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 → Kinetics ↓


Activation Energy and Catalyst


Introduction to dynamics

Chemical kinetics is the study of the speed or rate of chemical reactions. It also examines the factors that affect these rates, such as temperature, pressure, and concentration, as well as the mechanisms of reactions. Two key concepts in the study of chemical kinetics are activation energy and catalysis.

Understanding activation energy

First, let's find out what activation energy is. Activation energy is the minimum energy barrier that must be crossed for a chemical reaction to occur. This energy is required to break the bonds of the reactants and start the reaction process.

Visualization of activation energy

Transition state Feedback progress potential energy Reactants Products

The curve in the graph above shows the potential energy during the reaction process. At the peak of the curve, the system is in a transitional state between reactants and products. This point is known as the transition state, and the energy required to reach this state from the initial reactants is the activation energy (E a). The energy input is necessary to break the initial bonds and allow the reaction to proceed.

Calculating activation energy

The Arrhenius equation is commonly used in chemistry to calculate the activation energy of a reaction:

k = A * e-Ea / (RT)

Where:

  • k = rate constant of the reaction
  • A = pre-exponential factor (frequency of collisions with the correct orientation)
  • E a = activation energy
  • R = universal gas constant (8.314 J/mol K)
  • T = temperature in Kelvin

This equation highlights that as the activation energy E a increases, the rate constant k decreases when other factors remain constant, which means the rate of the reaction will slow down.

Example: Functional activation energy

Consider a simple reaction between hydrogen (H 2) and iodine (I 2) to form hydrogen iodide (HI):

H 2 + I 2 → 2HI

The process of forming HI from H 2 and I 2 involves breaking of HH and II bonds and forming new HI bonds. The energy required to break the initial bonds and reach the transition state is the activation energy of this reaction.

Role of catalysts

Catalysts are substances that increase the rate of a chemical reaction without being consumed in the process. They act by providing an alternative reaction pathway with a lower activation energy. Importantly, while they speed up the reaction rate, they do not change the thermodynamics of the reaction. The energy of the reactants and products remains unchanged.

Visualization of catalysts

induced not induced Feedback progress potential energy

As shown in the diagram above, the catalyzed reaction path takes a lower energy path than the uncatalyzed reaction. This lower path means that the reactants require less energy to reach the transition state. As a result, the overall reaction rate increases.

Types of catalysts

Catalysts can be classified into two main categories: homogenetic and heterogeneous.

  1. Homogeneous catalysts: These catalysts are in the same phase as the reactants. Typically, both the catalyst and the reactants are in solution. A common example of this is the use of sulfuric acid (H 2 SO 4) in the esterification of alcohols.
  2. Heterogeneous catalysts: These catalysts exist in a different phase than the reactants, often as a solid with liquid or gaseous reactants. A classic example of this is the use of platinum in catalytic converters to reduce vehicle emissions.

Example: Role of catalysts in lowering activation energy

Consider the decomposition of hydrogen peroxide (H 2 O 2) into water and oxygen. The reaction can be written as:

2H 2 O 2 → 2H 2 O + O 2

Without a catalyst, this reaction proceeds very slowly due to the high activation energy. However, if we use a catalyst such as manganese dioxide (MnO2), the decomposition occurs much faster.

This is because MnO2 provides an alternative reaction pathway with a lower activation energy, making the reaction faster. Remarkably, MnO2 remains unchanged at the end of the reaction.

Real-world applications of catalysts

The use of catalysts is widespread and extremely important in many industrial and biological processes.

Industrial applications

  • Petrochemical industry: Catalysts are used in petroleum refining to break down large hydrocarbon molecules into gasoline and other products.
  • Ammonia production: Iron catalysts are used to lower the activation energy in the Haber-Bosch process for the synthesis of ammonia (NH3) from nitrogen and hydrogen gases.

Biological catalysts

In living organisms, enzymes function as biological catalysts. Enzymes are complex proteins that catalyze biochemical reactions essential for life, including DNA polymerase in DNA replication and amylase in digestion.

Enzyme specificity is like a lock and key, where only specific substrates fit into the enzyme's active site, allowing for specific reactions with low activation energy.

Conclusion

Activation energy and catalysis play important roles in the study of chemical kinetics. Understanding these concepts enables chemists to manipulate reaction rates and develop new processes and materials. The principles of activation energy and catalysis have far-reaching implications, from industrial manufacturing to physiological functions.


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Activation Energy and Catalyst

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