Molecular Geometry and Bonding Theories

Physical and chemical properties of a molecule aredetermined by:

size and shapestrength 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 linesjoining the nuclei of the atoms in the molecule-must 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, lookat number of electron pairs attached to central atomand don’t distinguish between bonding electrons andlone 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 GeometryPairs Pairs

2 0 Linear

3 0 TrigonalPlanar

2 1 Angular

4 0 Tetrahedral 3 1 Trigonal

Pyramidal 2 2 Angular

5 0 TrigonalBipyramidal

4 1 Seesaw3 2 T-shaped2 3 Linear

6 0 Octahedral5 1 Square

Pyramidal4 2 Square

Planar

Determine the electron pair geometry and molecular geometry for the following:

SF4

IF5

ClF3

CO32-

H2S

http://www.dcu.ie/~pratta/jmgallery/JGALLERY.HTM

Non-bonding electrons exert greater repulsive forceson 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 aroundeach 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 individualbonds are polar, look at the symmetry of the molecule,if symmetric, non-polar, if asymmetric, polar. Lonepairs 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/vsepr/

• 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/essentialchemistry/flash/hybrv18.swf

Hybridization tutorial

How does valence bond theory explain molecules likeBeF2?

Be = 1s22s2 F = 1s22s22p5

In order for Be to bond with 2 F atoms, hybridizationor mixing of the orbitals occurs.

Hybrid orbitals require energy but they can overlapmore strongly resulting in a stronger bond. The energyreleased offsets the energy expended in the formationof the hybrid orbital.

Link electron pair geometry with hybridization:

linear sptrigonal planar sp2

tetrahedral sp3

trigonal bipyramidal sp3doctahedral sp3d2

Predict the hybridization of the following:

NH2- SF4 SO3

2- 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 facingeach other.

All single bonds are bonds.

bonds are formed by the side-ways overlap of porbitals. bonds are oriented perpendicular to theinternuclear axis.

Because there is less overlap in a bonds, generallythese bonds are weaker than bonds.

When multiple bonds are formed (such as double andtriple 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 areassociated with the two atoms forming the bond.

A molecule that does not have localized electrons isbenzene, C6H6. Benzene has two resonance structures resulting in 6 bonds of equal length. The3 bonds that form are said to be delocalized amongthe 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 orbitalscombine to form 2 molecular orbitals.

Orbital 1: Concentrates electron density between the twohydrogen 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 littleelectron density between nuclei.Destructive interference, atomic orbital canceleach other.Higher in energyDesignated 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 covalentbond.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 andbond order for the following:

C2 O2 F2 Ne2

Paramagnetism: substances that have one or moreunpaired electrons are attracted into a magneticfield and are said to be paramagnetic.

Diamagnetism: substances with no unpaired electronsare weakly repelled from a magnetic field and are saidto be diamagnetic.

Determine whether the above molecules are paramagnetic or diamagnetic.