HL STANDARD E

LEWIS STRUCTURESMOLCULAR GEOMETRY

& HYDRIDIZATION

14.1 MOLECULES w/ EXPANDED OCTETS

WHEN THE CENTRAL ATOMS IS AN ELEMENT FROM THE 3rd PERIOD OR BELOW, WE SOMETIMES FIND COMPOUNDS IN WHICH THERE ARE MORE THAN EIGHT ELECTRONS AROUND THE CENTRAL ATOM.

THIS ARRANGEMENT IS POSSIBLE BECAUSE THE d ORBITALS IN THE VALANCE SHELL OF THESE ATOMS HAVE ENERGY VALUES CLOSE TO THOSE OF THE p ORBITALS

TYPES OF COVALENT BONDSSINGLE BONDS• LONGEST OF THE 3 TYPES• WEAKEST OF THE 3 TYPES• CONTAINS ONE PAIR OF ELECTRONS (2 ELECTRONS)

DOUBLE BONDS• LENGTH IS GREATER THAN TRIPLE BUT LESS THAN SINGLE• STRENGTH IS GREATER THAN SINGLE BUT LESS THAN TRIPLE• CONTAINS TWO PAIRS OF ELECTRONS (4 ELECTRONS)

TRIPLE BONDS• SHORTEST OF THE 3 TYPES• STRONGEST OF THE 3 TYPES• CONTAINS THREE PAIRS OF ELECTRONS (6 ELECTRONS)

 LONGEST BONDS(SINGLE) ARE THE WEAKEST

SHORTEST BONDS(TRIPLE) ARE THE STRONGEST 

VSERP THEORYVALENCE SHELL ELECTRON REPLUSION THEORY

VSERP THEORY states that electron pairs found in the outer energy level or valence shell of atoms repel each other and thus position themselves as far apart as possible

The VSERP theory helps to predict the shape of the molecule. The following points should be considered in predicting shape (molecular geometry):

• Double or triples bonded electron pairs are oriented together and so behave in terms of repulsion as a single unit know as a ‘negative charged center’.

• The repulsion applies to both bonding and non- bonding (lone) pairs of electrons

• The total number of charge centers around the central atom will determine the geometrical arrangement of the electrons

• The shape of the molecule is determined by the angles between the bonded atoms

• Because non-bonded (lone) pairs are not shared between two atoms they cause more repulsion than bonded pairs

• The repulsion decrease in the following order:

lone-lone pair>lone-bonding pair>bonding-bonding pair

FOR EXAMPLE, IN THE LEWIS STRUCTURE OF WATER, THERE ARE TWO NON-BONDED (LONE) PAIRS OF ELECTRONS ON THE OXYGEN ATOM. THUS THE NON-BONDED (LONE) PAIR OF ELECTRONS REPEL MORE THAN THE LONE-BONDED PAIR BETWEEN THE OXYGEN AND HYDROGEN AND THE BONDED-BONDED PAIR OF ELECTRONS BETWEEN THE TWO HYDROGENS HAVE THE LEAST REPULSION.

NON-BONDED (LONE)&NON-BONDED (LONE) PAIR OF ELECTRONS

NON-BONDED(LONE)& BONDED PAIROF ELECTRONS

BONDED-BONDED PAIR OF ELECTRONS

BOND ANGLESDue to the fact that lone pair of

electrons repel more than bonded pairs, the bond angles will be different

in different shaped molecules

BEND(with 2 lone pair of e-) <TRIGONAL PYRAMIDAL<TETRAHEDRAL<BEND(with 1 lone pair of e-)<TRIGONAL PLANAR<LINEAR

****KNOW THE BOND ANGLES OF THIS SHAPES

LINEAR

TRIGONAL PLANAR

TETRAHEDRAL

BENT (w/ 2 lone pairs of e-)

TRIGONAL PYRAMIDAL

.. S // \:O: :O: ..

BENT (w/ 1 lone pairs of e-)

180o

120o

109.5o

109.5o

(107.5o ) 120o (119o)

109.5 o (104.5o)

ETHANOIC ACIDCH3COOH

104.5o

TETRAHEDRAL

TRIGONAL PLANAR

BENT

CAN A MOLECULE BE NONPOLAR AND STILL HAVE POLAR BONDS WITHIN THE MOLECULE???

FIRST THINK ABOUT WHAT MAKES A BOND POLAR!

NOW THINK ABOUT WHAT MAKES A MOLECULE NONPOLAR!

THE DIFFERENCE IN ELECTRONEGATIVITY VALUES (CHARGE SEPARATION) DETERMINE IF A BOND WILL BE POLAR OR NONPOLAR.

0-.3 difference is NONPOLAR.3-1.7 difference is POLAR

IF A MOLECULE IS SYMMETRICAL OR IF THERE IS NO NET DIPOLE MOMENT THEN IT IS NONPOLAR

EXAMPLE: CCl4

THE C-Cl BOND IS POLAR BUT THE MOLECULE IS TETRAHEDRAL SO IT IS SYMMETRICAL AND

THUS NONPOLAR

YES

WHICH OF THE FOLLOWING BONDS IS THE MOST POLAR

B-C or C-O or N-O or O-F

ANSWER:C-O

b/c they are further apart on the periodic table which means thatthey have the greatest charge separation.

HYBRIDIZATION

HYBRIDIZATION IS THE MIXING OF ORBITALS

THE NEW HYBRID ORBITALS FORMED MAY BE sp, sp2, sp3, (sp3d, sp3d2 not covered)

These new orbitals are called hybrid orbitals

The process is called hybridization

What this means is that both the s and one p orbital are involved in bonding to the connecting

atoms

Formation of sp hybrid orbitals

The combination of an s orbital and a p orbital produces 2 new orbitals called sp orbitals.

EXAMPLES OF MOLECULES WITH sp HYBRIDIZATION

ALL LINEAR MOLECULES

Formation of sp2 hybrid orbitals

EXAMPLES OF sp2 HYBRIDIZATION

ALL TRIGONAL PLANAR

AND BENT(one lone pair of e-)

.. S // \:O: :O: ..

Formation of sp3 hybrid orbitals

EXAMPLES OF sp3 HYBRIDIZATION

ALL TETRAHEDRAL

TRIGONAL PYRIMIDAL&

BENT (w/ two lone pair of e-)

Hybrid orbitals can be used to explain bonding and molecular geometry

Multiple Bonds

 Everything we have talked about so far has only dealt

with what we call sigma bonds Sigma bond (s) A bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms.

Pi bond (p) A bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis).

Example: H2C=CH2

Example: H2C=CH2

Example: HCCH

Delocalized p bonds

 When a molecule has two or more resonance structures,

the pi electrons can be delocalized over all the atoms that have pi bond overlap.

In general delocalized p bonding is present in all molecules where we can draw resonance structures with

the multiple bonds located in different places.

Benzene is an excellent example.  For benzene the p orbitals all overlap leading to a very delocalized electron

system

Example: C6H6 benzene