28 Sept 2024

CHEMICAL BONDING

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Chemical Bonding: A Comprehensive Overview

I. Primary Bond Types

  1. Ionic Bonds

    • Formed between oppositely charged ions.
    • Examples:
      • NaCl (sodium chloride)
      • KCl (potassium chloride)

  2. Covalent Bonds

    • Formed by the sharing of electrons between atoms.
    • Examples:
      • H₂ (hydrogen)
      • O₂ (oxygen)

  3. Metallic Bonds

    • Formed by the delocalization of electrons among metal atoms.
    • Examples:
      • Cu (copper)
      • Fe (iron)

II. Secondary Bond Types

  1. Hydrogen Bonds

    • Weak electrostatic attraction between hydrogen and highly electronegative atoms (oxygen, nitrogen, fluorine).
    • Examples:
      • H₂O (water)
      • NH₃ (ammonia)

  2. Van der Waals Forces

    • Weak intermolecular forces.
    • Types:
      • Dipole-Dipole Forces: Between polar molecules.
        • Example: HCl (hydrogen chloride)
      • London Dispersion Forces: Present in all molecules, including nonpolar ones.
        • Examples: Noble gases (He, Ne, Ar, etc.)

III. Other Bond Types

  1. Polar Covalent Bonds

    • Unequal sharing of electrons due to differences in electronegativity.
    • Examples:
      • H₂O (water)
      • CO₂ (carbon dioxide)

  2. Coordinate Covalent Bonds

    • One atom donates both electrons to form a bond.
    • Examples:
      • NH₄⁺ (ammonium ion)
      • H₃O⁺ (hydronium ion)

  3. Pi Bonds

    • Formed by sideways overlap of p-orbitals, contributing to double or triple bonds.
    • Examples:
      • Ethene (C₂H₄)
      • Benzene (C₆H₆)

  4. Sigma Bonds

    • Formed by end-to-end overlap of orbitals, representing a single bond.
    • Examples:
      • Ethane (C₂H₆)
      • Methane (CH₄)

IV. Intermolecular Forces

  1. Electrostatic Forces

    • Attraction or repulsion between charged particles.

  2. Induced Dipole Forces

    • Temporary dipoles induced in nonpolar molecules.

  3. Repulsive Forces

    • Forces that prevent molecules from getting too close together.

V. Bonding in Specific Element Groups

  1. Metals

    • Typically form metallic bonds.

  2. Nonmetals

    • Typically form covalent bonds.

  3. Metalloids

    • Can form both covalent and metallic bonds.

  4. Noble Gases

    • Primarily held together by London dispersion forces.

VI. Bonding Trends in the Periodic Table

  1. Electronegativity

    • Increases from left to right and top to bottom.

  2. Ionization Energy

    • Increases from left to right and top to bottom.

  3. Atomic Radius

    • Decreases from left to right and top to bottom.

Chemical Bonding: A Comprehensive Overview

I. Primary Bond Types

  1. Ionic Bonds

    • Electrostatic attraction between oppositely charged ions.
    • Examples:
      1. NaCl (sodium chloride)
      2. CaCO₃ (calcium carbonate)
      3. MgO (magnesium oxide)
      4. Al₂O₃ (aluminum oxide)
      5. KNO₃ (potassium nitrate)
      6. FeCl₃ (iron(III) chloride)
      7. CuSO₄ (copper(II) sulfate)
      8. BaCO₃ (barium carbonate)
      9. NH₄Cl (ammonium chloride)
      10. Pb(NO₃)₂ (lead(II) nitrate)

  2. Covalent Bonds

    • Sharing of electrons between atoms.
    • Examples:
      1. H₂ (hydrogen gas)
      2. O₂ (oxygen gas)
      3. N₂ (nitrogen gas)
      4. CH₄ (methane)
      5. C₂H₄ (ethylene)
      6. C₆H₁₂ (cyclohexane)
      7. CO₂ (carbon dioxide)
      8. H₂O (water)
      9. NH₃ (ammonia)
      10. CCl₄ (carbon tetrachloride)

