Graduate
  1. Physical Chemistry  1.1. Thermodynamics  1.1.1. Laws of Thermodynamics  1.1.2. Enthalpy and heat capacity  1.1.3. Entropy and free energy  1.1.4. Phase Equilibrium  1.1.5. Statistical thermodynamics  1.1.6. Thermodynamic cycle  1.1.7. Fugacity and activity  1.2. Quantum Chemistry  1.2.1. Principles of quantum mechanics  1.2.2. Schrödinger Equation  1.2.3. Particle in a box  1.2.4. Operators and eigenvalues  1.2.5. Atomic orbitals  1.2.6. Molecular orbital theory  1.2.7. Perturbation theory  1.2.8. Valence bond theory  1.2.9. Hybridization and chemical bonding  1.3. Chemical kinetics  1.3.1. Rate laws and reaction mechanisms  1.3.2. Collision theory  1.3.3. Transition state theory  1.3.4. Enzyme Kinetics  1.3.5. Chain Reactions and Polymerization  1.3.6. Photochemical reactions  1.4. Statistical mechanics  1.4.1. Partitioning functions  1.4.2. Molecular distribution functions  1.4.3. Boltzmann Distribution  1.4.4. Bose–Einstein and Fermi–Dirac statistics  1.5. Spectroscopy  1.5.1. Rotational Spectroscopy  1.5.2. Vibrational Spectroscopy  1.5.3. Electronic Spectroscopy  1.5.4. Nuclear magnetic resonance spectroscopy  1.5.5. Mass Spectrometry  1.5.6. Raman Spectroscopy  1.5.7. Electron paramagnetic resonance spectroscopy  1.6. Surface and colloidal chemistry  1.6.1. Absorption isotherm  1.6.2. Surface tension and wetting  1.6.3. Colloidal Stability  1.6.4. Catalysis  1.6.5. Emulsions and micelles  1.7. Electrochemistry  1.7.1. Nernst equation  1.7.2. Electrochemical cells  1.7.3. Conductivity and mobility  1.7.4. War  1.7.5. Fuel cells  1.7.6. Electrolysis
  2. Organic chemistry  2.1. Reaction mechanism  2.1.1. Nucleophilic substitution reactions  2.1.2. Electrophilic addition reactions  2.1.3. Elimination reactions  2.1.4. Rearrangement reactions  2.1.5. Radical reactions  2.1.6. Pericyclic reactions  2.2. Spectroscopy and structural determination  2.2.1. UV-Vis Spectroscopy  2.2.2. IR Spectroscopy  2.2.3. NMR Spectroscopy  2.2.4. Mass Spectrometry  2.2.5. X-ray crystallography  2.3. Stereoscopic  2.3.1. Chirality and optical activity  2.3.2. Structural analysis  2.3.3. Geometrical isomerism  2.3.4. Dynamic Stereochemistry  2.4. Organometallic Chemistry  2.4.1. Organolithium and organomagnesium reagents  2.4.2. Palladium-catalyzed cross-coupling reactions  2.4.3. Transition Metal Complexes  2.4.4. Metal–carbon bond  2.5. Polymer chemistry  2.5.1. Polymerization mechanism  2.5.2. Polymer Characteristics  2.5.3. Biodegradable Polymer  2.5.4. Conducting polymer  2.5.5. Supramolecular polymers  2.6. Medicinal chemistry  2.6.1. Drug design and development  2.6.2. Pharmacokinetics and pharmacodynamics  2.6.3. Structure-activity relationships  2.6.4. Molecular docking and drug screening
