chemistry 100 chapter 9 molecular geometry and bonding theories
TRANSCRIPT
Chemistry 100 Chapter 9
Molecular Geometry and Bonding Theories
Molecular Geometry
The three-dimensional arrangement of atoms in a molecule molecular geometry
Lewis structures can’t be used to predict geometry
Repulsion between electron pairs (both bonding and non-bonding) helps account for the molecular structure!
The VSEPR Model
Electrons are negatively charged, they want to occupy positions such that electron Electron interactions are minimized as much as possible
Valence Shell Electron-Pair Repulsion Model treat double and triple bonds as single domains resonance structure - apply VSEPR to any of them formal charges are usually omitted
Four Electron Domains – Three Different Geometries Replacement of bonding domains (B)
with nonbonding domains (E)results in a different molecular geometry.
AB4 AB3E AB2E2
Molecules With More Than One Central Atom
We simply apply VSEPR to each ‘central atom’ in the molecule.
• Carbon #1 – tetrahedral
• Carbon #2 – trigonal planar
Dipole Moments
The HF molecule has a bond dipole – a charge separation due to the electronegativity difference between F and H.
The shape of a molecule and the magnitude of the bond dipole(s) can give the molecule an overall degree of polarity ® dipole moment.
+H-F
Homonuclear diatomics ® no dipole moment (O2, F2, Cl2, etc)
Triatomic molecules (and greater). Must look at the net effect of all the bond dipoles.
In molecules like CCl4 (tetrahedral) BF3 (trigonal planar) all the individual bond dipoles cancel Þ no resultant dipole moment.
Bond Dipoles in Molecules
More Bond Dipoles
Valence Bond Theory and Hybridisation
Valence bond theory description of the covalent bonding and
structure in molecules.
Electrons in a molecule occupy the atomic orbitals of individual atoms.
The covalent bond results from the overlap of the atomic orbitals on the individual atoms
The Bonding in Diatomic Molecules
Hydrogen molecule a single bond between the
two H 1s orbitals a bond
Hydrogen Chloride a single bond from the
overlap of the Cl 3p orbital with the H 1s orbital
Chlorine molecule a single bond from the
overlap of the Cl 3p orbitals
Hybrid Atomic Orbitals
Look at the bonding picture in methane (CH4).
• Bonding and geometry in polyatomic molecules may be explained in terms of the formation of hybrid atomic orbitals
• Bonds overlap of the hybrid atomic orbitals on central atoms with appropriate half-filled atomic orbital on the terminal atoms.
The CH4 Molecule
The Formation of the sp3 Hybrids
Mix 3 “pure” p orbitals and a “pure” s orbital form an sp3
“hybrid” orbital. Rationalize the
bonding around the C central atom.
sp2 Hybridisation
Examine BH3 (a trigonal planar molecule)
sp Hybridisation
Examine BeF2 (a linear molecule). These sp hybrid orbitals have an
angle of 180 between them.
A Linear Molecule
The BeF2 molecule
Double Bonds
Look at ethene C2H4. Each central atom is an AB3 system,
the bonding picture must be consistent with VSEPR theory.
Sigma () Bonds
Sigma bonds are characterized by Head-to-head overlap. Cylindrical symmetry of electron density
about the internuclear axis.
Pi () Bonds
Pi bonds are characterized by Side-to-side
overlap. Electron density
above and below the internuclear axis.
Bond overlaps in C2H4
There are three different types of bonds
[sp2 (C ) – 1s (H) ] x 4 type
[sp2 (C 1 ) – sp2 (C 2 ) ] type
[2pz (C 1 ) – 2pz
(C 2 ) ] p type
The C2H4 Molecule
The Bond Angles in C2H4
Bond angles HCH = HCC 120. p bond is perpendicular
to the plane containing the molecule.
Double bonds – Rationalize by assuming
sp2 hybridization exists on the central atoms!
Any double bond one bond and a p bond
The Triple Bond in C2H2
Bond angles HCH = HCC = 180. p bonds are perpendicular to
the molecular plane.