  3. Metallic Bonds

    • Delocalized electrons shared among metal atoms.
    • Examples:
      1. Cu (copper)
      2. Fe (iron)
      3. Au (gold)
      4. Ag (silver)
      5. Al (aluminum)
      6. Zn (zinc)
      7. Ni (nickel)
      8. Pb (lead)
      9. Sn (tin)
      10. Hg (mercury)

II. Secondary Bond Types

  1. Hydrogen Bonds

    • Weak electrostatic attraction between hydrogen and highly electronegative atoms.
    • Examples:
      1. H₂O (water)
      2. NH₃ (ammonia)
      3. CH₃OH (methanol)
      4. C₆H₁₂O₆ (glucose)
      5. DNA double helix structure
      6. Proteins (e.g., collagen)
      7. Cellulose
      8. Glycogen
      9. Starch
      10. Hydrochloric acid (HCl)

  2. Van der Waals Bonds

    • Weak intermolecular forces.
    • Examples:
      1. Ar (argon)
      2. CH₄ (methane)
      3. CCl₄ (carbon tetrachloride)
      4. Xe (xenon)
      5. Graphite
      6. Diamond
      7. Fullerenes
      8. Hydrocarbons (e.g., hexane, octane)
      9. Fluorinated compounds (e.g., Teflon)
      10. Silicone oils

  3. Dipole-Dipole Forces

    • Electrostatic attraction between polar molecules.
    • Examples:
      1. CO₂ (carbon dioxide)
      2. HCl (hydrochloric acid)
      3. CH₃Cl (methyl chloride)
      4. C₆H₅Cl (chlorobenzene)
      5. H₂O (water)
      6. NH₃ (ammonia)
      7. CH₃OH (methanol)
      8. C₆H₁₂O₆ (glucose)
      9. Acetone
      10. Dimethyl ether

  4. London Dispersion Forces

    • Weak intermolecular forces.
    • Examples:
      1. Noble gases (e.g., He, Ne, Ar)
      2. Nonpolar molecules (e.g., CH₄, CCl₄)
      3. Hydrocarbons (e.g., hexane, octane)
      4. Fluorinated compounds (e.g., Teflon)
      5. Silicone oils
      6. Methane
      7. Ethane
      8. Propane
      9. Butane
      10. Pentane

III. Other Bond Types

  1. Polar Covalent Bonds

    • Unequal sharing of electrons due to differences in electronegativity.
    • Examples:
      1. H₂O (water)
      2. CO₂ (carbon dioxide)
      3. CH₃OH (methanol)
      4. NH₃ (ammonia)
      5. C₆H₅OH (phenol)
      6. HCl (hydrochloric acid)
      7. CH₃Cl (methyl chloride)
      8. C₆H₅Cl (chlorobenzene)
      9. Acetone
      10. Dimethyl ether

  2. Coordinate Covalent Bonds

    • Donation of electrons from one atom to another.
    • Examples:
      1. NH₃ (ammonia)
      2. H₂O (water)
      3. CO (carbon monoxide)
      4. CN⁻ (cyanide ion)
      5. SO₄²⁻ (sulfate ion)
      6. NO₂⁺ (nitronium ion)
      7. ClO⁻ (hypochlorite ion)
      8. O₃ (ozone)
      9. N₂O (nitrous oxide)
      10. XeF₂ (xenon difluoride)

  3. Pi Bonds

    • Sideways overlap of p-orbitals.
    • Examples:
      1. C₂H₄ (ethene)
      2. C₆H₆ (benzene)
      3. C₄H₆ (butadiene)
      4. C₂H₂ (acetylene)
      5. NO (nitric oxide)
      6. C₂H₂ (ethyne)
      7. C₅H₆ (cyclopentadiene)
      8. C₁₄H₁₀ (anthracene)
      9. C₁₀H₈ (naphthalene)
      10. C₁₀H₈ (azulene)

