Download Chapter 3 : Simple Bonding Theory Why do they make chemical

Chapter 3 : Simple Bonding Theory
N2
NH3
H2O
Why do they make chemical bonds ?
Chapter 3 : Simple Bonding Theory
Stabilization
1
Types of Chemical Bonds
Metallic Bond Ionic Bond Covalent Bond
Types of Chemical Bonds
Metallic Bond
Na: 1s22s22p63s1
2
Types of Chemical Bonds
Metallic Bond
Na: 1s22s22p63s1
e- Na+
e Na+
e e Na+
e
+
Na
Na+
Forming sea of electrons
(freely-moving valence electrons)
Structure ?
The way of packing
Minimizing energy
Packing efficiency
Close Packing of Spheres
1926 Goldschmidt proposed atoms could be considered as packing in
solids as hard spheres.
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Other Structures of Metals
Hexagonal-Close
Packing
(HCP)
Cubic-Close
Packing
(CCP, FCC)
FCC+1/2 of Td holes
Body-Centered
Cubic
(BCC)
Primitive
Cubic
Diamond Structure
Types of Chemical Bonds
Ionic Bond
NaCl
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Types of Chemical Bonds
Ionic Bond
Melting
conductivity
NaCl
No conductivity
Why conductivity?
and Cl- exit in NaCl.
Na+
Why Na+ and Cl- not Na and Cl ?
Simply, Na+ and Cl- are more stable than atomic
Na and Cl in NaCl
Na: 1s22s22p63s1
Na+: 1s22s22p6
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
Types of Chemical Bonds
Ionic Bond
NaCl
The way of packing
Minimizing energy
Packing efficiency
Lattice energy
NaCl(s) Æ Na+(g) + Cl-(g)
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Ionic Solids
Types of Chemical Bonds
Covalent Bond
In covalent bonding, electrons are “shared” between bonding partners.
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Types of Chemical Bonds
Covalent Bond
Bond energy
Types of Chemical Bonds
Covalent Bond
Think of the covalent bond as the
electron density existing
between the C and H atoms.
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Covalent Bond
Localized electron bonding models
•Lewis dot structure
•VSEPR (Valence shell electron pair repulsion)
•Valence bond theory (hybridization)
Delocalized electron bonding model
•Molecular orbital (MO) theory
Localized Model Limitations
• It is important to keep in mind that the models we are
discussing are just that…..models.
• We are operating under the assumption that when
forming bonds, atoms “share” electrons using atomic
orbitals
• Electrons involved in bonding: “bonding pairs”.
Electrons not involved in bonding: “lone pairs”.
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Lewis Dot Structures
• Developed by G. N. Lewis to serve as a
way to describe bonding in polyatomic
systems.
• Central idea: the most stable arrangement
of electrons is one in which all atoms have
a “noble” gas configuration. (Octet rule)
• Example: NaCl vs Na+ClNa: [Ne]3s1
Cl: [Ne]3s23p5
Na+: [Ne]
Cl-: [Ne]3s23p6 = [Ar]
Lewis Dot Structures
F
C
Lone Pair (6 x)
F
F
• Atoms are
represented by
atomic symbols
surrounded by
valence electrons.
• Electron pairs
between atoms
indicate bond
formation.
Bonding Pair
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Lewis Dot Structures
F
N–A=S
F
N: Total number of valence electrons when
considered as noble gases
A: Total number of actual valence electrons
S: Number of valence electrons shared
F
F
F2
F
16-14=2
F
Lewis Dot Structures
• An example: NO+
N
+
N
O
N-A=S
N
O
16-10=6
O
+
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Lewis Dot Structures
• An example: Cl2O
O
Cl
N-A=S
Cl
If one atom is
different from others,
position it in the
center of the
molecule.
24-20=4
Cl O Cl
Lewis Dot Structures
• An example: CH4
H H
H H
C
H
N-A=S
16-8=8
H C H
H
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Lewis Dot Structures
• An example: CO2
O
O
C
N-A=S
24-16=8
O C O
Octet Violation
O
CO double bond
C O
O
C O
Lewis Dot Structures
• An example: POCl3
P
O
Cl
N-A=S
Cl
40-32=8
O
Cl
P
Cl
O
Cl
Cl
Cl
P
Cl
Cl
If there are single
atoms of two
elements, position
one with the larger
atomic number in the
center of the
molecule.
Usually,
C family has 4 bonds
O family has 2 bonds
N family has 3 bonds
F family has 1 bonds
Later
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Resonance Structures
• The classic example: O3
O
O
O O
O
O
O
O
O
Both structures are correct!
Resonance Structures
• In this example, O3 has two resonance
structures:
O
O
O
• Conceptually, we think of the bonding
being an average of these two structures.
