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 ↓


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.


Chemical equilibrium occurs when a chemical reaction and its opposite reaction proceed at the same rate. In this situation, the concentrations of reactants and products remain constant over time, although they are not necessarily equal. This situation is characterized by the equilibrium constant, which provides valuable information about the extent of the reaction. In this detailed guide, we will understand the concepts and calculations related to the equilibrium constant in a simple, understandable way.

Understanding the equilibrium constant

The equilibrium constant, denoted as K, represents the ratio of the concentration of products to the concentration of reactants at equilibrium. The general form of a reversible chemical reaction is:

    AA + BB ⇌ CC + DD

where A and B are reactants, C and D are products, and a, b, c, d are their respective stoichiometric coefficients. The equilibrium constant expression for this reaction is given as:

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

Here, [A], [B], [C], [D] denote the molar concentrations of the respective species at equilibrium. The subscript c in K c indicates that the equilibrium constant is expressed in terms of concentration.

Visual example of equilibrium constant expression

Consider the response:

    N 2 + 3H 2 ⇌ 2NH 3

The equilibrium constant expression for this reaction will be:

    K c = ([NH 3 ] 2 )/([N 2 ][H 2 ] 3 )

This expression shows that the concentration of ammonia, NH 3, is square, while the concentration of hydrogen, H 2, is cubic, reflecting their stoichiometric coefficients in the balanced equation.

Different types of equilibrium constants

The equilibrium constant can be expressed in different terms, depending on which concentrations or activities are practically measurable. Here are the common types:

KC - concentration based

This is the most common type of equilibrium constant. It uses molar concentrations (mol/L) to express the equilibrium between gaseous or aqueous reactants and products.

KP - pressure based

For gaseous reactions, the equilibrium constant can also be represented in terms of partial pressures. This expression is similar to the concentration-based expression, but uses partial pressures (atm, bar):

    K P = ( P C C P D D ) / ( P A A P B B )

where P X represents the partial pressure of species X

Visual example of K p

If we consider the reaction:

    2SO 2 (g) + O 2 (g) ⇌ 2SO 3 (g)

The equilibrium constant in terms of partial pressure will be:

    K p = (P s O 3 2 ) / (P s O 2 2 P O 2 )

The use of K p is particularly useful when dealing with reactions involving gases at different pressure conditions.

Relation between K c and K p

The relationship between K c and K p for a gaseous reaction at constant temperature is given by the equation:

    K p = K c (RT) Δn

Where:

  • R is the universal gas constant (0.0821 L atm/mol K).
  • T is the temperature in Kelvin.
  • Δn is the change in moles of gas (moles of gaseous products - moles of gaseous reactants).

Visual example:

Consider the response:

    2NO 2 (g) ⇌ N 2 O 4 (g)

Here, Δn = 1 - 2 = -1. The relation between K c and K p is:

    K p = K c (rt) -1

Calculating the equilibrium constant

To calculate the equilibrium constant, you need to know the concentrations of the reactants and products at equilibrium. Here is a step-by-step method for calculating K c :

1. Start with a balanced chemical equation

This is important, because the coefficients in the balanced equation determine the form of the equilibrium constant expression.

2. Determine the initial concentration

Identify the initial concentrations of the reactants and products. In many cases involving gases, these can be provided as partial pressures.

3. Change in concentration

Use stoichiometry and the balanced equation to determine how the concentrations change when the system reaches equilibrium. Use the ICE (initial, change, equilibrium) table to organize these values.

           Initial change balance
  A(aq) [A] 0 -x [A] eq = [A] 0 - x
  B(aq) [B] 0 − Bx [B] eq = [B] 0 − Bx
  c(aq) [c] 0 + cx [c] eq = [c] 0 + cx

Here, x represents the change in concentration depending on the stoichiometry of the reaction.

4. Substitute the equilibrium concentrations in the expression

Insert the equilibrium concentrations obtained from the ICE table into the equilibrium constant expression to solve for K c.

Text example of calculation

Consider the response:

  H 2 (g) + I 2 (g) ⇌ 2HI(g)
  Initial: [H 2 ] = 1.00 M, [I 2 ] = 1.00 M, [HI] = 0.00 M
  At equilibrium: [HI] = 1.50 M

The changes when using the ICE table will be as follows:

           Initial change balance
  h 2 1.00 -x 1.00 - x
  i 2 1.00 -x 1.00 - x
  HI 0.00 +2x 1.50

From [HI] = 1.50 M, we get x = 0.75:

  [h 2 ] = 1.00 - 0.75 = 0.25 m
  [I 2 ] = 1.00 - 0.75 = 0.25 m
  [HI] = 1.50 m

Substitute these into the equilibrium expression:

  K c = ([HI] 2 )/([H 2 ][I 2 ])
        = (1.50 2 )/(0.25*0.25) = 36

The equilibrium constant K c is 36.

Factors affecting the equilibrium constant

While equilibrium constants have a set value for a given reaction at a certain temperature, they can be affected by temperature changes. Here's how temperature affects the equilibrium constant:

Exothermic reactions

For exothermic reactions, an increase in temperature will decrease the value of K, shifting the equilibrium towards the reactants.

Endothermic reactions

For endothermic reactions, an increase in temperature will increase the value of K, shifting the equilibrium towards the products.

Le Châtelier's principle can be used to predict the direction of the shift: an increase in temperature shifts the equilibrium toward absorption of the additional heat.

Le Chatelier's principle, visual example

Consider:

  N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g) + heat

If the temperature increases, the equilibrium shifts to the left, leading to a build-up of reactants, as the system will attempt to absorb the additional heat.

Applications and importance of equilibrium constant

The equilibrium constant is of practical importance in the field of chemistry and industry:

  • Prediction of the direction of the reaction: The magnitude of the equilibrium constant gives information about the direction in which the reaction is favored. A larger K value indicates a greater concentration of products at equilibrium.
  • Reaction extent: Small K values suggest that the reaction mixture contains more reactants than products at equilibrium.
  • Industrial processes: Reactions such as the synthesis of ammonia (Haber process) and sulfuric acid (Contact process) have equilibrium constants that determine efficient operating conditions in large-scale production.

Conclusion

Understanding how to calculate and interpret equilibrium constants is important in chemistry. These constants provide important data about the nature and extent of chemical reactions under equilibrium conditions, which is fundamental to both theoretical studies and practical applications. By mastering the calculations and concepts associated with equilibrium constants, one can predict the behavior of chemical systems in a variety of contexts.


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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.

https://www.chemistryelite.com/ug/1.8.4?ceal=en