  4. Sigma Bonds

    • End-to-end overlap of orbitals.
    • Examples:
      1. C₂H₆ (ethane)
      2. CH₄ (methane)
      3. NH₃ (ammonia)
      4. H₂O (water)
      5. HF (hydrogen fluoride)
      6. CHCl₃ (chloroform)
      7. CH₃OH (methanol)
      8. C₂H₅NH₂ (ethylamine)
      9. (CH₃)₂NH (dimethylamine)
      10. C₃H₈ (propane)

IV. Intermolecular Forces

  1. Electrostatic Forces

    • Attractions between charged particles.
    • Examples:
      1. NaCl (sodium chloride)
      2. CaCO₃ (calcium carbonate)
      3. MgO (magnesium oxide)
      4. Al₂O₃ (aluminum oxide)
      5. KNO₃ (potassium nitrate)
      6. FeCl₃ (iron(III) chloride)
      7. CuSO₄ (copper(II) sulfate)
      8. BaCO₃ (barium carbonate)
      9. NH₄Cl (ammonium chloride)
      10. Pb(NO₃)₂ (lead(II) nitrate)

  2. Induced Dipole Forces

    • Temporary dipoles induced by nearby molecules.
    • Examples:
      1. Noble gases (e.g., He, Ne, Ar)
      2. Nonpolar molecules (e.g., CH₄, CCl₄)
      3. Hydrocarbons (e.g., hexane, octane)
      4. Fluorinated compounds (e.g., Teflon)
      5. Silicone oils
      6. Methane
      7. Ethane
      8. Propane
      9. Butane
      10. Pentane

  3. Repulsive Forces

    • Forces that prevent molecules from getting too close.
    • Examples:
      1. All molecules, especially at short distances
      2. Ion-ion repulsion
      3. Electron-electron repulsion
      4. Nuclear-nuclear repulsion
      5. Pauli repulsion
      6. Steric repulsion
      7. Molecular orbital repulsion
      8. Electrostatic repulsion
      9. Exchange repulsion
      10. Consequences of Repulsive Forces


Factors Influencing Intermolecular Forces

I. Factors Influencing Intermolecular Forces

  1. Electronegativity

    • Definition: The ability of an atom to attract electrons towards itself.
    • Impact: Influences the strength of dipole-dipole interactions and hydrogen bonding.

  2. Polarizability

    • Definition: The ability of an atom or molecule to distort its electron cloud.
    • Impact: Affects the strength of London dispersion forces; larger, more polarizable molecules have stronger interactions.

  3. Size and Shape

    • Definition: The overall dimensions and configuration of a molecule.
    • Impact: Larger molecules have more electrons and a larger electron cloud, which can enhance London dispersion forces.

  4. Charge

    • Definition: The presence of positive or negative charges in ions or charged molecules.
    • Impact: Ions and charged molecules interact more strongly due to electrostatic attractions.

  5. Dipole Moment

    • Definition: A measure of the polarity of a molecule; polar molecules have a permanent electric dipole.
    • Impact: Determines the strength of dipole-dipole interactions and hydrogen bonding.

  6. Temperature

    • Definition: The measure of the average kinetic energy of particles.
    • Impact: Higher temperatures increase kinetic energy, reducing the effectiveness of intermolecular forces.

  7. Pressure

    • Definition: The force exerted by molecules in a given area.
    • Impact: Higher pressures increase molecular proximity, enhancing intermolecular forces.

  8. Molecular Weight

    • Definition: The mass of a molecule, often correlated with the number of atoms present.
    • Impact: Heavier molecules tend to have stronger intermolecular forces due to increased polarizability and electron cloud size.

II. Types of Intermolecular Forces Affected by These Factors

  1. London Dispersion Forces

    • Influenced by: Polarizability, size, and molecular weight.
    • Characteristics: Present in all molecules, stronger in larger or more polarizable ones.

  2. Dipole-Dipole Forces

    • Influenced by: Electronegativity, dipole moment, and temperature.
    • Characteristics: Stronger in polar molecules with significant dipole moments.

