28 Sept 2024

CHEMICAL BONDING

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

INORGANIC ACIDS

Here’s a well-organized list of commonly used inorganic acids with detailed categories and examples:

1. Halogen Acids:

  1. Hydrofluoric Acid - HF
  2. Hydrochloric Acid - HCl
  3. Hydrobromic Acid - HBr
  4. Hydroiodic Acid - HI
  5. Hydroastatic Acid - HAt

2. Oxyacids:

  1. Nitric Acid - HNO3
  2. Sulfuric Acid - H2SO4
  3. Phosphoric Acid - H3PO4
  4. Chromic Acid - H2CrO4
  5. Perchloric Acid - HClO4

3. Acidic Oxides:

  1. Carbonic Acid - H2CO3
  2. Silicic Acid - H2SiO3
  3. Boric Acid - H3BO3
  4. Germanic Acid - H2GeO3
  5. Telluric Acid - H6TeO6

4. Sulfur Containing Acids:

  1. Sulfurous Acid - H2SO3
  2. Thiosulfuric Acid - H2S2O3
  3. Dithionic Acid - H2S2O6
  4. Polythionic Acid - H2SxO6
  5. Sulfamic Acid - H3NSO3

5. Nitrogen Containing Acids:

  1. Hydrazoic Acid - HN3
  2. Cyanic Acid - HOCN
  3. Fulminic Acid - HCNO
  4. Isocyanic Acid - HNCO
  5. Cyanuric Acid - C3H3N3O3

6. Phosphorus Containing Acids:

  1. Hypophosphoric Acid - H4P2O6
  2. Pyrophosphoric Acid - H4P2O7
  3. Tripolyphosphoric Acid - H5P3O10
  4. Tetrapolyphosphoric Acid - H6P4O13
  5. Metaphosphoric Acid - HPO3

7. Chlorine Containing Acids:

  1. Hypochlorous Acid - HOCl
  2. Chlorous Acid - HClO2
  3. Chloric Acid - HClO3
  4. Perchloric Acid - HClO4
  5. Dichloroacetic Acid - C2H2Cl2O2

8. Bromine Containing Acids:

  1. Hypobromous Acid - HOBr
  2. Bromous Acid - HBrO2
  3. Bromic Acid - HBrO3
  4. Perbromic Acid - HBrO4
  5. Tribromacetic Acid - C2HBr3O2

9. Iodine Containing Acids:

  1. Hypoiodous Acid - HOI
  2. Iodous Acid - HIO2
  3. Iodic Acid - HIO3
  4. Periodic Acid - HIO4
  5. Diiodoacetic Acid - C2H2I2O2

10. Other Inorganic Acids:

  1. Arsenic Acid - H3AsO4
  2. Antimonic Acid - H3SbO4
  3. Bismuthic Acid - HBiO3
  4. Cadmium Acid - H2CdO3
  5. Stannic Acid - H2SnO3

11. Polyacids:

  1. Pyrophosphoric Acid - H4P2O7
  2. Tripolyphosphoric Acid - H5P3O10
  3. Tetrapolyphosphoric Acid - H6P4O13
  4. Silicotungstic Acid - H4SiW12O40
  5. Phosphotungstic Acid - H3PW12O40

12. Peroxo Acids:

  1. Peroxomonosulfuric Acid - H2SO5
  2. Peroxodisulfuric Acid - H2S2O8
  3. Peroxomonophosphoric Acid - H3PO5
  4. Peroxodiphosphoric Acid - H4P2O8

13. Oxohalogen Acids:

  1. Hypochlorous Acid - HOCl
  2. Hypobromous Acid - HOBr
  3. Hypoiodous Acid - HOI
  4. Chlorous Acid - HClO2
  5. Bromous Acid - HBrO2

========================================================================

1. Halogen Acids:

  • Uses: Disinfectants, metal cleaning, etching agents, catalysis.
  • Properties: Highly corrosive, strong acids, react with metals, colorless liquids/gases.
  • Examples:
    • Hydrochloric Acid (HCl): Used in steel pickling and pH regulation.
    • Hydrofluoric Acid (HF): Glass etching and metal cleaning.
    • Hydrobromic Acid (HBr): Organic bromination.
    • Hydroiodic Acid (HI): Reducing agent in organic synthesis.

2. Oxyacids:

  • Uses: Fertilizers, oxidizing agents, explosives, detergents.
  • Properties: Strong acids, highly reactive, form when nonmetal oxides dissolve in water.
  • Examples:
    • Nitric Acid (HNO3): Used in fertilizers, explosives.
    • Sulfuric Acid (H2SO4): Industrial chemical, battery acid.
    • Phosphoric Acid (H3PO4): Rust removal, fertilizers.
    • Chromic Acid (H2CrO4): Electroplating, cleaning glassware.

3. Acidic Oxides:

  • Uses: Drying agents, adsorbents, neutralizing agents.
  • Properties: Weak acids, react with water to form acids, nonmetal oxides.
  • Examples:
    • Carbonic Acid (H2CO3): Present in carbonated drinks.
    • Silicic Acid (H2SiO3): Used in desiccants, ceramics.
    • Boric Acid (H3BO3): Antiseptic, insecticide.

4. Sulfur Containing Acids:

  • Uses: Bleaching, reducing agents, textile industry, disinfectants.
  • Properties: Strong reducing agents, highly reactive, sulfur-oxygen combinations.
  • Examples:
    • Sulfurous Acid (H2SO3): Bleaching agent in the paper industry.
    • Thiosulfuric Acid (H2S2O3): Gold extraction.

5. Nitrogen Containing Acids:

  • Uses: Fertilizers, explosives, pharmaceuticals, dyes.
  • Properties: Highly reactive, involve nitrogen and oxygen.
  • Examples:
    • Hydrazoic Acid (HN3): Explosives, gas generation.
    • Cyanic Acid (HOCN): Used in organic synthesis.

6. Phosphorus Containing Acids:

  • Uses: Fertilizers, detergents, water treatment.
  • Properties: Polyprotic acids, phosphorus and oxygen compounds.
  • Examples:
    • Pyrophosphoric Acid (H4P2O7): Used in detergents.
    • Tripolyphosphoric Acid (H5P3O10): Water softener.

7. Chlorine Containing Acids:

  • Uses: Disinfectants, water treatment, bleaching.
  • Properties: Strong acids, reactive with metals.
  • Examples:
    • Hypochlorous Acid (HOCl): Disinfection in pools.
    • Perchloric Acid (HClO4): Strong oxidizer in rocket fuel.

8. Bromine Containing Acids:

  • Uses: Disinfection, organic synthesis.
  • Properties: Strong acids, reactive with metals.
  • Examples:
    • Bromic Acid (HBrO3): Oxidizer in organic synthesis.
    • Hypobromous Acid (HOBr): Used in water treatment.

9. Iodine Containing Acids:

  • Uses: Disinfectants, organic synthesis.
  • Properties: Strong acids, iodine-based.
  • Examples:
    • Iodic Acid (HIO3): Oxidizer in organic synthesis.
    • Periodic Acid (HIO4): Cleaving agents in biochemistry.

10. Other Inorganic Acids:

  • Uses: Industrial applications, metal processing.
  • Properties: Varies depending on the metal involved.
  • Examples:
    • Arsenic Acid (H3AsO4): Wood preservatives.
    • Stannic Acid (H2SnO3): Used in ceramics 

11. Polyacids:

  • Uses: Water softeners, detergents.
  • Properties: Complex acids with multiple acid groups.
  • Examples:
    • Silicotungstic Acid (H4SiW12O40): Catalysts in organic synthesis.
    • Phosphotungstic Acid (H3PW12O40): Staining agents in biology.

12. Peroxo Acids:

  • Uses: Bleaching, disinfectants.
  • Properties: Strong oxidizers, contain peroxide groups.
  • Examples:
    • Peroxomonosulfuric Acid (H2SO5): Bleaching agent.
    • Peroxodiphosphoric Acid (H4P2O8): Strong oxidizer.

13. Oxohalogen Acids:

  • Uses: Bleaching, disinfectants, oxidizers.
  • Properties: Highly reactive halogen-oxygen acids.
  • Examples:
    • Chlorous Acid (HClO2): Bleaching agent.
    • Bromous Acid (HBrO2): Oxidizer in organic chemistry.