Triple bond one bond and two p bonds
Triple bond rationalized by assuming sp hybridization exists on
the central atoms!
Bond Overlaps in C2H2
There are again three different types of bonds[sp (C ) – 1s (H) ] x 2 type [sp (C 1 ) – sp (C 2 ) ] type
[2py (C 1 ) – 2py
(C 2 ) ] p type [2pz
(C 1 ) – 2pz (C 2 ) ] p type
Bonding in H2O
Bonding Overlaps[sp3(O)–1s(H)] x 2
Bond Overlaps in H2CO
There are again three different types of bonds[sp (C) – 1s (H) ] x 2 type [sp2 (C) – sp2 (O) ] type [2p (C) – 2p (O) ] p type
Key Connection – VSEPR and Valence Bond Theory!!
sp3d Hybridisation
How can we use the hybridisation concept to explain the bonding picture PCl5.
There are five bonds between P and Cl (all stype bonds).
5 sp3d orbitals ® these orbitals overlap with the 3p orbitals in Cl to form the 5 s bonds with the required VSEPR geometry ® trigonal bipyramid.
Bond overlaps[sp3d (P ) – 3pz (Cl) ] x 5 type
sp3d2 Hybridisation
Look at the SF6 molecule. 6 sp3d2 orbitals ® these orbitals
overlap with the 2pz orbitals in F to form the 6 s bonds with the required VSEPR geometry ® octahedral.
Bond overlaps [sp3d2 (S ) – 2pz (F) ] x 6 type
Notes for Understanding Hybridization
Applied to atoms in molecules only Number hybrid orbitals = number of atomic
orbitals used to make them Hybrid orbitals have different energies and
shapes from the atomic orbitals from which they were made.
Hybridization requires energy for the promotion of the electron and the mixing of the orbitals ® energy is offset by bond formation.
Delocalised Bonding
Valence bond theory – bonding electrons have been totally associated with
the two atoms that form the bond they are localized. What about the bonding situation in benzene,
the nitrate ion, the carbonate ion?
Bonding in Aromatic Molecules
Benzene C-C s bonds are formed from the sp2 hybrid orbitals. Unhybridized 2pz orbital on adjacent C atoms overlap
(bonds).
Bonding in the Benzene Molecule The p bonds extend over the whole
molecule the p electrons bonds are delocalized – they are
free to move around the benzene ring. Resonance structures – delocalization of the
-electrons.
The Nitrate Anion
Three resonance structures Alternating single
and double bonds Blend resonance
structures Delocalized
bond over anion backbone
Molecular Orbital (M.O.) Theory
Valence bond and the concept of the hybridisation of atomic orbitals does not account for a number of fundamental observations of chemistry.
To reconcile these and other differences, we turn to molecular orbital theory (MO theory).
MO theory – covalent bonding is described in terms of molecular orbitals the combination of atomic orbitals that results in
an orbital associated with the whole molecule.
Constructive and Destructive Interference
+
+
Constructive
Destructive
ybonding = C1 ls (H 1) + C2 ls (H 2) yanti = C1 ls (H 1) - C2 ls (H 2)
Bonding Orbital ® a centro-symmetric orbital (i.e. symmetric about the line of symmetry of the bonding atoms).
Bonding M’s have lower energy and greater stability than the AO’s from which it was formed.
Electron density is concentrated in the region immediately between the bonding nuclei.
Anti-bonding orbital ® a node (0 electron density) between the two nuclei.
In an anti-bonding MO, we have higher energy and less stability than the atomic orbitals from
which it was formed. As with valance bond theory (hybridisation)
2 AO’s ® 2 MO’s
Bonding and Anti-Bonding M.O.’s from 1s atomic Orbitals
Energy
1s
1s
s1s
s*1s
The MO’s in the H2 Atom
The situation for two 2s orbitals is the same! The situation for two 3s orbital is the same.
Let’s look at the following series of moleculesH2, He2
+, He2
bond order = ½ {bonding - anti-bonding e-‘s}. Higher bond order º greater bond stability.