  3. Hydrogen Bonding

    • Influenced by: Electronegativity, dipole moment, and molecular shape.
    • Characteristics: Strong type of dipole-dipole interaction involving hydrogen bonded to electronegative atoms (e.g., N, O, F).

III. Consequences of Intermolecular Forces

  1. Physical Properties

    • Examples: Boiling point, melting point, viscosity, surface tension.
    • Impact: Stronger intermolecular forces generally lead to higher boiling/melting points and greater viscosity.

  2. Chemical Properties

    • Examples: Reactivity, solubility, phase behavior.
    • Impact: Intermolecular forces can determine how substances interact and dissolve in different solvents.

  3. Biological Processes

    • Examples: Protein folding, membrane structure, cell signaling.
    • Impact: The specific interactions between biomolecules are critical for their functions and stability.

IV. Real-World Applications

  1. Pharmaceuticals

    • Application: Designing drugs with optimal intermolecular forces to enhance efficacy and stability.

  2. Materials Science

    • Application: Creating materials with specific properties (e.g., strength, flexibility) tailored for particular uses.

  3. Energy Storage

    • Application: Optimizing battery performance by understanding and manipulating intermolecular forces for better energy retention and discharge.

Attractive Forces

I. Definition

Attractive forces are interactions between molecules or atoms that hold them together, shaping the physical and chemical properties of substances.

II. Types of Attractive Forces

  1. Ionic Bonds

    • Definition: Electrostatic attraction between oppositely charged ions.

  2. Covalent Bonds

    • Definition: Sharing of electrons between atoms.

  3. Hydrogen Bonds

    • Definition: Weak electrostatic attraction between hydrogen and electronegative atoms.

  4. Van der Waals Forces

    • Definition: Weak intermolecular forces, including:
      • London Dispersion Forces
      • Dipole-Dipole Forces

  5. Metallic Bonds

    • Definition: Delocalized electrons shared among metal atoms.

III. Characteristics of Attractive Forces

  1. Strength

    • Range: Varies from weak (van der Waals) to strong (covalent, ionic).

  2. Range

    • Short-range: (covalent, ionic)
    • Long-range: (van der Waals)

  3. Directionality

    • Directional: (covalent, hydrogen)
    • Non-directional: (ionic)

IV. Factors Influencing Attractive Forces

  1. Electronegativity

    • Impact: Affects ionic and covalent bond strength.

  2. Atomic Radius

    • Impact: Affects van der Waals forces.

  3. Molecular Shape

    • Impact: Influences hydrogen bonding and van der Waals interactions.

  4. Temperature

    • Impact: Affects molecular motion and intermolecular forces.

V. Consequences of Attractive Forces

  1. Physical Properties

    • Examples: Melting point, boiling point, viscosity.

  2. Chemical Properties

    • Examples: Reactivity, solubility, phase behavior.

  3. Biological Processes

    • Examples: Protein folding, membrane structure, cell signaling.

VI. Real-World Applications

  1. Materials Science

    • Application: Designing materials with specific properties.

  2. Pharmaceuticals

    • Application: Developing drugs with optimal binding affinity.

  3. Energy Storage

    • Application: Optimizing battery performance.

VII. Examples of Attractive Forces in Action

  1. Water's High Boiling Point

    • Due to hydrogen bonding.

  2. DNA's Double Helix Structure

    • Stabilized by hydrogen bonding.

  3. Metals' High Melting Points

    • Due to metallic bonding.

  4. Proteins' Complex Structures

    • Shaped by hydrogen bonding and van der Waals forces.

Repulsive Forces in Chemistry

I. Types of Repulsive Forces

  1. Ion-Ion Repulsion

    • Description: Occurs between two positively charged ions.
    • Cause: Electrostatic repulsion between like charges.
    • Characteristics:
      • Increases with decreasing distance between ions.
      • Important in ionic compounds (e.g., NaCl, CaCO₃).

