Bonding - Inorganic Chemistry

Unit 3
Chemical bonding:
Lewis structures, hybridisation and
molecular geometry
Inorganic 3
Slides prepared from information obtained from:
Cotton, F.A., Wilkinson, G. & Gaus, P.L. (1987). Basic inorganic chemistry. Third edition.
73-123.
Housecroft, C.E. & Sharpe, A.G. (2012) Inorganic Chemistry. Fourth edition. p31-59.
McMurry J.E. & Fay, R.C. (2012) Chemistry , six edition. 217-259.
D.F. Shriver; P.W. Atkins (2006) Inorganic Chemistry, 4th edition, Oxford University
Press.
BA. Averill; P. Eldredge, Principles of General Chemistry (v. 1.0M)
http://2012books.lardbucket.org/ Free online book
Atomic radii
Single bond covalent radius: typical contribution by an atom to the length of a
predominately covalent bond. E.g. covalent radius for fluorine is taken to be one
half the inter-nuclear distance in the homo-nuclear diatomic F2.
The Cl-Cl distance is 1.988 Å yields a covalent radius of 0.99 Å for the chlorine
atom.
Single bond covalent radius can be used to predict bond lengths for hetero-nuclear
diatomic molecules.
C-Si
0.77 Å + 1.17 Å =1.94 Å measured 1.87 Å
P-Cl
1.10 Å + 0.99 Å = 2.09 Å measured 2.04 Å
Table 1: Single bond covalent radii (Å)
H
0.28
N
0.70
O
0.66
C
0.77
P
1.10
Cl
0.99
Br
1.14
Si
1.17
F
0.64
Cotton and Wilkinson p60, 97
Three factors causing theoretical bond length to deviate
from predicted bond length
1 Multiple bonds: Multiple bonds are shorter than single bonds. (N2 triple bond
1.10 Å; N2 double bond 1.25 Å and N2 single bond 1.45 Å.
2 Hybridisation: Hybridisation can also affect the covalent radius since s
orbitals are more contracted than p orbitals. The radius will decrease with an
increase in the s character. C(sp3) 0.77 Å ; C(sp2) 0.73 Å ; C(sp) 0.70 Å.
3 Electronegativities: When there is great difference in electronegativities of
two atoms, the bond length is usually less than the sum of the covalent radii.
E.g. Si-F predicted distance 1.17 Å + 0.64 Å = 1.81 Å but actual distance is
1.54 Å. C-F predicted distance 0.77 Å + 0.64 Å = 1.44 Å but actual distance is
1.32 Å
Cotton and Wilkinson p 97
Van der Waals radii
Van der Waals radii is obtained from the non-bonded distance of closest
approach between atoms that are in contact with, but not bonded to, one
another.
Ionic covalent resonance
One resonance structure is covalent and the other resonance structure is ionic
A-B  A--B+
H-H  H--H+
http://www.calvin.edu/~dav4/Documents/ed100758t.pdf
Saturated molecules: number of valence e- available within molecule complex ion are sufficient
to allow Lewis diagram to be written with single bonds only achieving octet around each atom
(for all non-hydrogen atoms).
Unsaturated molecules: number of valence e- available within molecule complex ion are not
sufficient to allow Lewis diagram to be written with single bonds only.
Multiple bonds are required to achieve octet around each atom (for all non-hydrogen atoms).
Electron deficient molecules: All the valence e- used before an octet is achieved for all nonhydrogen atoms
e.g. H-Be-H
Hypervalent molecules: Can acquire more than an octet of electrons due to availability of dorbitals. The valence shell expands.
e.g.
LEWIS STRUCTURES
L
sp hybridization: linear
Averill and Eldredge et al.
sp2 hybridisation
Averill and Eldredge et al.
sp3 hybridisation
Averill and Eldredge et al.
sp3d hybridisation
trigonal bipyramidal
sp3d2 hybridisation
octahedral
Averill and Eldredge et al.
Hybridisation containing d orbitals and geometry
d2sp3 octahedral
s, px, py, pz, dz2, dx2-y2
dsp2 octahedral
s, px, py, dx2-y2
sd3 octahedral
s, dxy, dyz, dzx
dsp3 octahedral
s, px, py, pz, dz2
Do not memorise
dsp3 octahedral
s, px, py, pz, dx2-y2
Number of electrons
• Positive charge: Plus (+) means the
atoms/molecule has lost an electron
• Na  Na+ + e-
• Negative charge: Minus (-) means the
atoms/molecule has gained an electron
• Cl + e-  Cl• It works different to money in the bank
VSEPR Valance shell electron pair repulsion theory
ABxEy
B: atoms that are bonded
E: lone pairs
Number of charge clouds: x + y
Number of charge clouds determine electronic geometry
Molecular geometry (or just geometry) is derived from the electronic geometry
Cotton and Wilkinson p 86
ABxEy
Place lone pair
on equatorial
position
Place lone pair
on axial position
Averill and Eldredge et al.
