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Chemical Bonding
There are basically two types of
chemical bonds:
1. Covalent bonds—electrons are
shared by more than one nucleus
2. Ionic bonds—electrostatic attraction
between ions creates chemical bond
Hydrogen bonding and van der Waals’
bonding are subsets of electrostatic
bonding
The Octet Rule
Atoms want to have a filled valence
shell—for the main group elements, this
means having filled s and p orbitals.
Hence the name “octet rule” because
when the valence shell is filled, they
have a total of eight valence electrons.
We use “dot structures” to represent
atoms and their electrons.
Dots around the elemental symbol
represent the valence electrons.
Examples
Hydrogen (1s 1 ) H·
Carbon (1s^2 2s 2 2p 2 ) C
Chlorine ([Ne] 3s 2 3p 5 ) Cl
Lewis Dot Structures
Lewis Dot Structures
None of the atoms in the previous
example contained full valence shells.
When creating bonds, atoms may share
electrons in order to complete their
valence shells.
Lewis Dot Structures
Examples
H 2 molecule:
H needs two e - ’s to fill its valence shell.
Each hydrogen atom shares its electron
with the other in order to fill their
valence shells.
H· + ·H ⇒ H:H
Lewis Dot Structures
Examples
H 2 molecule: H:H
The result is a “covalent bond” in which
two electrons are shared between nuclei
and create a chemical bond in the
process.
When two e - ’s are shared, it makes a
“single” bond.
Lewis Dot Structures
Examples
F 2 molecule:
F has seven valence e - ’s, but wants eight
e - ’s to fill its valence shell.
:F· + ·F: ⇒ :F:F: ⇔ :F – F:
A line is often used to represent shared
electrons.
Lewis Dot Structures
Examples
CH 4 molecule:
C has four valence e - ’s; H has one
valence e -^.
H
·C· + 4 ·H ⇒ H· ·C· ·H ⇒ H : C : H
H
H
H
H – C – H
H
H
K
K
H
H
H· ·C· ·C· ·H ⇒ H : C : C : H
H
H
·^ ··
Lewis Dot Structures
Examples
C 2 H 4 molecule:
Is this complete— do all atoms have filled valence shells?
H : C C : H
H
H
Lewis Dot Structures
Examples
C 2 H 4 molecule:
H : C ··C : H
H
H
·· ⇒^ H – C^ – –C – H
H
H
K
K
Sharing of four electrons between two
nuclei results in a “double” bond.
Examples
NH 3 molecule:
N 3 H
H N H ⇒ H-N-H
Lewis Dot Structures
H H
||
Not all electrons in a Lewis dot structure need to be part of a chemical bond— some electrons may be in the form of “lone pairs”.
Examples
CO molecule:
Lewis Dot Structures
.O...
. C.. ⇒
.O ...
. C..
C has 5 e-’s; O has 7 e - ’s
. C..
..O..
C has 6 e-’s; O has 8 e - ’s
⇒ :C ...... O: ⇒^ :C-O:--
CO has a covalent triple bond—6 e-’s are shared to form the bond
Suppose you can draw two different Lewis structures that are equally valid (same formal charges). Which structure is correct? Example: NO 2
Resonance Structures
structure 1
each N-O bond is really a 1.5 bond, not a single or double bond
:O.. 2 :
N
:O 1 :
:O 2 :
N
:O 1 :
resonance structure
.. :O
..^2 :
N
:O 1 :
.. structure 2
Electronegativity and Bonding
When nuclei share electrons to form a
covalent chemical bond, the electrons are not necessarily shared equally—a shared electron may spend more time closer to one of the nuclei.
The electronegativity of the nuclei determines
how the electron is shared.
Electronegativity is a measure of how strongly
a “bound” electron participating in a chemical bond is attracted to a nucleus.
Electronegativity and Bonding
Electronegativity is related to electron affinity
and ionization energy.
Electronegativity (denoted by the greek
symbol χ) is highest for elements in the upper right hand side of the Periodic Table and increases from left to right and from bottom to top.
Electronegativity and Bonding
Polar Covalent Bonds
When two elements with different
electronegativities bond, the resulting covalent bond will be polar, i.e., the shared electrons will spend more time closer to the nucleus with the higher χ, so that end of the bond will be slightly more negative, and the other end will be slightly more positive.
H - F
χ=2.1 χ=4.
δ+ δ-
Polar Covalent Bonds
The molecule has a polar bond meaning the
electrical charge is not equally distributed between the nuclei involved in the chemical bond. The molecule also has a dipole moment—an uneven distribution of electrical charge.
H - F
χ=2.1 χ=4.
δ+ δ-
Polar Covalent Bonds
Other Examples Water: H 2 O is a bent molecule with two pairs of unshared e-’s in p orbitals. Is water a polar molecule?
O
H H
.. ..
χ(O) = 3.
χ(H) = 2.
δ+
δ- δ-
δ+
Polar Covalent Bonds
Examples Carbon monoxide: CO Is CO a polar molecule?
χ(O) = 3.
χ(C) = 2.
δ+ δ- :C ≡ O:
Polar Covalent Bonds
Examples Carbon dioxide: CO 2 is a linear triatomic molecule. Is CO 2 a polar molecule?
χ(O) = 3.
δ+ δ-^ χ(C) = 2.
O = C = O
..
.. ..
δ- δ+ ..
The bonding between atoms can have a
significant effect on the bond distance between atoms.
Multiple bonds between two atoms have
shorter bond lengths compared to single bonds involving the same elements:
Bond Lengths
Examples average bond lengths single bond C-C 154 pm C-O 143 pm double bond C=C 133 pm C=O 120 pm triple bond C≡C 120 pm C≡O 113 pm
Bond Lengths
The bond energy is the amount of
energy it takes to pull two atoms apart
and break the chemical bond.
