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 ↓


Law of mass action


The law of mass action is a fundamental principle in chemistry that helps us understand the behavior of chemical reactions, particularly in the context of chemical equilibrium. Formulated by Norwegian chemists Cato Guldberg and Peter Vaage in the late 19th century, this law provides a quantitative relationship between the concentrations of reactants and products at equilibrium. It is a cornerstone of physical chemistry and is essential to understanding how chemical reactions proceed.

Basic concepts

The law of mass action states that at a given temperature, the rate of a chemical reaction is directly proportional to the product of the molar concentrations of the reactants, each raised to the power corresponding to the number of molecules of that reactant participating in the reaction. This can be expressed mathematically for a general reaction as:

        AA + BB ⇌ CC + DD

In this reaction, A and B are reactants, C and D are products, and a, b, c and d are their stoichiometric coefficients, respectively.

Equilibrium constant

The law of mass action gives rise to the concept of the equilibrium constant, K, which is a measure of the state of equilibrium for a given reaction at a constant temperature. The equilibrium constant for the above reaction is expressed as:

        k = ([c]^c [d]^d) / ([a]^a [b]^b)

In this expression, the square brackets indicate the concentration of a species. The equilibrium constant is an important value in chemistry, because it indicates whether the forward or reverse reaction is favored. A large K value indicates that the products are favored at equilibrium, while a small K value indicates that the reactants are favored.

Understanding chemical equilibrium

Chemical equilibrium occurs when the rate of the forward reaction is equal to the rate of the reverse reaction, resulting in constant concentrations of reactants and products over time. This dynamic process can be better understood by examining a simple reversible reaction:

        N₂ + 3H₂ ⇌ 2NH₃

In this reaction, nitrogen gas reacts with hydrogen gas to form ammonia. At equilibrium, the rate at which nitrogen and hydrogen form ammonia is equal to the rate at which ammonia dissociates back into nitrogen and hydrogen. The equilibrium constant expression for this reaction is:

        K = ([NH₃]^2) / ([N₂][H₂]^3)

A visual representation

Looking at the law of mass action can help clarify these concepts. Consider a container where the reaction N₂ + 3H₂ ⇌ 2NH₃ occurs. The equilibrium represents the balance between the two sides of the reaction:

Reactants Products N₂ H₂ H₂ H₂ NH₃ NH₃

At equilibrium, there is a specific ratio of N₂, H₂, and NH₃, as determined by the equilibrium constant K

Text example for clarification

Consider a simple chemical reaction where the equilibrium constant is given. Suppose we have this reaction:

        a + 2b ⇌ c

Assume that the equilibrium constant K for this reaction is 10. If the concentration of A at equilibrium is 2 mol/L and B is 1 mol/L, you can calculate the concentration of C using the equilibrium constant expression:

        k = [c] / ([a][b]^2)

Insert known values:

        10 = [c] / (2 * 1^2)

Solve for [C]:

        [C] = 10 * 2 = 20 mol/L

Applications of the law of mass action

The law of mass action is not just an academic concept; it has real applications in various fields such as chemical engineering, environmental science, pharmacology and biochemistry. Here are some of the applications:

  • Chemical manufacturing: Industries use this law to optimize conditions for maximum yield of desired products.
  • Medicine: This principle is used to understand how drugs interact and reach equilibrium concentrations in the bloodstream.
  • Environmental science: Reactions in the atmosphere or hydrosphere are analyzed, often using equilibrium principles, to assess the levels and effects of pollution.

Factors affecting chemical equilibrium

Several factors can affect the position of the equilibrium, thereby affecting the equilibrium constant:

  • Temperature: An increase or decrease in temperature can change K because it affects the forward and reverse reaction rates differently.
  • Concentration: According to Le Chatelier principle a change in concentration may shift the equilibrium position.
  • Pressure: For gas-phase reactions, changes in pressure can shift the equilibrium position, especially if the number of moles of gas differs between the reactants and products.

Le Chatelier's principle and its relation to the law of mass action

Le Chatelier's principle is closely related to the law of mass action and states that if an external change is imposed on a system at equilibrium, the system adjusts itself partly to counteract the change and a new equilibrium is established. For example:

  • Addition of reactants: If more reactants are added to a system, the equilibrium will shift to the right, favoring the formation of products.
  • Removal of products: Removing products from the system shifts the equilibrium to the right, favoring product formation.
  • Change in temperature: Increasing the temperature favors the system in the endothermic reaction direction. Decreasing it favors the exothermic direction.

Conclusion

The law of mass action provides an important framework for understanding chemical equilibrium. Its mathematical formulation not only allows chemists to predict the outcome of reactions, but also to manipulate conditions to achieve desired results. Understanding this law is essential for anyone studying or working with chemical reactions and it helps to develop a deeper understanding of the dynamic processes in chemical systems. The equilibrium constant, although seemingly simple, is a comprehensive indicator of system behavior and guides us through the complex world of chemical reactions.


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Law of mass action

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