==============================================================

1. Halogen Acids:

  1. Hydrofluoric Acid - HF
  2. Hydrochloric Acid - HCl
  3. Hydrobromic Acid - HBr
  4. Hydroiodic Acid - HI
  5. Hydroastatic Acid - HAt
  6. Perchloric Acid - HClO4
  7. Chloric Acid - HClO3

2. Oxyacids:

  1. Nitric Acid - HNO3
  2. Sulfuric Acid - H2SO4
  3. Phosphoric Acid - H3PO4
  4. Chromic Acid - H2CrO4
  5. Perchloric Acid - HClO4
  6. Hypophosphorous Acid - H3PO2
  7. Selenic Acid - H2SeO4

3. Acidic Oxides:

  1. Carbonic Acid - H2CO3
  2. Silicic Acid - H2SiO3
  3. Boric Acid - H3BO3
  4. Germanic Acid - H2GeO3
  5. Telluric Acid - H6TeO6
  6. Stannic Acid - H2SnO3
  7. Oxalic Acid - H2C2O4

4. Sulfur Containing Acids:

  1. Sulfurous Acid - H2SO3
  2. Thiosulfuric Acid - H2S2O3
  3. Dithionic Acid - H2S2O6
  4. Polythionic Acid - H2SxO6
  5. Sulfamic Acid - H3NSO3
  6. Peroxomonosulfuric Acid - H2SO5
  7. Peroxodisulfuric Acid - H2S2O8

5. Nitrogen Containing Acids:

  1. Hydrazoic Acid - HN3
  2. Cyanic Acid - HOCN
  3. Fulminic Acid - HCNO
  4. Isocyanic Acid - HNCO
  5. Cyanuric Acid - C3H3N3O3
  6. Nitrous Acid - HNO2
  7. Dinitramic Acid - HN(NO2)2

6. Phosphorus Containing Acids:

  1. Hypophosphoric Acid - H4P2O6
  2. Pyrophosphoric Acid - H4P2O7
  3. Tripolyphosphoric Acid - H5P3O10
  4. Tetrapolyphosphoric Acid - H6P4O13
  5. Metaphosphoric Acid - HPO3
  6. Phosphorous Acid - H3PO3
  7. Orthophosphoric Acid - H3PO4

7. Chlorine Containing Acids:

  1. Hypochlorous Acid - HOCl
  2. Chlorous Acid - HClO2
  3. Chloric Acid - HClO3
  4. Perchloric Acid - HClO4
  5. Dichloroacetic Acid - C2H2Cl2O2
  6. Trichloroacetic Acid - CCl3COOH
  7. Chlorosulfuric Acid - HSO3Cl

8. Bromine Containing Acids:

  1. Hypobromous Acid - HOBr
  2. Bromous Acid - HBrO2
  3. Bromic Acid - HBrO3
  4. Perbromic Acid - HBrO4
  5. Tribromacetic Acid - C2HBr3O2
  6. Bromic Chloride - HBrClO4
  7. Bromosulfuric Acid - HBrSO3

9. Iodine Containing Acids:

  1. Hypoiodous Acid - HOI
  2. Iodous Acid - HIO2
  3. Iodic Acid - HIO3
  4. Periodic Acid - HIO4
  5. Diiodoacetic Acid - C2H2I2O2
  6. Periodic Acid (Ortho) - H5IO6
  7. Iodophosphoric Acid - H4IO4P

10. Other Inorganic Acids:

  1. Arsenic Acid - H3AsO4
  2. Antimonic Acid - H3SbO4
  3. Bismuthic Acid - HBiO3
  4. Cadmium Acid - H2CdO3
  5. Stannic Acid - H2SnO3
  6. Molybdic Acid - H2MoO4
  7. Vanadic Acid - H3VO4

11. Polyacids:

  1. Pyrophosphoric Acid - H4P2O7
  2. Tripolyphosphoric Acid - H5P3O10
  3. Tetrapolyphosphoric Acid - H6P4O13
  4. Silicotungstic Acid - H4SiW12O40
  5. Phosphotungstic Acid - H3PW12O40
  6. Polyphosphoric Acid - H6P4O13
  7. Polyboric Acid - H[B(OH)4]

12. Peroxo Acids:

  1. Peroxomonosulfuric Acid - H2SO5
  2. Peroxodisulfuric Acid - H2S2O8
  3. Peroxomonophosphoric Acid - H3PO5
  4. Peroxodiphosphoric Acid - H4P2O8
  5. Peroxonitric Acid - HNO4
  6. Peroxochloric Acid - HClO5
  7. Peroxoboric Acid - H2BO3

13. Oxohalogen Acids:

  1. Hypochlorous Acid - HOCl
  2. Hypobromous Acid - HOBr
  3. Hypoiodous Acid - HOI
  4. Chlorous Acid - HClO2
  5. Bromous Acid - HBrO2
  6. Iodous Acid - HIO2
  7. Perchloric Acid - HClO4
============================================================================

1. Halogen Acids

  1. Hydrofluoric Acid (HF):
    Uses: Glass etching, petroleum refining.
    Properties: Highly corrosive, weak acid, penetrates skin.

  2. Hydrochloric Acid (HCl):
    Uses: Metal cleaning, pH control, food processing.
    Properties: Strong acid, reacts with metals.

  3. Hydrobromic Acid (HBr):
    Uses: Organic synthesis, bromine production.
    Properties: Strong acid, highly corrosive.

  4. Hydroiodic Acid (HI):
    Uses: Organic reductions, pharmaceuticals.
    Properties: Strong reducing agent, colorless.

  5. Hydroastatic Acid (HAt):
    Uses: Rare, research in halogen chemistry.
    Properties: Highly radioactive, corrosive.

  6. Perchloric Acid (HClO4):
    Uses: Rocket fuel, explosives manufacturing.
    Properties: Strong acid, powerful oxidizer.

  7. Chloric Acid (HClO3):
    Uses: Explosives, bleaching agent.
    Properties: Strong oxidizer, highly unstable.

2. Oxyacids

  1. Nitric Acid (HNO3):
    Uses: Fertilizer production, explosives.
    Properties: Strong acid, powerful oxidizer.

  2. Sulfuric Acid (H2SO4):
    Uses: Battery acid, chemical manufacturing.
    Properties: Strong acid, dehydrating agent.

  3. Phosphoric Acid (H3PO4):
    Uses: Fertilizers, food additives.
    Properties: Moderate acid, non-volatile.

  4. Chromic Acid (H2CrO4):
    Uses: Metal plating, wood preservation.
    Properties: Strong oxidizer, highly toxic.

  5. Perchloric Acid (HClO4):
    Uses: Rocket propellants, analytical chemistry.
    Properties: Strong acid, very reactive.

  6. Hypophosphorous Acid (H3PO2):
    Uses: Reducing agent in plating, chemical reactions.
    Properties: Weak acid, reducing properties.

  7. Selenic Acid (H2SeO4):
    Uses: Oxidizing agent, electroplating.
    Properties: Strong acid, powerful oxidizer.

3. Acidic Oxides

  1. Carbonic Acid (H2CO3):
    Uses: Carbonated beverages, pH buffering.
    Properties: Weak acid, unstable in solution.

  2. Silicic Acid (H2SiO3):
    Uses: Ceramics, water purification.
    Properties: Weak acid, forms silica gels.

  3. Boric Acid (H3BO3):
    Uses: Antiseptic, insecticide.
    Properties: Weak acid, mild antiseptic.

  4. Germanic Acid (H2GeO3):
    Uses: Semiconductor manufacturing.
    Properties: Weak acid, reactive with metals.

  5. Telluric Acid (H6TeO6):
    Uses: Catalyst in chemical reactions.
    Properties: Strong acid, oxidizing agent.

  6. Stannic Acid (H2SnO3):
    Uses: Tin plating, corrosion inhibitor.
    Properties: Weak acid, amphoteric behavior.

  7. Oxalic Acid (H2C2O4):
    Uses: Rust removal, dyeing textiles.
    Properties: Strong organic acid, toxic.

4. Sulfur Containing Acids

  1. Sulfurous Acid (H2SO3):
    Uses: Bleaching agent, preservative.
    Properties: Weak acid, reducing agent.

  2. Thiosulfuric Acid (H2S2O3):
    Uses: Photography, gold extraction.
    Properties: Reducing agent, unstable in solution.

  3. Dithionic Acid (H2S2O6):
    Uses: Strong oxidizer in chemical processes.
    Properties: Powerful oxidizing agent.

  4. Polythionic Acid (H2SxO6):
    Uses: Found in volcanic gases, chemical research.
    Properties: Weak acid, varies in sulfur content.

