Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Chemical equilibrium ↓


Le Chatelier's principle


Le Chatelier's principle is a fundamental concept in chemistry, particularly in the study of chemical equilibrium. Named after French chemist Henri Louis Le Chatelier, it provides valuable information about how a chemical system at equilibrium reacts when subjected to changes in concentration, temperature, or pressure. Let's dive into this important principle that helps predict the direction of the reaction when any change is applied to a system at equilibrium.

What is chemical equilibrium?

Chemical equilibrium is the state of a reversible reaction where the rate of the forward reaction is equal to the rate of the backward reaction. This means that the concentrations of reactants and products remain constant over time. Equilibrium can be represented by a double arrow in a chemical equation:

 a + b ⇌ c + d

Here, the forward reaction is A + B → C + D and the reverse reaction is C + D → A + B At equilibrium, the rates of these reactions are equal, and there is no net change in the concentration of the species involved.

Le Chatelier's principle explained

Le Chatelier's principle states that if changing conditions disturbs dynamic equilibrium, the equilibrium state shifts to counteract the change. In simple terms, if you disturb a system in equilibrium, it will adjust itself to minimize the disturbance and restore a new equilibrium.

Factors affecting balance

Several factors can disturb the equilibrium of a chemical system:

  • Concentration: Change in the concentration of reactants or products.
  • Temperature: Change in the temperature of the system.
  • Pressure: Change in pressure, usually for gases.

Effect of concentration changes

Consider the general equilibrium reaction:

 a + b ⇌ c + d

Let us examine each change in concentration:

  • Adding more of A or B: If we add more of reactant A or B, the system will adjust to reduce this amount. Thus, it will shift the equilibrium to the right, producing more products C and D.
  • Removal of some A or B: The system will try to replace the lost reactants by shifting the equilibrium to the left, causing more A and B to be formed.
  • Adding excess C or D: Adding products will shift the equilibrium to the left, causing the added products to be eliminated and more reactants to be formed.
  • Removal of some C or D: The equilibrium position will shift to the right and the lost products will be replaced by forming more C and D.

These reactions help the system maintain equilibrium, which implements Le Chatelier's principle.

Effect of temperature change

The effect of temperature change depends on the nature of the reaction (endothermic or exothermic):

Exothermic reactions

For exothermic reactions (where heat is released), an increase in temperature will shift the equilibrium toward the reactants. Conversely, a decrease in temperature will shift the equilibrium toward the products.

Example:

 N 2 + 3H 2 ⇌ 2NH 3 + heat

If you increase the temperature, the system absorbs this "extra" heat, thereby affecting the reaction that consumes heat (here the opposite reaction occurs). Thus, the equilibrium shifts to the left.

Endothermic reactions

For endothermic reactions (where heat is absorbed), an increase in temperature shifts the equilibrium toward the products, while a decrease in temperature shifts it toward the reactants.

Example:

 Heat + N 2 O 4 (g) ⇌ 2NO 2 (g)

By increasing the temperature, the system promotes the forward reaction, consuming additional heat, and so the equilibrium shifts to the right.

Effect of pressure changes

Changes in pressure affect the equilibrium involving gases. The response of a system to a change in pressure can be understood based on the number of moles of gas on each side of the equation.

  • Increase in pressure: The system will shift the equilibrium position towards where there are fewer moles of gas.
  • Decrease in pressure: The system shifts the equilibrium towards where there are more moles of gas.

Example:

 N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g)

Here, there are 4 moles of gas on the reactant side and 2 moles of gas on the product side. Increasing the pressure shifts the equilibrium to the right (towards ammonia), to the side where there are fewer moles of gas.

Visual representation of Le Chatelier's principle

To understand the concepts visually, let's use a graphic to depict dynamic changes in equilibrium:

Reactants Products Change Shift

In this representation, the blue rectangle represents the reactants, and the red rectangle represents the products. The orange line through the middle symbolizes the point where the equilibrium is at its initial equilibrium. When a change occurs, represented by the top label "change," the equilibrium shifts in one direction, represented by the bottom label "shift."

Real life applications of Le Chatelier's principle

Le Chatelier's principle is not just a classroom concept. It is widely used in industrial processes:

  • Haber process: This process creates ammonia, an important ingredient for fertilizers. By manipulating temperature and pressure according to Le Chatelier's principle, industrial plants can optimize ammonia production.
  • Contact process: In the manufacture of sulfuric acid, equilibrium considerations ensure maximum yield of products under various pressure and temperature conditions.

Conclusion

Le Chatelier's principle is an important tool for chemists, helping to predict how equilibrium systems respond to changes. Understanding this principle provides insight into the dynamics of chemical reactions, allowing chemists to profitably use equilibrium shifts, especially in industrial applications.

Whether adjusting factors for optimum yield in production or understanding natural processes, this principle is indispensable for effectively interpreting and managing chemical equilibrium.


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Le Chatelier's principle

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