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 → Chemical reactions ↓


Types of Chemical Reactions


Chemical reactions are processes where one or more substances are transformed into different substances. They are a fundamental concept in chemistry and are classified into different types based on various factors such as the nature of the reactants and products or the chemical change they undergo. Understanding these types helps us predict the products of a reaction and understand the principles that govern chemical changes. We will dive deeper into some of the most common types: combination reactions, decomposition reactions, displacement reactions, redox reactions, and more, each of which has its own unique characteristics and examples.

1. Combination reactions

Combination reactions, also called synthesis reactions, occur when two or more reactants combine to form a single product. These reactions are characterized by the general formula:

 A + B → AB

Here, A and B are the reactants, and AB is the product formed by their combination. A classic example of this is the synthesis of water from hydrogen and oxygen:

 2H 2 + O 2 → 2H 2 O

Visual example:

,

In the SVG example, the blue color represents hydrogen atoms and the red color represents oxygen. Together, they form a purple water molecule.

2. Decomposition reactions

Decomposition reactions are the opposite of combination reactions. Here, a single compound breaks down into two or more simpler substances. Its general form is:

 AB → A + B

A common decomposition reaction is the breakdown of water into hydrogen and oxygen gases:

 2H 2 O → 2H 2 + O 2

Visual example:

,

In the SVG example, the purple water molecule disintegrates into blue hydrogen and red oxygen atoms.

3. Displacement reactions

Displacement reactions occur when one element replaces another element in a compound, forming a new element and a new compound. It can be further classified into single and double displacement reactions. The general form of a single displacement reaction is:

 a + bc → ac + b

For example, when zinc reacts with copper sulfate:

 4Zn + CuSO4 → ZnSO4 + Cu

Visual example:

, ,

In the SVG example, the gray zinc atom displaces the blue copper in copper sulfate.

4. Double displacement reactions

In a double displacement reaction, there is an exchange of ions between two compounds, leading to the formation of two new compounds. Its general form is:

 AB + CD → AD + CB

A classic example of this is the reaction between silver nitrate and sodium chloride:

 AgNO 3 + NaCl → AgCl + NaNO 3

Visual example:

, ,

In this SVG example, blue and green rectangles swap partners with yellow and purple circles.

5. Redox reactions

Redox reactions involve the transfer of electrons between two substances. In these reactions, one species is oxidized by losing electrons and the other is reduced by gaining electrons. The redox process can be represented by the following half-reactions:

        Oxidation: A → A n+ + ne -
        Reduction: B + ne - → B n-

For example, in the reaction between hydrogen and fluorine:

 H 2 + F 2 → 2HF

Hydrogen is oxidized and fluorine is reduced. The visual representation for redox can be quite extensive, so we'll keep it conceptual here.

6. Neutralization reactions

Neutralization reactions are a subtype of double displacement reaction where an acid and a base react to form water and a salt. The general equation is:

 HX + MOH → MX + H 2 O

A common example of this is the reaction between hydrochloric acid and sodium hydroxide:

 HCl + NaOH → NaCl + H 2 O

These reactions are common in daily life and are important for a variety of biological and industrial processes.

7. Combustion reactions

Combustion reactions involve the burning of a substance, usually in the presence of oxygen, and the release of energy in the form of heat and light. The general equation for the complete combustion of a hydrocarbon can be written as:

 C x H y + O 2 → CO 2 + H 2 O

For example, the combustion of methane:

 CH 4 + 2O 2 → CO 2 + 2H 2 O

Combustion reactions are responsible for powering countless engines and are vital in the energy sector.

8. Precipitation reactions

Precipitation reactions are a form of double displacement reactions where two aqueous solutions react to form one or more insoluble products, called precipitates. The general equation is:

 AB(aq) + CD(aq) → AD(s) + CB(aq)

A simple example of this is the reaction between potassium iodide and lead(II) nitrate, in which lead(II) iodide forms as a precipitate:

 2KI(aq) + Pb(NO 3 ) 2 (aq) → 2KNO 3 (aq) + PbI 2 (s)

The solid precipitate is often observed as a cloud in solution or as a separate solid.

Conclusion

Chemical reactions are central to our understanding of chemistry and matter. Each type of reaction provides essential information about how substances interact, combine, and transform into new materials. Recognizing and predicting the products of these reactions is a vital skill in chemistry. This understanding enables chemists to develop new materials, create energy efficiently, and understand the complex reactions that occur in the world around us.


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Types of Chemical Reactions

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