  5. Sulfamic Acid (H3NSO3):
    Uses: Descaler, cleaning agent.
    Properties: Weak acid, stable in solid form.

  6. Peroxomonosulfuric Acid (H2SO5):
    Uses: Strong oxidizer, bleaching agent.
    Properties: Powerful oxidizing acid.

  7. Peroxodisulfuric Acid (H2S2O8):
    Uses: Polymerization initiator.
    Properties: Strong oxidizer, decomposes in water.

5. Nitrogen Containing Acids

  1. Hydrazoic Acid (HN3):
    Uses: Explosives, gas-generating agent in airbags.
    Properties: Highly explosive, toxic.

  2. Cyanic Acid (HOCN):
    Uses: Organic synthesis.
    Properties: Weak acid, unstable.

  3. Fulminic Acid (HCNO):
    Uses: Explosive, research in isomers.
    Properties: Extremely reactive, unstable.

  4. Isocyanic Acid (HNCO):
    Uses: Intermediate in organic chemistry.
    Properties: Unstable, colorless liquid.

  5. Cyanuric Acid (C3H3N3O3):
    Uses: Pool water treatment, stabilizer for chlorine.
    Properties: Weak acid, forms stable complexes.

  6. Nitrous Acid (HNO2):
    Uses: Diazotization reactions in organic synthesis.
    Properties: Weak acid, decomposes easily.

  7. Dinitramic Acid (HN(NO2)2):
    Uses: Energetic material, research in propellants.
    Properties: Highly explosive, unstable.

6. Phosphorus Containing Acids

  1. Hypophosphoric Acid (H4P2O6):
    Uses: Reducing agent in various reactions.
    Properties: Weak acid, forms complex salts.

  2. Pyrophosphoric Acid (H4P2O7):
    Uses: Detergents, food additives.
    Properties: Moderate acid, hydrolyzes in water.

  3. Tripolyphosphoric Acid (H5P3O10):
    Uses: Water softeners, detergents.
    Properties: Strong chelating agent.

  4. Tetrapolyphosphoric Acid (H6P4O13):
    Uses: Industrial water treatment, food processing.
    Properties: Forms soluble complexes with metals.

  5. Metaphosphoric Acid (HPO3):
    Uses: Drying agent, food additive.
    Properties: Highly hygroscopic, non-volatile.

  6. Phosphorous Acid (H3PO3):
    Uses: Fungicides, plastic stabilizers.
    Properties: Reducing agent, hygroscopic.

  7. Orthophosphoric Acid (H3PO4):
    Uses: Food industry, fertilizers.
    Properties: Weak acid, non-volatile.

7. Chlorine Containing Acids

  1. Hypochlorous Acid (HOCl)
    Uses: Disinfectant, bleaching agent.
    Properties: Weak acid, strong oxidizer.

  2. Chlorous Acid (HClO2)
    Uses: Bleaching paper, disinfectant.
    Properties: Strong oxidizer, unstable.

  3. Chloric Acid (HClO3)
    Uses: Explosives, oxidizing agent.
    Properties: Strong acid, highly reactive.

  4. Perchloric Acid (HClO4)
    Uses: Rocket fuel, analytical chemistry.
    Properties: Strong acid, powerful oxidizer.

  5. Dichloroacetic Acid (C2H2Cl2O2)
    Uses: Chemical synthesis, medical research.
    Properties: Toxic, carcinogenic.

  6. Trichloroacetic Acid (CCl3COOH)
    Uses: Cosmetic treatments, chemical peels.
    Properties: Strong acid, corrosive.

  7. Chlorosulfuric Acid (HSO3Cl)
    Uses: Sulfonation agent in organic synthesis.
    Properties: Strong acid, reactive with water.

8. Bromine Containing Acids

  1. Hypobromous Acid (HOBr)
    Uses: Disinfectant, oxidizer.
    Properties: Weak acid, unstable in light.

  2. Bromous Acid (HBrO2)
    Uses: Bleaching agent, oxidation.
    Properties: Strong oxidizer, unstable.

  3. Bromic Acid (HBrO3)
    Uses: Explosives, oxidizing agent.
    Properties: Strong acid, highly reactive.

  4. Perbromic Acid (HBrO4)
    Uses: Oxidizing agent in organic chemistry.
    Properties: Very strong acid, powerful oxidizer.

  5. Tribromacetic Acid (C2HBr3O2)
    Uses: Organic synthesis, chemical research.
    Properties: Strong acid, corrosive.

  6. Bromic Chloride (HBrClO4)
    Uses: Research in halogen chemistry.
    Properties: Highly reactive and corrosive.

  7. Bromosulfuric Acid (HBrSO3)
    Uses: Industrial chemical processing.
    Properties: Strong acid, reactive with metals.

9. Iodine Containing Acids

  1. Hypoiodous Acid (HOI)
    Uses: Disinfectant, organic synthesis.
    Properties: Weak acid, oxidizer.

  2. Iodous Acid (HIO2)
    Uses: Oxidizing agent in chemical reactions.
    Properties: Unstable, decomposes easily.

  3. Iodic Acid (HIO3)
    Uses: Disinfectants, analytical chemistry.
    Properties: Strong oxidizer, stable in solution.

  4. Periodic Acid (HIO4)
    Uses: Organic synthesis, protein structure analysis.
    Properties: Strong acid, powerful oxidizer.

  5. Diiodoacetic Acid (C2H2I2O2)
    Uses: Organic synthesis, research in iodine chemistry.
    Properties: Strong acid, highly corrosive.

  6. Periodic Acid (Ortho) (H5IO6)
    Uses: Research in analytical chemistry.
    Properties: Strong oxidizing agent, stable.

  7. Iodophosphoric Acid (H4IO4P)
    Uses: Organic synthesis, catalyst.
    Properties: Strong acid, highly reactive.

10. Other Inorganic Acids

  1. Arsenic Acid (H3AsO4)
    Uses: Wood preservatives, insecticides.
    Properties: Highly toxic, strong acid.

  2. Antimonic Acid (H3SbO4)
    Uses: Flame retardants, paints.
    Properties: Toxic, reacts with bases.

  3. Bismuthic Acid (HBiO3)
    Uses: Oxidizing agent, chemical research.
    Properties: Strong acid, unstable.

  4. Cadmium Acid (H2CdO3)
    Uses: Pigments, electroplating.
    Properties: Toxic, stable in solution.

  5. Stannic Acid (H2SnO3)
    Uses: Tin plating, catalyst.
    Properties: Amphoteric, forms salts with metals.

  6. Molybdic Acid (H2MoO4)
    Uses: Catalysts, metal corrosion inhibitor.
    Properties: Weak acid, forms complex salts.

  7. Vanadic Acid (H3VO4)
    Uses: Catalyst in organic synthesis.
    Properties: Strong oxidizer, reacts with metals.

11. Polyacids

  1. Pyrophosphoric Acid (H4P2O7)
    Uses: Water softeners, food additives.
    Properties: Moderate acid, hydrolyzes in water.

  2. Tripolyphosphoric Acid (H5P3O10)
    Uses: Detergents, water treatment.
    Properties: Strong chelating agent.

  3. Tetrapolyphosphoric Acid (H6P4O13)
    Uses: Industrial water treatment.
    Properties: Forms stable complexes with metals.

  4. Silicotungstic Acid (H4SiW12O40)
    Uses: Catalysts, chemical processing.
    Properties: Strong acid, highly reactive.

  5. Phosphotungstic Acid (H3PW12O40)
    Uses: Catalysts, biological stains.
    Properties: Strong acid, water-soluble.

  6. Polyphosphoric Acid (H6P4O13)
    Uses: Fertilizers, industrial processes.
    Properties: Strong acid, chelating agent.

  7. Polyboric Acid (H[B(OH)4])
    Uses: Antiseptic, neutron absorber.
    Properties: Weak acid, complexing agent.

12. Peroxo Acids

  1. Peroxomonosulfuric Acid (H2SO5)
    Uses: Bleaching agent, strong oxidizer.
    Properties: Powerful oxidizing acid.

  2. Peroxodisulfuric Acid (H2S2O8)
    Uses: Polymerization initiator, bleach.
    Properties: Strong oxidizer, decomposes in water.

  3. Peroxomonophosphoric Acid (H3PO5)
    Uses: Strong oxidizing agent in chemical reactions.
    Properties: Strong acid, reactive.

