The best examples of Le Chatelier's principle: examples and applications

Before getting lost in symbols and equations, start with things you actually see and touch. Some of the best examples of Le Chatelier’s principle are hiding in your kitchen, your lungs, and even your soda can.

Carbonated drinks and CO₂ pressure

A classic example of Le Chatelier’s principle is a carbonated drink. Inside a sealed can, the equilibrium is:

\[ \text{CO}_2(\text{g}) \rightleftharpoons \text{CO}_2(\text{aq}) \]

High pressure above the liquid forces more CO₂ to dissolve. Open the can, pressure drops, and the system is disturbed. According to Le Chatelier’s principle, the equilibrium shifts to oppose the change—so it produces more gas to raise the pressure again. Translation: bubbles.

Warm soda goes flat faster because the dissolution of CO₂ is exothermic. Raising temperature favors the endothermic direction (gas formation), shifting the equilibrium toward CO₂(g). This is a clean, intuitive example of how temperature and pressure changes interact with equilibrium constants and reaction quotients.

Blood CO₂, breathing, and pH balance

Your blood chemistry provides powerful real examples of Le Chatelier’s principle: examples and applications in physiology. The key equilibrium is:

\[ \text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^- \]

When you hold your breath, CO₂ builds up in your blood. The system responds by shifting right, producing more H⁺ and lowering pH (more acidic). When you hyperventilate, you exhale more CO₂, decreasing its concentration. The equilibrium shifts left, reducing H⁺ and raising pH.

Medical sources, including the National Institutes of Health, describe this CO₂–bicarbonate buffer as a central part of acid–base regulation in humans (NIH). It’s not just a textbook concept; hospitals literally use this equilibrium when interpreting blood gas results.

Industrial-scale examples of Le Chatelier’s principle: examples and applications in manufacturing

If you want examples of Le Chatelier’s principle: examples and applications with serious economic impact, look at industrial chemistry. Here, small shifts in equilibrium can mean millions of dollars.

The Haber process for ammonia

The Haber–Bosch process is the poster child. The reaction is:

\[ \text{N}_2(\text{g}) + 3\,\text{H}_2(\text{g}) \rightleftharpoons 2\,\text{NH}_3(\text{g}) + \text{heat} \]

A few key Le Chatelier insights drive the design:

• Pressure: Four moles of gas on the left, two on the right. Increasing pressure shifts equilibrium toward fewer moles of gas → more NH₃.
• Temperature: The forward reaction is exothermic. Lowering temperature favors ammonia formation, but too low and the reaction rate crawls. Industrial plants compromise at moderate temperatures (about 400–500 °C) and very high pressures.
• Concentration: Removing ammonia as it forms (by cooling and liquefying it) pulls the equilibrium to the right, increasing yield.

This is one of the best examples of Le Chatelier’s principle being used deliberately. Without these optimizations, modern fertilizer production—and modern agriculture—would look very different.

Contact process for sulfuric acid

Another industrial example of Le Chatelier’s principle is the contact process for sulfuric acid production. The key equilibrium is:

\[ 2\,\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightleftharpoons 2\,\text{SO}_3(\text{g}) + \text{heat} \]

Again, the forward reaction is exothermic and reduces the total number of gas moles. High pressure and moderate temperature favor SO₃ formation. In practice, industries use catalysts and recycle unreacted SO₂ and O₂, constantly disturbing the system in a way that keeps equilibrium shifted toward product.

Laboratory examples of Le Chatelier’s principle: examples and applications you’ll actually see in class

Chemistry instructors love colorful examples of Le Chatelier’s principle because you can watch equilibrium shifts happen in real time.

Cobalt complex and color change

One of the most famous lab demonstrations uses cobalt(II) chloride:

\[ \text{[Co(H}_2\text{O)}_6]^{2+} + 4\,\text{Cl}^- \rightleftharpoons \text{[CoCl}_4]^{2-} + 6\,\text{H}_2\text{O} \]

• The pink hexaaquacobalt(II) complex is favored at lower temperatures.
• The blue tetrachlorocobaltate(II) complex is favored at higher temperatures.

