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


Chemical kinetics is a branch of physical chemistry that deals with the study of the rates of chemical reactions. Unlike thermodynamics, which indicates whether a reaction is possible, kinetics provides information about the speed or rate at which a reaction occurs and how different variables affect this rate.

Basic concepts and definitions

The core of chemical kinetics involves understanding how quickly or slowly reactions proceed. The rate of a reaction can be affected by many factors, including the concentration of the reactants, temperature, the presence of catalysts, and the physical state of the reactants.

Rate of reaction

The rate of a chemical reaction refers to how fast the reactant is consumed or how fast the product is formed. It can be expressed as the change in concentration of the reactant or product over time. Mathematically, it is given as:

Rate = -Δ[Reactant]/Δt = Δ[Product]/Δt

Where Δ[reactant] and Δ[product] represent the change in concentration of reactant and product respectively over time interval Δt. Negative sign is used for reactant because its concentration decreases with time.

Factors affecting the reaction rate

Several factors affect how fast a chemical reaction occurs:

  • Concentration: Generally, increasing the concentration of reactants increases the reaction rate. This is due to an increase in the number of reactant particles that can collide to form products.
  • Temperature: Increasing the temperature generally increases the rate of the reaction. Higher temperatures impart more kinetic energy to the reactant molecules, increasing the frequency and energy of collisions.
  • Catalyst: Catalysts are substances that speed up a reaction without being consumed in the process. They provide an alternative pathway with a lower activation energy for the reaction to occur.
  • Physical state: The physical state of the reactants (solid, liquid, gas) can affect the reaction rate. For example, gases generally react more quickly than solids because their particles collide more often.
  • Surface area: For reactions involving solids, increasing the surface area (e.g., by grinding into a powder) allows more particles to interact and react simultaneously, increasing the reaction rate.

Reaction rate law

The reaction rate law or rate equation describes how the rate of a reaction depends on the concentration of the reactants. For a general reaction:

aA + bB → cC + dD

The rate law can be expressed as:

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

Here, k is the rate constant, and [A] and [B] are the concentrations of the reactants. The exponents m and n are the reaction orders, which are determined experimentally and indicate how the rate depends on the concentration of each reactant.

Order of reaction

The overall order of the reaction is the sum of the exponents in the rate law. For example, if m = 1 and n = 1, then the reaction is second-order overall. Reaction orders can be zero, first, second, etc., and may even be partial or negative in some complex reactions.

Integrated rate law

Integrated rate laws relate the concentrations of reactants or products to time. They help determine the age of a reaction or predict concentrations at future times. There are different integrated rate laws for zero-order, first-order, and second-order reactions.

Zero-order reactions

In zero-order reactions, the rate is independent of the concentration of the reactants. The integrated rate law is:

[A] = [A]0 - kt

where [A]0 is the initial concentration, k is the zero-order rate constant, and t is time.

First order reactions

For first-order reactions, where the velocity is proportional to the concentration of one reactant, the integrated velocity law is:

ln([A]/[A]0) = -kt

This can be rearranged to find the concentration as a function of time:

[A] = [A]0e-kt

Second order reactions

Second-order reactions depend either on the square of the concentration of one reactant or on the concentrations of two different reactants. The integrated rate law for a second-order reaction involving one reactant is:

1/[A] = 1/[A]0 + kt

Collision theory and transition state theory

Two major theories explain how reactions occur at the molecular level: collision theory and transition state theory.

Collision theory

Collision theory states that for molecules to react, they must collide with sufficient energy and in the proper orientation. The minimum energy required for a reaction to occur is known as the activation energy (Ea).

EAReactantsProducts

In the example above, the reactants must overcome an energy barrier (Ea) to form products. The higher Ea is, the slower the reaction, because fewer molecules have enough energy to react.

Transition state theory

Transition state theory focuses on the idea of an activated complex, a transient configuration during a reaction when bonds are simultaneously being broken and formed. This complex, also known as the transition state, occurs at the energy peak of the reaction.

Transition stateReactantsProductsActive Complex

Understanding these principles helps chemists control reaction conditions, such as temperature and pressure, and use catalysts that alter the reaction pathway, lowering the energy barrier without changing the initial and final states.

Catalysts and enzymes

Catalysts play an important role in increasing reaction rates by providing alternative reaction pathways with lower activation energies. Enzymes, biological catalysts, are particularly important in biochemical reactions.

Types of catalysts

Catalysts can be classified as homologous or heterologous:

  • Homogeneous catalysts: These catalysts exist in the same phase as the reactants, usually in solution. An example of this is the acid-catalyzed esterification of carboxylic acids.
  • Heterogeneous catalysts: These catalysts are in a different state from the reactants, often solid catalysts with liquid or gaseous reactants. The use of platinum in catalytic converters for cars is an example of this.

Enzymes are remarkable in their specificity and efficiency as biological catalysts. They often operate under mild conditions (e.g., body temperature, neutral pH), controlling complex biochemical pathways.

Mechanism of reactions

The step-by-step sequence of elementary reactions by which the overall chemical change occurs is known as the reaction mechanism. Mechanisms provide information about which bonds are broken and formed during the reaction, what is the order of these events, and what intermediates are involved.

Primary reactions

Elementary reactions are simple reactions that describe a single molecular event, such as a bond-breaking or bond-forming step. These often involve one or two molecules, and their stoichiometry matches the molecularities:

  • Monomer: A single molecule undergoes decomposition or isomerization. Example: A2 → 2A
  • Bimolecular: It involves collision between two molecules. Example: A + B → Products
  • Termolecular: This involves three molecules, although this is rare due to the low probability of simultaneous collisions. Example: 2A + B → product

Rate determination stage

The rate determining step (RDS) is the slowest step in the mechanism and controls the overall rate of the reaction. It acts as a bottleneck, limiting how fast the product can form. Identifying the RDS is necessary to obtain the rate law from the mechanism.

Case studies and examples

Let's look at some classic examples of chemical kinetics:

Haber process

The Haber process synthesizes ammonia from nitrogen and hydrogen gases. The reaction is as follows:

N2 (g) + 3H2 (g) → 2NH3 (g)

This process is catalyzed by iron, which increases the rate and makes production on an industrial scale possible. The rate of ammonia formation is affected by temperature, pressure, and the presence of a catalyst.

Enzyme kinetics: Michaelis-Menten mechanism

This mechanism describes how enzymes catalyze reactions with substrate molecules. The simplified reaction is:

E + S ↔ ES → E + P

where E is the enzyme, S is the substrate, ES is the enzyme-substrate complex, and P is the product. The rate equation is expressed as:

Vmax[S]=KmRate

v = (Vmax[S])/(Km + [S])

Here, Vmax is the maximum reaction rate, and Km is the Michaelis constant, both of which help in understanding enzyme efficiency.

Conclusion

Chemical kinetics is a powerful tool in understanding reaction rates and mechanisms. It helps chemists design and optimize reactions for industrial processes, biological systems, or the development of new materials. By analyzing rates and exploring factors such as catalysts and reaction mechanisms, kinetic studies provide insight and control over chemical transformations important for technological and scientific advances.


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Chemical kinetics

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