Part 1: Engage (Anchoring Phenomenon)

Ammonia (NH₃) is essential for producing the fertilizers that feed roughly half the world’s population. The industrial method for making ammonia — the Haber process (N₂ + 3H₂ ⇌ 2NH₃ + heat) — never achieves 100% yield, no matter how long the reaction runs. Chemical engineers must carefully control conditions to maximize ammonia output while balancing energy costs and equipment safety.

1. Observations and Questions:

• Why does the Haber process stop short of converting all nitrogen and hydrogen into ammonia?
• If you could change one condition inside the reaction chamber (temperature, pressure, or concentration), what would you predict happens to the amount of ammonia produced?
• Generate at least two “need to know” questions about how changing conditions might shift the equilibrium toward more product.

Part 2: Explore (Simulation Investigation)

Open the Le Chatelier’s Principle simulation. Set the temperature to 400 K, ensure the volume/pressure slider is in the middle, and press Run to observe the system at equilibrium. Notice the Shift Indicator (which shows which direction the reaction is favored) and the Conc vs Time graph.

2. Data Collection:

Concentration Experiments:

• Run the simulation at equilibrium. Click Add N₂ to inject more nitrogen gas. Watch the Shift Indicator and the Conc vs Time graph. Record what happens to the concentrations of N₂, H₂, and NH₃ as the system re-establishes equilibrium.
• Reset and click Add H₂ to inject more hydrogen. Record the changes in all three species.
• Reset and click Remove NH₃ to simulate removing product as it forms. Record the shift direction and concentration changes.

Volume/Pressure Experiment:

• Reset. Move the volume/pressure slider to the left to decrease the volume (increase pressure). Observe and record the shift direction and the final equilibrium concentrations of N₂, H₂, and NH₃. Note: the reaction has 4 moles of gas on the left (1 N₂ + 3 H₂) and 2 moles on the right (2 NH₃).
• Repeat with the slider moved to the right to increase the volume (decrease pressure).

Temperature Experiments:

• Reset. Set the temperature to 200 K (low). Run and record the equilibrium concentrations of NH₃ and the shift direction.
• Increase temperature to 600 K, then 800 K. Record NH₃ concentration and shift direction at each step. Remember: the reaction is exothermic (releases 92 kJ of heat).

Part 3: Explain (Sensemaking)

3. Explaining Equilibrium Shifts:

• Use your data to explain why adding a reactant (N₂ or H₂) pushed the equilibrium toward products. What is happening at the molecular level inside the chamber?
• Use your temperature data to explain why the Haber process cannot simply be run at a very low temperature to maximize ammonia yield. Consider both equilibrium position (Le Chatelier’s Principle) and reaction rate (collision theory).

4. Patterns and Relationships:

• Look at your volume/pressure data. Why does changing the pressure affect this particular reaction differently than it would affect a reaction with equal moles of gas on both sides? What is the relationship between the mole ratio (4 vs 2) and the shift direction?
• Based on your data, describe the tension between thermodynamic favorability (equilibrium position) and kinetic feasibility (reaction rate). Why can’t chemical engineers just maximize ammonia yield by choosing the single best temperature or pressure?

Part 4: Elaborate/Evaluate (Argumentation & Modeling)

5. Engineering Design Proposal: Imagine you are a chemical engineer designing a new Haber process plant. Write a proposal recommending the optimal operating conditions (temperature range, pressure range, and concentration management strategy) for producing ammonia.

Your proposal must:

• Specify a target temperature range (from the 200–800 K range you tested) and justify it by balancing equilibrium yield against reaction rate.
• Specify whether you would operate at high pressure or low pressure, and explain why the mole ratio of gases supports your choice.
• Describe how you would manage reactant concentrations (N₂ and H₂) and product removal to continuously drive the reaction forward.
• Discuss at least one real-world tradeoff (e.g., energy cost, equipment safety, catalyst effectiveness) and how your proposed conditions address it.
• Use evidence from your simulation data to support every claim.