperature

False. Gas pressure is directly related to temperature, as described by the gas laws. Specifically, Gay-Lussac's Law states that the pressure of a gas is directly proportional to its absolute temperature when the volume and the number of moles are kept constant. Mathematically, this is expressed as:

P₁/T₁ = P₂/T₂

Where:

• P₁ = Initial pressure
• T₁ = Initial absolute temperature (in Kelvin)
• P₂ = Final pressure
• T₂ = Final absolute temperature (in Kelvin)

Therefore, if you increase the temperature of a gas in a closed container, the pressure will increase proportionally.

4. All Gases are Ideal Under All Conditions

False. While the concept of an ideal gas is a useful simplification for many calculations, it's important to remember that real gases deviate from ideal behavior, especially at high pressures and low temperatures.

An ideal gas is a hypothetical gas that perfectly obeys the assumptions of the kinetic molecular theory. In reality, gas particles do have some volume, and intermolecular forces do exist, particularly when the particles are close together (high pressure) or moving slowly (low temperature).

The van der Waals equation is a modified version of the ideal gas law that accounts for these deviations:

(P + a(n/V)²) (V - nb) = nRT

Where:

• P = Pressure
• V = Volume
• n = Number of moles
• R = Ideal gas constant
• T = Temperature
• a = A constant that accounts for intermolecular attractions
• b = A constant that accounts for the volume of gas particles

The van der Waals constants (a and b) are specific to each gas and reflect the strength of the intermolecular forces and the size of the gas particles.

5. The Average Kinetic Energy of Gas Particles is the Same for All Gases at the Same Temperature

True. This statement is a direct consequence of the kinetic molecular theory. The theory states that the average kinetic energy of gas particles is directly proportional to the absolute temperature of the gas. The equation for average kinetic energy is:

KE = (3/2)kT

Where:

• KE = Average kinetic energy
• k = Boltzmann constant (1.38 x 10⁻²³ J/K)
• T = Absolute temperature (in Kelvin)

Notice that the equation doesn't depend on the identity of the gas. Therefore, at the same temperature, all gases will have the same average kinetic energy. This means that lighter gas particles will have higher average speeds than heavier gas particles at the same temperature.

6. Gas Particles Travel in Straight Lines Until They Collide with Something

True. This statement is another cornerstone of the kinetic molecular theory. In the absence of external forces, gas particles move in straight lines. Their paths are only disrupted when they collide with other gas particles or the walls of the container. These collisions are assumed to be perfectly elastic, meaning that no kinetic energy is lost during the collision.

7. Increasing the Volume of a Gas at Constant Temperature Decreases the Pressure

True. This statement is a manifestation of Boyle's Law, which states that the pressure of a gas is inversely proportional to its volume when the temperature and the number of moles are kept constant. Mathematically, this is expressed as:

P₁V₁ = P₂V₂

Where:

• P₁ = Initial pressure
• V₁ = Initial volume
• P₂ = Final pressure
• V₂ = Final volume

Therefore, if you increase the volume of a gas, the pressure will decrease proportionally, and vice versa.

8. Gases Cannot Be Liquefied

False. Gases can be liquefied under appropriate conditions of temperature and pressure. The process of liquefaction involves decreasing the kinetic energy of the gas particles so that intermolecular forces become significant enough to hold the particles together in a liquid state.

Lowering the temperature reduces the kinetic energy of the gas particles, allowing attractive forces to dominate. Increasing the pressure forces the gas particles closer together, increasing the effectiveness of intermolecular forces.

The critical temperature is the temperature above which a gas cannot be liquefied, no matter how much pressure is applied. Each gas has its own critical temperature.