  2. Electron-Electron Repulsion

    • Description: Occurs between two electrons in an atom or molecule.
    • Cause: Electrostatic repulsion between like charges.
    • Characteristics:
      • Increases with decreasing distance between electrons.
      • Important in atomic physics and chemistry (e.g., electron configuration).

  3. Nuclear-Nuclear Repulsion

    • Description: Occurs between two positively charged nuclei.
    • Cause: Electrostatic repulsion between like charges.
    • Characteristics:
      • Increases with decreasing distance between nuclei.
      • Important in nuclear physics (e.g., nuclear reactions).

  4. Pauli Repulsion

    • Description: Occurs when two electrons with the same spin occupy the same orbital.
    • Cause: Pauli exclusion principle.
    • Characteristics:
      • Increases with decreasing distance between electrons.
      • Important in atomic physics and chemistry (e.g., electron configuration).

  5. Steric Repulsion

    • Description: Occurs when two molecules or groups of atoms are too close.
    • Cause: Repulsion between electron clouds.
    • Characteristics:
      • Increases with decreasing distance between molecules or groups.
      • Important in organic chemistry (e.g., molecular shape, reactivity).

  6. Molecular Orbital Repulsion

    • Description: Occurs when two molecular orbitals overlap.
    • Cause: Repulsion between electrons in overlapping orbitals.
    • Characteristics:
      • Increases with decreasing distance between molecules.
      • Important in molecular physics and chemistry (e.g., chemical bonding).

  7. Electrostatic Repulsion

    • Description: Occurs between two charged particles (ions, electrons, or molecules).
    • Cause: Electrostatic repulsion between like charges.
    • Characteristics:
      • Increases with decreasing distance between charged particles.
      • Important in physics and chemistry (e.g., ionic compounds, electrochemistry).

  8. Exchange Repulsion

    • Description: Occurs when two electrons exchange places in an atom or molecule.
    • Cause: Repulsion between electrons in different orbitals.
    • Characteristics:
      • Increases with decreasing distance between electrons.
      • Important in atomic physics and chemistry (e.g., electron configuration).

  9. Correlation Repulsion

    • Description: Occurs when two electrons interact with each other.
    • Cause: Repulsion between electrons in different orbitals.
    • Characteristics:
      • Increases with decreasing distance between electrons.
      • Important in atomic physics and chemistry (e.g., electron correlation).

II. Factors Influencing Repulsive Forces

  1. Distance between particles or molecules.
  2. Charge and Electronegativity of particles or molecules.
  3. Size and Shape of particles or molecules.
  4. Spin and Orbital Orientation of electrons.
  5. Temperature and Pressure conditions.

III. Consequences of Repulsive Forces

  1. Molecular Shape and Structure: Determine the geometric arrangement of atoms in a molecule.
  2. Chemical Reactivity and Bonding: Influence how and why molecules interact with each other.
  3. Physical Properties: Affect characteristics such as boiling point and viscosity.
  4. Biological Processes: Play a crucial role in processes like protein folding and membrane structure.

Types of Attractive Forces

1. Ionic Bonds

  • Definition: Electrostatic attraction between oppositely charged ions.
  • Examples:
    1. NaCl (sodium chloride)
    2. CaCO3 (calcium carbonate)
    3. MgO (magnesium oxide)
    4. Al2O3 (aluminum oxide)
    5. KNO3 (potassium nitrate)
    6. FeCl3 (iron(III) chloride)
    7. CuSO4 (copper(II) sulfate)
    8. BaCO3 (barium carbonate)
    9. NH4Cl (ammonium chloride)
    10. Pb(NO3)2 (lead(II) nitrate)

2. Covalent Bonds

  • Definition: Sharing of electrons between atoms.
  • Examples:
    1. H2 (hydrogen gas)
    2. O2 (oxygen gas)
    3. N2 (nitrogen gas)
    4. CH4 (methane)
    5. C2H4 (ethylene)
    6. C6H12 (cyclohexane)
    7. CO2 (carbon dioxide)
    8. H2O (water)
    9. NH3 (ammonia)
    10. CCl4 (carbon tetrachloride)

