Download VSEPR Theory - Crestwood Local Schools

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
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
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 LOCATIONToo far from one nucleus
Orbital Hybridization
Overlap of two p-orbitals
GOOD SPOT
between both
nuclei
BAD location far from
other nucleus
One atom & its
p-orbital
represents the nucleus
BAD location far from
other nucleus
The other atom &
its p-orbital
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
sp2
hybrids
3 EP
sp3 hybrids
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:
 Lone Pairs:
 Electron Domains:
 Bond Angle:
 Example:
Image:
2
0
2

180°
CO2
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:
 Lone Pairs:
 Electron Domains:
 Bond Angle:
 Example:
Image:

3
0
3
120°
BF3
Trigonal Planar
Carbonate Ion (CO32-)
Nitrate Ion (NO3-)
Bent or Angular
Bonding Pairs:
 Lone Pairs:
 Electron Domains:
 Bond Angle:
 Example:
Image:

2
1
3
120° (119°)
SO2
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:
 Lone Pairs:
 Electron Domains:
 Bond Angle:
 Example:
Image:

4
0
4
109.5°
CH4
Tetrahedral
Silicon Tetrachloride (SiCl4)
Methane (CH4)
Trigonal Pyramidal
Bonding Pairs:
 Lone Pairs:
 Electron Domains:
 Bond Angle:
 Example:
Image:

3
1
4
109.5° (107.5°)
NH3
Trigonal Pyramidal
Hydronium Ion (H3O+)
Ammonia (NH3)
Bent or Angular (Ver. 2)
Bonding Pairs:
 Lone Pairs:
 Electron Domains:
 Bond Angle:
 Example:
Image:

2
2
4
109.5° (104.5°)
H2O
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