LEWIS STRUCTURES MOLCULAR GEOMETRY & HYDRIDIZATION
WHEN THE CENTRAL ATOMS IS AN ELEMENT FROM THE 3 rd 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 BONDS SINGLE 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 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
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 PAIR OF ELECTRONS BONDED- BONDED PAIR OF ELECTRONS
Due 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-)
LINEAR TRIGONAL PLANAR TETRAHEDRAL BENT (w/ 2 lone pairs of e-) TRIGONAL PYRAMIDAL.. S // \ :O:.. BENT (w/ 1 lone pairs of e-) 180 o 120 o o o (107.5 o ) 120 o (119 o ) o (104.5 o )
ETHANOIC ACID CH 3 COOH o 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 difference is NONPOLAR difference is POLAR IF A MOLECULE IS SYMMETRICAL OR IF THERE IS NO NET DIPOLE MOMENT THEN IT IS NONPOLAR EXAMPLE: CCl 4 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 that they have the greatest charge separation.
HYBRIDIZATION IS THE MIXING OF ORBITALS THE NEW HYBRID ORBITALS FORMED MAY BE sp, sp 2, sp 3, (sp 3 d, sp 3 d 2 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 sp 2 hybrid orbitals
EXAMPLES OF sp 2 HYBRIDIZATION ALL TRIGONAL PLANAR AND BENT(one lone pair of e-).. S // \ :O:..
Formation of sp 3 hybrid orbitals
EXAMPLES OF sp 3 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 bondA bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms. Sigma bond ( ) A bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms.
Pi bondA bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis). Pi bond ( ) A bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis).
Example: H 2 C=CH 2
Example: HC CH
Delocalized 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 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 orbitals all overlap leading to a very delocalized electron system Example: C 6 H 6 benzene