Chemistry · Grade 11 · Gases and Atmospheric Chemistry · Term 3

Real Gases vs. Ideal Gases

Students will explore the conditions under which real gases deviate from ideal gas behavior and the reasons why.

• Ontario Curriculum Expectations
• HS-PS1-3

About This Topic

The ideal gas law PV = nRT assumes gas particles are point masses with no volume and no intermolecular forces between them. Real gases deviate from ideal behavior under high pressures and low temperatures. At high pressures, particles occupy significant space relative to the container, reducing the effective volume available and causing measured pressure to exceed ideal predictions after corrections. At low temperatures, particles move slowly enough for attractive forces to pull them back from container walls, lowering pressure below ideal values. Grade 11 students graph the compressibility factor Z = PV/nRT, where Z ≠ 1 signals deviations, and apply the van der Waals equation (P + an²/V²)(V - nb) = nRT to quantify corrections.

This topic extends kinetic molecular theory into atmospheric chemistry applications, such as gas behavior in extreme conditions like the upper atmosphere or industrial compression. Students predict and compare how intermolecular forces and particle volume alter pressure-volume relationships, honing analytical skills for data interpretation.

Active learning excels with this topic because molecular-scale phenomena are invisible. Simulations and data plotting let students manipulate variables to observe deviations emerge, while group discussions link observations to explanations, turning abstract theory into intuitive understanding.

Key Questions

1. Explain the molecular reasons why real gases deviate from ideal gas behavior at high pressures and low temperatures.
2. Compare the behavior of an ideal gas to that of a real gas under extreme conditions.
3. Predict how the intermolecular forces and particle volume of a real gas affect its pressure and volume compared to an ideal gas.

Learning Objectives

• Explain the molecular basis for deviations of real gases from ideal gas behavior under specific conditions of high pressure and low temperature.
• Compare the pressure-volume relationships of an ideal gas and a real gas when subjected to extreme conditions.
• Analyze the impact of intermolecular forces and finite particle volume on the compressibility factor (Z) of real gases.
• Predict how changes in temperature and pressure will affect the deviation of a real gas from ideal behavior using the van der Waals equation.

Before You Start

Kinetic Molecular Theory of Gases

Why: Students need to understand the assumptions of the KMT regarding particle motion, volume, and forces to grasp why real gases deviate.

The Ideal Gas Law (PV=nRT)

Why: Students must be familiar with the ideal gas law before they can understand the corrections and deviations applied to real gases.

Key Vocabulary

Ideal Gas Law	A theoretical model describing the behavior of an ideal gas, assuming particles have negligible volume and no intermolecular forces.
Compressibility Factor (Z)	A ratio (PV/nRT) that indicates how much a real gas deviates from ideal gas behavior; Z=1 for an ideal gas.
Intermolecular Forces	Attractive or repulsive forces between neighboring molecules, such as van der Waals forces, which are significant in real gases at low temperatures.
Particle Volume	The finite space occupied by gas molecules themselves, which becomes significant relative to the container volume at high pressures.
Van der Waals Equation	A modified ideal gas law that includes correction terms for intermolecular attractions and finite particle volume to better describe real gas behavior.

Watch Out for These Misconceptions

Common Misconception

Real gases deviate from ideal behavior under all conditions.

What to Teach Instead

Deviations are negligible at low pressures and high temperatures, where ideal assumptions hold well for lab work. Graphing activities with real data help students identify moderate condition zones, building judgment on model applicability through pattern recognition.

Common Misconception

Particle volume causes pressure decreases in real gases.

What to Teach Instead

Finite volume reduces free space at high P, leading to higher-than-ideal pressure after accounting for attractions. Simulations toggling volume correction clarify this counterintuitive effect, as students watch Z rise above 1 and connect it to crowding.

Common Misconception

Intermolecular forces increase measured pressure.

What to Teach Instead

Attractions reduce pressure by pulling particles from walls, dominant at low T. Peer demos contrasting cooled vs heated gases reveal this, with discussions reinforcing how speed modulates force impact.

Active Learning Ideas

PhET Simulation: Real Gas Deviations

Direct pairs to the PhET Gas Properties simulation set to real gas mode. Have them adjust pressure from 1 to 100 atm and temperature from 100 K to 500 K, recording Z values in a table. Pairs then graph Z vs P and identify conditions of largest deviation.

35 min·Pairs

Graphing Stations: Compressibility Curves

Prepare stations with data for CO2, N2, and He at varying P and T. Small groups plot Z vs P graphs at fixed T, rotating stations. Conclude with whole-class share-out on trends.

45 min·Small Groups

Van der Waals Calculations: Pairs Challenge

Provide van der Waals constants a and b for several gases. Pairs calculate corrected P and V for given conditions, compare to ideal values, and predict which gas deviates most under high P low T.

30 min·Pairs

Demo Discussion: Dry Ice Pressure

Demonstrate pressure buildup from dry ice in a sealed flask at room T vs cooled. Students predict ideal vs real behavior beforehand, then explain observations in whole class using molecular terms.

25 min·Whole Class

Real-World Connections

• Chemical engineers designing high-pressure storage tanks for gases like hydrogen or natural gas must account for real gas behavior to ensure safety and efficiency, preventing tank rupture.
• Atmospheric scientists studying gas behavior in the upper atmosphere, where temperatures are extremely low and pressures can vary significantly, use real gas models to predict atmospheric composition and reactions.
• Industrial processes involving gas liquefaction, such as the production of liquid nitrogen or oxygen, rely on understanding how real gases deviate from ideal behavior at very low temperatures and high pressures.

Assessment Ideas

Quick Check

Present students with a graph showing the compressibility factor (Z) versus pressure for several gases at a constant low temperature. Ask: 'Which gas shows the most deviation from ideal behavior at high pressures? Explain your reasoning based on intermolecular forces and particle volume.'

Discussion Prompt

Pose the following scenario: 'Imagine a gas cylinder containing helium at room temperature and atmospheric pressure, and another cylinder containing ammonia under high pressure and low temperature. Which gas is more likely to deviate significantly from ideal gas behavior? Justify your answer by discussing the specific properties of each gas.'

Exit Ticket

Ask students to write two sentences explaining why real gases behave differently from ideal gases at high pressures, and two sentences explaining why they behave differently at low temperatures. They should use at least two key vocabulary terms in their answers.

Frequently Asked Questions

Why do real gases deviate from ideal gas behavior at high pressure and low temperature?

High pressure makes particle volume significant, shrinking effective container volume and raising pressure beyond ideal. Low temperature slows particles, allowing intermolecular attractions to reduce wall collisions and lower pressure. Students quantify this via Z factor graphs or van der Waals terms, seeing molecular causes directly tied to conditions.

How does the van der Waals equation account for real gas behavior?

The equation corrects ideal PV=nRT by adding an attraction term a(n/V)² to P, accounting for reduced pressure, and subtracting particle volume nb from V. For 1 mol, (P + a/V²)(V - b) = RT. Pairs calculations under extreme conditions show improved accuracy over ideal law for gases like CO2.

How can active learning help students grasp real vs ideal gases?

Hands-on simulations like PhET let students change P and T to see Z deviate live, matching predictions to molecular ideas. Group graphing of real data reveals trends like Z<1 at low T, while demos with dry ice make effects visible. These build deep intuition over rote learning.

What are molecular reasons for real gas deviations?

Ideal gases ignore particle volume and forces; real particles exclude volume at high P, causing excess pressure, and experience attractions at low T, causing deficit. Kinetic theory explains: slow speeds enhance attractions, crowding enhances volume effects. Applications to atmospheric gases contextualize for students.