VSEPR Theory

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

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

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 Theory Terms and Definitions

VSEPR – Molecular Shape

Multiple covalent bonds around the same atom determine the shape

Negative e- pairs (same charge) repel each other

Repulsions push the pairs as far apart as possible

Bond Angle:

• Angle formed by any two terminal (outside) atoms and a central atom

• Caused by the repulsion 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

Overlap of two s-orbitals

NOT A GOOD LOCATION- Too far from one nucleus

Note: shared in this overlap the e- pair would spend most of the time in an unfavorable location

GOOD SPOT

between both

nuclei

Orbital Hybridization

Overlap of two p-orbitals

One atom & its p-orbital

The other atom & its p-orbital

represents the nucleus

BAD location far from other nucleus

GOOD SPOT between both

nuclei

BAD location far from other 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

2 EP 4 EP 3 EP

sp2 hybrids sp3 hybrids

Orbital Hybridization

sp-Hybridization

sp2 -Hybridization

sp3 -Hybridization

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

Hybridization – Key Points

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

Nitrogen Gas (N2)

Carbon Dioxide (CO2)

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

Carbonate Ion (CO32-)

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

Methane (CH4)

Silicon Tetrachloride (SiCl4)

Trigonal Pyramidal

Bonding Pairs: 3

Lone Pairs: 1

Electron Domains: 4

Bond Angle: 109.5° (107.5°)

Example: NH3

Image:

Trigonal Pyramidal

Ammonia (NH3)

Hydronium Ion (H3O+)

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