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Crystal field theory  3.1.2. Ligand field theory  3.1.3. Spectrochemical Series  3.1.4. Chelation and stability  3.2. Organometallic Chemistry  3.2.1. Metal Carbonyls  3.2.2. Catalysis by organometallic complexes  3.2.3. Metallocenes  3.3. Bio-inorganic chemistry  3.3.1. Metalloproteins and enzymes  3.3.2. The role of metals in biological systems  3.3.3. Metal ion transport and storage  3.4. Solid state chemistry  3.4.1. Crystal Structures  3.4.2. Band theory of solids  3.4.3. Superconductors  3.4.4. Defects in the crystal  3.5. Lanthanides and Actinides  3.5.1. Electronic configuration in lanthanides and actinides  3.5.2. Coordination chemistry of the lanthanides  3.5.3. Magnetic Properties of Lanthanides
  4. Analytical chemistry  4.1. Chromatography  4.1.1. Gas Chromatography  4.1.2. High Performance Liquid Chromatography  4.1.3. Thin Layer Chromatography  4.2. Spectroscopic Techniques  4.2.1. Atomic absorption spectroscopy  4.2.2. X-ray diffraction  4.2.3. Inductively coupled plasma spectroscopy  4.3. Electroanalytical methods  4.3.1. Potentiometry  4.3.2. Voltammetry  4.3.3. Colometry  4.4. Mass Spectrometry  4.4.1. Ionization Technique  4.4.2. Fragmentation Pattern  4.5. Chemometrics  4.5.1. Multidisciplinary analysis  4.5.2. Machine Learning in Chemistry
  5. Theoretical and Computational Chemistry  5.1. Molecular dynamics simulation  5.1.1. Force fields and energy minimization  5.1.2. Monte Carlo simulation in molecular dynamics simulations  5.2. Quantum chemical methods  5.2.1. Hartree–Fock theory  5.2.2. Density functional theory  5.2.3. Semi-empirical methods  5.3. Computational drug design  5.3.1. Molecular Docking  5.3.2. QSAR Modeling  5.3.3. Virtual Screening
  6. Biochemistry  6.1. Enzyme Kinetics  6.1.1. Michaelis-Menten kinetics  6.1.2. Inhibition mechanism  6.2. Metabolism and Bioenergetics  6.2.1. Glycolysis  6.2.2. Citric acid cycle  6.2.3. Electron transport chain  6.3. molecular Biology  6.3.1. DNA replication and repair  6.3.2. Protein synthesis  6.3.3. Gene Regulation  6.4. Structural Biochemistry  6.4.1. Protein Folding  6.4.2. Membrane Biophysics
  7. Environmental Chemistry  7.1. Atmospheric Chemistry  7.1.1. Greenhouse gases  7.1.2. Ozone depletion  7.1.3. Air pollutants and smoke  7.2. Water Chemistry  7.2.1. pH and water hardness  7.2.2. Wastewater treatment  7.2.3. Aquatic Chemistry  7.3. Soil Chemistry  7.3.1. Heavy metal contamination  7.3.2. Soil pH and Buffering  7.3.3. Nutrient cycling  7.4. Toxicology and Chemical Safety  7.4.1. Toxicokinetics  7.4.2. Risk Assessment in Toxicology and Chemical Safety in Environmental Chemistry  7.4.3. Effects of pollutants on the environment