• Electrons are delocalized between the
oxygens such that on average the bond
strength is equivalent to 1.5 O-O bonds.
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Resonance Structures
• An example: CO322-
2-
2-
O
O
O
C
C
C
O
O
O
O
O
O
116 pm
143 pm
Experimentally, 129 pm
Resonance Structures
• An example: CO322-
O
2-
2-
O
O
O
C
C
C
O
O
O
O
O
The electronic energy is lowered in a molecule
with several resonance structures.
### Think about particle in a box ###
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Formal Charge
A
B
Half the electron(s) of the bonding
Then
Formal charge = net charge of the atom
0
2-
0
O
N
C
O
-1
O
O
0
O
+1
+
Cl
0
-1
P
0
Cl
Cl 0
0
Sum of formal charges = net charge of the
molecule or ion
Formal Charge
Why formal charge?
Help in assigning bonding when there are several
possibilities of Lewis structure.
1. Structures with small FC (-2,+2 or less) are more likely.
2. Nonzero FCs on adjacent atoms are usually of opposite sign.
3. More electronegative atoms should have negative FC.
4. FCs of opposite signs separated by large distance are unlikely.
5. The largest sum of the electronegativity differences for
adjacent atoms. Ex: HOCl more stable than HClO
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Beyond the Octet Rule
(Expanded Shell in 3rd and higher period)
F
F
F
Cl
F
F
F
F
F
S
F
Old : using d orbitals Æ new : not necessarily (MO theory)
Beyond the Octet Rule
(Expanded Shell in 3rd and higher period)
0
-1
O
0
Cl
P
0
Cl
O
Cl
Cl
P
+1
Cl
0
Cl
Formal charge favors
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Beyond the Octet Rule
(Expanded Shell in 3rd and higher period)
Beyond the Octet Rule
(Be and B)
• Some atoms (Be and B in particular)
undergo bonding, but will form stable
molecules that do not fulfill the octet rule.
-2
+1
F=Be=F Æ network
Cl=Be=Cl Æ network or dimerize
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Beyond the Octet Rule
(Be and B)
F
F
B
F
F
F
B
…
F
FC favors, but B-F single bond : 152 pm,
actually, 131 pm So….
H
BH3 Æ dimerize
H
B
H
H
B
H
H
Beyond the Octet Rule
• Finally, one can encounter odd electron
systems where full pairs will not exist.
• Example: Chlorine Dioxide.
O Cl O
Unpaired electron
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Covalent Bond
Localized electron bonding models
•Lewis dot structure
•VSEPR (Valence shell electron pair repulsion)
•Valence bond theory (hybridization)
Delocalized electron bonding model
•Molecular orbital (MO) theory
VSEPR
• The Lewis Dot Structure approach provided some
insight into molecular structure in terms of bonding,
but what about geometry?
• Recall from last lecture that we had two types of
electron pairs: bonding and lone.
• Valence Shell Electron Pair Repulsion (VSEPR):
3D structure is determined by minimizing
repulsion of electron pairs.
=> Position valence electrons as far from each other
as possible.
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VSEPR
• AXmEn
A: central atom
X: atom or group of atom surrounding A
E: a lone pair of electrons
• Steric number: SN = m + n
Bear in mind that VSEPR provides approximate
shapes for molecules , not a complete picture of
bonding.
VSEPR
SN = 2
Linear Structure
180°
• Example: BeF2, CO2
F Be F
O
C
O
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VSEPR
SN = 3
120°
Trigonal (planar triangular)
• BF3, SO3
O
F
F
B
S
F
O
O
VSEPR
SN = 4
H
109.5°
H C H
H
Lewis Structure
VSEPR Structure
Tetrahedral structure
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VSEPR
SN = 5
Trigonal bipyramidal structure: 120° in
plane, and two orbitals at 90° to plane
• PCl5:
Cl
Cl
Cl
Cl
P
Cl
120o
Cl
90o
P
Cl
Cl
Cl
Cl
VSEPR
SN = 6
• Octahedral structure: all angles are 90°.
