Hybridization

Hybridization 

• Atom can take its atomic orbitals and hybridize them into molecular orbitals 

  • Used to make covalent bonds which are overlapping of such orbitals

• Hybridization is chosen specifically to give you the number of σ bonds and lone pairs needed

  • π bonds are not hybridized 

• Once atomic orbitals are hybridized, those orbitals no longer exist

• Involves end-to-end overlapping 

  • Hence why π bonds are not hybridized

• Only one bonded atom = not hybridized 

  • Uses its atomic orbital

Types of bonds

1. σ bonds

• Cylindrical symmetrical - symmetrical along bond 

• All single covalent bonds

• Freely rotatable

  • This is because the electron density is concentrated along the bond axis

2. π bonds 

• Symmetrical along the nodal axis

• Internuclear axis

• Two lobes

• “Locks” rotation, no rotation allowed 

  • would break the side-to-side overlap and thus the bond

σ bond formation 

• Overlap between two orbitals happens along the axis connecting the two nuclei

  • End-to-end overlap 

• Formed by different types of overlaps

  • s + s (e.g., in H₂).

  • s + p (e.g., in CH₃).

  • p + p (e.g., in O₂).

  • Hybrid orbitals (such as sp, sp², sp³) formed from the combination of s and p orbitals (e.g., in methane, CH₄).

π bond formation 

• Overlap of p orbitals happen above and below the internuclear axis

  • Side-to-side overlap

• Formed by 

  • P orbitals

  • Double and triple bonds

Molecular structure 

• Single bonds

  • 1 σ bonds

• Double bonds

  • 1 σ bond, 1 π bond

• Triple bonds

  • 1 σ bond, 2 π bonds

Orbitals 

Sp 

• Made from an s and a p orbital

• Linear 

Spp = sp²

• Made from an s and two p orbitals

• Trigonal planar

• Energy of sp² orbital is between that of s and p orbital and higher than an sp orbital 

Sppp = sp³

• Made from an s and three p orbitals 

• Tetrahedral 

Spppd = sp³d

• Made from an s, three p, and one d orbital

• Trigonal bipyramidal 

Spppdd = sp³d²

• Made from an s, three p, and two d orbitals 

• Octahedral