3. Hydrogen Bonds

  • Definition: Weak electrostatic attraction between hydrogen and electronegative atoms.
  • Examples:
    1. H2O (water)
    2. NH3 (ammonia)
    3. CH3OH (methanol)
    4. C6H12O6 (glucose)
    5. DNA double helix structure
    6. Proteins (e.g., collagen)
    7. Cellulose
    8. Glycogen
    9. Starch
    10. Hydrochloric acid (HCl)

4. Van der Waals Forces

  • Definition: Weak intermolecular forces.
  • Examples:
    1. Ar (argon)
    2. CH4 (methane)
    3. CCl4 (carbon tetrachloride)
    4. Xe (xenon)
    5. Graphite
    6. Diamond
    7. Fullerenes
    8. Hydrocarbons (e.g., hexane, octane)
    9. Fluorinated compounds (e.g., Teflon)
    10. Silicone oils

5. Metallic Bonds

  • Definition: Delocalized electrons shared among metal atoms.
  • Examples:
    1. Cu (copper)
    2. Fe (iron)
    3. Au (gold)
    4. Ag (silver)
    5. Al (aluminum)
    6. Zn (zinc)
    7. Ni (nickel)
    8. Pb (lead)
    9. Sn (tin)
    10. Hg (mercury)

6. Dipole-Dipole Forces

  • Definition: Electrostatic attraction between polar molecules.
  • Examples:
    1. CO2 (carbon dioxide)
    2. HCl (hydrochloric acid)
    3. CH3Cl (methyl chloride)
    4. C6H5Cl (chlorobenzene)
    5. H2O (water)
    6. NH3 (ammonia)
    7. CH3OH (methanol)
    8. C6H12O6 (glucose)
    9. Acetone
    10. Dimethyl ether

7. Pi Bonds

  • Definition: Sideways overlap of p-orbitals.
  • Examples:
    1. Ethene (C2H4)
    2. Benzene (C6H6)
    3. Butadiene (C4H6)
    4. Acetylene (C2H2)
    5. Nitric oxide (NO)
    6. Ethyne (C2H2)
    7. Cyclopentadiene (C5H6)
    8. Anthracene (C14H10)
    9. Naphthalene (C10H8)
    10. Azulene (C10H8)

8. Sigma Bonds

  • Definition: End-to-end overlap of orbitals.
  • Examples:
    1. Ethane (C2H6)
    2. Methane (CH4)
    3. Ammonia (NH3)
    4. Water (H2O)
    5. Hydrogen fluoride (HF)
    6. Chloroform (CHCl3)
    7. Methanol (CH3OH)
    8. Ethylamine (C2H5NH2)
    9. Dimethylamine ((CH3)2NH)
    10. Propane (C3H8)

Repulsive Forces

I. Definition

Repulsive forces are interactions that prevent molecules or atoms from getting too close to one another, influencing the structure and stability of substances.

II. Types of Repulsive Forces and Examples

1. Ion-Ion Repulsion

Definition: Occurs between two positively charged ions.

  • Examples:
    1. Na⁺ and Ca²⁺ ions in aqueous solution
    2. Al³⁺ and Fe³⁺ ions in ionic compounds
    3. K⁺ and Na⁺ ions in potassium sodium tartrate
    4. Mg²⁺ and Ca²⁺ ions in magnesium calcium carbonate
    5. Cu²⁺ and Zn²⁺ ions in copper zinc alloys
    6. Li⁺ and Na⁺ ions in lithium sodium batteries
    7. Fe³⁺ and Cr³⁺ ions in stainless steel
    8. Ni²⁺ and Co²⁺ ions in nickel cobalt alloys
    9. Pb²⁺ and Sn²⁺ ions in lead tin solders
    10. Sr²⁺ and Ba²⁺ ions in strontium barium titanate

2. Electron-Electron Repulsion

Definition: Occurs between two electrons in an atom or molecule.