Cotton and Wilkinson p 86 - 89
Place lone pair
on equatorial
position
They are all dsp3 hybridised
Get hybridisation of central
atom by counting number of
lone pairs and atoms
surrounding central atom.
Hypervalent molecules
Cotton and Wilkinson p 86 - 89
Place lone pair on
axial position
They are all dsp2 hybridised
Get hybridisation of central
atom by counting number of
lone pairs and atoms
surrounding central atom.
Organic carbonyls AB3
• sp2 hybridized central carbon atom
• Trigonal planar molecular geometry
• Groups and bonds around central
carbon equivalent due to resonance
• O-C-O thus ideal angles of 120°
Cotton and Wilkinson p 91
Organic carbonyls AB3
• sp2 hybridized central atoms
• Trigonal planar molecular geometry
• Groups around central carbon non-equivalent causing
deviation from the ideal 120°
• The electrons of the C=O double bond require more room in
the structure
• The more electron negative groups causes smaller angles
between them (e.g. angle smaller for F-C-F than H-C-H)
Qs: discuss the hybridisation, geometry and bond angles for H2C=O vs Cl2C=O vs F2C=O
Why is the angle 120° for CO32- and less than 120° for H2C=O?
Cotton and Wilkinson p 91
Bond angles: NO2- vs NO2 vs NO2+
• AB2E for NO2- and NO2
• sp2 hybridisation for nitrogen atom; molecular geometry is bent
• ONO angle less than 120° because of the larger repulsion and larger
volume require by the lone pair electrons on the nitrogen atom
• Removal of electron from NO2- to give NO2 causes increase in OCO angle
• Less repulsion and less volume required for one electron, thus larger angle
• AB2 for NO2+ Angle 180°
• sp hybridisation for nitrogen atom; molecular geometry is linear
Qs: Why is the angle increasing between the ONO going from NO2- to NO2 to NO2+?
Cotton and Wilkinson p 91
AH3E1 with A from group 5A
The XAX angle is the smallest in the molecule :AX3 where the
atoms X are the most electronegative
Qs: Why is the bond angle (XAX) smaller for NF3 than for NCl3
Cotton and Wilkinson p 94
AH3E1 with A from group 5A
•
•
•
•
AB3E1
sp3 hybridized central atoms
pyramidal molecular geometry
Lone pair causes deviations from the ideal 109.5° angles expected
for perfect sp3 hybrid sets
• Lone pair causes a decrease in the angle
• HAH angle is the smallest in the molecules :AH3 where the central
atom A is the largest
• Angles close to 90° indicates that choice of sp3 hybridisation is
inappropriate when the central atoms are antimony (Sb) and
arsenic (As)
Qs: discuss and compare the geometry, angles and hybridisation of NH3, PH3, AsH3 and SbH3.
Cotton and Wilkinson p 94
AH2E2 with A from group 6A
•
•
•
•
AB2E2
sp3 hybridized central atoms
bent molecular geometry
Repulsion between the two lone pairs causes deviations from the
ideal 109.5° angles expected for perfect sp3 hybrid sets
• Lone pairs causes a decrease in the angles
• HAH angle is the smallest in the molecules :AH2 where the central
atom A is the largest
• Angles close to 90° indicates that choice of sp3 hybridisation is
inappropriate when the central atoms are selenium (Se) and
tellurium (Te)
Cotton and Wilkinson p 94
Polarity
Averill and Eldredge et al.
Averill and Eldredge et al.
Unit 3
Chemical bonding:
Molecular Orbital theory
Inorganic 3
Slides prepared from information obtained from:
Cotton, F.A., Wilkinson, G. & Gaus, P.L. (1987). Basic inorganic chemistry. Third edition.
73-123.
Housecroft, C.E. & Sharpe, A.G. (2012) Inorganic Chemistry. Fourth edition. p31-59.
McMurry J.E. & Fay, R.C. (2012) Chemistry , six edition. 252-255.
D.F. Shriver; P.W. Atkins (2006) Inorganic Chemistry, 4th edition, Oxford University
Press.
BA. Averill; P. Eldredge, Principles of General Chemistry (v. 1.0M)
http://2012books.lardbucket.org/ Free online book
Molecular orbitals are constructed as linear combinations of atomic orbitals
Sigma () molecular orbitals
A bonding orbital arises from the constructive interference of neighbouring
atomic orbitals;
an antibonding orbital arises from their destructive interference, as indicated by a
node between the atoms.
