Study Guides
Big Picture
Lewis structures are two-dimensional diagrams, but molecules are actually three-dimensional. The VSEPR (Valence
Shell Electron Pair Repulsion) theory allows us to visualize a molecule’s shape and the relative locations of a molecule’s
electrons. The size, shape, and orientation of these electron orbitals play an important role in an atom’s bonding
patterns. Orbital hybridization and molecular orbitals explain the nature of these orbitals.
Key Terms
Steric Number: The number of atoms bonded to the central atom plus the number of lone pairs.
Chemistry
VSEPR Theory & Hybridization
Hybridization: When two or more orbitals “mix” to form two or more orbitals of an “average” size and energy.
sp-hybridized orbital: One of two orbitals created by hybridizing one s orbital with one p orbitals, producing two
orbitals at 180° degrees apart. Atoms forming two double bonds will hybridize in this fashion, leaving two unhybridized p orbitals.
sp -hybridized orbital: One of three orbitals created by hybridizing one s orbital with two p orbitals, producing
2
3 orbitals at 120° degrees apart. Atoms forming one double bond will hybridize in this fashion (leaving one
un-hybridized p orbital).
sp3-hybridized orbital: One of four orbitals created by hybridizing one s orbital with three p orbitals, producing 4
orbitals at 109.5° degrees apart. Most atoms forming only single bonds will form this type of orbital.
sp3d-hybridized orbital: One of five orbitals created by hybridizing one s orbital, 3 p orbitals, and 1 d orbital,
producing three orbitals 120° degrees apart and two more orbitals 180° degrees apart on the vertical.
sp3d2-hybridized orbital: One of six orbitals created by hybridizing one s orbital, 3 p orbitals, and 2 d orbitals,
creating 6 orbitals 90° degrees apart.
Sigma Bond: A chemical bond involving s orbitals, p orbitals, or hybridized orbitals.
Pi Bond: A chemical bond involving un-hybridized p orbitals.
Molecular Geometry
The VSEPR theory helps to predict the molecular geometry. In VSEPR theory, the valence-electron pairs move as far
away from each other as possible in order to minimize repulsion between electron pairs (see table below).
• The geometry is determined by the steric number.
• A lone pair repels more strongly than paired electrons.
sp
sp2
sp3d
sp3d2
This guide was created by Steven Lai, Rory Runser, and Jin Yu. To learn more
about the student authors, visit http://www.ck12.org/about/about-us/team/
interns.
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v1.1.12.2012
Disclaimer: this study guide was not created to replace
your textbook and is for classroom or individual use only.
Image Credit: CK-12 Foundation, CC-BY-NC-SA 3.0
sp3
Chemistry
VSEPR Theory & Hybridization
cont .
Hybridization
Hybridization explains the formation of
bonds as the overlap of atomic orbitals and
is a central part of the valence bond theory.
Hybridized orbitals are formed by mixing
atomic orbitals.
• The number of hybridized orbitals equals
the number of atomic orbitals used to
form the hybridized orbitals.
• Each
hybridized orbital contains two
electrons.
Types of Hybridized Orbitals
Orbitals Hybridized
Hybridized
Orbital Names
Number of Hybridized
Orbitals
s+p
sp
2
s+p+p
2
sp
3
s+p+p+p
sp3
4
s+p+p+p+d
3
sp d
5
s+p+p+p+d+d
sp3d2
6
Consider the molecule CH4 with the electron configuration 1s22s22p2. The carbon atom forms four bonds, but in the
ground state electron configuration, carbon only has two half-filled orbitals (the 2p orbitals). In order to form four
bonds, one of the electrons in the 2s orbital must be promoted to the empty 2p orbital, creating four half-filled orbitals.
Image Credit: CK-12 Foundation, CC-BY-NC-SA 3.0
If the four hydrogen atoms bonded to the atomic orbitals as is, then one of the bonds would be different from the
others. In reality, all the bonds are identical. Hybridization explains this observation by creating four identical sp3
orbitals. This is beneficial for an atom because it lowers the energy state of the higher-energy orbitals and makes all
orbitals have an equal energy level, thus assisting bonding patterns (each orbital can make one chemical bond).
