Transcript Molecular Geometry and Bonding Theories

Molecular Geometry and
Bonding Theories
Physical and chemical properties of a molecule are
determined by:
size and shape
strength and polarity of bonds
Lewis structures do not indicate shapes of molecules,
simply the number and types of bonds.
To translate Lewis structures into three dimensions,
bond angles- the angles made by lines
joining the nuclei of the atoms in the moleculemust be used.
Valence Shell Electron Pair
Repulsion Theory (VSEPR)
• Best arrangement of electron pairs is the
one that minimizes repulsions.
• Arrangement of electron pairs around a
central atom is called electron-pair
geometry. No distinction is made between
bonding and nonbonding electrons.
• Molecular geometry is the arrangement of
atoms in space; distinction between
bonding and nonbonding electrons.
How do electron pairs affect shape?
2 electron pairs = linear = 180o
3 e.p. = trigonal planar = 120o
4 e.p. = tetrahedral = 109.5o
5 e.p. = trigonal bipyramidal =
120o and 90o
6 e.p. = octahedral = 90o
In order to determine electron pair geometry, look
at number of electron pairs attached to central atom
and don’t distinguish between bonding electrons and
lone pairs.
Molecular geometry takes into account lone pairs which
influence the shape of the molecule. Multiple bonds
are treated the same as single bonds.
Bonding Non bonding
Pairs
Pairs
Geometry
2
0
Linear
3
0
2
1
Trigonal
Planar
Angular
4
3
0
1
2
2
Tetrahedral
Trigonal
Pyramidal
Angular
5
0
4
3
2
1
2
3
6
5
0
1
4
2
Trigonal
Bipyramidal
Seesaw
T-shaped
Linear
Octahedral
Square
Pyramidal
Square
Planar
Determine the electron pair geometry and
molecular geometry for the following:
SF4
IF5
ClF3
CO32H2S
http://www.dcu.ie/~pratta/jmgallery/JGALLERY.HTM
Non-bonding electrons exert greater repulsive forces
on adjacent pairs and compress angles.
CH4, NH3, H2O
all have a tetrahedral electron pair
geometry, but their molecular
geometry is dictated by the presence
of lone pairs on the central atom.
Bond angles are 109.5o, 107o, and
104.5o respectively.
In absence of central atom, predict geometry around
each atom of backbone.
Polar molecules
-degree of polarity measured by dipole moment
-dipole moment of molecule depends on the
polarities of the individual bonds and the
geometry of the molecule.
Once you have established whether the individual
bonds are polar, look at the symmetry of the molecule,
if symmetric, non-polar, if asymmetric, polar. Lone
pairs on the central atom = polar.
VSEPR Theory
• Provides simple means for predicting
shapes of molecules.
• Does not explain why bonds exist or form.
• http://www.chem.purdue.edu/gchelp/vsep
r/
• Molecular shapes
Valence Bond Theory
• Combines Lewis’s idea of electron pair
bonds with atomic orbitals.
• Atomic orbital of one atom merges with that
of another atom
• Orbitals share a region of space or overlap.
http://www.mhhe.com/physsci/chemistry/essential
chemistry/flash/hybrv18.swf
Hybridization tutorial
How does valence bond theory explain molecules like
BeF2?
Be = 1s22s2
F = 1s22s22p5
In order for Be to bond with 2 F atoms, hybridization
or mixing of the orbitals occurs.
Hybrid orbitals require energy but they can overlap
more strongly resulting in a stronger bond. The energy
released offsets the energy expended in the formation
of the hybrid orbital.
Link electron pair geometry with hybridization:
linear
trigonal planar
tetrahedral
trigonal bipyramidal
octahedral
sp
sp2
sp3
sp3d
sp3d2
Predict the hybridization of the following:
NH2-
SF4
SO32-
SF6
 bonds are formed by the overlap of two s orbitals.
Concentration is symmetric on internuclear axis.
 bonds can form by the overlap of two s orbitals,
an s and a p orbital or two p orbitals that are facing
each other.
All single bonds are  bonds.
 bonds are formed by the side-ways overlap of p
orbitals.  bonds are oriented perpendicular to the
internuclear axis.
Because there is less overlap in a  bonds, generally
these bonds are weaker than  bonds.
When multiple bonds are formed (such as double and
triple bonds), the first bond is a  bond and the
remaining bonds are  bonds.
Predict the hybridization and the number of  and 
bonds in formaldehyde: H2CO.
 and  bonds are considered localized: electrons are
associated with the two atoms forming the bond.
A molecule that does not have localized electrons is
benzene, C6H6. Benzene has two resonance
structures resulting in 6 bonds of equal length. The
3  bonds that form are said to be delocalized among
the 6 carbon atoms.
It is important to understand that wherever
resonance occurs with multiple bonds, the  bonds
that form will be considered delocalized.
General Conclusions
• Every pair of bonded atoms shares one or
more pairs of electrons. In every bond at
least one pair of electrons is localized.
• The electrons in a  bonds are localized.
• When atoms share more than one pair of
electrons, the additional pairs form 
bonds.
• Electrons in  bonds that extend over more
than two atoms are delocalized.
Molecular Orbital (MO) Theory
• Molecular orbitals form from a
combination of atomic orbitals.
• Contain a maximum of two electrons.
• Two atomic orbitals form two molecular
orbitals.
When 2 hydrogen atoms combine, the two 1s orbitals
combine to form 2 molecular orbitals.
Orbital 1:
Concentrates electron density between the two
hydrogen nuclei.
Considered constructive interference.
Orbital is lower in energy, more stable.
Designated a 1s bonding orbital.
Orbital 2:
Atomic orbitals combine and lead to very little
electron density between nuclei.
Destructive interference, atomic orbital cancel
each other.
Higher in energy
Designated 1s* antibonding orbital.
Order of filling molecular orbitals:
1s 1s* 2s 2s* 
2p

2p
2p* 2p*
Each  molecular orbital holds 2 electrons; each 
orbital holds 4 electrons.
Bond order is a measure of the stability of a covalent
bond.
B.O. = 1/2(number of bonding electrons - number of
antibonding electrons)
Bond order of 1 = single bond
2 = double bond
3 = triple bond
0 = molecule doesn’t exist
Determine the molecular orbital configurations and
bond order for the following:
C2
O2
F2
Ne2
Paramagnetism: substances that have one or more
unpaired electrons are attracted into a magnetic
field and are said to be paramagnetic.
Diamagnetism: substances with no unpaired electrons
are weakly repelled from a magnetic field and are said
to be diamagnetic.
Determine whether the above molecules are
paramagnetic or diamagnetic.