PHD
  1. Inorganic chemistry  1.1. Coordination chemistry  1.1.1. Ligand field theory  1.1.2. Crystal field theory  1.1.3. Molecular Orbital Theory for Coordination Compounds  1.1.4. Stability of Coordination Compounds  1.1.5. Electronic spectra of coordination compounds  1.1.6. Isomerism in Coordination Compounds  1.1.7. Kinetics and mechanism of coordination reactions  1.2. Organometallic Chemistry  1.2.1. Metal Carbonyls  1.2.2. Metal alkyl and aryl  1.2.3. Fischer and Schrock carbonaceans  1.2.4. Oxidative Addition and Reductive Elimination  1.2.5. Metal Hydrides  1.2.6. Catalysis by organometallic compounds  1.2.7. Metallocenes and Sandwich Complexes  1.2.8. CH activation  1.3. Solid state chemistry  1.3.1. Crystal structures and lattices  1.3.2. Band theory and electrical properties  1.3.3. Defects in the crystal  1.3.4. Synthesis of solid state materials  1.3.5. Magnetic and optical properties of solids  1.3.6. Superconductivity  1.3.7. Solid state ionics  1.4. Bio-inorganic chemistry  1.4.1. Metalloproteins and enzymes  1.4.2. Metal ion transport and storage  1.4.3. The role of metals in medicine  1.4.4. Bioinspired catalysis  1.4.5. Metal–DNA interactions  1.4.6. Metal Complexes in Medical Science  1.5. Lanthanides and Actinides  1.5.1. Electronic configuration and oxidation states in lanthanides and actinides  1.5.2. Spectral and Magnetic Properties in Lanthanides and Actinides  1.5.3. Separation and Extraction of Lanthanides and Actinides  1.5.4. Coordination chemistry of f-block elements  1.6. Main Group Chemistry  1.6.1. Chemistry of boron compounds  1.6.2. Chemistry of Silicon and Germanium Compounds  1.6.3. Chemistry of Phosphorus and Sulfur Compounds  1.6.4. Organosilicon chemistry  1.6.5. Noble Gas Chemistry
  2. Organic chemistry  2.1. Reaction mechanism  2.1.1. Nucleophilic substitution reactions  2.1.2. Electrophilic addition and substitution reactions  2.1.3. Elimination reactions  2.1.4. Rearrangement reactions  2.1.5. Radical reactions  2.1.6. Pericyclic Reactions  2.1.7. Photochemical reactions  2.2. Stereoscopic  2.2.1. Chirality and optical activity  2.2.2. Structural analysis  2.2.3. Stereoelectronic effects  2.2.4. Asymmetric synthesis  2.2.5. Dynamic Stereochemistry  2.3. Organometallic Chemistry in Organic Synthesis  2.3.1. Organolithium and organomagnesium reagents  2.3.2. Transition metal catalyzed coupling reactions  2.3.3. Palladium-catalyzed reactions  2.3.4. Olefin Metathesis  2.3.5. Gold and silver catalysis  2.4. Natural products chemistry  2.4.1. Alkaloids and terpenoids  2.4.2. Steroids and Flavonoids  2.4.3. Total synthesis of natural products  2.4.4. Biosynthesis of natural products  2.5. Supramolecular Chemistry  2.5.1. Host-Guest Chemistry  2.5.2. Self-assembly and molecular recognition  2.5.3. Supramolecular Catalysis  2.5.4. Molecular Machines
  3. Physical Chemistry  3.1. Thermodynamics  3.1.1. Laws of Thermodynamics  3.1.2. Chemical Potential and Equilibrium  3.1.3. Statistical thermodynamics  3.1.4. Non-equilibrium thermodynamics  3.2. Quantum Chemistry  3.2.1. Schrödinger equation and its applications  3.2.2. Molecular orbital theory  3.2.3. Density functional theory  3.2.4. Post-Hartree–Fock methods  3.3. Chemical kinetics  3.3.1. Reaction rate theory  3.3.2. Mechanism and rate laws  3.3.3. Catalysis and enzyme kinetics  3.3.4. Reaction dynamics  3.4. Spectroscopy and molecular structure  3.4.1. Infrared Spectroscopy  3.4.2. UV-visible spectroscopy  3.4.3. Nuclear Magnetic Resonance Spectroscopy  3.4.4. Mass Spectrometry  3.4.5. Raman Spectroscopy  3.5. Surface and Colloid Chemistry  3.5.1. Absorption and Catalysis  3.5.2. Surfactants and micelles  3.5.3. Colloidal Stability  3.5.4. Nanomaterials and Interfaces
  4. Analytical chemistry  4.1. Chromatography  4.1.1. Gas Chromatography  4.1.2. High Performance Liquid Chromatography  4.1.3. Thin Layer Chromatography  4.1.4. Ion Chromatography  4.2. Electroanalytical techniques  4.2.1. Voltammetry and Polarography  4.2.2. Conductometry and potentiometry  4.2.3. Colometry  4.3. Spectroscopic Methods  4.3.1. Atomic absorption spectroscopy  4.3.2. Fluorescence Spectroscopy  4.3.3. X-ray diffraction  4.3.4. Electron paramagnetic resonance spectroscopy  4.4. Chemometrics  4.4.1. Data processing and statistical methods  4.4.2. Multidisciplinary analysis  4.4.3. Machine Learning in Analytical Chemistry
  5. Theoretical and Computational Chemistry  5.1. Molecular dynamics simulation  5.2. Quantum chemistry methods  5.3. Computational drug design  5.4. Machine Learning in Chemistry  5.5. Ab initio molecular dynamics
  6. Biophysics and Medicinal Chemistry  6.1. Drug discovery and design  6.2. Enzyme kinetics and inhibition  6.3. Chemical biology approach  6.4. Protein-ligand interactions  6.5. Bioorthogonal chemistry
  7. Materials chemistry  7.1. Polymer chemistry  7.1.1. Polymerization mechanism  7.1.2. Functional Polymers  7.1.3. Biodegradable Polymer  7.1.4. Conducting and semiconducting polymers  7.2. Nano Chemistry  7.2.1. Nanoparticles and nanostructures  7.2.2. Carbon-based nanomaterials  7.2.3. Quantum dots  7.3. Energy and Environmental Chemistry  7.3.1. Battery and Fuel Cell Chemistry  7.3.2. Green chemistry and sustainable materials  7.3.3. Photocatalysis