  4. Peroxodiphosphoric Acid (H4P2O8)
    Uses: Bleaching, strong oxidizer.
    Properties: Highly reactive, strong acid.

  5. Peroxonitric Acid (HNO4)
    Uses: Explosives, oxidizer.
    Properties: Strong acid, powerful oxidizer.

  6. Peroxochloric Acid (HClO5)
    Uses: Research in chemical synthesis.
    Properties: Powerful oxidizer, unstable.

  7. Peroxoboric Acid (H2BO3)
    Uses: Oxidizing agent in organic chemistry.
    Properties: Strong acid, reacts with metals.

13. Oxohalogen Acids

  1. Hypochlorous Acid (HOCl)
    Uses: Disinfectant, bleaching agent.
    Properties: Weak acid, strong oxidizer.

  2. Hypobromous Acid (HOBr)
    Uses: Disinfectant, oxidizer.
    Properties: Weak acid, unstable.

  3. Hypoiodous Acid (HOI)
    Uses: Organic synthesis, disinfectant.
    Properties: Weak acid, reactive.

  4. Chlorous Acid (HClO2)
    Uses: Oxidizer in bleaching, disinfecting.
    Properties: Unstable, decomposes easily.

  5. Bromous Acid (HBrO2)
    Uses: Oxidizing agent, bleaching.
    Properties: Unstable, strong oxidizer.

  6. Iodous Acid (HIO2)
    Uses: Research in iodine chemistry.
    Properties: Strong oxidizer, unstable.

  7. Perchloric Acid (HClO4)
    Uses: Rocket propellants, analytical chemistry.
    Properties: Strong acid, powerful oxidizer.

ORGANIC ACIDS


Here’s the detailed list of organic acids categorized by industry applications, including their molecular names:

1. Food and Beverage Industry

  • Food Preservation:

    1. Citric Acid - C6H8O7
    2. Lactic Acid - C3H6O3
    3. Acetic Acid - CH3COOH
    4. Sorbic Acid - C6H8O2
    5. Benzoic Acid - C7H6O2

  • Flavor Enhancement:

    1. Malic Acid - C4H6O5
    2. Tartaric Acid - C4H6O6
    3. Ascorbic Acid - C6H8O6
    4. Fumaric Acid - C4H4O4
    5. Glycolic Acid - C2H4O3

  • Beverage Production:

    1. Phosphoric Acid - H3PO4
    2. Citric Acid - C6H8O7
    3. Tartaric Acid - C4H6O6
    4. Lactic Acid - C3H6O3
    5. Acetic Acid - CH3COOH

  • Wine and Brewery:

    1. Tartaric Acid - C4H6O6
    2. Malic Acid - C4H6O5
    3. Lactic Acid - C3H6O3
    4. Citric Acid - C6H8O7
    5. Acetic Acid - CH3COOH

  • Dairy Products:

    1. Lactic Acid - C3H6O3
    2. Citric Acid - C6H8O7
    3. Acetic Acid - CH3COOH
    4. Tartaric Acid - C4H6O6
    5. Propionic Acid - C3H6O2

2. Pharmaceutical Industry

  • APIs Synthesis:

    1. Lactic Acid - C3H6O3
    2. Ascorbic Acid - C6H8O6
    3. Citric Acid - C6H8O7
    4. Salicylic Acid - C7H6O3
    5. Acetic Acid - CH3COOH

  • Drug Formulation:

    1. Citric Acid - C6H8O7
    2. Tartaric Acid - C4H6O6
    3. Lactic Acid - C3H6O3
    4. Glycolic Acid - C2H4O3
    5. Phosphoric Acid - H3PO4

  • Excipients:

    1. Citric Acid - C6H8O7
    2. Tartaric Acid - C4H6O6
    3. Malic Acid - C4H6O5
    4. Fumaric Acid - C4H4O4
    5. Sorbic Acid - C6H8O2

  • Vitamin Production:

    1. Ascorbic Acid - C6H8O6
    2. Nicotinic Acid - C6H6N2O
    3. Folic Acid - C19H19N7O6
    4. Pantothenic Acid - C9H17NO5
    5. Biotin - C10H16N2O3

3. Cosmetics and Personal Care

  • Skin Care Products:

    1. Glycolic Acid - C2H4O3
    2. Salicylic Acid - C7H6O3
    3. Lactic Acid - C3H6O3
    4. Citric Acid - C6H8O7
    5. Malic Acid - C4H6O5

  • Hair Care Products:

    1. Lactic Acid - C3H6O3
    2. Citric Acid - C6H8O7
    3. Glycolic Acid - C2H4O3
    4. Tartaric Acid - C4H6O6
    5. Oleic Acid - C18H34O2

  • Oral Care Products:

    1. Citric Acid - C6H8O7
    2. Phosphoric Acid - H3PO4
    3. Lactic Acid - C3H6O3
    4. Tartaric Acid - C4H6O6
    5. Malic Acid - C4H6O5

  • Fragrances and Perfumes:

    1. Benzoic Acid - C7H6O2
    2. Salicylic Acid - C7H6O3
    3. Acetic Acid - CH3COOH
    4. Lactic Acid - C3H6O3
    5. Citric Acid - C6H8O7

4. Agriculture and Animal Feed

  • Fertilizers:

    1. Citric Acid - C6H8O7
    2. Malic Acid - C4H6O5
    3. Fumaric Acid - C4H4O4
    4. Oxalic Acid - C2H2O4
    5. Acetic Acid - CH3COOH

  • Animal Feed Additives:

    1. Citric Acid - C6H8O7
    2. Fumaric Acid - C4H4O4
    3. Lactic Acid - C3H6O3
    4. Propionic Acid - C3H6O2
    5. Butyric Acid - C4H8O2

  • Veterinary Medicine:

    1. Acetic Acid - CH3COOH
    2. Lactic Acid - C3H6O3
    3. Citric Acid - C6H8O7
    4. Benzoic Acid - C7H6O2
    5. Salicylic Acid - C7H6O3

5. Chemical Industry

  • Polymer Production:

    1. Acrylic Acid - C3H4O2
    2. Methacrylic Acid - C4H6O2
    3. Maleic Acid - C4H4O4
    4. Fumaric Acid - C4H4O4
    5. Adipic Acid - C6H10O4

  • Resin Manufacturing:

    1. Acrylic Acid - C3H4O2
    2. Methacrylic Acid - C4H6O2
    3. Citric Acid - C6H8O7
    4. Tartaric Acid - C4H6O6
    5. Maleic Acid - C4H4O4

  • Adhesives and Coatings:

    1. Acrylic Acid - C3H4O2
    2. Citric Acid - C6H8O7
    3. Maleic Acid - C4H4O4
    4. Fumaric Acid - C4H4O4
    5. Tartaric Acid - C4H6O6

  • Cleaning Agents:

    1. Citric Acid - C6H8O7
    2. Acetic Acid - CH3COOH
    3. Lactic Acid - C3H6O3
    4. Phosphoric Acid - H3PO4
    5. Glycolic Acid - C2H4O3

6. Textile and Leather Industry

  • Dyeing and Printing:

    1. Citric Acid - C6H8O7
    2. Acetic Acid - CH3COOH
    3. Tartaric Acid - C4H6O6
    4. Malic Acid - C4H6O5
    5. Lactic Acid - C3H6O3

  • Leather Tanning:

    1. Acetic Acid - CH3COOH
    2. Formic Acid - HCOOH
    3. Lactic Acid - C3H6O3
    4. Tartaric Acid - C4H6O6
    5. Oxalic Acid - C2H2O4

7. Paper and Pulp Industry

  • Pulp Bleaching:

    1. Peracetic Acid - C2H4O3
    2. Citric Acid - C6H8O7
    3. Acetic Acid - CH3COOH
    4. Formic Acid - HCOOH
    5. Phosphoric Acid - H3PO4

  • Water Treatment:

    1. Citric Acid - C6H8O7
    2. Acetic Acid - CH3COOH
    3. Phosphoric Acid - H3PO4
    4. Sulfuric Acid - H2SO4 5


Here’s the organized list of 100 commonly used organic acids with their molecular names and purposes:

Carboxylic Acids

  1. Acetic Acid (CH₃COOH) - Used as a preservative and flavoring agent in food.
  2. Citric Acid (C₆H₈O₇) - Acts as a preservative and flavoring in beverages and candies.
  3. Lactic Acid (C₃H₆O₃) - Used in food preservation and as a pH regulator.
  4. Malic Acid (C₄H₆O₅) - Enhances flavor in food and beverages.
  5. Tartaric Acid (C₄H₆O₆) - Used in baking and as a stabilizing agent in cream of tartar.
  6. Glycolic Acid (C₂H₄O₃) - Commonly used in skincare products for exfoliation.
  7. Salicylic Acid (C₇H₆O₃) - Used in acne treatments and as a pain reliever.
  8. Adipic Acid (C₆H₁₀O₄) - Used in the production of nylon and food additives.
  9. Fumaric Acid (C₄H₄O₄) - Used as a food additive and in chemical synthesis.
  10. Succinic Acid (C₄H₆O₄) - Used in food and beverage flavoring.

