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Le Châtelier’s Principle and Shifting Equilibrium Mark as Favorite (0 Favorites)

DEMONSTRATION in Le Châtelier's Principle, Establishing Equilibrium, Equilibrium Constants. Last updated July 23, 2024.


In this demonstration, students will be introduced to the concepts of Le Châtelier’s Principle and reversible reactions through the formation of a copper-ammonia complex ion.

Grade Level

High School

AP Chemistry Curriculum Framework

This demonstration supports the following unit, topics, and learning objectives:

  • Unit 7: Equilibrium
    • Topic 7.1: Introduction to Equilibrium
      • TRA-6.A: Explain the relationship between the occurrence of a reversible chemical or physical process, and the establishment of equilibrium, to experimental observations.
    • Topic 7.9: Introduction to Le Châtelier’s Principle
      • TRA-8.A: Identify the response of a system at equilibrium to an external stress using Le Châtelier's principle.

NGSS Alignment

This demonstration will help prepare your students to meet the performance expectations in the following standard:

  • HS-PS1-6: Refine the design of a chemical system by specifying a change in conditions that would produce increased amounts of products at equilibrium.
  • Scientific and Engineering Practices:
    • Engaging in Argument from Evidence


By the end of this demonstration, students should be able to:

  • Understand how adding ammonium impacts the equilibrium.
  • Explain how adding hydrochloric acid affects the equilibrium.
  • Demonstrate with evidence and reasoning how a chemical reaction can achieve equilibrium in the forwards or reverse direction.

Chemistry Topics

This demonstration supports students’ understanding of:

  • Equilibrium
  • Le Châtelier's Principle
  • Equilibrium Constants
  • Establishing Equilibrium


Teacher Preparation: 30 minutes
Lesson: 30 minutes


  • 100mL volumetric flask
  • Scale
  • Weighing boat
  • 250 mL Erlenmeyer flask
  • 100 mL graduated cylinders (2)
  • Stirring rod (or stir bar magnet and stir plate)
  • Water
  • 2.5g copper sulfate pentahydrate, CuSO4·5H2O
  • 30mL of household ammonia, NH3
  • 60 mL 0f 1M hydrochloric acid, HCl


  • Always wear safety goggles when handling chemicals in the lab.
  • Students should wear proper safety gear during chemistry demonstrations. Safety goggles and lab apron are required.
  • When working with acids, if any solution gets on the skin, it should be thoroughly flushed with water.
  • When diluting acids, always add acid to water.
  • Consult the Safety Data Sheets for hydrochloric acid and ammonia before conducting the demonstration.

Teacher Notes

  • This demonstration is suitable for any level of chemistry to show students how equilibrium is established by running the reaction forward and then how to drive the reaction in reverse.
  • Students will complete a student observation sheet during the demonstration. They are prompted with several questions that correspond to steps in the procedure and are also encouraged to make predictions. Be sure that adequate wait time is provided so that students can accomplish this.
  • Background:
    • Le Châtelier's Principle: When a system at equilibrium is disrupted by a stress, the system will shift to alleviate that stress.
    • When copper(II) sulfate pentahydrate is dissolved in water, a pale blue solution will form. If ammonia is added to the solution, then the dark blue copper-ammonia complex ion will form.
    • This equilibrium can be described be the equation: Cu2+ (aq)  + 4NH3 (aq)    ↔ Cu(NH3)42+(aq)
    • When 60 mL of 1 M hydrochloric acid is added to the complex ion solution, the equilibrium shifts in reverse, back towards the reactants, and the solution will turn pale blue again due to the presence of Cu2+ ions.
    • If you want to shift the reaction forward again, add 75 mL of household ammonia and the dark blue copper-ammonia complex ion will return.
  • Demonstration Procedures:
    • Dissolve 2.5 g of copper(II) sulfate pentahydrate in water in order to make 100 mL of 0.1 M CuSO4·5H2O. The color of this solution should be pale blue. Add this to a 250 mL Erlenmeyer flask.
    • Then add 30 mL of household ammonia, NH3, to the copper(II) sulfate solution. Stir this mixture by hand or using a stir plate and magnetic stir bar. This shifts the equilibrium to the right. As a result, the dark blue tetraaminecopper(II) complex ion, Cu(NH3)42+, forms.
    • Then add 60 mL of 1 M hydrochloric acid, HCl. Stir this mixture by hand or using a stir plate and magnetic stir bar. This will shift the reaction in reverse and breaks down the complex ion to form pale blue Cu2+ ions. The HCl is not concentrated enough to form the tetrachlorocopper(II) complex ion, CuCl42-, and resulting precipitated cloud of ammonium chloride, NH4Cl.
    • When students see a clear solution of household ammonia poured into a pale blue copper(II) sulfate solution, most students wouldn’t expect the dark blue copper-ammonia complex ion.
  • Alternatively, this activity can be completed as a smaller-scale lab activity performed by the students themselves.
  • The photograph on the left shows the pure, pale blue 0.1 M copper(II) sulfate pentahydrate solution.
  • The photograph on the right shows the solution after ammonia is added, a dark blue copper-ammonia complex ion forms.

For the Student


Student Observation Sheet

Cu2+(aq)  +  4 NH3(aq) ⟷ Cu(NH3)42+(aq)
Initial observation

Make a prediction about what will happen when household NH3 is added.
Observation after ammonia is added

What evidence do you have indicating that a chemical reaction took place?
Make a prediction about what will happen when HCl is added.

Observation after HCl is added

What evidence do you have indicating that a chemical reaction took place?

Which way did the equilibrium shift?

Extension Questions

  1. What does it mean when a reaction is at equilibrium?
  2. What is Le Châtelier’s Principle?
  3. Why does adding HCl shift the reaction in reverse?
  4. Write the equilibrium constant expression for the reaction.