VSEPR Theory VSEPR Theory Shapes of Molecules
Molecular Structure or Molecular Geometry The 3-dimensional arrangement of the atoms that make-up a molecule. Determines several properties of a substance, including: reactivity, polarity, phase of matter, color, magnetism, and biological activity. The chemical formula has no direct relationship with the shape of the molecule. VSEPR Theory Shapes of Molecules
Molecular Structure or Molecular Geometry The 3-dimensional shapes of molecules can be predicted by their Lewis structures. Valence-shell electron pair repulsion (VSEPR) model or electron domain (ED) model: Used in predicting the shapes. The electron pairs occupy a certain domain. They move as far apart as possible. Lone pairs occupy additional domains, contributing significantly to the repulsion and shape. VSEPR Theory Terms and Definitions
Bonding Pairs (AX) Electron pairs that are involved in the bonding. Lone Pairs (E) – aka non-bonding pairs or unshared pairs Electrons that are not involved in the bonding. They tend to occupy a larger domain. Electron Domains (ED) Total number of pairs found in the molecule that contribute to its shape. VSEPR – Molecular Shape
Multiple covalent bonds Bond Angle: around the same atom • Angle formed by any determine the shape two terminal (outside) Negative e- pairs (same atoms and a central charge) repel each other atom • Caused by the repulsion Repulsions push the pairs as far apart as possible of shared electron pairs. Hybridization
What’s a hybrid? • Combining two of the same type of object and contains characteristics of both • Occurs to orbitals during bonding Orbital hybridization • Process in which atomic orbitals are mixed to form new hybrid orbitals • Each hybrid orbital contains one electron that it can share with another atom Carbon is most common atom to undergo hybridization • Four hybrid orbitals from 1 s and 3 p orbitals • Hybrid = sp3 orbital Orbital Hybridization
Atomic orbitals such as s and p are not well suited for overlapping and allowing two atoms to share a pair of electrons The best location of shared pair is directly between two atoms e- pair spends little time in best location • With overlap of two s-orbital • With overlap of two p-orbitals Orbital Hybridization
Hybrid orbitals (cross of atomic orbitals) •Shape more suitable for bonding One large lobe and one very small lobe Large lobe oriented towards other nucleus •Angles more suitable for bonding Angles predicted from VSEPR Orbital Hybridization Overlap of two s-orbitals Note: shared in this overlap the e- pair would spend most of the time in an unfavorable location
GOOD SPOT between both nuclei
NOT A GOOD LOCATION- Too far from one nucleus Orbital Hybridization
Overlap of two p-orbitals
GOOD SPOT BAD location far from BAD location far from other nucleus between both other nucleus nuclei One atom & its The other atom & p-orbital its p-orbital
represents the nucleus Orbital Hybridization
Hybrid orbitals yield more favorable shape for overlap • Atomic orbitals are not shaped to maximize attractions nor minimize repulsions Hybrid orbital shape • One large lobe oriented towards other atom • Notice the difference in this shape compared to p-orbital shape Orbital Hybridization
Hybrid orbitals create more favorable angles for overlap, too. Atomic orbitals are not shaped to maximize attractions nor minimize repulsions BUT the angles are also not favorable p-orbitals are oriented at 90 to each other Other angles are required: 180, 120, or 109.5 Orbital Hybridization
- Each e pair requires a hybrid orbital If two hybrid orbitals required than two atomic orbitals must be hybridized, an s and a p orbital forming two sp orbitals at 180
sp hybrids sp2 hybrids sp3 hybrids
2 EP 3 EP 4 EP sp-Hybridization sp2 -Hybridization sp3 -Hybridization Hybridization – Key Points
The number of hybrid (molecular) orbitals obtained equals the number of atomic orbitals combined. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed. Examples: • 1 s + 1 p = 2 sp orbitals • 1 s + 2 p = 3 sp2 orbitals • 1 s + 3 p = 4 sp3 orbitals Electron-Pair Geometry vs Molecular Geometry
Electron-pair geometry • Where are the electron pairs • Includes bonding pairs (BP) = shared between 2 atoms nonbonding pairs (NBP) = lone pair Molecular geometry • Where are the atoms • Includes only the bonding pairs
2 Electron Domains (ED) around central atom
Two clouds pushed as far apart as possible • Greatest angle possible 180 • LINEAR shape
Linear
Bonding Pairs: 2
Lone Pairs: 0
Electron Domains: 2
Bond Angle: 180°
Example: CO2 Image: Linear
Carbon Dioxide (CO2)
Nitrogen Gas (N2) 3 Electron Domains (ED) around central atom
Three electron clouds pushed as far apart as possible • Greatest angle possible = 120 • TRIGONAL (3) PLANAR (flat) shape
Examples of 3 ED
3 Bonded Pairs + 0 Non-Bonded Pairs • 3 ED = Electron Pair Geometry is trigonal planar • All locations occupied by atoms, • So Molecular Geometry is also trigonal planar 2 Bonded Pairs + 1 Non-Bonded Pair • 3 ED = Electron Pair Geometry is trigonal planar • Only two bonding pairs • One of the locations is only lone pair of e- • So molecular geometry is bent Trigonal Planar
Bonding Pairs: 3
Lone Pairs: 0
Electron Domains: 3
Bond Angle: 120°
Example: BF3 Image: Trigonal Planar
2- Carbonate Ion (CO3 )
- Nitrate Ion (NO3 ) Bent or Angular
Bonding Pairs: 2
Lone Pairs: 1
Electron Domains: 3
Bond Angle: 120° (119°)
Example: SO2 Image: Bent or Angular
- Nitrite Ion (NO2 ) 4 Electron Domains (ED) around central atom
Four clouds pushed as far apart as possible • Greatest angle no longer possible in two dimensions • Requires three-dimensional • TETRAHEDRAL shape Examples of 4 ED
4 Bonded Pairs + 0 Non-Bonded Pairs • 4 ED: Both Electron Pair Geometry and Molecular Geometry are tetrahedral 3 Bonded Pairs + 1 Non-Bonded Pair • 4 ED: Electron Pair Geometry is tetrahedral Molecular Geometry is TRIGONAL PYRAMIDAL No atom at top location 2 Bonded Pairs + 2 Non-Bonded Pairs • 4 ED: Electron Pair Geometry is tetrahedral Molecular geometry is BENT No atoms at two locations Tetrahedral
Bonding Pairs: 4
Lone Pairs: 0
Electron Domains: 4
Bond Angle: 109.5°
Example: CH4 Image: Tetrahedral
Silicon Tetrachloride (SiCl4)
Methane (CH4) Trigonal Pyramidal
Bonding Pairs: 3
Lone Pairs: 1
Electron Domains: 4
Bond Angle: 109.5° (107.5°)
Example: NH3 Image: Trigonal Pyramidal
+ Hydronium Ion (H3O )
Ammonia (NH3) Bent or Angular (Ver. 2)
Bonding Pairs: 2
Lone Pairs: 2
Electron Domains: 4
Bond Angle: 109.5° (104.5°)
Example: H2O Image: Bent or Angular (Ver. 2)
Chlorine Difluoride (ClF2) Summary of 4 Electron Domain Shapes Exceptions to Octet Rule
Reduced Octet • H only forms one bond only one pair of e- • Be tends to only form two bonds only two pair of e- • B tends to only form three bonds only three pair of e- Expanded Octet • Empty d-orbitals can be used to accommodate extra e- • Elements in the third row and lower can expand • Up to 6 pairs of e- are possible Lewis Structures in Which the Central Atom Exceeds an Octet Summary: Molecular Geometry of Expanded Octets