Chemistry 100 Chapter 9
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Transcript of Chemistry 100 Chapter 9
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Chemistry 100 Chapter 9
Molecular Geometry and Bonding Theories
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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!
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The VSEPR Model
Electrons are negatively charged, they want to occupy positions such that electron Electron interactions are minimised 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
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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
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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
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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.
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Bond Dipoles in Molecules
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More Bond Dipoles
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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
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The Bonding in H2
Hydrogen molecule a single bond indicating the overlap of the
1s orbitals on the individual atoms cylindrical symmetry with respect to the
line joining the atomic centres, i.e., a bond
H H
Overlap Region
1s (H1) – 1s(H2) bond
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The Bonding in H2
H H
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The Cl2 Molecule
In the chlorine molecule, we observe a single bond indicating the overlap of the 3p orbitals on the individual atoms.
Cl Cl
Bonding description 3pz (Cl 1) – 3pz (Cl 2)
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Is This a Bond?
Cl Cl
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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.
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The CH4 Molecule
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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.
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sp2 Hybridisation
What if we try to rationalise the bonding picture in the BH3 (a trigonal planar molecule)?
We mix 2 “pure” p orbitals and a “pure” s orbital to form “hybrid” or mixed sp2 orbitals.
These three sp2 hybrid orbitals lie in the same plane with an angle of 120 between them.
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A Trigonal Planar Molecule
Overlap regions
B
H
H H
Overlap region
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sp Hybridisation
What if we try to rationalize the bonding picture in the BeH2 species (a linear molecule)?
We mix a single “pure” p orbital and a “pure” s orbital to form two “hybrid” or mixed sp orbitals
These sp hybrid orbitals have an angle of 180 between them.
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A Linear Molecule
The BeH2 molecule
BeH H
Overlap Regions
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Double Bonds
Look at ethene C2H4. Each central atom is an AB3 system,
the bonding picture must be consistent with VSEPR theory.
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The Bond
Additional feature an unhybridized p
orbital on adjacent carbon atoms.
Overlap the two parallel 2pz orbitals (a -bond is formed).
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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 ) ] type
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The C2H4 Molecule
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The Bond Angles in C2H4
Bond angles HCH = HCC 120. 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 bond
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The Triple Bond in C2H2
Bond angles HCH = HCC = 180. bonds are perpendicular to
the molecular plane. Triple bond one bond and
two bonds
Triple bond rationalized by assuming sp hybridization exists on
the central atoms!
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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 ) ] type [2pz
(C 1 ) – 2pz (C 2 ) ] type
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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) ] type
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Key Connection – VSEPR and Valence Bond Theory!!
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sp3d Hybridisation
How can we use the hybridisation concept to explain the bonding picture PCl5.
There are five bonds between P and Cl (all type bonds).
5 sp3d orbitals these orbitals overlap with the 3p orbitals in Cl to form the 5 bonds with the required VSEPR geometry trigonal bipyramid.
Bond overlaps[sp3d (P ) – 3pz (Cl) ] x 5 type
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sp3d2 Hybridisation
Look at the SF6 molecule. 6 sp3d2 orbitals these orbitals
overlap with the 2pz orbitals in F to form the 6 bonds with the required VSEPR geometry octahedral.
Bond overlaps [sp3d2 (S ) – 2pz (F) ] x 6 type
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Notes for Understanding Hybridisation 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.
Hybridisation requires energy for the promotion of the electron and the mixing of the orbitals energy is offset by bond formation.
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Delocalised Bonding
In almost all the cases where we described the bonding n the molecule, the bonding electrons have been totally associated with the two atoms that form the bond they are localised.
What about the bonding situation in benzene, the nitrate ion, the carbonate ion?
In benzene, the C-C bonds are formed from the sp2 hybrid orbitals. The unhybridised 2pz orbital on C overlaps with another 2pz orbital on the adjacent C atom.
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Three bonds are formed. These bonds extend over the whole molecule (i.e. the bonds are delocalised).
The electrons are free to move around the benzene ring.
Any species where we had several resonance structures, we would have delocalisation of the -electrons.
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Delocalised Electrons in Molecules
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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.
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Recall the wave properties of electrons.
constructive interference the two e- waves interact favourably; loosely analogous to a build-up of e- density between the two atomic centres. destructive interference unfavourable interaction of e- waves; analogous to the decrease of e- density between two atomic centres.
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Constructive and Destructive Interference
+
+
Constructive
Destructive
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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.
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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
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Bonding and Anti-Bonding M.O.’s from 1s atomic Orbitals
Energy
1s
1s
1s
*1s
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The MO’s in the H2 Atom
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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.