For some diatomic gas phase species,
we know the bond energies exactly
through measurement.
Bond Energies
Reaction Energies
Energy
reacants
C 2 H 2 + O 2
CO 2 + H 2 O products
2435 kJ
Exothermic reaction (^) Example:
6 H 2 O + 2 N 2 → 4 NH 3 + 2 O 2 Which bonds are broken? 12 x O-H 12 (460 kJ mol-1^ ) = 5520 kJ 2 x N≡N 2 (945 kJ mol-1^ ) = 1890 kJ total energy needed = 7410 kJ
Reaction Energies
Example: 6 H 2 O + 2 N 2 → 4 NH 3 + 3 O 2 Which bonds are formed? 12 x N-H 12 (390 kJ mol-1^ ) = 4680 kJ 3 x O=O 3 (495 kJ mol-1^ ) = 1485 kJ total energy released = 6165 kJ ΔE (^) rxn = ΣBEreact - ΣBE (^) prod = 7410 kJ – 6165 kJ = 1245 kJ (1267 kJ actual)
Reaction Energies Reaction Energies
Energy reacants
H 2 O + N 2
NH 3 + O 2 products
1245 kJ
Endothermic reaction
Molecular Orbitals
When atomic orbitals (AO’s) overlap to create a covalent bond, the result is the formation of molecular orbitals (MO’s). Molecular orbitals define the region of space most likely to contain bonding electrons— MO’s are drawn as 90% electron density contours just as AO’s are drawn 90% electron density contours in atoms.
Molecular Orbitals
Because electrons can be described as waves, when the AO’s overlap, the waves may either constructively interfere or destructively interfere. Constructive interference between the AO’s results in a bonding MO—destructive interference between AO’s results in an anti- bonding MO with a node along the internuclear axis. Anti-bonding orbitals are higher in energy than bonding orbitals.
Molecular Orbitals
Overlap of s-type atomic orbitals to form either bonding or anti-bonding molecular orbitals. Anti-bonding orbitals are designated with an asterisk (*).
Molecular Orbitals
Molecular Orbitals
Molecular Orbitals
A molecular orbital diagram can be constructed from the atomic orbitals of the bonding elements.
H 2 : H· ·H ⇒ H – H
H H
1s 1s σ
↑ σ* ↑ ↑↓
The 1s e - ’s of each H atom go into the lower energy σ MO resulting in the formation of a single σ bond. Configuration: (σ) 2
Molecular Orbitals
He 2 : He: :He
The 1s e - ’s of each He atom fill both the σ MO and the σ* MO. The bonding effect of the σ MO is offset by the anti-bonding effect of the σ* MO—no net bonding is observed. Configuration: (σ)^2 (σ*) 2
He He 1s 1s σ
σ* ↑↓ ↑↓ ↑↓
↑↓
Molecular Orbitals
Bond order is defined in the following way: bond order ≡ ½ (# e - ’s in bonding MO’s)
- ½ (# e - ’s in anti-bonding MO’s) H 2 : bond order = ½ (2) – ½ (0) = 1 H 2 has a single bond He 2 : bond order = ½ (2) – ½ (2) = 0 He 2 has no bond ⇒ compound does not form
F 2 σ* anti-bonding orbital made from s atomic
orbitals
F 2 σ bonding orbital made from s atomic orbitals
Problems with Valence Bond Theory and Lewis Dot Structures
Valence Bond Theory (including Lewis structures and hybrid orbital theories) does an excellent job at explaining the bonding in many chemical systems. It fails miserably in describing delocalized bonding systems and some very simple molecules like O 2.
Problems with Valence Bond Theory and Lewis Dot Structures
Molecular oxygen, O 2 Properties O 2 has a double bond O 2 is paramagnetic (has two unpaired electrons) Predicted Lewis structure:
O = O.. ..
.. .. The Lewis structure correctly predicts a
double bond, but there are no unpaired electrons.
Molecular Orbitals
O 2 :
O
↑↓ 2s
O
↑↓ 2s σs
σs*
↑ 2p
↑ 2p
σp
σp * π*
π
↑↓ ↑ ↑↓↑
Molecular Orbitals
O 2 :
O
↑↓ 2s
O
↑↓ 2s σs
σs*
↑ 2p
↑ 2p
σp
σp * π*
π
↑↓
↑↓
↑↓ ↑↓ ↑↓
↑↓ ↑↓
Bond order = ½(8 – 4) = 2 (double bond)
↑ ↑ ↑ ↑
MO Description of Molecular
Oxygen
Electron configuration of O 2 is:
(σs )^2 (σs *)^2 (σs )^2 (σs )^2 (σp)^2 (π)^4 (π)^2
1s 2s 2p
MO theory correctly predicts that O 2 has a
double bond (bond order = 2) and that O 2
is paramagnetic.
Molecular Orbital Diagram for Heteronuclear
Diatomic Molecule
CO:
C
↑↓ 2s
O
↑↓ 2s σs
σs*
↑ 2p ↑ 2p
σp
σp * π*
π
↑↓ ↑
↑
Molecular Orbital Diagram for Heteronuclear
Diatomic Molecule
CO:
C
↑↓ 2s
O
↑↓ 2s σs
σs*
↑ 2p ↑ 2p
σp
σp * π*
π
↑↓ ↑
↑
↑↓ ↑↓
Molecular Orbital Diagram for Heteronuclear
Diatomic Molecule
CO:
C
↑↓ 2s
O
↑↓ 2s σs
σs*
↑ 2p ↑ 2p
σp
σp * π*
π
↑↓ ↑
↑
↑↓
↑↓
↑↓ ↑↓ ↑↓
bond order = ½(8 – 2) = 3 (triple bond)
