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 ↓


Reaction quotient


The concept of the reaction quotient, represented by the symbol Q, is an essential part of understanding chemical equilibrium. It allows chemists to determine the direction in which a chemical reaction will proceed at any time other than equilibrium. This understanding is important for both predicting and controlling chemical reaction processes, especially in industrial and laboratory settings. In this comprehensive guide, we will explore the reaction quotient, its calculation, significance, and applications through text explanations and visual examples.

Understanding reaction quotient

In simple terms, the reaction quotient (Q) is a measure of the relative amounts of products and reactants present during a reaction at a particular time. It is used to estimate the direction in which the reaction must proceed to reach equilibrium. The mathematical expression of the reaction quotient is similar to that of the equilibrium constant (K), but they differ in one important aspect:

  • Equilibrium constant, K: calculated using the concentrations of reactants and products when the system is at equilibrium.
  • Reaction quotient, Q: calculated using the concentrations of reactants and products at any given time, not necessarily at equilibrium.

For a general chemical reaction:

    AA + BB ⇌ CC + DD

The reaction quotient (Q) is given by the following expression:

    Q = ([C]^c [D]^d) / ([A]^a [B]^b)

Where:

  • [A], [B], [C], [D] are the molar concentrations of reactants and products.
  • a, b, c, d are the coefficients in a balanced chemical equation.

Comparison of Q and K

The value of Q can be compared with the equilibrium constant K to determine which direction the reaction will shift in order to reach equilibrium:

  • If Q < K: The reaction will proceed in the forward direction, and reach equilibrium by converting reactants into products.
  • If Q = K: The system is in equilibrium, and there will be no net change in the concentrations of reactants and products.
  • If Q > K: The reaction will proceed in the reverse direction, converting products back into reactants to reach equilibrium.

Example response

Let us consider the following reaction as an example:

    2 NO₂(g) ⇌ N₂O₄(g)

Suppose that at a given time you have the following concentrations:

  • [NO₂] = 0.20 M
  • [N₂O₄] = 0.10 M

And the equilibrium constant K for this reaction at a specific temperature is 0.50.

The reaction quotient is calculated as:

    Q = ([N₂O₄]) / ([NO₂]^2)
    Q = (0.10) / (0.20)^2
    Q = 0.10 / 0.04
    Q = 2.5

Given that Q (2.5) is greater than K (0.50), the reaction will shift in the opposite direction to reach equilibrium, converting N₂O₄ back to NO₂.

Visual representation

Consider the following graphical representation of the reaction and observe how Q changes:

Reactants Products Balance point (K) Q < K, proceed Q > K, move backward

Practical applications

The reaction quotient is important to chemists in several ways:

Chemical synthesis

In industrial processes involving the synthesis of chemicals, knowing the reaction quotient allows engineers to adjust conditions to optimize product yield.

Environmental monitoring

Environmental scientists use Q to assess the status of chemical reactions in natural processes, such as the dissolution of gases in water bodies.

Laboratory experiment

In the lab, researchers calculate Q to analyze the progress of reactions, ensuring they reach the desired stage in time.

Examples in analytical chemistry

Consider the following equilibrium in a typical acid-base reaction:

    CH₃COOH(aq) + OH⁻(aq) ⇌ CH₃COO⁻(aq) + H₂O(l)

If the initial concentration is given:

  • [CH₃COOH] = 1.0 M
  • [OH⁻] = 0.5 M
  • [CH₃COO⁻] = 0.2 M

To find Q we omit water because it is a pure liquid:

    Q = ([CH₃COO⁻]) / ([CH₃COOH][OH⁻])
    Q = (0.2) / (1.0 * 0.5)
    Q = 0.4

If K is given as 0.50, since Q (0.4) < K (0.50), the reaction will proceed in the forward direction and more CH₃COO⁻ will be formed.

Conclusion

The reaction quotient is a versatile and powerful tool in chemistry, helping to predict the behavior and progress of reactions. By understanding Q in relation to the equilibrium constant K, chemists can influence and control reactions to efficiently achieve desired chemical transformations. This understanding lays the foundation for more advanced chemical studies and practical applications in both industry and research.


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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.

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Reaction quotient

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