• Example: SF6:
F
F
F
F
F
F
S
F
F
S
F
F
F
F
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VSEPR
SN = 7
• Pentagonal bipyramidal structure: 72° in plane,
and two orbitals at 90° to plane
F
• Example: IF7:
F
F
F
F
F
I
F
F
F
I
F
F
F
F
F
SN = 8
VSEPR
• Square antiprismatic
• Example: TaF83-:
F
70.5o
F
F
F
99.6o
Ta
F
F
109.5o
F
F
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Refinement (I) of VSEPR
bonding pair
lone pair
Closer to X and spead out
around X
Lp – Lp >> Lp-Bp > Bp-Bp
Refinement (I) of VSEPR
SN = 3
Sn
Cl
95o
Cl
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Refinement (I) of VSEPR
SN = 4
109.5°
106.6°
Tetrahedral
Trigonal pyramidal
104.5°
Vent
Refinement (I) of VSEPR
SN = 5
SF4
F
120o
F
F
F
90o
101.6o
S
S
F
F
F
S
F
F
173o
F
Lp-Bp: 3x 90o + 1x 180o
Bp-Bp: 3x 90o + 3x 120o
F
F
seesaw
Lp-Bp: 2x 90o + 2x 120o
Bp-Bp: 4x 90o + 1x 120o
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Refinement (I) of VSEPR
SN = 5
F
169.8 pm
ClF3
F
87.5
o
Cl
F
F
Cl
159.8 pm
F
F
F
F
Cl
F
F
distorted T
Cl
F
F
Refinement (I) of VSEPR
SN = 5
XeF2
Xe
F
F
F
Xe
F
linear
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Refinement (I) of VSEPR
SN = 6
SF6
IF5
XeF4
F
F
F
F
F
F
S
F
F
Octahedral
F
I
F
F
F
Xe
F
81.9o
F
F
~ square pyramidal
Square planar
Refinement (II) of VSEPR
Multiple bonds have slightly greater repulsive effects
than single bonds because of the repulsive effect of
π electrons
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Refinement (II) of VSEPR
Refinement (II) of VSEPR
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Electronegativity
A measure of the tendency of an atom to
attract a bonding pair of electrons.
Electronegativity
Pauling's Electronegativity
The bond energy E(AB) in a molecule AB is always greater than the mean of
the bond energies E(AA) + E(BB) in the homonuclear species AA and BB.
Pauling argued that in an "ideal" covalent bond E(AB) should equal this mean,
and that the "excess" bond energy is caused by electrostatic attraction
between the partially charged atoms in the heternuclear species AB. In effect,
he was saying that the excess bond energy arises from an ionic contribution to
the bond.
He managed to treat this ionic contribution by the equation
E(AB) = (1/2)[E(AA)+E(BB)] + 96.48(χA - χB)2
in which E(AB) is expressed in kJ mol-1 (1 electron volt, 1eV, = 96.48 kJ mol1) and χA - χB represents the difference in "electronegativity" between the two
elements.
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Electronegativity
Milliken's Electronegativity
Average of ionization energy and electron affinity
χM = (IEv + EAv)/2
χM = 3.48[(IEv + EAv)/2 - 0.602] in Pauling scale
Alled-Rochow Electronegativity
A scale of electronegativity based upon the electrostatic force of
attraction between the nucleus and the valence electrons.
χAR = 0.744 + 0.359Zeff/r2 in Pauling scale
r = covalent radius
Electronegativity
Allen's Electronegativity
The average one-electron energy of the valence shell electrons in the ground
state free atoms. This means that values of the Allen electronegativity can be
calculated from spectroscopic data.
χspec = (mεp + nεs)/(m + n)
m = number of p-electrons
n = number of s-electrons
εp = p ionization energy
εs = s ionization energy
The values obtained correlate well with Pauling electronegativity and with
Allred-Rochow electronegativity.
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Electronegativity
Electronegativity
Bond Polarity and Polar Molecule
µ = Qr (Dipole molment)
H
C
H
measured by dielectric constant
H
H
zero net dipole
Which force derives liquefaction of this?
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Electronegativity
Hydrogen Bonding
BP
Evidence for H-bonding
Known H-bondings
H2O : max 4 H-bonds
HF : max 2 H-bonds
NH3 : not certain
Electronegativity
Life much depends on Hydrogen Bonding
Ice
DNA
Protein
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Refinement (III) of VSEPR
Electronegativity and Atomic Size Effects
AXn : χ↓-> angle ↑,
size ↑and bond length ↑-> angle ↑
AXn : χ, size counter-balance
AXn : size ↑and bond length ↑-> angle ↓, χ↓-> angle ↓
Refinement (IV) of VSEPR
Ligand Close-Packing
For a series of molecules with the same central atom, the nonbonded distances between the outer atoms are consistent.
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Covalent Bond
Localized electron bonding models
•Lewis dot structure
•VSEPR (Valence shell electron pair repulsion)
•Valence bond theory (hybridization)
Delocalized electron bonding model
•Molecular orbital (MO) theory
Valence Bond Theory (VBT)
First attempt of quantum mechanical
explanation of chemical bonding
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Valence Bond Theory (VBT)
Ψ = φΑ(1)φΒ(2)+φΑ(2)φΒ(1)
Ψ = φΑ(1)φΒ(2)
Valence Bond Theory (VBT)
Ψ = φΑ(1)φΒ(2)+φΑ(2)φΒ(1)
Ψ = φΑ(1)φΒ(2)
Each electron is free to migrate to the other atom.