Amino Acids

  1. Glycine (C₂H₅NO₂) - Used as a dietary supplement and in food processing.
  2. Alanine (C₃H₇NO₂) - Important for protein synthesis and energy metabolism.
  3. Valine (C₅H₁₁NO₂) - Used in nutritional supplements for muscle recovery.
  4. Leucine (C₆H₁₃NO₂) - Essential for protein synthesis and muscle repair.
  5. Isoleucine (C₆H₁₃NO₂) - Used for muscle metabolism and immune function.

Hydroxyacids

  1. Lactic Acid (C₃H₆O₃) - See above.
  2. Glycolic Acid (C₂H₄O₃) - See above.
  3. Malic Acid (C₄H₆O₅) - See above.
  4. Tartaric Acid (C₄H₆O₆) - See above.
  5. Citric Acid (C₆H₈O₇) - See above.

Keto Acids

  1. Pyruvic Acid (C₃H₄O₃) - Key intermediate in metabolic pathways.
  2. Oxaloacetic Acid (C₄H₄O₅) - Important in the citric acid cycle.
  3. Alpha-Ketoglutaric Acid (C₅H₆O₅) - Involved in amino acid synthesis.
  4. Beta-Ketoglutaric Acid (C₅H₆O₅) - Used in metabolic processes.

Fatty Acids

  1. Acrylic Acid (C₃H₄O₂) - Used in producing polymers and coatings.
  2. Methacrylic Acid (C₄H₆O₂) - Important in the manufacture of plastics.
  3. Oleic Acid (C₁₈H₃₄O₂) - Used in cosmetics and food products.
  4. Palmitic Acid (C₁₆H₃₂O₂) - Commonly found in animal fats and used in cosmetics.
  5. Stearic Acid (C₁₈H₃₆O₂) - Used in soaps and cosmetics.

Aromatic Acids

  1. Benzoic Acid (C₇H₆O₂) - Used as a food preservative.
  2. Salicylic Acid (C₇H₆O₃) - See above.
  3. Anthranilic Acid (C₇H₇NO₂) - Used in dyes and fragrances.
  4. p-Aminobenzoic Acid (C₇H₈NO₂) - Important in sunscreen formulations.

Heterocyclic Acids

  1. Nicotinic Acid (C₆H₅NO₂) - Used in dietary supplements and to treat cholesterol.
  2. Picolinic Acid (C₆H₅NO₂) - Acts as a chelator in nutritional applications.
  3. Quinolinic Acid (C₇H₅NO₄) - Involved in neurotransmitter metabolism.

Sulfonic Acids

  1. Methanesulfonic Acid (CH₄SO₃) - Used in chemical synthesis and as a solvent.
  2. Ethanesulfonic Acid (C₂H₅SO₃) - Used in pharmaceuticals and industrial applications.
  3. Propane Sulfonic Acid (C₃H₇SO₃) - Employed in surfactants and detergents.

Phosphonic Acids

  1. Phosphoric Acid (H₃PO₄) - Used in fertilizers and food flavoring.
  2. Glyceral Phosphoric Acid (C₃H₇O₆P) - Involved in energy metabolism.
  3. Phosphonoacetic Acid (C₂H₅O₅P) - Used in herbicides and insecticides.

Other Organic Acids

  1. Oxalic Acid (C₂H₂O₄) - Used in cleaning and bleaching agents.
  2. Maleic Acid (C₄H₄O₄) - Important in the production of resins and coatings.
  3. Glutaric Acid (C₅H₈O₄) - Used in plasticizers and pharmaceuticals.
  4. Pimelic Acid (C₇H₁₂O₄) - Used in the synthesis of polymers.
  5. Suberic Acid (C₈H₁₄O₄) - Used in the production of nylon.
  6. Azelaic Acid (C₉H₁₆O₄) - Used in skin treatments for acne and rosacea.
  7. Sebacic Acid (C₁₀H₁₈O₄) - Used in the production of lubricants and plastics.
  8. Undecylenic Acid (C₁₁H₂₀O₂) - Used in antifungal treatments.

Additional 50 Organic Acids

  1. Butyric Acid (C₄H₈O₂) - Used in flavoring and as a preservative.
  2. Caproic Acid (C₆H₁₂O₂) - Used in the production of esters for flavors and fragrances.
  3. Caprylic Acid (C₈H₁₆O₂) - Used in food additives and antimicrobial applications.
  4. Lauric Acid (C₁₂H₂₄O₂) - Used in soaps and detergents, as well as in food.
  5. Myristic Acid (C₁₄H₂₈O₂) - Used in cosmetics and as a food additive.
  6. Palmitoleic Acid (C₁₆H₃₀O₂) - Used in dietary supplements and skin care.
  7. Linoleic Acid (C₁₈H₃₂O₂) - An essential fatty acid used in food and supplements.
  8. Linolenic Acid (C₁₈H₃₀O₂) - An essential fatty acid important for human health.
  9. Arachidonic Acid (C₂₀H₃₂O₂) - Used in nutrition and as a precursor for eicosanoids.
  10. Eicosapentaenoic Acid (C₂₀H₃₀O₂) - Used in omega-3 fatty acid supplements.
  11. Docosahexaenoic Acid (C₂₂H₃₂O₂) - Important for brain health and found in fish oil.
  12. Cyclohexanecarboxylic Acid (C₇H₁₂O₂) - Used in organic synthesis and as a plasticizer.
  13. Cyclopropanecarboxylic Acid (C₄H₆O₂) - Used in synthetic organic chemistry.
  14. Dichloroacetic Acid (C₂H₂Cl₂O₂) - Used in the treatment of certain medical conditions.
  15. Trichloroacetic Acid (C₂HCl₃O₂) - Used in chemical peels and as a reagent in laboratories.
  16. Chloroacetic Acid (C₂H₃ClO₂) - Used in the synthesis of various organic compounds.
  17. Bromoacetic Acid (C₂H₃BrO₂) - Used in organic synthesis and as a reagent.
  18. Iodoacetic Acid (C₂H₃IO₂) - Used in biochemical research as a reagent.
  19. Glyoxylic Acid (C₂H₂O₃) - Used in organic synthesis and as a reagent in laboratories.
  20. Oxamic Acid (C₂H₃NO₃) - Used in organic synthesis and medicinal chemistry.
  21. Methylmalonic Acid (C₄H₆O₄) - Used in biochemical research and diagnostics.
  22. Ethylmalonic Acid (C₅H₈O₄) - Involved in metabolic pathways and studied in metabolism.
  23. Propylmalonic Acid (C₆H₁₀O₄) - Used in organic synthesis.
  24. Butylmalonic Acid (C₇H₁₂O₄) - Used in chemical research and synthesis.
  25. Pentylmalonic Acid (C₈H₁₄O₄) - Used in organic synthesis and research.
  26. Hexylmalonic Acid (C₉H₁₆O₄) - Used in organic chemistry applications.
  27. Heptylmalonic Acid (C₁₀H₁₈O₄) - Used in organic synthesis.
  28. Octylmalonic Acid (C₁₁H₂₀O₄) - Used in chemical synthesis and research.
  29. Nonylmalonic Acid (C₁₂H₂₂O₄) - Used in organic chemistry.
  30. Decylmalonic Acid (C₁₃H₂₄O₄) - Used in synthesis of organic compounds.
  31. Undecylmalonic Acid (C₁₄H₂₆O₄) - Used in organic synthesis.
  32. Dodecylmalonic Acid (C₁₅H₂₈O₄) - Used in chemical research.
  33. Tridecylmalonic Acid (C₁₆H₃₀O₄) - Used in organic chemistry.
  34. Tetradecylmalonic Acid (C₁₇H₃₂O₄) - Used in synthetic applications.
  35. Pentadecylmalonic Acid (C₁₈H₃₄O₄) - Used in chemical research.
  36. Hexadecylmalonic Acid (C₁₉H₃₆O₄) - Used in organic synthesis.
  37. Heptadecylmalonic Acid (C₂₀H₃₈O₄) - Used in organic chemistry.
  38. Octadecylmalonic Acid (C₂₁H₄₀O₄) - Used in synthetic chemistry.
  39. Nonadecylmalonic Acid (C₂₂H₄₂O₄) - Used in research and synthesis.
  40. Eicosylmalonic Acid (C₂₃H₄₄O₄) - Used in organic chemistry.
  41. Heneicosylmalonic Acid (C₂₄H₄₆O₄) - Used in organic synthesis.
  42. Docosylmalonic Acid (C₂₅H₄₈O₄) - Used in chemical research.
  43. Tricosylmalonic Acid (C₂₆H₅₀O₄) - Used in organic synthesis.
  44. Tetracosylmalonic Acid (C₂₇H₅₂O₄) - Used in organic chemistry.
  45. Pentacosylmalonic Acid (C₂₈H₅₄O₄) - Used in chemical synthesis.
  46. Hexacosylmalonic Acid (C₂₉H₅₆O₄) - Used in organic chemistry.
  47. Heptacosylmalonic Acid (C₃₀H₅₈O₄) - Used in research and synthesis.
  48. Octacosylmalonic Acid (C₃₁H₆₀O₄) - Used in organic chemistry.
  49. Nonacosylmalonic Acid (C₃₂H₆₂O₄) - Used in synthetic applications.
  50. Triacontylmalonic Acid (C₃₃H₆₄O₄) - Used in organic synthesis.

