Chapter 9: Molecular Geometry and Bonding Theories

Chapter 9: Molecular Geometry and Bonding Theories
9.1 Molecular Shapes

Lewis structures do not indicate shape of molecule
bond angles – angles made by the lines joining the nuclei of the atoms in the molecule; determines shape of molecule along with bond length
AB_n molecule, A is central atom bonded to n B atoms

9.2 The VSEPR Model

best arrangement of a given number of electron pairs is one that minimizes the repulsion among them

9.2.1 Predicting Molecular Geometries

two types of valence shell electron pairs 1)bonding pairs and 2)nonbonding pairs
electron-pair geometry – arrangement of electron pairs about the central atom of an AB_n
molecular geometry – arrangement of atoms in space
when describing shapes give molecular geometry rather than electron-pair geometry
steps to predict molecular geometries with VSEPR
1) sketch Lewis structure of the molecule or ion
2) count total number of electron pairs around central atom and arrange to minimize electron-pair repulsion
3) describe molecular geometry in terms of angular arrangement of the bonding pairs
4) double or triple bond counts as one bonding pair

9.2.2 The Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles

bond angles decrease as number of nonbonding electron pairs increase
nonbonding electron pairs exert greater repulsive forces on adjacent electron pairs ® compress angles between bonding pairs
electrons in multiple bonds exert a greater repulsive force on adjacent electron pairs than do single bonds

9.2.3 Molecules with Expanded Valence Shells

most stable electron-pair geometry for five electron pairs is trigonal bipyramid
two geometrically distinct electron pairs – axial pairs and equatorial pairs
equatorial pairs feel less repulsion than axial pairs
nonbonding pairs always equatorial
most stable electron-pair geometry for six electron pairs is octahedron
all angles are 90 or 190 degrees

9.2.4 Molecules with More than One Central Atom
9.3 Polarity of Molecules

polar if centers of negative and positive charge do not coincide
"d +" and "d -" used to indicate partial positive and negative charges
or an arrow that shows a shift in electron density
polar molecules align themselves in an electric field
dipole – two electrical charges of equal magnitude but opposite sign are separated by a distance
dipole moment – size of dipole, m

if equal magnitude, Q+ and Q-, separated by a distance r then dipole moment is product of Q and r
m = Qr
unit = debyes (D) = 3.34x10^-30 columb-meters (C-m)
charge of molecules measured in units of electronic charge e, 1.60x10^-19 C. and distance Å
dipole moments provide information about charge distributions in molecules

9.3.1Dipole Moments of Polyatomic Molecules

bond dipole – dipole moment due only to the two atoms bonded together
bond dipoles and dipole moments are vector quantities
overall dipole moment of a polyatomic molecule is sum of its bond dipoles

9.4 Covalent Bonding and Orbital Overlap

VSEPR does not explain why bonds exist between atoms
Valence-bond theory – extension of Lewis’s notion of electron-pair bonds
In Lewis theory, covalent bonding occurs when atoms share electrons
Valence-bond theory, covalent bonding occurs when valence atomic orbital of one atom merges with that of another atom

Orbitals are said to overlap

as distance between the atoms decreases, the overlap between their 1s orbitals increases
increase in electron density à decrease in potential energy of system
strength of bond increases ® decrease in energy
as atoms come close together ® energy increases rapidly

increase due to electrostatic repulsion between nuclei
observed bond length is the distance at which the attractive forces between unlike charges (electrons and nuclei) balanced by repulsive forces of like charges (electron-electron and nucleus-nucleus)

9.5 Hybrid Orbitals

9.5.1 sp Hybrid Orbitals

hybrid orbitals – orbitals formed by mixing two or more atomic orbitals on an atom
formed by hybridization
promotes one s electron to the p orbital

9.5.2 sp² and sp³ Hybrid Orbitals

mixing one 2s and 2p orbitals yields two equivalent sp hybrid orbitals that point in opposite directions
s orbital can mix with all three p orbitals which form sp³ hybrid orbitals
each sp³ hybrid orbital has a large lobe that points toward a vertex of a tetrahedron

9.5.3 Hybridization Involving d Orbitals

mix one s orbital, three p orbitals, and one d orbital ® five sp³d orbitals
mix one s orbital, three p orbitals, and two d orbitals ® siz sp³d² hybrid orbitals (directed toward vertices of an octahedron
corresponds to notion of expanded valence shells

9.5.4 Summary

predicting hybrid orbitals:
1) draw Lewis structure for molecule or ion
2) determine the electron-pair geometry using the VSEPR model
3) specify hybrid orbitals needed to accommodate electron pairs based on their geometrical arrangement

