Section 6.5 – Molecular Geometry

The properties of molecules depend on the bonding and the molecular geometry, the 3-dimensional arrangement of the atoms in space.

Molecular Polarity

This is the uneven distribution of molecular charge, and it is determined by the polarity of each bond, along with the geometry of the molecule.

Two Theories – based on evidence

VSEPR Theory: Accounts for molecular bond angles.

Hybridization: Describes the orbitals that contain the valence electrons of a molecule’s atoms.

VSEPR Theory ValenceShellElectronPairRepulsionRepulsion between the

valence-shell electrons surrounding an atom causes these sets to be oriented as far apart as possible.

CH4 NH3 H2O

• Draw the Lewis structures for all 3 molecules• Sets of electrons that remain together in

bonds or in lone pairs are referred to as electron domains. Electron domains prefer to be as far apart as possible from each other within a molecule.

• How many electron domains are located around the central atom of each molecule?

Activity

CH4 NH3 H2O

• Collect supplies – 1 set of gumdrops and toothpicks for each pair of students

• Build a 3-D model of each molecule – SHOW ME YOUR MODEL!!!!

• Remember – electron domains want to be as far apart as possible.

• How do lone pairs affect the shape of each molecule?

Activity

Diatomic MoleculesDiatomic molecules are

composed of two atoms, so the geometry is always linear

the molecular polarity is determined by the electronegativity differences between the atoms. H2: non-polar HCl: polar

Shorthand for DescribingFor molecules containing more

than two atoms, we can use the following symbols with subscripts:A – the central atomB – number of bonds

on the central atomE – number of lone

pairs on central atom(for atoms that have double or

triple bonds, it is treated as a single B for geometry)

The Basis For VSEPR Theory

That one must consider the locations of all electron pairs of the valence electrons in the molecule.

Polyatomic ions are treated the same way.

The Basis For VSEPR Theory

The following examples do not have lone pairs that influence the geometry of the molecule.

Linear – AB2

Central atom with two single bonds, no lone pairs. Because the valence electron pairs in the bonds repel each other, the bonds are as far apart as possible (180°).

Ex: BeH2

Trigonal Planar – AB3

The 3 A-B bonds stay furthest apart by pointing to the corners of an equilateral triangle, giving 120° angles between the bonds. Ex.: BH3

Tetrahedral – AB4

Octet rule is followed here. The distance between the A-B bonds is maximized if each bond points to the corners of a tetrahedron, giving bond angles of 109.5° between the bonds.

Ex.: CH4

Trigonal-bipyramidal – AB5

120° angles between bonds within the trigonal plane, 90° bond angles between the axial bond and those in the plane.

Ex.: PCl5

Octahedral – AB6

6 bonds to the central atom, all equidistant from each other. 90° bond angles.

Ex.: SF6

Lone Pairs Do Occupy Space and Influence Geometry

But our description of the molecular geometry refers to the positions of the atoms only.

A summary of the shapes of various molecules is in Table 6-5, p. 186.

Different sizes of B groups may distort some bond angles that are given in the table.

VSEPR and Unshared Electron Pairs

One must always write out the Lewis structure for a molecule to decide on the proper geometry, the chemical formula of something does not tell you about lone pairs around the central atom.

Bent – AB2E

2 bonds to central atom with one lone pair. The lone pair bends what one would expect to be linear. The lone pair takes up more space than a bond and shoves the bonded atoms closer together than the 120° for trigonal planar.

Bent – AB2E2

The addition of a second lone pair forces the bonding atoms even closer together than what one expects from tetrahedral.

Trigonal Pyramidal – AB3EWith 3 bonds one

expects trigonal planar with 120° between the bonds, but the lone pair bends the plane away from the pair, forcing the atoms closer together.

Molecular Polarity

Reflecting on the geometries, we can now see why lone pairs on the central atom make a molecule polar – it changes the

geometry of the molecule and creates an uneven “tug of war” across the molecule.

Intermolecular ForcesThese are the forces of

attraction that occur between molecules. They vary in strength but are generally weaker than regular bonding (ionic, covalent, or metallic).

Melting and Boiling PointsUsually are a good measure of the force of attraction. The

higher the boiling point, the stronger the forces between particles.

Dipole Force

The strongest intermolecular forces exist between polar molecules. Each one acts as a dipole, created by equal but opposite charges that are separated by a short distance.

The direction of the arrow is pointed to the negative pole, the crossed tail indicates the + side

Dipole-dipole forcesThese are the forces of

attraction between polar molecules.

Example:

bp for F2 is -188°C

bp for HF is 20°C

bp for HCl is -85°C

Which is the stronger

dipole-dipole force?

For Molecules With More Than 2 AtomsThe molecular polarity

depends on both the polarity and the orientation of each bond.

Induced DipolesThe electrons of a nonpolar

molecule can be temporarily attracted by a polar molecule. This is weaker than a regular dipole-dipole force.

This can be very important in thesolubility of gases in water.

Hydrogen BondingA very special type of

dipole-dipole force in which a hydrogen atom is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom of another molecule.

Usually Represented ByDotted Lines

Explains High bp’s of Some CompoundsGives the H atom a large

positive charge, and it’s small size allows it to come very close to the unshared pair of electrons on an adjacent molecule.

Extremely Important in Biochemistry

StereoisomersIsomers – same chemical formula, different

structures (isopropyl alcohol vs 2-propanol)Stereoisomers – structures are mirror

images.

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London Dispersion ForcesThis is a weak attractive

force resulting from the imbalance of electrons and the creation of an instantaneous dipole. Important for noble gases and nonpolar molecules.

Increased ForceWith the increased number

of electrons in the interacting atoms or molecules, thus with increasing atomic or molar mass.

Assignment – Due Wed. EOP

Section 6.5 Worksheet

Mixed Review Worksheet

Molecular Geometry Worksheet

6.5 Textbook Problems