Chapter 1: Orbitals And Bonding
Chapter 1 Topics: Bonding Concepts
Look back at “Chemistry, The Central Science” by Brown, Lemay, Bursten!
Ionization Potential (Ch. 8)
Ionic Bonds (Ch. 8)
Hund’s Rule (Ch. 8)
Aufbau Principle (Ch. 8)
Lewis Structures (Ch. 8)
Electronegativity (Ch. 8)
Octet Rule (Ch. 8)
Bond Dissociation Energy (Ch. 8)
Covalent Bonds (Ch. 8/9)
Wave Functions (Ch. 9)
Quantum Numbers (Ch. 8)
Electronic Configuration (Ch. 8)
Pauli Exclusion Principle (Ch. 8)
Atomic Orbitals (Ch. 8)
Dipole Moment (Ch. 8)
Valence Electrons (Ch. 8)
Resonance Structures (Ch. 8)
Formal Charge (Ch. 8)
Nodes (Ch. 8/9)
Molecular Orbitals (Ch. 9)
Orbital Nomenclature
Organic Chemists usually don’t use quantum numbers – but we have to remember the correlations:
Higher Energy
principal quantum
number (n)
azimuthal quantum
number (l)
Shell 1:
One S orbital
(1s)
Shell 2:
One S, three P
(2s, 2p)
Shell 3:
One S, three P, five D
Shell 4:
One S, three P, five D, seven F (4s, 4p, 4d, 4f)
(3s, 3p, 3d)
The three P orbitals are designated Px, Py, Pz (different spacial orientations).
magnetic quantum
number (m)
Ionic Bonding
Ionic bonds: One atom transfers electron to another. Molecule held together by electostatic (magnetic)
forces. Formed between two atoms of very different electonegativities (>2.0 electronegativity difference)
Li
Loss of one electron will
lead to a completely
empty valence shell
F
Li
F
or
LiF
Addition of one electron
will lead to a completely
filled valence shell (full
octet)
Atoms are especially stable when all of the valence orbitals are either completely filled or
completely empty (the "noble gas" configuration). This has been adapted to the octet rule: (most)
atoms are stable when there are 8 electrons in their outermost (valence) shell. For this course: 1st &
2nd row atoms can never have more than 8 valence electrons and/or 4 valence orbitals!!!
Electronegativity
Electronegativity and Percent Covalency
Covalent Bonding
Covalent bonds: Two atoms share electrons. Both atoms can count the shared electrons toward their octet. This type
of bond is formed between two atoms of similar electonegativities (<2.0 electronegativity difference)
H
+
Sharing one additional
electron will lead to a
completely filled valence
shell
H
Sharing one additional
electron will lead to a
completely filled valence
shell
H H
or
H
H
or
A line (bond) signifies
2 shared electrons
Both hydrogens have a filled
valence shell (shared
electrons count for both atoms)
H2
Polar Covalent Bonding
H F
or
H
!+
H
Electronegativity 2.1
or
F
HF
!–
F
Electronegativity 4.0
Fluorine has a higher electronegativity, and will "pull" electrons toward itself causing bond polarization. This creates
a dipole along the bond axis.
Identical atoms will share electrons equally: a nonpolar covalent bond.
Nonidentical atoms will not share electrons equally: a polar covalent bond.
Writing Lewis Structures
Lewis Structures: Represent connectivity of a chemical species. Dots reach represent one
electron; lines represent a shared electron pair; atomic symbols represent the nucleus and
all non-valence electrons. Nonbonding electron pairs are frequently omitted!
Writing Lewis Structures
Lewis Structures: Represent connectivity of a chemical species. Dots reach represent one
electron; lines represent a shared electron pair; atomic symbols represent the nucleus and
all non-valence electrons. Nonbonding electron pairs are frequently omitted!
H
H
H
NH3
N
ammonia
H
N H
H
H
or
N
H
H
Writing Lewis Structures
Lewis Structures: Represent connectivity of a chemical species. Dots reach represent one
electron; lines represent a shared electron pair; atomic symbols represent the nucleus and
all non-valence electrons. Nonbonding electron pairs are frequently omitted!
H2CCH2
H
C
C
H
H C C H
H H
ethylene
H
H
H
H C
H
C H
H
or
H
C
H
C
H
Writing Lewis Structures
Lewis Structures: Represent connectivity of a chemical species. Dots reach represent one
electron; lines represent a shared electron pair; atomic symbols represent the nucleus and
all non-valence electrons. Nonbonding electron pairs are frequently omitted!
O
O
H
H3CS(O)CH3
H
dimethylsulfoxide
H
C
S
S
H
C
H
H
H
H
H
O
C
H
H C S C H
H
H
H
H
C
H
H
H
H
C
O
S
H
Formal Charge = Valence e– – Nonbonding e– – 1/2 Bonding e–
H
C
H
H
Atomic Orbitals: A Brief Review
node (nodal plane)
A 1s orbital
A 2p orbital
Atomic Orbitals: A Brief Review
In General, electrons are lower in energy if:
1. They are closer to a positive charge
2. They are in an orbital with fewer nodes
3. They have a larger orbital space to occupy
Molecular Orbitals: A Brief Review
H
H
H
or
H
H
H
Mixing orbitals with same phase: leads
to a bonding interaction that is
stabilizing (lower in energy than the
atomic orbitals)
H
H
Mixing orbitals of opposite phase: leads
to an antibonding interaction that is
destabilizing (higher in energy than the
atomic orbitals)
Molecular Orbitals: A Brief Review
node (nodal plane)
!* H–H
H
H
H
a !*-orbital (antibonding)
H
H
H
! H–H
a !-orbital (bonding)
Molecular Orbitals: A Brief Review
Energy
!* H–H
H
H
H
a !*-orbital (antibonding)
H
52 kcal/mol
H
H
a !-orbital (bonding)
! H–H
Electrons are 52 kcal/mol (per electron) more stable in a σ H–H orbital (larger orbital space)
than in a hydrogen 1s orbital.
Molecular Orbitals: A Brief Review
+
H
H
H
bond formation
H
H
bond cleavage
(BDE)
H
exothermic by 104 Kcal/mol (!H = –104 Kcal/mol)
H
+
H
endothermic by 104 Kcal/mol (!H = +104 Kcal/mol)
The bond dissociation energy (BDE) is the energy required to break a bond homolytically
(into a diradical). BDE (H2) = 104 Kcal/mol
Molecular Orbitals: A Brief Review
C
C
C
or
C
C
Mixing orbitals with same phase: leads
to a bonding interaction that is
stabilizing (lower in energy than the
atomic orbitals)
C
C
C
Mixing orbitals of opposite phase: leads
to an antibonding interaction that is
destabilizing (higher in energy than the
atomic orbitals)
Molecular Orbitals: A Brief Review
!* C–C
C
C
C
a !*-orbital (antibonding)
C
a !-orbital (bonding)
C
C
! C–C
All mutiple bonds that we will encounter will be !-type
Curved Arrow Notation
Curved arrows are used to designate the movement or flow of electrons.
the electrons
start here (a
filled orbital)
H
O
+
correct
H
H
O
H
O
H
this atom has
the empty
orbital to
receive the
electrons
X
H
O
+
incorrect
H
H
Resonance Structures
Resonance Structures allow Lewis Structures to describe multicenter bonding (more than 2
atoms sharing electrons). There are also situations where resonance structures are used to
show bond polarization.
O3
O
O
O
O
O
O
O
O
O
O
O
O
ozone
O
O
O
O
O
O
+1
A more accurate single representation:
–1/2
O
O
–1/2
O