httpswww.campusvirtual.uniovi.espluginfile.php814526mod_resourcecontent1Lesson203.1.pdf.pdf

Lesson 3. Molecules: Electronic Structure and Bonding
Lesson 3.1
• Introduction to Chemical Bonding. Types of chemical bonding
• Theory of Lewis. Resonance. Exceptions to the Octet Rule
• Valence-shell electron-pair repulsion theory
• Polar covalent bonds. Electronegativity
• Order, length and strength of chemical bonds
Lesson 3.2
• Valence Bond Theory
• Molecular Orbital Theory
Contents

Introduction to Chemical Bonding. Types of chemical bonding
Types of Chemical Bonding
“Overlapping”
“Sharing electrons”
Covalent Bonds
“oppositely charged ions”
“Networks”
attraction cation-anion releases energy
and stabilize network
Ionic Bonds
“Sea of electrons”
atomic cores = positive nuclei
+
inner shell of electrons
Valence shell electron
Metallic Bonds
Lesson 3

+ +
+
+
+
+
Covalent Bond
Energy
(released)
Chemical Bonding (Covalent)

Distances Binding Enegies Force constant
Measurement of chemical bonding strength

G. Lewis (1875-1946) Langmuir (1881-1957)
N. P. 1932
W. Kopssel (1888-1956)
Lewis Bonding Theory. Covalent Bonds
1) Chemical bonding ↔ Electrons of the valence shell (outermost)
2) Covalent bond ↔ electrons are share between atoms.
3) Atom → Noble gas configuration (eight valence electrons ↔ octet).
Fundamental ideas
Lewis Theory

Lewis Symbols
1) Chemical symbol: nucleus + core electrons (inner-shell)
2) Dots: valence shell electrons.
Si
Si: [Ne] 3s2, 3p2

Lewis structures
Single Bonds
Multiple Bonds

Lewis Bonding Theory. Bond Order (b.o.) and Bond Length
Bond order
• Number of covalent bonds between a pair of given atoms.
• single (bond order = 1), double (bond order = 2), triple (bond order = 3), etc.
• The higher the bond order that is, the more tightly the atoms are held together.
A B A B
A B
Bond Order (A-B) 1 3
2
Bond length d(A-B)single < d(A-B)double < d(A-B)triple
Bond Strength (A-B)single < (A-B)double < (A-B)triple

Lewis Bonding Theory. Polar covalent Bonds. Electronegativity
Electron Density Map
1. Molecule A-A (atoms are identical): bond pairs electrons are equally shared
Non-polar Covalent Bonds
δ-
δ+
Polar Covalent Bonds
2. Molecule A-B (Non identical Atoms): bond pairs electrons are NOT equally shared
Electron Density Map
δ-
δ+

Lewis Bonding Theory. Polar covalent Bonds. Electronegativity
Ability of an atom to attract toward itself the electrons in a chemical bond
Electronegativity (χ)
χ(A) : Relates with Ip(A) and EA(A)
Pauling scale
χA (0.7 – 4.0)
χA increases from left to right across a period in the periodic table (decreasing metallic character ).
χA decreases with increasing atomic number in a group.
Note that the transition metals do not follow these trends.

Lewis Bonding Theory. Writing Lewis Structures.
Writing Lewis Structures
1) Count Valence-Shell electrons
CO2
Atom Valence Shell electrons
C [He]2s2 2p2 → 4 electrons
O(x2) [He]2s2 2p4 → 6 (x2) = 12 electrons
total 16 valence electrons
2) Draw Skeletal structure i) Less electronegative → Central atom
ii) More electronegative → Terminal atom
iii) H is always terminal and C is central
iv) Place 2 electrons (single bond) between atoms
C O
O
χ(C) < χ(O) → C: Central atom
O: Terminal Atom
4 electrons ↔ 12 electrons left
IONS
+
-
= - 1 electron
= + 1 electron

3) Firstly, complete the OCTET of TERMINAL atoms
C O
O
: :
:
:
:
:
3.1. Hydrogen octet = 2 electrons
3.2. Check if all valence electrons are placed in the structure
3.3. If not, place remaining electrons on central atom
16 electrons

3) Check if all atoms meets OCTET criteria
C O
O
: :
:
:
:
:
8 e-
8 e-
4 e-
- If NOT → Use MULTIPLE bonds to complete OCTET
C O
O
: :
:
:
:
8 e-
8 e-
6 e-
C O
O
: :
:
:
8 e-
8 e-
8 e-
C O
O
: :
:
:
8 e-
8 e-
8 e-
Formal Charges

