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


Kinetics, in the context of chemistry, deals with the rate of chemical reactions—how quickly or slowly a reaction occurs. This understanding is important for a wide range of scientific and industrial applications, from pharmaceuticals to materials science and environmental chemistry. By studying kinetics, chemists can determine which factors influence reaction rates and develop models to predict the progress of chemical reactions.

Basic concepts

The study of kinetics involves some key concepts, including reaction rates, rate laws, and the mechanism of reactions. In addition, it considers the effect of various conditions on the speed of reactions.

Response rates

The rate of a chemical reaction is a measure of how quickly reactants convert into products over time. It is usually expressed as the concentration of a certain amount of reactant or product over a certain period of time. Mathematically, the rate can be defined as:

Rate = -Δ[reactants]/Δt = Δ[products]/Δt

Here, [Reactant] and [Product] represent the concentrations of reactants and products and Δ represents the change over time interval Δt. Note the negative sign for reactants, which shows that their concentrations decrease with time.

Rate law and order of reaction

Rate laws express the relationship between the reaction rate and the concentration of reactants. The general form of the rate law is:

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

In the above equation, k is the rate constant, which is specific to a given reaction at a particular temperature. The exponents m and n are known as the order of the reaction with respect to reactants A and B, respectively. The sum m + n represents the overall order of the reaction.

A B

Determination of the rate constant and order of reaction

The values of the rate constant k and the orders m and n are usually determined through experimental data. To obtain these values, several experimental approaches can be used, such as the method of initial rates, which consists of measuring the initial rate of the reaction for different initial concentrations of reactants.

Example calculation

Consider a hypothetical reaction:

A + 2B → C

Suppose we conduct an experiment to measure the initial rate of a reaction at several different initial concentrations:

Experiment [A] (m) [B] (m) Initial rate (m/s)
     1 0.10 0.10 0.005
     2 0.20 0.10 0.010
     3 0.10 0.20 0.020

From this data, we can estimate the order of the reaction with respect to each reactant. For example, doubling the concentration of A doubles the rate (from Experiment 1 to 2), indicating that the reaction is first-order with respect to A. Similarly, doubling the concentration of B also doubles the rate (from Experiment 1 to 3), indicating that the reaction is also first-order with respect to B. Therefore, the rate law is:

Rate = k[A][B]

Reaction mechanism

A detailed description of the step-by-step process by which reactants are converted into products is known as a reaction mechanism. Each successive step in this process is called an elementary step.

Primary reactions

Elementary reactions occur in a single step and involve a small number of molecules, usually one or two. For example, in the reaction mechanism:

A + B → I (slow) I + B → C (fast)

Here, the first step is slow and involves the direct interaction of molecules A and B to form an intermediate I. The second step is fast, where intermediate I reacts with another molecule of B to form C.

Factors affecting the reaction rate

Several factors can affect the rate of a chemical reaction:

Concentration

Increasing the concentration of reactants generally increases the rate of the reaction. This happens because higher concentrations mean there are more reactant particles available to collide and react.

Temperature

As the temperature rises, the kinetic energy of the reactant molecules increases, leading to more frequent and energetic collisions. This generally leads to higher reaction rates.

Presence of catalyst

Catalysts are substances that increase the rate of a reaction without costing themselves. They work by providing an alternative pathway with a lower activation energy for the reaction to take place.

Surface area

For reactions involving solids, greater surface area allows for more collisions between reactant molecules. This can be achieved by using finely divided powders or by stirring the solution.

Temperature dependence and the Arrhenius equation

The effect of temperature on the rate constant k for most reactions can be expressed mathematically using the Arrhenius equation:

k = a * e^(-ea/rt)

In the above equation, A is the pre-exponential factor or frequency factor, e is the base of the natural logarithm, R is the gas constant, T is the temperature in Kelvin, and Ea is the activation energy of the reaction. By taking the natural logarithm of both sides, we can write:

ln(k) = ln(A) - Ea/Rt

This linear form allows us to determine the activation energy from a plot of ln(k) against 1/T, which gives a straight line with a slope of -Ea/R.

Conclusion

Understanding kinetics provides us with valuable information about how various factors affect the speed of chemical reactions, from simple laboratory procedures to large-scale industrial applications. By comprehensively studying reaction rates, mechanisms, and factors such as concentrations and temperatures, chemists can optimize conditions and make informed decisions in their work.


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Kinetics

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