Chemical Bonding and Molecular Structure

• Kössel-Lewis approach to chemical bonding:

  • Kössel and Lewis provided a logical explanation of valence based on the inertness of noble gases
  • Lewis pictured the atom with a positively charged 'Kernel' and an outer shell that could hold a maximum of eight electrons
  • Atoms achieve a stable octet when linked by chemical bonds
  • Sodium and chlorine bond by electron transfer, while molecules like Cl2, H2, F2 bond by sharing a pair of electrons
  • Lewis symbols represent valence electrons in an atom, with the number of dots indicating the number of valence electrons
• Octet Rule:

  • Developed by Kössel and Lewis in 1916
  • Atoms combine by transferring or sharing valence electrons to achieve an octet in their valence shells
• Covalent Bond:

  • Langmuir refined Lewis postulations and introduced the term covalent bond
  • Covalent bonds involve sharing electron pairs between atoms to achieve noble gas configurations
  • Lewis dot structures represent covalent bonds, with each bond formed by sharing an electron pair
• In molecules where atoms share three electron pairs, like in N2 and C2H2, a triple bond is formed
• Lewis dot structures provide a picture of bonding in molecules and ions in terms of shared pairs of electrons and the octet rule
• Steps for writing Lewis dot structures:

  • Obtain the total number of electrons required by adding the valence electrons of the combining atoms
  • For anions, each negative charge means the addition of one electron; for cations, each positive charge results in the subtraction of one electron
  • Distribute the total number of electrons as bonding shared pairs between the atoms
  • The least electronegative atom usually occupies the central position in the molecule/ion
  • Ensure each bonded atom gets an octet of electrons
• In Lewis structures, the formal charge of an atom is the difference between the number of valence electrons in the free atom and the electrons assigned in the Lewis structure
• Limitations of the Octet Rule:

  • Incomplete octet of the central atom: Elements with less than four valence electrons may have fewer than eight electrons surrounding the central atom
  • Odd-electron molecules: Molecules with an odd number of electrons may not satisfy the octet rule for all atoms
  • Expanded octet: Elements beyond the third period may have more than eight valence electrons around the central atom due to the availability of 3d orbitals for bonding
• Formation of ionic compounds depends on the ease of formation of positive and negative ions from neutral atoms and the arrangement of ions in the solid lattice
• In the gas phase, when an atom in its ground state gains an electron, the electron gain process may be exothermic or endothermic
• Ionization is always endothermic
• Electron affinity is the negative of the energy change accompanying electron gain
• Ionic bonds are formed more easily between elements with low ionization enthalpies and elements with high negative values of electron gain enthalpy
• Most ionic compounds have cations derived from metallic elements and anions from non-metallic elements
• Ionic compounds in the crystalline state consist of orderly three-dimensional arrangements of cations and anions held together by coulombic interaction energies
• The lattice enthalpy of an ionic solid is the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions
• Bond length is the equilibrium distance between the nuclei of two bonded atoms in a molecule
• Bond angle is the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion
• Bond enthalpy is the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state
• Bond order is given by the number of bonds between two atoms in a molecule
• Resonance structures are used when a single Lewis structure is inadequate for the representation of a molecule
• Resonance is introduced to deal with difficulties in accurately depicting structures of molecules like O3
• According to the concept of resonance:

  • When a single Lewis structure cannot describe a molecule accurately, multiple structures with similar energy, positions of nuclei, bonding, and non-bonding pairs of electrons are considered as canonical structures of the hybrid
  • For O3, the canonical structures or resonance structures are shown, with the hybrid (III structure) representing the molecule more accurately
• Resonance in the CO3

  2– ion:

  • The single Lewis structure with two single bonds and one double bond is inadequate as it represents unequal bonds
  • Experimental findings show that all carbon to oxygen bonds in CO3

  2– are equivalent, making it best described as a resonance hybrid of canonical forms I, II, and III
• Resonance in the CO2 molecule:

  • Experimentally determined carbon to oxygen bond length in CO2 is 115 pm, lying between the values for a carbon to oxygen double bond (121 pm) and a carbon to oxygen triple bond (110 pm)
  • The structure of CO2 is best described as a hybrid of canonical or resonance forms I, II, and III
• Polarity of Bonds:

  • No bond or compound is completely covalent or ionic in reality
  • Covalent bonds between similar atoms result in nonpolar covalent bonds, while bonds between different atoms can lead to polar covalent bonds
• In general, regarding resonance:

  • Resonance stabilizes the molecule as the energy of the resonance hybrid is lower than any single canonical structure
  • Resonance averages the bond characteristics as a whole
• The VSEPR theory explains the shape of molecules based on the repulsive interactions of electron pairs in the valence shell of atoms
• Main postulates of VSEPR theory:

  • The shape of a molecule depends on the number of valence shell electron pairs around the central atom
  • Electron pairs in the valence shell repel each other due to their negatively charged electron clouds
  • Electron pairs tend to occupy positions in space that minimize repulsion and maximize distance between them
  • Valence shell is considered a sphere with electron pairs localizing on the spherical surface at maximum distance from each other
  • A multiple bond is treated as a single electron pair, and multiple electron pairs of a bond are treated as a single super pair
• In VSEPR theory, lone pairs of electrons occupy more space compared to bonding pairs, leading to greater repulsion between lone pairs
• VSEPR theory categorizes molecules into two types:

  • Molecules with a central atom having no lone pair
  • Molecules with a central atom having one or more lone pairs
• Table 4.6 shows the arrangement of electron pairs around a central atom without any lone pairs and the geometries of some molecules/ions of type AB
• Table 4.7 displays the shapes of some simple molecules and ions where the central atom has one or more lone pairs of electrons
• Table 4.8 explains the reasons for distortions in the geometry of molecules containing bond pairs and lone pairs
• When a bond is formed between two hydrogen atoms, the resulting hydrogen molecule is more stable than isolated hydrogen atoms
• The energy released in the formation of a bond is called bond enthalpy, corresponding to the minimum in the potential energy curve
• 435.8 kJ of energy is required to dissociate one mole of an H2 molecule
• In the formation of a hydrogen molecule, there is a minimum energy state when two hydrogen atoms are so near that their atomic orbitals undergo partial interpenetration, known as overlapping of atomic orbitals
• The strength of a covalent bond is determined by the extent of overlap between atomic orbitals
• Covalent bonds can be classified into sigma (σ) bonds and pi (π) bonds based on the types of overlapping