9. The Rate of Diffusion of a Gas is Inversely Proportional to its Molar Mass

True. This statement is related to Graham's Law of Diffusion, which states that the rate of diffusion of a gas is inversely proportional to the square root of its molar mass. Mathematically, this is expressed as:

Rate₁/Rate₂ = √(M₂/M₁)

Where:

• Rate₁ = Rate of diffusion of gas 1
• Rate₂ = Rate of diffusion of gas 2
• M₁ = Molar mass of gas 1
• M₂ = Molar mass of gas 2

Diffusion is the process by which gas particles spread out and mix due to their random motion. Lighter gases diffuse faster than heavier gases because they have higher average speeds at the same temperature.

10. Gases Always React Completely with Each Other

False. Gases do not always react completely with each other. The extent to which gases react depends on several factors, including:

• The chemical nature of the gases: Some gases are more reactive than others. For example, hydrogen and oxygen will react readily to form water under the right conditions, while nitrogen and helium are relatively inert.
• The temperature: Higher temperatures generally increase the rate of chemical reactions.
• The pressure: Pressure can affect the rate and equilibrium of gas-phase reactions.
• The presence of a catalyst: A catalyst can speed up a reaction without being consumed itself.
• Equilibrium: Many gas-phase reactions are reversible and reach a state of equilibrium where the rates of the forward and reverse reactions are equal. At equilibrium, the reaction may not go to completion, meaning that some reactants will remain.

11. The Volume of a Gas is Directly Proportional to the Number of Moles of Gas at Constant Temperature and Pressure

True. This statement is Avogadro's Law. Avogadro's Law states that equal volumes of all gases, at the same temperature and pressure, contain the same number of molecules. This implies that the volume of a gas is directly proportional to the number of moles of gas when the temperature and pressure are held constant.

Mathematically, Avogadro's Law can be represented as:

V₁/n₁ = V₂/n₂

Where:

• V₁ = Initial volume
• n₁ = Initial number of moles
• V₂ = Final volume
• n₂ = Final number of moles

This law is crucial for understanding stoichiometry in gas-phase reactions.

12. Intermolecular Forces in Gases are Strong

False. One of the fundamental assumptions of the kinetic molecular theory of gases is that intermolecular forces (attractive or repulsive forces between gas molecules) are negligible. This is largely true for ideal gases. In real gases, intermolecular forces do exist, but they are significantly weaker than those in liquids and solids. These weak intermolecular forces in real gases become more important at high pressures and low temperatures.

13. Gases Conduct Heat Very Well

False. Gases are generally poor conductors of heat compared to liquids and solids. This is because the molecules in a gas are much farther apart than in liquids or solids, making it more difficult for them to transfer energy through collisions. The thermal conductivity of a gas depends on factors such as the gas's density, specific heat, and the average speed of its molecules.

14. The Color of a Gas is Always Invisible

False. Many gases are indeed invisible, but not all. Some gases have distinct colors. For example:

• Chlorine gas (Cl₂) is greenish-yellow.
• Nitrogen dioxide (NO₂) is reddish-brown.
• Ozone (O₃) has a pale blue tint.

The color of a gas arises from the way its molecules interact with light. Certain molecules absorb specific wavelengths of light, and the color we perceive is the result of the remaining wavelengths that are transmitted or reflected.

Summary of True Statements

To recap, here are the statements about gases that are true based on the principles of chemistry and physics:

• Gas particles are in constant, random motion.
• The average kinetic energy of gas particles is the same for all gases at the same temperature.
• Gas particles travel in straight lines until they collide with something.
• Increasing the volume of a gas at constant temperature decreases the pressure.
• The rate of diffusion of a gas is inversely proportional to its molar mass.
• The volume of a gas is directly proportional to the number of moles of gas at constant temperature and pressure.

Conclusion

Understanding the properties and behavior of gases is essential in many scientific and engineering fields. By knowing the true statements about gases and the underlying principles, we can better understand the world around us and make accurate predictions about gas behavior in various situations. From the air we breathe to the fuels that power our engines, gases play a crucial role in our lives. Remember to consider the conditions under which these statements hold true, particularly when dealing with real gases, where deviations from ideal behavior can occur.