GURUJI


🤷 Something to figure out or answer is, how a rock from asteroid belt propel towards earth. What causes it come our way? That one rock wiped out dinosaurs.

🤷 This is what is happening. But the question is Why is it happening only to mercury but not earth?

🏋️ Research on Vulcan planet. It's hypothetical

😇 You may find clue for mercury

🤯 That was one of the theory before Einstein

💆 Vulcan planet

https://phys.org/news/2008-10-solar-young-twin-asteroid-belts.amp

https://science.howstuffworks.com/asteroid-belt.htm

https://www.universetoday.com/130136/far-asteroid-belt-earth/amp/

https://wisp.physics.wisc.edu/astro104/lecture19/lec19g.html

http://www.cbat.eps.harvard.edu/special/rocknroll/0003834.html

The first minor planet to be discovered was Ceres in 1801, though it was called a 'planet' at the time and an 'asteroid' soon after; the term minor planet was not introduced until 1841, and was considered a subcategory of 'planet' until 1932.

https://en.wikipedia.org/wiki/Minor_planet#:~:text=The%20first%20minor%20planet%20to,of%20'planet'%20until%201932.





*Noted*

Uranus is slightly larger in diameter than its neighbor Neptune, yet smaller in mass. It is the second least dense planet; Saturn is the least dense of all.

Uranus gets its blue-green color from methane gas in the atmosphere. Sunlight passes through the atmosphere and is reflected back out by Uranus' cloud tops. Methane gas absorbs the red portion of the light, resulting in a blue-green color.  ( Composites Color ).🤔🤨😨💆


https://en.wikipedia.org/wiki/Orbital_eccentricity#/media/File:Animation_of_Orbital_eccentricity.gif



WHO ARE YOU { 😎 vulcan 🧭🌐 }🙇 ‼️⁉️

the gravitational constant


The central body in an orbital system can be defined as the one whose mass (M) is much larger than the mass of the orbiting body (m), or M ≫ m.



Notes for Vulcan Planet:

The idea of Vulcan as a planet that existed between the sun and Mercury was proposed between 1859 and 1915 to explain anomalies in Mercury's orbit. 

The International Astronomical Union still reserves the name "Vulcan" for the hypothetical planet and for any hypothetical asteroids that may exist within Mercury's orbit. 
 

Mercury is denser than Venus because it has a larger metallic core and is made of heavier elements: 
 
Density: 
Mercury's density is 13.59 grams per cubic centimeter, while 
Venus's average density is 5.2 grams per cubic centimeter. 
 
Core: 
Mercury's core is large and metallic, with a radius of about 1,289 miles (2,074 kilometers), which is about 85% of the planet's radius. 
 
Elements:
Mercury is made of heavier elements than Venus. 
 

Earth is the densest object in our solar system, even though it's made of lighter materials and has a substantial atmosphere. This is because Earth has enough mass to compress itself significantly due to gravity. 


27 Sept 2024

Ionic Radius

Group-wise Ionic Radii

Group 1: Alkali Metals (Cations)

  • Atomic Number 3: Li⁺ - 76 pm
  • Atomic Number 11: Na⁺ - 102 pm
  • Atomic Number 19: K⁺ - 138 pm
  • Atomic Number 37: Rb⁺ - 152 pm
  • Atomic Number 55: Cs⁺ - 167 pm
  • Atomic Number 87: Fr²⁺ - 270 pm (estimated)

Group 2: Alkaline Earth Metals (Cations)

  • Atomic Number 4: Be²⁺ - 45 pm
  • Atomic Number 12: Mg²⁺ - 72 pm
  • Atomic Number 20: Ca²⁺ - 100 pm
  • Atomic Number 38: Sr²⁺ - 113 pm
  • Atomic Number 56: Ba²⁺ - 114 pm
  • Atomic Number 88: Ra²⁺ - 130 pm (estimated)

Group 13: Boron Group (Cations)

  • Atomic Number 5: B³⁺ - 27 pm (very small)
  • Atomic Number 13: Al³⁺ - 53 pm
  • Atomic Number 31: Ga³⁺ - 135 pm
  • Atomic Number 49: In³⁺ - 100 pm
  • Atomic Number 81: Tl³⁺ - 156 pm

Group 14: Carbon Group (Cations)

  • Atomic Number 6: C⁴⁻ - 182 pm (as carbide, larger anion)
  • Atomic Number 14: Si⁴⁺ - 40 pm

Group 15: Nitrogen Group (Anions and Cations)

  • Atomic Number 7: N³⁻ - 140 pm
  • Atomic Number 15: P³⁻ - 175 pm
  • Atomic Number 33: As³⁻ - 197 pm

Group 16: Chalcogens (Anions)

  • Atomic Number 8: O²⁻ - 140 pm
  • Atomic Number 16: S²⁻ - 184 pm
  • Atomic Number 34: Se²⁻ - 198 pm
  • Atomic Number 52: Te²⁻ - 206 pm
  • Atomic Number 86: Po²⁻ - 150 pm (estimated)

Group 17: Halogens (Anions)

  • Atomic Number 9: F⁻ - 133 pm
  • Atomic Number 17: Cl⁻ - 181 pm
  • Atomic Number 35: Br⁻ - 196 pm
  • Atomic Number 53: I⁻ - 220 pm
  • Atomic Number 85: At⁻ - 202 pm (estimated)

Summary

  • Cations (from Groups 1, 2, 13, and some in 14 and 15) typically have smaller ionic radii compared to their neutral atoms due to the loss of electrons.
  • Anions (from Groups 15, 16, and 17) have larger ionic radii due to the gain of electrons.

These values are approximate and can vary based on specific compounds and conditions.