9.6 Multiple Bonds

internuclear axis – line joining nuclei
sigma bonds (s ) – overlap of two s orbitals, s and p orbital, two p orbitals, p orbital with sp hybrid
pi bond – overlap of two p orbitals oriented perpendicularly to the internuclear axis
covalent bond in which the overlap regions lie above and below the internuclear axis
no probability of finding electron on internuclear axis
pi bonds weaker than sigma bonds
single bonds are sigma bonds, double bond consists of one sigma bond and one pi bond, triple bond consists of one sigma bond and two pi bonds
cannot experimentally observe a pi bond directly

pi bonds create rigidity in molecules

double and triple bonds are more common in molecules with small atoms

9.6.1 Delocalized pi bonding

localized bonding – sigma and pi electrons are associated totally with the two atoms that form the bond
delocalized molecules have pi bonds and more than one resonance structure

9.6.2 General Conclusions

1) in every bond at least one pair of electrons is localized in space between atoms in a sigma bond. Appropriate set of hybrid orbitals used to form sigma bonds between atom and neighbors determined by observed geometry of the molecule
2) electrons in sigma bonds are localized in region between two bonded atoms and do not make a significant contribution to the bonding between any other atom
3) when atoms share more than one pair of electrons, additional pairs are pi bonds
4) molecules with two or more resonance structures can have pi bonds that extend over more than two bonded atoms

9.7 Molecular Orbitals

molecular orbital theory – another way to describe bonding in molecules
electrons allowed in energy states called molecular orbitals

9.7.1 the Hydrogen Molecule

when two atomic orbitals overlap, two molecular orbitals form
bonding molecular orbital – the lower energy orbital concentrates charge density in the region between the nuclei
antibonding molecular orbital – excludes electrons form the region between the nuclei
atomic orbitals cancel each other in the antibonding molecular orbital
excludes electrons from region that bonds have to be formed
electron repelled from boning region
sigma molecular orbitals – orbitals formed from the combination of bonding and antibonding molecular orbitals
energy-level diagram (molecular orbital diagram) – shows the interaction between two 1s orbitals to form s _1s and s _1s* molecular orbitals

9.7.2 Bond Order

an bond order of 1 = a single bond
bond order of 2 = double bond
bond order of 3 = triple bond
½, 3/2 or 5/2 also possible
bond order of 0 = no bond

9.8 Second-Row Diatomic Molecules

homonuclear diatomic molecules – composed of two identical atoms
1) number of molecular orbitals formed = number of atomic orbitals combined
2) atomic orbitals combine most effectively with other atomic orbitals of similar energy
3) effectiveness of combining atomic orbitals proportional to overlap with one another
4) each molecular orbital can have up to 2 electrons (pauli exclusion principle)
5) Hund’s Rule

9.8.1 Molecular Orbitals for Li₂ and Be₂

core electrons usually do not contribute significantly to bonding in molecule formation
in Be₂ there is an equal number of bonding and antibonding electrons so bond order equals 0 therefore Be₂does not exist
Li₂ has a bond order of 1 = 1 single bond

9.8.2 Molecular Orbitals from 2p Atomic Orbitals

p orbitals that are perpendicular to the internuclear axis form pi molecular orbitals
head to head 2p orbitals form s _2p and s _2p* orbitals
overlap sideways forms p _2p* and p _2p orbitals
s _2p lower in energy than p _2p
s _2p* higher energy than p _2p*

9.8.3 Electron Configurations for B₂ Through Ne₂

1) 2s atomic orbitals lower than 2p atomic orbitals
2) s _2p lower in energy than p _2p and s _2p* higher energy than p _2p*
3) both p _2p and p _2p* molecular orbitals are doubly degenerate, there are two degenerate molecular orbitals of each type
For B₂, C₂, and N₂ the s _2p molecular orbital is above the p _2p molecular orbitals in energy. For O₂, F₂, and Ne₂ the s _2p molecular orbital below the p _2p molecular orbitals

9.8.4 Electron Configurations and Molecular Properties

paramagnetism – attraction of unpaired electrons in a magnetic field
diamagnetism – substances with no unpaired electrons are weakly repelled in magnetic field
weaker than paramagnetism
test for paramagnetism and diamagnetism is to weigh substance without presence of magnetic field and then with a magnetic field
substance will appear to weigh more in paramagnetic in magnetic field, or would appear to weigh less if diamagnetic