AP Chemistry 3.6 Deviation from Ideal Gas Behavior Notes - New Syllabus 2024-2025
AP Chemistry 3.6 Deviation from Ideal Gas Behavior Notes- New syllabus
AP Chemistry 3.6 Deviation from Ideal Gas Behavior Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Explain the relationship among non-ideal behaviors of gases, interparticle forces, and/or volumes.
Key Concepts:
- Non-Ideal Behavior of Gases
Relationship Among Non-Ideal Gas Behaviors, Interparticle Forces, and Particle Volumes
The Ideal Gas Law, \( \mathrm{PV = nRT} \), assumes that gas particles have no volume and no intermolecular forces. In reality, these assumptions break down under certain conditions, and gases exhibit non-ideal behavior. Real gases deviate most significantly at high pressures (small volume) and low temperatures (strong attractive forces).
1. Causes of Non-Ideal Behavior:
a) Interparticle Attractions:
- Real gas molecules experience weak attractive forces (dispersion, dipole–dipole, or hydrogen bonding).
- When these attractions become significant, gas particles collide with container walls less forcefully → the observed pressure is lower than predicted by \( \mathrm{PV = nRT} \).
- These effects are strongest at low temperatures and moderate pressures, where molecules move more slowly and can interact longer.
b) Finite Particle Volume:
- Gas molecules occupy physical space; at high pressures, the volume of particles is no longer negligible compared to the container volume.
- This causes the available free space for motion to be smaller than the measured container volume → the observed volume is larger than predicted by the ideal gas law.
2. Conditions for Ideal vs. Non-Ideal Behavior:
Condition | Gas Behavior | Explanation |
---|---|---|
High Temperature, Low Pressure | Nearly Ideal | Particles are far apart; IMF effects negligible; volume negligible. |
Low Temperature | Non-Ideal (Attractive Forces Important) | Particles move slowly; IMF effects cause reduced pressure. |
High Pressure | Non-Ideal (Volume Effects Important) | Particles are crowded; finite size reduces free volume. |
3. van der Waals Equation (Correction for Real Gases):
To account for real gas behavior, the van der Waals equation modifies the ideal gas law by introducing correction terms for intermolecular attractions (\(a\)) and finite volume (\(b\)):
- \( \mathrm{a} \): accounts for attractive forces (adds to pressure).
- \( \mathrm{b} \): accounts for finite volume of particles (subtracts from available volume).
Interpretation:
- The term \( \mathrm{a\frac{n^2}{V^2}} \) corrects for intermolecular attractions that reduce pressure.
- The term \( \mathrm{nb} \) corrects for finite particle volume that reduces free space for motion.
4. Graphical Representation of Deviations:
a) PV/nRT vs. P Plot:
- For an ideal gas: the plot is a horizontal line at 1.0.
- For a real gas:
- At low pressure → ratio < 1 (due to attractive forces lowering P).
- At very high pressure → ratio > 1 (due to finite molecular volume).
Comparison of Ideal vs. Real Gas Behavior
Feature | Ideal Gas | Real Gas |
---|---|---|
Particle Volume | Negligible | Finite; significant at high pressures |
Intermolecular Forces | None | Attractive and repulsive forces exist |
Pressure Behavior | Exact prediction from \( \mathrm{PV=nRT} \) | Lower due to IMF; higher at very high P (volume effects) |
Conditions of Validity | High T, Low P | Low T, High P (deviations large) |
Example :
Explain why \(\mathrm{CO_2}\) gas deviates from ideal behavior more than \(\mathrm{He}\) under the same conditions of temperature and pressure.
▶️ Answer/Explanation
Step 1: The degree of non-ideality depends on intermolecular forces and molecular size.
Step 2: \(\mathrm{CO_2}\) molecules are larger and exhibit stronger London dispersion forces due to more electrons and greater polarizability.
Step 3: Helium atoms are small, nonpolar, and have very weak dispersion forces.
Step 4: Thus, \(\mathrm{CO_2}\) molecules experience greater attractions, reducing their measured pressure more significantly at low temperatures or moderate pressures.
Final Answer: \(\mathrm{CO_2}\) deviates more from ideal gas behavior because of stronger intermolecular attractions and larger molecular volume compared to \(\mathrm{He}\).