Atomic NumberElementIonic Radius (pm)
1H120
3Li152
4Be112 (Be²⁺)
5B88 (B³⁺)
6C77 (C⁴⁻)
7N75 (N³⁻)
8O73 (O²⁻)
9F72 (F⁻)
11Na186
12Mg160 (Mg²⁺)
13Al143 (Al³⁺)
14Si118 (Si⁴⁻)
15P110 (P³⁻)
16S104 (S²⁻)
17Cl99 (Cl⁻)
19K227
20Ca197 (Ca²⁺)
21Sc162 (Sc³⁺)
22Ti147 (Ti⁴⁺)
23V133 (V⁵⁺)
24Cr133 (Cr³⁺)
25Mn120 (Mn²⁺)
26Fe130 (Fe²⁺)
27Co125 (Co²⁺)
28Ni124 (Ni²⁺)
29Cu128 (Cu²⁺)
30Zn88 (Zn²⁺)
31Ga135 (Ga³⁺)
32Ge124 (Ge⁴⁺)
33As114 (As³⁻)
34Se116 (Se²⁻)
35Br114 (Br⁻)
36Kr88
37Rb248
38Sr215 (Sr²⁺)
39Y162 (Y³⁺)
40Zr160 (Zr⁴⁺)
41Nb146 (Nb⁵⁺)
42Mo139 (Mo³⁺)
43Tc136 (Tc⁷⁺)
44Ru127 (Ru⁴⁺)
45Rh135 (Rh³⁺)
46Pd125 (Pd²⁺)
47Ag155 (Ag⁺)
48Cd96 (Cd²⁺)
49In165 (In³⁺)
50Sn118 (Sn⁴⁺)
51Sb133 (Sb³⁻)
52Te134 (Te²⁻)
53I133 (I⁻)
54Xe140
55Cs262
56Ba214 (Ba²⁺)
57La126 (La³⁺)
58Ce113 (Ce³⁺)
59Pr108 (Pr³⁺)
60Nd106 (Nd³⁺)
61Pm103 (Pm³⁺)
62Sm96 (Sm³⁺)
63Eu95 (Eu³⁺)
64Gd93 (Gd³⁺)
65Tb92 (Tb³⁺)
66Dy89 (Dy³⁺)
67Ho87 (Ho³⁺)
68Er86 (Er³⁺)
69Tm85 (Tm³⁺)
70Yb83 (Yb³⁺)
71Lu79 (Lu³⁺)
72Hf175 (Hf⁴⁺)
73Ta146 (Ta⁵⁺)
74W137 (W⁶⁺)
75Re139 (Re⁶⁺)
76Os133 (Os⁴⁺)
77Ir137 (Ir⁴⁺)
78Pt133 (Pt²⁺)
79Au144 (Au⁺)
80Hg96 (Hg²⁺)
81Tl156 (Tl³⁺)
82Pb120 (Pb²⁺)
83Bi120 (Bi³⁺)
84Po135 (Po²⁻)
85At202 (At⁻)
86Rn120
87Fr270
88Ra215 (Ra²⁺)
89Ac186 (Ac³⁺)
90Th206 (Th⁴⁺)
91Pa187 (Pa⁴⁺)
92U196 (U⁶⁺)
93Np193 (Np⁶⁺)
94Pu195 (Pu⁶⁺)
95Am197 (Am³⁺)
96Cm196 (Cm³⁺)
97Bk195 (Bk³⁺)
98Cf197 (Cf³⁺)
99Es198 (Es³⁺)
100Fm197 (Fm³⁺)
101Md196 (Md³⁺)
102No195 (No³⁺)
103Lr193 (Lr³⁺)
104Rf200 (Rf⁴⁺)
105Db197 (Db⁵⁺)
106Sg195 (Sg⁶⁺)
107Bh193 (Bh⁷⁺)
108Hs191 (Hs⁶⁺)
109Mt189 (Mt⁸⁺)
110Ds187 (Ds⁵⁺)
111Rg185 (Rg⁴⁺)
112Cn183 (Cn²⁺)
113Nh181 (Nh³⁺)
114Fl180 (Fl⁴⁺






Atomic NumberElementIonic Radius (pm)
1Hydrogen (H)120 (H⁺)
3Lithium (Li)152 (Li⁺)
4Beryllium (Be)112 (Be²⁺)
5Boron (B)88 (B³⁺)
6Carbon (C)77 (C⁴⁻)
7Nitrogen (N)75 (N³⁻)
8Oxygen (O)73 (O²⁻)
9Fluorine (F)72 (F⁻)
11Sodium (Na)186 (Na⁺)
12Magnesium (Mg)160 (Mg²⁺)
13Aluminum (Al)143 (Al³⁺)
14Silicon (Si)118 (Si⁴⁻)
15Phosphorus (P)110 (P³⁻)
16Sulfur (S)104 (S²⁻)
17Chlorine (Cl)99 (Cl⁻)
19Potassium (K)227 (K⁺)
20Calcium (Ca)197 (Ca²⁺)
21Scandium (Sc)162 (Sc³⁺)
22Titanium (Ti)147 (Ti⁴⁺)
23Vanadium (V)133 (V⁵⁺)
24Chromium (Cr)133 (Cr³⁺)
25Manganese (Mn)120 (Mn²⁺)
26Iron (Fe)130 (Fe²⁺)
27Cobalt (Co)125 (Co²⁺)
28Nickel (Ni)124 (Ni²⁺)
29Copper (Cu)128 (Cu²⁺)
30Zinc (Zn)88 (Zn²⁺)
31Gallium (Ga)135 (Ga³⁺)
32Germanium (Ge)124 (Ge⁴⁺)
33Arsenic (As)114 (As³⁻)
34Selenium (Se)116 (Se²⁻)
35Bromine (Br)114 (Br⁻)
36Krypton (Kr)88
37Rubidium (Rb)248 (Rb⁺)
38Strontium (Sr)215 (Sr²⁺)
39Yttrium (Y)162 (Y³⁺)
40Zirconium (Zr)160 (Zr⁴⁺)
41Niobium (Nb)146 (Nb⁵⁺)
42Molybdenum (Mo)139 (Mo³⁺)
43Technetium (Tc)136 (Tc⁷⁺)
44Ruthenium (Ru)127 (Ru⁴⁺)
45Rhodium (Rh)135 (Rh³⁺)
46Palladium (Pd)125 (Pd²⁺)
47Silver (Ag)155 (Ag⁺)
48Cadmium (Cd)96 (Cd²⁺)
49Indium (In)165 (In³⁺)
50Tin (Sn)118 (Sn⁴⁺)
51Antimony (Sb)133 (Sb³⁻)
52Tellurium (Te)134 (Te²⁻)
53Iodine (I)133 (I⁻)
54Xenon (Xe)140
55Cesium (Cs)262 (Cs⁺)
56Barium (Ba)214 (Ba²⁺)
57Lanthanum (La)126 (La³⁺)
58Cerium (Ce)113 (Ce³⁺)
59Praseodymium (Pr)108 (Pr³⁺)
60Neodymium (Nd)106 (Nd³⁺)
61Promethium (Pm)103 (Pm³⁺)
62Samarium (Sm)96 (Sm³⁺)
63Europium (Eu)95 (Eu³⁺)
64Gadolinium (Gd)93 (Gd³⁺)
65Terbium (Tb)92 (Tb³⁺)
66Dysprosium (Dy)89 (Dy³⁺)
67Holmium (Ho)87 (Ho³⁺)
68Erbium (Er)86 (Er³⁺)
69Thulium (Tm)85 (Tm³⁺)
70Ytterbium (Yb)83 (Yb³⁺)
71Lutetium (Lu)79 (Lu³⁺)
72Hafnium (Hf)175 (Hf⁴⁺)
73Tantalum (Ta)146 (Ta⁵⁺)
74Tungsten (W)137 (W⁶⁺)
75Rhenium (Re)139 (Re⁶⁺)
76Osmium (Os)133 (Os⁴⁺)
77Iridium (Ir)137 (Ir⁴⁺)
78Platinum (Pt)133 (Pt²⁺)
79Gold (Au)144 (Au⁺)
80Mercury (Hg)96 (Hg²⁺)
81Thallium (Tl)156 (Tl³⁺)
82Lead (Pb)120 (Pb²⁺)
83Bismuth (Bi)120 (Bi³⁺)
84Polonium (Po)135 (Po²⁻)
85Astatine (At)202 (At⁻)
86Radon (Rn)120
87Francium (Fr)270 (Fr⁺)
88Radium (Ra)215 (Ra²⁺)
89Actinium (Ac)186 (Ac³⁺)
90Thorium (Th)206 (Th⁴⁺)
91Protactinium (Pa)187 (Pa⁴⁺)
92Uranium (U)196 (U⁶⁺)
93Neptunium (Np)193 (Np⁶⁺)
94Plutonium (Pu)187

CMB

https://map.gsfc.nasa.gov/resources/cmbtheory.html

https://map.gsfc.nasa.gov/resources/cmblinks.html

26 Sept 2024

ENCYCLOPEDIA

https://www.sciencedirect.com/browse/journals-and-books

Browse 5,231 journals and 35,993 books




Encyclopedia of Mathematical Physics


Encyclopedia of Materials: Science and Technology

Encyclopedia of the World's Biomes

Encyclopedia of Materials: Metals and Alloys

Encyclopedia of Energy

Encyclopedia of Geology

Encyclopedia of Materials: Electronics

Encyclopedia of Biodiversity

Encyclopedia of Ecology


https://onlinelibrary.wiley.com/ 

https://link.springer.com/#browse-by-subjects

HAWKING RADIATION

https://www.sciencedirect.com/topics/physics-and-astronomy/hawking-radiation

23 Sept 2024

HOW COLORS CREATED

The sun emits radiation across a wide range of wavelengths, with the majority of its energy falling within a narrow bandwidth. The percentage of solar radiation emitted at different wavelengths is: 
 
Ultraviolet (UV)
About 7% of the sun's radiation is in the UV range, from 100–400 nanometers (nm). 
 
