Gas Laws: Real Gas vs. Ideal Gas Sandbox Task

Task Overview

In this task, you will explore the differences between ideal gases and real gases using the “Real Gas vs. Ideal Gas Sandbox” simulation. You will manipulate parameters such as gas type, temperature, volume, and amount to investigate how intermolecular forces and particle volume cause deviations from the Ideal Gas Law.

Learning Objectives

Phenomenon

Why do high-pressure gas cylinders (like scuba tanks or industrial oxygen tanks) hold slightly more or slightly less gas than the Ideal Gas Law predicts? Under extreme conditions, the simple assumption that gas particles have no volume and do not interact breaks down.

Instructions

Part 1: Exploring Ideal Behavior

  1. Open the Real Gas vs. Ideal Gas Sandbox simulation.
  2. Set the Gas Type to “Ideal Gas”.
  3. Set Temperature to 300 K, Volume to 2.0 L, and Amount to 1.0 mol.
  4. Record the Calculated Pressure. What is the Compressibility Factor (Z)?
  5. Prediction: If you decrease the volume to 1.0 L while keeping temperature and amount constant, what should happen to the pressure according to Boyle’s Law?
  6. Test your prediction. Decrease the Volume to 1.0 L. Record the new pressure. Did it match your prediction?

Part 2: Introducing Real Gases

  1. Change the Gas Type to “Carbon Dioxide (CO₂)”.
  2. Set Temperature to 300 K, Volume to 2.0 L, and Amount to 1.0 mol.
  3. Compare the Ideal Gas Law pressure with the Real Gas (Van der Waals) pressure. Are they the same? Which one is higher/lower?
  4. Look at the Compressibility Factor (Z). Is it greater than 1, equal to 1, or less than 1?
  5. Observe the Particle View. How does the behavior of the CO₂ particles differ from the Ideal Gas particles (hint: look for clustering or “bonds” forming temporarily).

Part 3: Investigating Intermolecular Forces (The ‘a’ constant)

The van der Waals constant ‘a’ represents the strength of intermolecular attractive forces.

  1. Set Volume to 2.0 L, Amount to 1.0 mol, and Temperature to 150 K (a low temperature).
  2. Compare the Real Gas pressure and Compressibility Factor (Z) for Helium (He) and Water Vapor (H₂O).
    • Helium (a = 0.0346): Real Pressure = __, Z = ____
    • Water Vapor (a = 5.536): Real Pressure = __, Z = ____
  3. Analysis: Based on your observations, how does a larger ‘a’ value affect the real pressure compared to the ideal pressure? Explain why this happens at the microscopic level using the particle view.

Part 4: Investigating Particle Volume (The ‘b’ constant)

The van der Waals constant ‘b’ represents the physical volume occupied by the gas particles themselves.

  1. Set Temperature to 800 K (a high temperature, where attractive forces are less significant), Amount to 5.0 mol (a large amount of gas), and Volume to 0.5 L (a very small volume).
  2. Compare the Ideal Gas Law pressure with the Real Gas (Van der Waals) pressure for Nitrogen (N₂).
    • Nitrogen (b = 0.0387): Ideal Pressure = __, Real Pressure = ____
  3. Look at the Compressibility Factor (Z). Is it greater than 1 or less than 1?
  4. Analysis: Why does the real pressure become higher than the ideal pressure under conditions of high pressure and small volume? Explain this using the concept of the ‘b’ constant.

Part 5: The Mystery Gas Challenge

  1. Select “Mystery Gas X”.
  2. Design and conduct a brief investigation to determine under what conditions (High/Low Temp, High/Low Volume, High/Low Moles) Mystery Gas X deviates the most from ideal behavior.
  3. Claim: Mystery Gas X deviates most from ideal behavior when temperature is ___ and volume is _______.
  4. Evidence: Provide specific data points (Temperature, Volume, Ideal Pressure, Real Pressure, Z factor) to support your claim.
  5. Reasoning: Based on its van der Waals constants (‘a’ and ‘b’), explain why Mystery Gas X behaves this way under these specific conditions.