Æ Probability to find 2e-’s between two nuclei Is high.
Æ bonding
Overlap of atomic orbitals
Valence
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Valence Bond Theory (VBT)
H + Cl
H
Cl + Cl
H Cl
Cl
1s
3s
Cl
3p
3p
1s
Cl
3p
3p
σ m.o.
σ m.o.
VBT
But what about CH4?
H
H
C H
H
tetrahedral,
4 equivalent bonds
H
4H + C Æ H C H
H
C
2s
2p
promote
electron
2s
2p
???
36
VBT
But what about CH4?
H
H
C
2s
promote
electron
2p
tetrahedral,
4 equivalent bonds
C H
H
hybridize
2s
2p
H
4H
sp3 hybrid a.o.s:
H
C(sp3)
tetrahedral
sp3
hybrid
a.o.s
C H
H
σ(sp3C + 1sH)
VBT
BeH2 facts: H
Be
H 2 equivalent bonds
linear, 180°
Be
2s
2p
promote
electron
atomic configuration
can't form two bonds
sp hybrid a.o.s:
hybridize
2s
2p
two bonds, but not
equal, not 180°
and
2p
180° !
= H Be
H(1s) Be(sp) H(1s)
(linear)
atomic orbitals
sp
hybrid
a.o.s
H
σ(spBe + 1sH)
molecular orbitals
37
VBT
H
BH3 facts:
H
B
2s
B
trigonal planar,
3 equivalent bonds
H
promote
electron
2p
hybridize
2p
2s
sp2
hybrid
a.o.s
sp2 hybrid a.o.s:
2p
H
3H
=
H
B(sp2)
σ(sp2
trig. plan.
B
B
H
+ 1sH)
VBT
But what about CH4?
H
H
C
sp3
2s
2p
promote
electron
tetrahedral,
4 equivalent bonds
C H
H
hybridize
2s
2p
H
4H
hybrid a.o.s:
C(sp3)
tetrahedral
H
C H
H
sp3
hybrid
a.o.s
σ(sp3C + 1sH)
38
VBT
NH3
N
2p
2s
sp3
lone pair in sp3 a.o.
sp3 hybridized
N H
H
H
H2O
O
σ(sp3N + 1sH)
2p
2s
sp3 hybridized
sp3
lone pairs in
O H
H
σ(sp3
O
sp3
a.o.s
+ 1sH)
VBT
PCl5
3p
3s
sp3d
3s
Cl
Cl P Cl
Cl
hybrid a.o.s:
3s
3p
3s
sp3d
Cl
trigonal
bipyramidal
SF6
3d
3p
3d
3p
sp3d2
F
sp3d2
F
F
hybrid a.o.s:
octahedral
S
F
F
F
39
VBT
Summary:
e– pair geometry
linear
sp
trigonal planar
sp2
tetrahedral
sp3
trigonal bipyramidal
sp3d
octahedral
sp3d2
VBT
Multiple bonds
hybridization
trigonal planar = sp2
C2H4 facts: H
double bonds
H
H
1s
C
H
H
1s
sp2
2p
C
H
H
H
1s
1s
2p
sp2
π bond
2p
H
all six atoms lie
in same plane
H
C C
C C
H
H
overlap
p orbitals
σ(sp2C + 1sH)
2
σ(sp C + sp2C)
H
H
C C
H
H
H
all atoms coplanar
for p orbital overlap
H
C
=
H
C
H
double bond =
1 π bond +
1 σ bond
40
VBT
Multiple bonds
C2H2 facts: H
triple bonds
H
C H
H
1s
C
2p
sp
C
linear = sp
1s
2p
sp
2 π bonds
2p
H
C
C C
C
H
H
C
H = H
σ(spC + 1sH)
C
C
H
triple bond =
2 π bonds +
1 σ bond
σ(spC + spC)
VBT
What is the hybridization of each indicated atom in the following
molecule? How many sigma and pi bonds are in the molecule?
H
O
H
C O C H
C C
H
H
C
N
41
VBT
What is the hybridization of each indicated atom in the following
molecule? How many sigma and pi bonds are in the molecule?
sp3
sp2
sp2
H
O
H
sp3
C O C H
C C
H
H
C
sp2
N
sp
12 σ bonds and 4 π bonds
How many lone pair electrons?
VBT
Delocalized π bonding
all sp2, trigonal planar
(120º, planar hexagon)
benzene, C6H6
H
H
H
H
H
H
=
H
H
H
H
H
H
6 e– in 6 p orbitals
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