Lewis Bonding Theory. Formal charges
Apparent charges ↔ atoms have not contributed with equal numbers of electrons to the
covalent bonds.
Formal Charges
number of valence
electrons in the free
(un-combined) atom
number of electrons
assigned to that atom
in the Lewis structure
Formal Charge = -
Lone Pairs
Belonging enterely to atom Bond Pairs
Equally shared between atoms
(half contribution)
C O
O
: :
:
:
C O
O
: :
:
:
(6-7 = -1)
(6-5 = +1) (4-4 = 0)
(6-6 = 0)
(6-6 = 0) (4-4 = 0)
(+1) (0) (-1) (0)
(0)
(0)

General rules to determine the plausibility of a Lewis structure based on its
formal charges:
i. The sum of the formal charges in a Lewis structure must equal zero for a neutral molecule
and must equal the magnitude of the charge for a polyatomic ion.
ii. Where formal charges are required, they should be as small as possible.
iii. Negative formal charges usually appear on the most electronegative atoms; positive
formal charges, on the least electronegative atoms.
iv. Structures having formal charges of the same sign on adjacent atoms are unlikely.
Lewis Bonding Theory. Formal charges
Formal Charges

Lewis Bonding Theory. Resonance
δ- δ+ δ-
Resonance Molecule with different “valid” (identical contribution) Lewis struscture
dO-O = 1.28 Å
O O O
Average Bond order = between double and single = 1.5

• Average bond order = [1/3 (1+1+2)] = 1.33
• 3 identical C-O bond distance = 129 pm. Intermediate between a C-O single bond (143
pm) and a C=O double bond (120 pm)
b.o. (C-O) = 1
b.o. (C-O) = 2
Resonance
δ-
δ+
δ-
δ-

Lewis Bonding Theory. Coordinative Covalent Bonds
The Lewis theory of bonding describes a covalent bond as the sharing of a pair of electrons, but this does
not necessarily mean that each atom contributes an electron to the bond. A covalent bond in which a single
atom contributes both of the electrons to a shared pair is called a coordinate covalent bond
Coordinative Covalent Bond
Note that once the coordinative bond has formed, it is impossible to say which of the four bonds is the coordinate covalent bond.
Thus, a coordinate covalent bond is indistinguishable from a regular covalent bond.
δ+ δ-
δ-
δ+
δ+
δ+

Lewis Bonding Theory. Resonance and Formal charges
Sometimes, resonance structures do not contribute equally.
For example, consider the azide anion, N3
-, for which three resonance structures are given
below.
We can decide which resonance structure likely contributes most to the hybrid by applying the
general rules for formal charges.
Major contribution to the Resonance Hybrid

Lewis Bonding Theory. Exceptions to octet Rule
1. Odd-electron Molecules
Lewis theory deals with electron pairs and does not tell us where to put the unpaired electron. Thus, we locate the
unpaired electron following the Formal Charges Rules.
δ- δ+
δ-

2. Incomplete Octet Molecules
In some compounds, the number of electrons surrounding the central atom in a stable molecule is fewer
than eight.
For example, beryllium, [Be] = 1s22s2; it has two valence electrons in the 2s orbital. In the gas phase, beryllium hydride
(BeH2) exists as discrete molecules. The Lewis structure of BeH2 is
Elements in Group 13, particularly boron (B) and aluminium (Al), also tend to form compounds in which they are
surrounded by fewer than eight electrons.
B: 1s22s22p1, it has a total of three valence electrons
δ-
δ+
δ+
δ+
Octet Rule

3. Expanded Valence Shell (Octet)
i) Atoms of the second-period elements cannot have more than eight valence electrons around the central atom
ii) Atoms of elements of the third period and beyond, form compounds in which more than eight electrons surround the
central atom.
In addition to the 3s and 3p orbitals, elements in the third period also have 3d orbitals that can be used in bonding. These
orbitals enable an atom to form an expanded octet (the use of the 3d orbitals for valence-shell expansion is a matter
of scientific dispute)
iii) Molecules with expanded valence shells typically involve:
(a) Central atom: Non-metal atoms of the third period and beyond
(b) Terminal atom: highly electronegative atoms.
For example: PCl5, SbF6, etc.
iv) Expanded valence shells have also been used in cases where they appear to give a better Lewis structure that meet the
Formal Charges Rules
Ej. [SO4]-2

Molecular Geometry: Valence-Shell Electron-Pair Repulsion Model – VSEPR Model
Molecular geometry is the three-dimensional arrangement of atoms in a molecule
The shape of a molecule is established by experiment or by a quantum mechanical calculation confirmed by
experiment. The results of these experiments and calculations are generally in good agreement with the
valence-shell electron-pair repulsion theory (VSEPR). In VSEPR model, we focus on pairs of electrons in the
valence electron shell of a central atom in a structure
Electron Pair Domains
Electron Pair Regions
Electron Pair Groups
Lone Pairs (LP)
Bond Pairs (BP)
Multiple Bond Pairs (MBP)
VSEPR
Molecular geometry is predicted on
the base of the number of electron
pairs GROUPS surround the central
atom
Linear
2 electron groups 3 electron groups 4 electron groups
Trigonal planar Tetrahedral
Trigonal Bipyramidal Octahedral
5 electron groups 6 electron groups