Sigma and Pi Bonds
Hybridized and unhybridized orbitals can then combine to form molecular orbitals that belong to the entire molecule.
The two types of molecular orbitals are sigma bonds and pi bonds.
Figure: Sigma and pi bonds
Image Credit: Rory Runser, CC-BY-SA 3.0
• Each molecular orbital still holds only two electrons.
• Sigma (σ) bonds form from the head-on-head overlap of s, p, or hybridized orbitals. All covalently bonded molecules
contain sigma bonds.
• Pi (�) bonds form from the side-by-side overlap of unhybridized p orbitals.
• Double bonds have 1 sigma and 1 pi bond, while triple bonds have 1 sigma and 2 pi bonds.
• Sigma bonds are stronger than pi bonds.
Polar Molecules
The shape of the molecule can be used to determine if the molecule is polar. To see if a molecule is polar or nonpolar,
1.Draw the molecular geometry
2.Label the bonds with arrows that point toward the more electronegative atom
3.If all the arrows cancel out, then the molecule is nonpolar. Otherwise, the molecule is polar.
Imagine the arrows are pushing on the molecule. If the overall effect of all the arrows cause the molecule to
move in one direction, the molecule is polar. If the molecule doesn’t move, the molecule is nonpolar.
Notes
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Common Problems
• Determining the shape of a molecule.
• Draw the Lewis structure for the molecule.
• Determine the number of electron pairs (bonded
• Determining the hybridization state of a molecule.
• Draw the Lewis structure for the molecule.
• Find the number of electron pairs, and match it
and unbonded) around the central atom.
with the correct hybridization state
• Use
the table (or your memory) to find the
correct shape.
• Determining the polarity of a molecule.
Drawing VSEPR Structures
Basic VSEPR
Example: Determine the VSEPR Structure of ammonia (NH3).
1.Draw the Lewis structure for the molecule.
2.Determine the number of electron pairs (bonded and unbonded) around the central atom.
3.Use the table (or your memory) to find the correct shape.
Figure: (left) Ammonia Lewis structure (right) Ammonia VSEPR shape. The top orbital has a
pair of electrons that fill the orbital.
Ammonia has 4 orbitals, one of which is filled. Thus, it is trigonal pyramidal.
Advanced VSEPR
Example: Determine the VSEPR structure of sulfur tetrafluoride (SF4).
1.Count up the number of valence electrons in the entire molecule
• Sulfur = 6, Fluorine = 7
• 6 + 4 · 7 = 34
2.Determine how many electron pairs are necessary for the non-central atom. Since sulfur is the central atom, each
fluorine will have 6 non-bonding electrons (3 electron pairs). Deduct these from the electron count.
• 34 – 6 · 4 = 34 – 25 = 10
3.There are now 10 electrons associated with the central atom, fluorine. Four bonds will form, one between the sulfur
and each fluorine. Deduct 2 electrons for each bond formed.
• 10 – 2 · 4 = 10 – 8 =2
4.There are thus 2 non-bonding electrons on the sulfur. Using the VSEPR Chart, 1 electron pair and 4 bonds means
this is seesaw-shaped.
Figure: VSEPR structure of the seesaw-shaped sulfur tetrafluoride.
Hybridization & VSEPR
Example: What is the hybridization state
of xenon in XeF4?
• Follow
the steps for determining the
VSEPR structure of an atom.
• Xenon
is square planar with 2 filled
orbitals.
• Look up on the chart which hybridization
this corresponds to. This means xenon
has a hybridization of sp3d2.
Polarity & VSEPR
Example: Is OF2 polar?
1.Draw the VSEPR structure of the molecule following the steps
above.
2.Determine the electronegativities of each atom.
Fluorine = 4, oxygen = 3.4. Draw dipoles along the bonds, with
the positive ends at the atom with the lower electronegativity.
3.If the arrows cancel each other out, the molecule is not polar. In
this case, OF2 is polar.
Figure: Polarity of oxygen difluoride
Page 3 of 3
Chemistry
VSEPR Problem Guide