Valence Bond Theory Flashcards

Definitions and Characteristics of Scientific Theories:
- A scientific theory is defined as a strongly supported explanation for observed natural laws or extensive bodies of experimental data.
- To be accepted, a theory must fulfill two primary criteria: it must explain experimental data and accurately predict future behavior.
VSEPR Theory vs. Valence Bond Theory:
- The Valence Shell Electron Pair Repulsion (VSEPR) theory is widely accepted due to its ability to predict three-dimensional molecular shapes consistent with data from thousands of molecules.
- Despite its success in predicting shape, VSEPR theory lacks an explanation for chemical bonding itself. L
- While quantum mechanics predicts regions where electrons are likely found around free atoms (e.g., spherical $s$ orbitals, dumbbell-shaped $p$ orbitals), these atomic orbitals insufficient for describing electron locations within molecules.
Covalent Bond Formation:
- A covalent bond is defined as the sharing of a pair of electrons by two atoms, where the pair is simultaneously attracted by the nuclei of both atoms.
- Valence Bond Theory and hybridization provide a more complete model for understanding these electron distributions in molecules.

Description of a Covalent Bond:
- Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals, where each orbital contains a single electron. This overlap yields a pair of electrons shared between the bonded atoms.
Conditions for Bond Formation:
1. An orbital on one atom must overlap with an orbital on a second atom (occupying the same I region of space).
2. The single electrons in each respective orbital must combine to form an electron pair.
Mechanism of Physical Linking:
- The two atoms are physically linked by the force of the mutual attraction between the negatively charged electron pair and the two positively charged nuclei.
Bond Strength and Extent of Overlap:
- The strength of a covalent bond is directly dependent on the extent to which the orbitals overlap.
- Extensive orbital overlap results in stronger bonds, whereas less overlap results in weaker bonds.

Energy States During Atomic Approach (Figure 8.2):
- Large Separation: When atoms are far apart, there is no orbital overlap and no interaction. By convention, the sum of the energies is set at $0$ .
- Initial Interaction: As atoms approach, orbitals begin to overlap. Electrons begin to feel the attraction of the other atom's nucleus. Simultaneously, repulsions occur between electrons and between the two nuclei.
- Decreasing Energy: While atoms are widely separated, the attractive forces are slightly l stronger than the repulsive forces, causing the system's energy to decrease and a bond to begin forming.
- Optimum Distance: The energy reaches its lowest, most stable value at a specific distance determined by the atoms involved. This is defined as the bond distance.
- Destabilization: If the nuclei move closer than the bond distance, repulsions between nuclei and between confined electrons exceed the attractive forces, causing the system's energy to rise sharply.
Quantitative Example: The $H_2$ Molecule:
- The minimum energy state for the $H-H$ bond occurs at a value of $-7.24 \times 10^{-19}$ † $J$ .
- The bond length (internuclear distance) for $H_2$ is $74\,pm$ .

Geometry of Overlap:
- Other than spherical $s$ orbitals, the orientation of orbitals significantly affects overlap.
- Overlap is maximized when orbitals are oriented along a direct line between the two nuclei (end-to-end overlap).
- For example, two $p$ orbitals have the greatest overlap when directed end-to-end. Any other arrangement or angle results in less overlap and therefore a weaker bond (Figure 8.3).

Sigma ( $\sigma$ ) Bonds:
- Definition: A covalent bond where electron density is concentrated in the region along the internuclear axis. A line passing through the nuclei also passes through the center of the overlap region.
- Formation Examples:
  - Overlap of two $s$ orbitals (e.g., $H_2$ ).
  - Overlap of one $s$ orbital and one $p$ orbital (e.g., $HCl$ ).
  - End-to-end overlap of two $p$ orbitals (e.g., $Cl_2$ ).
- Lewis Structure Correlation: All single bonds in Lewis structures are described as $\sigma$ bonds in valence bond theory.
Pi ( $\pi$ ) Bonds:
- Definition: A covalent bond resulting from the side-by-side overlap of two $p$ orbitals.
- Geometry: The regions of orbital overlap are located on opposite sides of the internuclear axis.
- The Node: Along the actual internuclear axis, there is a node, which is a plane where the probability of finding an electron is zero.

Sequence of Bond Formation:
- The first bond formed between any two atoms will always be a $\sigma$ bond.
- There can only be one $\sigma$ bond between any two atoms.
Bond Composition:
- Single Bond: Consists of one $\sigma$ bond.
- Double Bond: Consists of one $\sigma$ bond and one $\pi$ bond (e.g., Oxygen, $O_2$ ).
- Triple Bond: Consists of one $\sigma$ bond and two $\pi$ bonds (e.g., Nitrogen, $N_2$ ).

Bond Data (Bond Length in $pm$ , Energy in $kJ\,mol^{-1}$ ):
- $H-F$ : $91.7\,pm$ , $565\,kJ\,mol^{-1}$
- $H-N$ : $101.5\,pm$ , $391\,kJ\,mol^{-1}$
- $H-O$ : $97.5\,pm$ , $467\,kJ\,mol^{-1}$
- $H-Cl$ : $127.5\,pm$ , $431\,kJ\,mol^{-1}$
- $H-Br$ : $141.4\,pm$ , $366\,kJ\,mol^{-1}$
- $H-I$ : $160.9\,pm$ , $298\,kJ\,mol^{-1}$
- $C-C$ : $150.6\,pm$ , $347\,kJ\,mol^{-1}$
- $C=C$ : $133.5\,pm$ , $614\,kJ\,mol^{-1}$
- $C \equiv C$ : $120.8\,pm$ , $839\,kJ\,mol^{-1}$
- $C-N$ : $142.1\,pm$ , $305\,kJ\,mol^{-1}$
- $C=N$ : $130.0\,pm$ , $615\,kJ\,mol^{-1}$
- $C \equiv N$ : $116.1\,pm$ , $891\,kJ\,mol^{-1}$
- $C=O$ : $113.7\,pm$ , $1072\,kJ\,mol^{-1}$
- $O-O$ : $148\,pm$ , $146\,kJ\,mol^{-1}$
- $O=O$ : $120.8\,pm$ , $498\,kJ\,mol^{-1}$
- $F-F$ : $141.2\,pm$ , $159\,kJ\,mol^{-1}$
- $Cl-Cl$ : $198.8\,pm$ , $243\,kJ\,mol^{-1}$