Electron Pair Domains
Electron Pair Regions
Electron Pair Groups
Electron-group geometry
resulting of bond and lone pairs
Molecular geometry
resulting of bond pairs (but influence by electron-
group geometry)
predicted H-N-H angle = 109.5º
experimental H-N-H angle = 107º
Electron-group geometry and Molecular geometry coincide when the central atom has no lone pairs
VSEPR Symbols (Nomenclature)
A = Central atom
X = Ligand (substituent) bonding to central atom
E = Lone pair

VSEPR when Central Atom has no lone pairs electrons
2 electron groups 3 electron groups 4 electron groups
5 electron groups 6 electron groups
AX2
AX3
AX4
AX5
AX6

Predicting Molecular Shapes by VSEPR
1. Write Lewis structure
2. Find the VSEPR Symbols of the structure
3. Find the Electron-group Geometry (bond pairs and lone pairs)
4. Find the Molecular Geometry (bond pairs)
5. Distortions of Molecular Geometry (see later)

Molecular Geometry: VSEPR Model. Applications
2 electron groups (AX2)
3 electron groups (AX3; AX2E)

4 electron groups (AX4; AX3E; AX2E2 )

5 electron groups (AX5; AX4E; AX3E2; AX2E3 )

6 electron groups (AX6; AX5E; AX4E2)

Molecular Geometry: VSEPR Model. Distortions and Limitations
Second Order Effects on VSEPR model
• Repulsion between electron groups: Bond Pairs (BP), Lone Pairs (LP)
and Multiple Bond Pairs (MBP)
• Electronegativity of Substituents
• Combined Effects
• Non-equivalent positions

Molecular Geometry: VSEPR Model. Distortions
Repulsion between electron groups
1. The closer together two electron domains (BP, LP, MBP) are forced, the stronger the
repulsion between them.
2. Lone Pair (LP) electrons spread out more than do Bond Pair (BP). Thus, the ORDER OF
REPULSIVE FORCES, from strongest to weakest:
LP-LP > LP-BP > BP-BP
3. Multiple Bond Pairs (MBP) electrons spread out more than Bond Pairs (BP) but less than Lone
Pairs (LP)

Repulsion between electron groups. Examples
BP
LP
More Repulsions
AX3E AX4
LP
BP
AX3
BP
MBP
AX3
BP-LP
BP-MBP

Repulsion between electron groups. Examples
Electronegativity A-B
A = Central atom; B = Terminal atom
χA < χB
χA ≈ χB
χA > χB
Lower repulsive interactions between BP
Higher repulsive interactions between BP
χF > χCl > χBr > χI
χN > χP > χAs
χF > χH

Non-equivalent axial and equatorial position in Trigonal Bipyramidal
187°
101.4°
1.646 Å
1.545 Å
AX4E

Non-equivalent axial and equatorial position in Trigonal Bipyramidal
AX3E2
T-Shape
AX2E3

Non-equivalent axial and equatorial position in Octahedral
AX5E1 All octahedral positions are equivalent (90º)
AX4E2
AX4E2
Square planar

Molecules with multi-central atoms (skeletal structures)
Determine the geometry around each of the central. The set of all the observed geometries to each
central atom will be combined leading to the description of the geometry of the complete molecule.

POLARITY: Bonding and Molecules. DIPOLE MOMENTS
H F
𝛿+
𝛿−
𝝁 = 𝑸 𝒙 𝒓
Q Partial Charges located loacted on each atomo (δ)
r Distance between centres of positive and negative charges
μ Dipolar moment. Debye (D). 1 D = 3.34 10-30 C m (Coulomb x metre)

Determination of molecular polarity:
1. The polarity of the bonds in the molecule
2. The geometry of the molecule
= 0
Polar molecules: μ ≠ 0
Non-polar molecules: μ = 0

= 0 ≠ 0

resonance structures
b.o. (S-O) = 2
b.o. (S-O) = 1
S
O
O
O
O
2-
Resonance Hybrid
𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑏. 𝑜. (𝑆 − 𝑂) =
𝑁𝑢𝑚𝑏𝑒𝑟 𝑆 − 𝑂 𝑏𝑜𝑛𝑑 𝑝𝑎𝑖𝑟𝑠 𝑖𝑛 𝐿𝑒𝑤𝑖𝑠 𝑠𝑡𝑟𝑢𝑐𝑡𝑢𝑟𝑒
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑆 − 𝑂 𝑠𝑖𝑛𝑔𝑙𝑒 𝑐𝑜𝑛𝑛𝑒𝑐𝑡𝑖𝑜𝑛𝑠
𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑏. 𝑜. 𝑆 − 𝑂 =
6
4
= 1.5
1.49 Å
S=O ~ 1.43 Å
S-O ~ 1.54 Å
Resonance

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