Heat the solution, and it shifts blue; cool it, and it shifts pink. Add chloride ions, and the equilibrium moves toward the blue complex to consume the added Cl⁻. This is a visually striking example of Le Chatelier’s principle: examples and applications in coordination chemistry.

Chromate–dichromate equilibrium

Another popular example of equilibrium shifting is the chromate–dichromate system:

\[ 2\,\text{CrO}_4^{2-} + 2\,\text{H}^+ \rightleftharpoons \text{Cr}_2\text{O}_7^{2-} + \text{H}_2\text{O} \]

• Chromate (CrO₄²⁻) is yellow.
• Dichromate (Cr₂O₇²⁻) is orange.

Add acid (H⁺), and the equilibrium shifts right: the solution turns more orange. Add base to remove H⁺, and it shifts left: the solution turns yellow. This is a clean pH-driven case where you can literally see Le Chatelier’s principle in action.

Gas-phase examples of Le Chatelier’s principle: examples and applications with pressure and volume

Gas equilibria are perfect for illustrating how pressure and volume changes influence the reaction quotient (Q) and equilibrium position.

Nitrogen dioxide and dimerization

The brown color of nitrogen dioxide (NO₂) and the colorless dimer dinitrogen tetroxide (N₂O₄) provide another elegant example of Le Chatelier’s principle:

\[ 2\,\text{NO}_2(\text{g}) \rightleftharpoons \text{N}_2\text{O}_4(\text{g}) \]

• Fewer moles of gas on the right.
• Lower temperatures favor N₂O₄ (exothermic dimerization).

Compress the gas (increase pressure), and the equilibrium shifts right to reduce the number of gas molecules. The color fades as more colorless N₂O₄ forms. Warm it up, and the brown color intensifies as the equilibrium shifts back to NO₂.

Ozone formation and atmospheric chemistry

In the atmosphere, ozone (O₃) is formed and destroyed through a series of equilibria. A simplified step is:

\[ \text{O}_2 + \text{O} \rightleftharpoons \text{O}_3 \]

Changes in sunlight intensity, temperature, and pollutant concentrations disturb these equilibria. While the actual atmospheric chemistry is more complex, the general idea follows Le Chatelier’s logic: systems shift to oppose imposed changes. Research organizations such as NASA and the EPA discuss how temperature and emissions affect ozone levels and related equilibria (EPA).

Biochemical examples of Le Chatelier’s principle: examples and applications in your body

Biochemistry is full of real examples of Le Chatelier’s principle: examples and applications that keep you alive, whether you’re thinking about it or not.

Hemoglobin and oxygen binding

Hemoglobin (Hb) binding to oxygen can be treated as an equilibrium:

\[ \text{Hb} + \text{O}_2 \rightleftharpoons \text{HbO}_2 \]

In the lungs, high O₂ concentration and relatively lower CO₂ favor the right side, promoting oxygen loading. In tissues, lower O₂ and higher CO₂ push the equilibrium left, encouraging oxygen release.

CO₂ also affects pH through the bicarbonate buffer system mentioned earlier. Changes in pH shift the equilibrium position for oxygen binding (the Bohr effect), a topic widely discussed in physiology texts and medical resources such as those hosted by the National Library of Medicine (NLM).

ATP hydrolysis and coupled reactions

Cells use Le Chatelier’s principle implicitly when they couple ATP hydrolysis to otherwise unfavorable reactions. The hydrolysis equilibrium is:

\[ \text{ATP} + \text{H}_2\text{O} \rightleftharpoons \text{ADP} + \text{P}_i + \text{H}^+ \]

By continuously consuming ADP and inorganic phosphate (Pᵢ) in other pathways, cells keep disturbing this equilibrium and effectively pull it to the right, maintaining a high ATP/ADP ratio. The concept of reaction coupling is standard in biochemistry courses and is discussed in detail in open educational resources such as those from MIT and other universities (MIT OpenCourseWare).

Connecting Le Chatelier’s principle to equilibrium constants and Q

So far, we’ve focused on examples of Le Chatelier’s principle: examples and applications you can visualize. Under the hood, every one of these systems is governed by the relationship between the equilibrium constant (K) and the reaction quotient (Q).

• K is fixed at a given temperature. It tells you the ratio of product to reactant concentrations at equilibrium.
• Q is the same ratio, but for whatever concentrations you have right now.

