Chemical Bonding and Molecular Structure mind map

Chemical Bonding and Molecular Structure: what the chapter covers

The main idea of each topic with its first key points. Open a branch above for everything, or read the full notes.

1. 1Chapter Overview

   Chemical Bonding and Molecular Structure explains why atoms combine, how bonds form, and why molecules have definite shapes and properties. Atoms bond mainly to attain lower energy and more stable electronic arrangements, often resembling noble gas configurations. Ionic bonding involves electron transfer and electrostatic attraction, while covalent bonding involves sharing of electron pairs. Bond parameters such as bond length, bond angle, bond enthalpy, bond order and polarity help compare molecules. VSEPR theory predicts shapes using electron-pair repulsions. Valence bond theory explains overlap and hybridization, while molecular orbital theory explains bonding using delocalized molecular orbitals and predicts magnetic behavior. Hydrogen bonding, metallic bonding and weak interactions explain many physical properties important for NEET.

   • The driving force for bonding is lowering of potential energy and attaining stability.
   • Octet rule is useful but not universal; many molecules show incomplete, expanded or odd-electron octets.
   • Molecular shape depends on electron-pair geometry and the number of lone pairs on the central atom.
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2. 2Ionic & Covalent Bonding

   Ionic and covalent bonding are the two basic ways atoms combine. Ionic bonding involves transfer of electrons from an electropositive atom to an electronegative atom, producing oppositely charged ions held by electrostatic attraction. It is favored by low ionization enthalpy, high electron gain enthalpy and high lattice enthalpy. Covalent bonding involves sharing of electron pairs between atoms, usually non-metals, to attain stable valence-shell arrangements. Lewis dot structures show valence electrons, bonding pairs and lone pairs. The octet rule is useful but has exceptions such as BF3, PCl5, SF6 and NO. Formal charge helps choose the best Lewis structure, while resonance explains delocalization when one structure is insufficient. Polar covalent bonds arise due to unequal sharing of electrons.

   • Ionic bond formation is favorable when the energy released by lattice formation compensates energy required for ion formation.
   • Covalent bonds may be single, double or triple depending on number of shared electron pairs.
   • A coordinate bond is a covalent bond in which both shared electrons are donated by one atom.
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3. 3Bond Parameters

   Bond parameters are measurable quantities that describe the nature, strength and geometry of chemical bonds. Bond length is the equilibrium distance between two bonded nuclei and generally decreases as bond order increases. Bond angle is the angle between two bonds around a central atom and is strongly influenced by hybridization and lone-pair repulsion. Bond enthalpy measures the energy required to break a bond in the gaseous state and indicates bond strength. Bond order represents the number of bonds between atoms or, in molecular orbital theory, half the difference between bonding and antibonding electrons. Dipole moment measures polarity and depends on both bond polarity and molecular shape. NEET often asks trend comparisons using these parameters.

   • Bond length is affected by atomic size, bond order, resonance and hybridization.
   • More s-character in hybrid orbitals usually gives shorter and stronger bonds.
   • Bond enthalpy is always positive for bond dissociation because energy is required to break bonds.
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4. 4VSEPR Theory & Molecular Shapes

   VSEPR theory, or Valence Shell Electron Pair Repulsion theory, predicts molecular shapes by assuming that electron pairs around a central atom repel each other and arrange as far apart as possible. Both bonding pairs and lone pairs are counted as electron domains, but molecular shape is described using only the positions of atoms. The strength of repulsion follows lone pair-lone pair greater than lone pair-bond pair greater than bond pair-bond pair. Therefore, lone pairs compress bond angles and distort ideal geometries. Two electron domains give linear geometry, three give trigonal planar, four give tetrahedral, five give trigonal bipyramidal and six give octahedral electron-pair geometry. NEET strongly focuses on shapes of BeCl2, BF3, CH4, NH3, H2O, PCl5, SF6, XeF2 and XeF4.

   • Multiple bonds count as one electron domain in VSEPR shape prediction.
   • Lone pairs occupy more space than bonding pairs because they are attracted by only one nucleus.
   • The actual bond angle is often less than ideal when lone pairs are present.
   Read VSEPR Theory & Molecular Shapes notes

5. 5Valence Bond Theory & Hybridization

   Valence bond theory explains covalent bond formation by overlap of half-filled atomic orbitals. Greater overlap gives stronger bonds, and the bond lies in the region where electron density is concentrated between nuclei. End-to-end overlap forms a sigma bond, while sidewise overlap forms a pi bond. Hybridization is the mixing of atomic orbitals of similar energy on the same atom to form equivalent hybrid orbitals with definite geometry. sp, sp2 and sp3 hybridizations explain linear, trigonal planar and tetrahedral arrangements respectively. Expanded geometries such as trigonal bipyramidal and octahedral are often described using sp3d and sp3d2 in the NCERT-level model. Hybridization also explains bond angles, equivalent bonds and the relation between s-character and bond strength.

   • Only orbitals of comparable energy and proper orientation combine effectively in hybridization.
   • Hybrid orbitals form sigma bonds or hold lone pairs; unhybridized p orbitals form pi bonds.
   • Hybridization is predicted from steric number in many NEET questions.
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6. 6Molecular Orbital Theory

   Molecular orbital theory describes bonding by combining atomic orbitals to form molecular orbitals spread over the entire molecule. When atomic orbitals combine constructively, a lower-energy bonding molecular orbital forms; when they combine destructively, a higher-energy antibonding molecular orbital forms. Electrons fill molecular orbitals according to Aufbau principle, Pauli exclusion principle and Hund’s rule. Bond order is calculated as half the difference between bonding and antibonding electrons. A positive bond order indicates stability, while zero bond order means the molecule is unstable. MOT successfully explains bond order, relative bond length, bond enthalpy and magnetic behavior. Its most famous NEET application is explaining why O2 is paramagnetic, which Lewis theory and simple VBT fail to explain.

   • Molecular orbitals belong to the whole molecule, not to a single atom.
   • Only atomic orbitals of comparable energy and proper symmetry combine effectively.
   • Filling order differs for lighter and heavier second-period diatomic molecules.
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7. 7Hydrogen Bonding & Metallic Bonding

   Hydrogen bonding is a strong dipole-dipole interaction that occurs when hydrogen is bonded to a highly electronegative and small atom such as fluorine, oxygen or nitrogen, and is attracted to a lone pair on another electronegative atom. It may be intermolecular, as in water and HF, or intramolecular, as in o-nitrophenol. Hydrogen bonding explains high boiling points of H2O, HF and NH3 compared with similar hydrides, the open structure of ice, solubility of alcohols and biological structures such as proteins and DNA. Metallic bonding is explained by metal cations immersed in a sea of delocalized electrons. This model explains electrical conductivity, thermal conductivity, malleability, ductility and metallic lustre.

   • Hydrogen bonding is stronger than ordinary dipole-dipole forces but weaker than covalent or ionic bonds.
   • F, O and N are effective because they are small and highly electronegative.
   • Intermolecular hydrogen bonding increases molecular association, viscosity and boiling point.
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