Visible light
About 42–43% of the sun's radiation is in the visible light range, from 380–780 nm. 
 
Near infrared (NIR)
About 52–55% of the sun's radiation is in the NIR range, from greater than 780 nm. 
 
Infrared
About 49% of the sun's radiation is in the infrared range, from 700 nm–1 mm. 
 
X-rays, gamma rays, and radio waves
Less than 1% of the sun's radiation is emitted as these types of radiation. 
 
The ozone layer absorbs most of the sun's UV radiation, protecting Earth. The hole in the ozone layer over the Antarctic allows more UV radiation to reach the Earth, which can be harmful to habitats.

*The sun emits all colors of visible light*, and, *in fact, emits all frequencies of electromagnetic waves*. This includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays. 

The sun emits all of the colors of the visible spectrum because it is a hot thermal body that emits light through the process of thermal radiation.

*Why is the sun yellow color ?*

It may be tempting to examine the color content of sunlight and identify the brightest color (the peak frequency) as the actual color of the sun.

The sun's actual color is white, but it appears to have different colors due to the Earth's atmosphere: 
 
White light
The sun emits light across the entire visible spectrum of colors, from red to violet, in roughly equal amounts. This combination is perceived as white light. 
 
Scattering
The Earth's atmosphere scatters blue light more than red light, which makes the sun appear yellow. This phenomenon is called *Rayleigh scattering*. 
 
Sun position
The sun appears white at noon because it's directly overhead, so its light has less air to travel through and less scattering from dust and other particles. 
 
Sunrises and sunsets
During sunrises and sunsets, the sun's light passes through more atmosphere, causing even more blue light to scatter. This means that the sun's spectrum has a greater percentage of red light during these times.

Close examination of the visible-light spectrum from our Sun.

Patterns are also evident in a graph of an object's reflectance. Elements, molecules, and even cell structures have unique signatures of reflectance. 

A graph of an object's reflectance across a spectrum is called a spectral signature. 

Spectral signatures of different Earth features within the visible light spectrum ARE shown below.

*This is how light interacts with atoms*

https://dept.harpercollege.edu/chemistry/chm/100/dgodambe/thedisk/spec/5back4.htm

*Absorption* and *Emission*

An atom changes from a ground state to an excited state by taking on energy from its surroundings in a process called absorption.

 The electron absorbs the energy and jumps to a higher energy level. 

In the reverse process, emission, the electron returns to the ground state by releasing the extra energy it absorbed.

The test tube in the figure contains an orange solution. The solar spectra is white light. When sunlight shines through an orange solution, the violet, blue and green *wavelengths are absorbed*. The other colors pass through. The transmitted light or reflected is the light we see, and *it looks orange.*

When sunlight is shined on a green leaf, the violet, red and orange wavelengths are absorbed. The reflected wavelengths appear green. ( Second image above mentioned )


*Why black and white in light ?*

Black and white objects are just the extremes of colored objects. Black objects absorb all the light shined on them. There is no reflected light, so we see black (the absence of color). If all of the light is reflected, we see all the wavelengths, which means we see white light

Then you may ask.. what aboUt other colors happens 🤔


*Why do objects emit complementary colors?*

The excited electrons absorb certain wavelengths of light. What humans see is the complementary color of the absorbed wavelengths,

 i.e., the remaining wavelengths of light that are not absorbed. For example, if an object absorbs the red wavelengths of light, we will perceive it as green (red's complementary color).


*Complementary colors* are *pairs of colors* which, when combined or mixed, cancel each other out (lose hue) by producing a grayscale color like white or black. When placed next to each other, they create the strongest contrast for those two colors. Complementary colors may also be called "opposite colors".


*Ok simple question:*

Why does the pink color happen in the sun's refection of an object?

 *Notice*: that we also don't see pink as a color on the visible light spectrum. 

This is because pink is also a mixture of colors. *We see pink when both red and violet light are reflected.*

When we see color, we see the wavelength of light reflected, or not absorbed, by the object.


https://dept.harpercollege.edu/chemistry/chm/100/dgodambe/thedisk/spec/complem.htm

Thermal radiation happens because excited state electrons jump ground states and emit thermal energy.


*Is molecules radiating or reflecting the colors ?*

*In a general sense :*

Molecules reflect light to create color, but molecules themselves do not have color: 
 
Reflection

The electrons on the outside of an atom or molecule absorb light and then emit some of it back. The color we see is determined by the wavelengths of light that are reflected. For example, an apple appears red because its surface reflects red wavelengths and absorbs the rest.

*Deep sense:*

Light reflects off molecules when the light waves' frequencies don't match the natural vibration frequencies of the molecules' atoms.

When this happens, the electrons in the atoms vibrate briefly, then re-emit the energy as a reflected light wave


Here are some other things to know about light and molecules: 
 
*Light absorption*
When light waves' frequencies match the energy levels of the electrons in a material, the electrons absorb the light's energy and change their energy state. The energy can be re-emitted as a photon, or it can be retained by the material, which heats up. 
 
*Light scattering*
Light scatters in all directions when it hits small particles, like gas molecules, or larger particles, like dust or water droplets. The amount of scattering depends on the size of the particle compared to the wavelength of light. 
 
*Chromophores*
Chromophores are regions in molecules where the energy difference between two molecular orbitals is within the visible spectrum. When light hits a chromophore, an electron can be excited from its ground state to an excited state. 
 
*Angle of incidence and reflection*
When light reflects off a surface, the angle at which it comes in (the angle of incidence) always equals the angle at which it reflects.


*Visible Light Reflection and Transmission*

Reflection and transmission of light waves occur because the frequencies of the light waves do not match the natural frequencies of vibration of the objects. When light waves of these frequencies strike an object, the electrons in the atoms of the object begin vibrating. But instead of vibrating in resonance at a large amplitude, the electrons vibrate for brief periods of time with small amplitudes of vibration; then the energy is reemitted as a light wave. 

If the object is transparent, then the vibrations of the electrons are passed on to neighboring atoms through the bulk of the material and reemitted on the opposite side of the object. Such frequencies of light waves are said to be transmitted. If the object is opaque, then the vibrations of the electrons are not passed from atom to atom through the bulk of the material. Rather the electrons of atoms on the material's surface vibrate for short periods of time and then reemit the energy as a reflected light wave. Such frequencies of light are said to be reflected.


21 Sept 2024

NON-BARYONIC DARK MATTER

Non-baryonic dark matter comes in two basic forms: hot and cold. These terms refer to the ability for this dark matter to cool to non-relativistic velocities and clump into smaller units

https://ned.ipac.caltech.edu/level5/Bothun2/Bothun4_6_4.html


20 Sept 2024

DARK ENERGY

https://science.nasa.gov/universe/the-universe-is-expanding-faster-these-days-and-dark-energy-is-responsible-so-what-is-dark-energy/

https://www.energy.gov/science/articles/our-expanding-universe-delving-dark-energy

COSMIC RAYS

hyperphysics.phy-astr.gsu.edu/hbase/Astro/cosmic.html

NEUTRINOS

https://www.researchgate.net/figure/aCosmic-Ray-shower-in-earth-atmosphere_fig1_369014844

PROTON

There are a few ways that magnetic fields and proton beams interact: 
 
Proton spin
Protons have a magnetic moment due to a quantum property called spin, which makes them act like tiny bar magnets. When placed in an external magnetic field, a proton's spin can either align with the field or flip to oppose it. 
 
Proton movement in a magnetic field
When a proton moves in a magnetic field, it experiences a force that causes it to move in circles. The magnetic field "steers" the charged particle, but it doesn't make it speed up or slow down. 
 
Proton beam and magnetic field direction
If a proton beam moves parallel to a magnetic field, no force acts on it, and it continues moving in a straight line.

PHOTONS

Yes, photons have spin: 
 
Spin angular momentum (SAM)
Photons have SAM, which comes from the rotation of their electric field. SAM can only point forward or backward relative to the direction of the beam. 
 
Spin number
Photons can have a spin of either -1 or +1. 
 
Spin and polarization
The polarization of light is often considered its "intrinsic" spin degree of freedom. 
 
Spin and bosons
Photons are bosons because they have an integer spin. By the spin-statistics theorem, all bosons follow Bose–Einstein statistics. 
 
Spin and 3D movies
The spin property of photons is used to create 3D movies. In a 3D movie theater, one image is projected with positive-spin photons and the other with negative-spin photons. The 3D glasses filter out either the positive-spin or negative-spin photons.