PHD → Physical Chemistry → Chemical kinetics ↓


Mechanism and rate laws


Chemical kinetics is a branch of physical chemistry that involves the study of the rates of chemical processes. The ability to predict and understand reaction rates is important for controlling chemical processes, efficiently synthesizing materials, and understanding natural phenomena.

Kinetic mechanism

A reaction mechanism is the step-by-step sequence of elementary reactions by which an overall chemical transformation occurs. Each step (also called an elementary step or reaction) has its own microscopic reaction rate and balance of mass and energy.

Early stages

An elementary step refers to a reaction that occurs in a single kinetic step at the molecular level. These steps cannot be further divided and usually involve a maximum of two or three reactive species.

Example of mechanism

Consider the following reaction mechanism:

    Step 1: A + B → C (slow)
    Step 2: C + B → D (fast)

Overall reaction: A + 2B → D

In this mechanism, the first step is the rate determining step because it is the slowest and thus controls the rate of the overall reaction.

Rate laws

A rate law is a mathematical equation that relates the reaction rate to the concentration of reactants. Rate laws are determined experimentally and cannot be derived from the stoichiometric coefficients of the balanced equation. Their general form is:

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

Where:

  • k is the rate constant and it is specific for a particular reaction at a given temperature.
  • [A] and [B] are the molar concentrations of reactants A and B.
  • m and n are the reaction orders with respect to each reactant, and they show how the rate is affected by the concentration of the reactants.

Example of a rate law

Consider the decomposition of hydrogen peroxide:

2H 2 O 2 → 2H 2 O + O 2

Let the experimentally determined rate law for this reaction be:

Rate = k [H 2 O 2 ]^1

This shows that the reaction is first order with respect to hydrogen peroxide.

Determination of mechanisms via rate laws

To extract a possible reaction mechanism from a rate law, one must ensure that the proposed mechanism can lead to the same rate law form as observed experimentally. Typically, the rate-determining step is used to extract the rate law. Consider the mechanism:

    Step 1: A + B → C (slow)
    Step 2: C → D (fast)

Given that the first step is slow, we can write the rate law as:

Rate = K [A] [B]

This rate law corresponds to the slow (i.e., rate-determining) step. Consistency with the experimental rate law can validate the proposed mechanism.

Example with rate law consistency

Consider the response:

2NO3 + O2 → 2NO3

Experimental rate law: Rate = k [NO]^2 [O 2 ]

Mechanism:

    Step 1: NO + O 2 ⇌ NO 3 (rapid equilibrium)
    Step 2: NO 3 + NO → 2NO 2 (slow)

For the rapid equilibrium phase, we can say:

[NO 3 ] = [NO][O 2 ]

Substituting the slow step, we get:

Rate = k [NO 3 ] [NO] = k' [NO] [O 2 ] [NO] = k' [NO]^2 [O 2 ]

This reflects the experimentally observed rate law.

Visualization of reaction mechanisms

Mechanisms can often be visualized through reaction coordinate diagrams that show the energy changes during the reaction.

A+B D Transition state C

In this diagram, the Y-axis represents potential energy, while the X-axis represents the progress of the reaction. The peaks represent transition states, and the valleys represent intermediate states.

Endothermic vs. exothermic

Whether a reaction is endothermic or exothermic can be judged from the energy levels of the reactants and products. If the products are lower in energy, the reaction is endothermic, releasing heat into the surroundings. Conversely, if the products are higher in energy than the reactants, the reaction is endothermic.

Experimental determination of rate laws

Determining rate laws experimentally involves measuring the reaction rate at different reactant concentrations, often using the method of initial rates. By varying the concentration of one reactant while keeping the others constant, one can estimate the order with respect to that reactant.

Initial rate method

This method involves measuring the rate of the reaction under different concentrations, immediately after it starts. Consider a hypothetical reaction:

A + B → Product

Perform several experiments with different initial concentrations of A and B. Then, measure the initial reaction rates:

[A] (M) [B] (M) Starting rate (m/s)
0.1 0.1 2.0 × 10-3
0.2 0.1 4.0 × 10-3
0.1 0.2 2.0 × 10-3

From these results, the reaction order with respect to A and B can be estimated. Here, doubling the concentration of A doubles the rate, which shows first-order dependence on A, while doubling B does not cause any change in the rate, which shows zero-order dependence on B. Thus, the rate law is:

Rate = k [A]^1 [B]^0 = k [A]

Factors affecting the reaction rate

The rate of a chemical reaction is influenced by several factors:

  • Concentration: Increasing reactant concentration generally increases the reaction rate due to more frequent collisions.
  • Temperature: Increasing temperature increases the kinetic energy, resulting in more frequent and more energetic collisions.
  • Catalyst: Catalysts increase the reaction rate by providing an alternative pathway with a lower activation energy, without being consumed.
  • Surface area: Greater surface area of the solid reactant increases the reaction rate.

Chemical reactions are intricately linked to the world we live in. Understanding and using the principles of reaction mechanisms and rate laws provide chemists with powerful tools for innovation and discovery in a variety of fields.


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Mechanism and rate laws

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