explains how atoms form covalent bonds through orbital . It's all about shared electrons and the strength of their connections. This theory helps us understand why molecules have specific shapes and bond strengths.
Sigma and pi bonds are the building blocks of molecular structures. Sigma bonds form through head-on orbital overlap, while pi bonds result from sideways overlap. These concepts are crucial for grasping how atoms join to create complex molecules.
Valence Bond Theory
Atomic orbital overlap in covalent bonds
Top images from around the web for Atomic orbital overlap in covalent bonds
Single Covalent Bonds | Introduction to Chemistry View original
Covalent bonds form when atomic orbitals overlap enabling electrons to be shared between atoms
Orbitals must have matching symmetry and orientation to overlap effectively (s orbitals, p orbitals)
Overlapping orbitals must also have comparable energies to form stable bonds
Stronger covalent bonds result from greater overlap between orbitals
Orbital overlap depends on the size and shape of the involved orbitals (s, p, d, f)
Shared electrons in the overlapping region are attracted to both atomic nuclei lowering the overall energy and stabilizing the bond
cloud forms in the region of orbital overlap where the shared electrons are typically located
The extent of orbital overlap influences the and overall structure of the molecule
Sigma vs pi bonds
Sigma (σ) bonds
Formed by head-on overlap of atomic orbitals (end-to-end)
Can involve overlap of s orbitals, p orbitals, or a combination of s and p orbitals (sp, sp², sp³ hybrid orbitals)
Electron density concentrated along the between the bonded atoms
Stronger than pi bonds due to more extensive orbital overlap
Pi (π) bonds
Formed by sideways overlap of p orbitals (parallel)
Require unhybridized p orbitals perpendicular to the internuclear axis
Electron density concentrated above and below the internuclear axis
Weaker than sigma bonds due to less orbital overlap
Multiple bonds between atoms consist of one and one or more pi bonds
: one sigma bond and one (C=C in )
: one sigma bond and two pi bonds (C≡C in )
The number and type of bonds between atoms determine the , which affects bond strength and length
Bond energy and orbital interactions
is the energy required to break a chemical bond
Stronger bonds have higher bond energies (H-H 436 kJ/mol, C-C 348 kJ/mol, C=C 614 kJ/mol, C≡C 839 kJ/mol)
Higher bond energy and stronger bonds result from greater orbital overlap
Orbital overlap depends on the size, shape, and orientation of the involved orbitals
As the distance between bonded atoms increases, orbital overlap decreases leading to lower bond energy and a weaker bond
Optimal bond length occurs when attractive and repulsive forces between atoms are balanced maximizing orbital overlap and bond energy
Bond energy is also influenced by the difference between bonded atoms
Polar covalent bonds have slightly lower bond energies than pure covalent bonds due to larger differences (H-Cl vs H-H)
The bond energy of a molecule can be estimated by summing the average bond energies of all the bonds present
D0(ABn)=n1∑D0(A−B), where D0(ABn) is the molecule's bond energy, D0(A−B) is each A-B bond's average bond energy, and n is the number of A-B bonds in the molecule
Advanced concepts in bonding
: the mixing of atomic orbitals to form new hybrid orbitals with different shapes and energies
: predicts molecular geometry based on electron pair repulsion
Resonance: occurs when multiple valid Lewis structures can be drawn for a molecule, resulting in a hybrid structure that better represents the actual electron distribution