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


Chemical equilibrium is an essential concept in chemistry that provides insight into how chemical reactions occur and reach a state of equilibrium. It is not only central to the mastery of chemistry, but is also important in a variety of applications, from industrial processes to biological systems. This comprehensive guide will explore the key aspects of chemical equilibrium, providing a comprehensive understanding of how it works and its implications.

Understanding chemical equilibrium

In general chemistry, chemical equilibrium refers to the state of a reversible reaction where the rate of the forward reaction is equal to the rate of the backward reaction. At this point, the concentrations of reactants and products remain constant over time.

Reactants Products KF KB

Dynamic nature of equilibrium

Chemical equilibrium is dynamic, meaning that reactions continue to occur at the molecular level. Even if there is no gross change in concentration, molecules of reactants continue to transform into products and vice versa.

For example, consider the simple synthesis of water:

2H2(g) + O2(g) ⇌ 2H2O(g)

At a given temperature and pressure in a closed container, the rate at which hydrogen and oxygen dissociate to form water is equal to the rate at which water dissociates to form hydrogen and oxygen. Thus, there is no net change in the concentration of either species.

Equilibrium constant

The equilibrium constant, denoted as K, is a numerical value that relates the concentrations of reactants and products at equilibrium in a reversible reaction. It is derived from the law of mass action and helps to predict the position of equilibrium.

Manifestations of K

For a hypothetical reaction:

aA + bB ⇌ cC + dD

The expression of equilibrium constant is given as:

K = [C]c[D]d / [A]a[B]b

where [A], [B], [C], and [D] are the concentrations of the chemical species at equilibrium, and a, b, c, and d are their stoichiometric coefficients.

Factors affecting the equilibrium constant

The equilibrium constant is only affected by changes in temperature. Changes in temperature can shift the equilibrium constant, causing the value of K to vary. In contrast, changes in concentration and pressure do not affect K

Le Chatelier's principle

Le Chatelier's principle is fundamental in predicting how a change in conditions can alter a state of chemical equilibrium. It states that if a change in conditions disturbs dynamic equilibrium, the state of equilibrium shifts to counteract that change and reestablish equilibrium.

Effects of concentration

Adding or removing a reactant or product shifts the equilibrium so that the equilibrium concentrations are reestablished. For example, adding more reactant will drive the reaction toward the product.

a + b ⇌ c + d +B right shift

Effects of pressure

Changes in pressure affect the equilibrium involving gases. When pressure is increased by decreasing the volume, the equilibrium shifts to the side with fewer gas molecules.

Consider the response:

N2(g) + 3H2(g) ⇌ 2NH3(g)

Increasing the pressure will benefit the product side, which will have fewer moles of gas (4:2 ratio).

Effect of temperature

A change in temperature can shift the equilibrium, depending on whether the reaction is exothermic or endothermic. If the reaction is exothermic, an increase in temperature shifts the equilibrium to the left, toward the reactants.

Applications of chemical equilibrium

Chemical equilibrium principles are applied in many industrial processes, environmental systems, and even the human body.

Industrial applications

The Haber process for ammonia synthesis is a classic example of equilibrium application. By manipulating temperature, pressure, and concentration, ammonia production is optimized:

N2(g) + 3H2(g) ⇌ 2NH3(g)

The conditions chosen are often a compromise to achieve economically viable production rates while maintaining significant yields.

Biological balance

In biological systems, balancing processes regulate vital functions, such as the binding and release of oxygen by hemoglobin:

Hb + O2 ⇌ HbO2

This balance is dynamic and changes depending on the oxygen concentration, facilitating the transport and distribution of oxygen in the body.

Conclusion

Chemical equilibrium is important in understanding reaction mechanisms and system dynamics in chemistry. By examining the equilibrium constant and Le Chatelier's principle, scientists and engineers can control and optimize reactions for desired results.


Undergraduate → 1.8


U
username
0%
completed in Undergraduate

← Prev (1.7.5)
Partition coefficient
Next (1.8.1) →
Law of mass action

Comments 



Chemical equilibrium

https://www.chemistryelite.com/ug/1.8?ceal=en