Graduate → Physical Chemistry → Chemical kinetics ↓


Rate laws and reaction mechanisms


Chemical kinetics is a branch of physical chemistry that studies the speed or rate at which chemical reactions occur and the mechanisms by which these reactions occur. Two fundamental concepts in the study of chemical kinetics are rate laws and reaction mechanisms. Understanding these concepts provides insight into the kinetics of chemical reactions and helps control conditions to optimize specific reactions. In this lesson, we will explore these concepts in depth using simple language and examples.

Rate laws

The rate law of a chemical reaction is an equation that relates the rate of the reaction to the concentrations of the reactants. The rate law is determined experimentally and can help understand how changes in conditions affect the speed of the reaction. It is important to note that the rate law cannot be deduced from the stoichiometry of the overall reaction, but rather is determined from experimental observations.

General form of the rate law

For a general reaction where reactants A and B form a product, the rate law can be expressed as:

Rate = k [A]^m [B]^n

Here:

  • Rate is the reaction rate.
  • k is the rate constant, specific to the reaction at a given temperature.
  • [A] is the concentration of reactant A
  • [B] is the concentration of reactant B
  • m and n are the relative reaction orders of A and B, respectively. These are determined experimentally.

Reaction order

The reaction order shows the dependence of the rate on the concentration of each reactant. Let's look at some common scenarios:

  • Zero-order reactions: The rate is independent of the concentration of the reactants.
    The rate law is Rate = k.
  • First order reactions: The rate is directly proportional to the concentration of a reactant.
    The rate law is Rate = k [A].
  • Second order reactions: The rate is proportional to the square of the concentration of a reactant, or proportional to the product of two reactants.
    The rate law can be Rate = k [A]^2 or Rate = k [A][B].

Rate fixing law

Rate laws are typically determined by the method of initial rates, which involves conducting a series of experiments to measure initial reaction rates at different concentrations. This method helps establish how the reaction rate depends on the concentrations of the reactants.

Example: Simple decomposition reaction

Consider a simple decomposition reaction where compound X decomposes to form Y:

X → Y

Suppose experimental data show that doubling the concentration of X doubles the reaction rate. This suggests a first-order reaction with respect to X, which is described by the rate law:

Rate = k [X]

Reaction mechanism

The reaction mechanism provides a detailed description of the steps through which reactants are transformed into products. It consists of a series of elementary reactions, each of which represents a single step in the overall reaction.

Primary reactions

Elementary reactions are single-step processes with a definite molecularity, which indicates the number of reactant molecules involved. Common types include:

  • Unimolecular reactions: It involves one molecule. Example: A → Products
  • Bimolecular reactions: It involves two molecules. Example: A + B → Products
  • Termolecular reactions: Rarely do they involve more than three molecules. Example: A + B + C → Products

Rate determination stage

The rate determining step is the slowest step in the reaction mechanism that determines the overall reaction rate. Understanding this step is important in proposing a valid reaction mechanism that aligns with the observed rate law.

Example: Reaction mechanism of N2O5 decomposition

Consider the decomposition of dinitrogen pentoxide (N2O5):

2 N2O5 → 4 NO2 + O2

A possible mechanism may involve the following steps:

  • Step 1: N2O5 → NO2 + NO3 (fast)
  • Step 2: N2O5 + NO3 → 3 NO2 (slow)

The second step is the slowest and is therefore the rate-determining step. The rate law corresponding to this mechanism is:

Rate = k [N2O5]

Graphical representation of the response profile

Reaction profiles depict the energy changes during a chemical reaction. They are valuable tools for understanding the energy barriers associated with each step of the mechanism. Consider the following example:

Reactants intermediate Products Transition state 1 Transition state 2 energy

This reaction profile represents a multi-step mechanism with two transition states and an intermediate state. The vertical axis indicates the energy levels, while the horizontal axis shows the reaction coordinates.

Experimental validation of the mechanisms

Experimental validation of a proposed reaction mechanism involves comparing the predicted rate law based on the mechanism to the observed rate law. Consistency between the two supports the mechanism, although it cannot conclusively prove its validity due to the possible existence of alternative mechanisms giving the same rate law.

Example: Verification process

For the N2O5 mechanism above, suppose experiments determined that the reaction is first-order with respect to N2O5. The proposed mechanism predicting a first-order rate law is thus consistent and potentially valid.

Conclusion

Understanding rate rules and reaction mechanisms is essential in chemical kinetics, as they provide important insights into the dynamics of chemical reactions. By conducting experiments to determine rate rules and propose mechanisms, chemists can predict and control how reactions proceed, allowing the optimization of industrial processes and the synthesis of desired products efficiently.

Through experimental determination and verification of rate laws and mechanisms, chemists can gain a deeper understanding of reaction kinetics, which contributes to progress in chemical research and industry.


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Rate laws and reaction mechanisms

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