Sigma () molecular orbitals
Averill and Eldredge et al.
Diamagnetic
Averill and Eldredge et al.
Molecular orbital theory
Diatomic, homo-neuclear
Bond order = (1-0)/2=0.5
Paramagnetic
Bond order = (2-1)/2=0.5
Paramagnetic
Bond order = (2-2)/2=0
Unstable
Not observed
Averill and Eldredge et al.
Bond order = (2-0)/2=1
Diamagnetic
Bond order = (4-2)/2=1
Diamagnetic
Bond order = (2-2)/2=0
Unstable
Not observed
Housecroft and Sharpe
Molecular orbital theory
Diatomic, homoneuclear
Higher bond order: more stable molecule, stronger bond, more energy required to
break bond, shorter bond length
Averill and Eldredge et al.
Bond order = (2-0)/2=1
Diamagnetic
e.g. Li2
Averill and Eldredge et al.
Bond order = (2-2)/2=0
Diamagnetic
e.g. Be2
Unstable
Sigma () and Phi ( ) molecular orbitals
Averill and Eldredge et al.
Sigma () and Phi ( ) molecular orbitals
Delta molecular orbitals
The Greek letter δ in their name refers to d orbitals, since the orbital symmetry of the
delta bond is the same as that of the usual (4-lobed) type of d orbital when seen
down the bond axis. This type of bonding is observed in atoms that have occupied d
orbitals with low enough energy to participate in covalent bonding, for example, in
organometallic species of transition metals.
http://en.wikipedia.org/wiki/Delta_bond
Cotton and Wilkinson p105
Brown et al, Central Science
Averill and Eldredge et al.
Molecular orbital theory
Diatomic, homoneuclear
Housecroft and Sharpe
Housecroft and Sharpe
Bond order = (8-6)/2=1
Diamagnetic
Averill and Eldredge et al.
Bond order = (8-4)/2=2
Paramagnetic
Different magnetic properties for
Lewis diagram vs molecular orbital
diagram
Experiment: paramagnetic
O=O
diamagnetic
Averill and Eldredge et al.
Because the O2 molecule has two unpaired electrons, it is paramagnetic.
Consequently, it is attracted into a magnetic field, which allows it to remain
suspended between the poles of a powerful magnet until it evaporates
*2p
*2p
2p
2p
*2s
2s
Averill and Eldredge et al.
Bond order = (8-3)/2=2.5
Paramagnetic
Averill and Eldredge et al.
NO- weaker bond than NO
Bond order = (8-4)/2=2
Paramagnetic
NO
NO- Averill and Eldredge et al.
Bond order = (8-2)/2=3
Diamagnetic
NO+ stronger bond than NO
NO+
Averill and Eldredge et al.
Isoelectric molecules:
same number of electrons
Valence isoelectric molecules:
same number of valence electrons
For example: CO, N2, CN- and NO+ each has two nuclei and 10 valence electrons
O2 and NO- each has two nuclei and 12 valence electrons
Not to be confused with:
Isostructural: same molecular geometry
But not same number of electrons
[BH4]-; CH4; [NH4]+ tetrahedral geometry
Valence isoelectric molecules
2 *2p
π*2p
2 2p
2p
2p
π2p
1 *2s
2s
2s
12s
N
10 valence electrons
N2
N
Bond order = (8-2)/2=3
Bond order = (8-2)/2=3
Diamagnetic
10 valence electrons
Bond order = (8-2)/2=3
Paramagnetic
NO-
10 valence electrons
Better molecular orbital bonding model for CO, but complex
10 valence electrons
Housecroft and Sharpe
No nonbonding electrons
Simpler molecular orbital bonding model for CO, but not so correct
3  Has antibonding and
bonding character
1  Not 100%
nonbonding
It has bonding and
non-bonding
character
C
O
10 valence electrons
Housecroft and Sharpe
N2 vs CO, isoelectric but different molecular orbital diagrams
2 *2p
π*2p
2 2p
2p
2p
π2p
1 *2s
2s
2s
12s
N
N2
N
• Oxygen orbitals lower than carbon orbitals
because it is more electronegative
• 2s-2p energy separation for O greater than C
• Resulting in 2s from C overlap 2p from O
giving 2s-2p mixing
• 3 anti-bonding character, CO+ slightly
stronger bond than CO
Non bonding electrons do not participate in bonding therefore not part of bond order
calculation
Bond order = (2-0)/2=1
Housecroft and Sharpe
Non bonding electrons do not participate in bonding therefore not part of bond order
calculation
Bond order = (2-0)/2=1
Averill and Eldredge et al.