  • Examples:
    1. Electrons in atomic orbitals (e.g., 1s, 2s, 2p)
    2. Electrons in molecular orbitals (e.g., σ, π, δ)
    3. Electron pairs in covalent bonds (e.g., H₂, O₂)
    4. Lone pair electrons in molecules (e.g., NH₃, H₂O)
    5. Electrons in conjugated systems (e.g., benzene, butadiene)
    6. Electrons in aromatic compounds (e.g., toluene, naphthalene)
    7. Electrons in anti-bonding molecular orbitals
    8. Electrons in transition metal complexes
    9. Electrons in lanthanide and actinide compounds
    10. Electrons in radicals and carbanions

3. Nuclear-Nuclear Repulsion

Definition: Occurs between two positively charged nuclei.

  • Examples:
    1. Nuclear reactions (e.g., fusion, fission)
    2. Radioactive decay (e.g., alpha, beta, gamma)
    3. Nuclear scattering experiments
    4. Particle accelerators
    5. Cosmic ray interactions
    6. Stellar nucleosynthesis
    7. Big Bang nucleosynthesis
    8. Nuclear reactions in stars
    9. Nuclear reactions in supernovae
    10. Nuclear reactions in black holes

4. Pauli Repulsion

Definition: Occurs when two electrons with the same spin occupy the same orbital.

  • Examples:
    1. Electron configuration of atoms (e.g., Aufbau principle)
    2. Molecular orbital theory
    3. Hund's rule of maximum multiplicity
    4. Pauli exclusion principle
    5. Exchange interactions in molecules
    6. Electron correlation effects
    7. Quantum chemistry calculations
    8. Electron spin resonance spectroscopy
    9. Nuclear magnetic resonance spectroscopy
    10. Electron paramagnetic resonance spectroscopy

5. Steric Repulsion

Definition: Occurs when two molecules or groups of atoms are too close.

  • Examples:
    1. Molecular shape and structure
    2. Conformational analysis
    3. Molecular mechanics simulations
    4. Steric hindrance in chemical reactions
    5. Enzyme-substrate interactions
    6. Protein-ligand binding
    7. Membrane transport
    8. Molecular recognition
    9. Supramolecular chemistry
    10. Crystal packing

6. Molecular Orbital Repulsion

Definition: Occurs when two molecular orbitals overlap.

  • Examples:
    1. Molecular orbital theory
    2. Electronic spectroscopy
    3. Photochemistry
    4. Electron transfer reactions
    5. Conducting polymers
    6. Organic electronics
    7. Quantum chemistry calculations
    8. Molecular simulations
    9. Chemical reactivity
    10. Catalysis

7. Electrostatic Repulsion

Definition: Occurs between two charged particles.

  • Examples:
    1. Ionic compounds (e.g., NaCl, CaCO₃)
    2. Electrochemistry
    3. Colloidal suspensions
    4. Electrophoresis
    5. Electrostatic precipitation
    6. Particle interactions
    7. Surface chemistry
    8. Biophysical chemistry
    9. Nanoparticle interactions
    10. Biological membranes

8. Exchange Repulsion

Definition: Occurs when two electrons exchange places in an atom or molecule.

  • Examples:
    1. Electron exchange reactions
    2. Exchange interactions in molecules
    3. Electron correlation effects
    4. Quantum chemistry calculations
    5. Molecular orbital theory
    6. Valence bond theory
    7. Electron spin resonance spectroscopy
    8. Nuclear magnetic resonance spectroscopy
    9. Electron paramagnetic resonance spectroscopy
    10. Chemical reactivity

9. Correlation Repulsion

Definition: Occurs when two electrons interact with each other.

  • Examples:
    1. Electron correlation effects
    2. Quantum chemistry calculations
    3. Molecular orbital theory
    4. Valence bond theory
    5. Electron spin resonance spectroscopy
    6. Nuclear magnetic resonance spectroscopy
    7. Electron paramagnetic resonance spectroscopy
    8. Chemical reactivity
    9. Catalysis
    10. Materials science
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