When you disturb a system—by changing concentration, pressure, or temperature—you change Q. The system then shifts in whichever direction brings Q back to K.

For instance, in the Haber process:

\[ \text{N}_2 + 3\,\text{H}_2 \rightleftharpoons 2\,\text{NH}_3 \]

Adding more N₂ or H₂ makes Q smaller than K. The system responds by forming more NH₃ until Q = K again. Removing NH₃ also lowers Q, again shifting right.

Temperature is a bit different. Changing temperature actually changes K itself, because K is related to the reaction’s enthalpy via the van ’t Hoff equation. For exothermic reactions, raising temperature lowers K, favoring reactants. For endothermic reactions, raising temperature increases K, favoring products. That’s why the optimal industrial conditions for systems like the Haber and contact processes are a negotiated balance between kinetics and thermodynamics.

If you want a rigorous mathematical treatment of K, Q, and temperature effects, standard university general chemistry resources and textbooks (for example, those hosted by major universities and the American Chemical Society) go into the derivations in depth.

Modern trends: Le Chatelier’s principle in climate and environmental chemistry

As of 2024–2025, some of the most discussed real examples of Le Chatelier’s principle: examples and applications are in environmental and climate chemistry.

Ocean acidification

Rising atmospheric CO₂ dissolves in seawater and participates in a series of equilibria:

\[ \text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^- \rightleftharpoons 2\,\text{H}^+ + \text{CO}_3^{2-} \]

As atmospheric CO₂ increases, more CO₂ dissolves, shifting equilibria toward more bicarbonate and hydrogen ions. The increased H⁺ lowers pH—this is ocean acidification. Organizations like NOAA and the EPA track and explain this process and its impact on marine ecosystems.

Here, Le Chatelier’s principle helps frame the narrative: adding CO₂ to the atmosphere doesn’t just sit there; it pushes a chain of equilibria in the oceans, with measurable shifts in pH documented over recent decades.

Carbonate equilibria in freshwater and drinking water treatment

Water treatment plants manage carbonate equilibria to control corrosion and scaling in pipes. Adjusting pH, dissolved inorganic carbon, and hardness shifts equilibria among CO₂, bicarbonate, and carbonate, influencing whether minerals like CaCO₃ dissolve or precipitate. The U.S. Environmental Protection Agency provides technical guidance on water treatment chemistry and corrosion control in distribution systems (EPA).

Again, the principle is simple: change conditions, shift equilibria. But the applications are very 2020s—aging infrastructure, lead and copper concerns, and the need for fine-tuned water chemistry control.

FAQ: Common questions about examples of Le Chatelier’s principle

Q: What are some everyday examples of Le Chatelier’s principle?
Everyday examples include carbonated drinks losing fizz when opened, blood maintaining pH through the CO₂–bicarbonate buffer, and temperature-dependent color changes in equilibrium systems like the cobalt chloride complex. These examples of Le Chatelier’s principle: examples and applications show up in cooking, breathing, and even climate discussions.

Q: Can you give a simple example of Le Chatelier’s principle involving temperature?
A simple temperature-based example of Le Chatelier’s principle is the NO₂/N₂O₄ equilibrium. Cooling the mixture favors the exothermic formation of colorless N₂O₄, so the brown color fades. Heating it shifts the equilibrium back toward brown NO₂.

Q: How do industrial processes use examples of Le Chatelier’s principle?
Processes like the Haber and contact processes adjust pressure, temperature, and concentrations to push equilibria toward desired products. These are some of the best examples of Le Chatelier’s principle being used strategically to maximize yield while balancing cost and reaction rate.

Q: Do catalysts change the position of equilibrium in these examples?
No. Catalysts speed up both the forward and reverse reactions equally, so they help a system reach equilibrium faster but do not change the equilibrium position or the equilibrium constant K. They matter for practicality, not for where the equilibrium finally settles.

Q: How are equilibrium constants related to the examples of Le Chatelier’s principle we see?
In all these examples, the system shifts because the disturbance makes the reaction quotient Q different from the equilibrium constant K. The system then moves in the direction that brings Q back to K. Le Chatelier’s principle is the qualitative summary of that quantitative relationship.