Olmsted Chapter 8 and 9
Compounds and their Bonding
Chemical Bond Formation
• When a chemical reaction occurs
between two atoms, their valence
electrons are reorganized so that a net
attractive force – a chemical bond –
occurs between the atoms
– Ionic – typically metal with a nonmetal
– Covalent – typically a two nonmetals
Ionic bond
• Sodium metal reacts with chlorine gas in a
violently exothermic reaction to produce NaCl
2Na(s) + Cl2(g) -> 2NaCl(s)
• The chlorine has a high affinity for electrons,
and the sodium has a low ionization potential.
• Thus the chlorine gains an electron from the
sodium atom.
• This can be represented using electron-dot
symbols
1
• The arrow indicates the transfer of the electron
from sodium to chlorine to form the Na+ metal
ion and the Cl- chloride ion.
• Each ion now has an octet of electrons in its
valence shell:
Na+ [He] 2s22p6
Cl- [Ne] 3s23p6
• The “bond” is the attractive force between the
positive and negative ions.
The Energy of Ion-Pair Formation
Usually exothermic
– Ionization of atom to form cation (Na+)
– Ionization of atom to form anion (Cl-)
– Pairing of the two to form ion-pair
Na(g) Æ Na+(g) + 1e- ΙΕ = 502 kJ/mol
Cl(g) + 1e-Æ Cl-(g) -ΕΑ = -349 kJ/mol
Na+(g) + Cl-(g) Æ NaCl (g) Eion pair = -552 kJ/mol
Na (g) + Cl (g) Æ NaCl (g) ∆E = -399 kJ/mol
The ions are drawn together, energy is released,
ions form solid lattice
Eion pair = k
(q1)(q2 )
r
• Where
–
–
–
–
k is a constant
r is the distance between the two ions
q1 is the electronic charge one ion
q2 is the electronic charge the other ion
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Lattice Energy
Lattice Energy - the energy required to form a
solid ionic compound from its gaseous ions.
Na+(g) + Cl-(g) Æ NaCl (s) Elattice = -786 kJ/mol
– It is a measure of just how much stabilization results
from the arranging of oppositely charged ions in an
ionic solid.
– The magnitude of the lattice energy depends upon
the charges of the ions, their size and the particular
lattice arrangement.
Non-Polar Covalent Bond
Equal sharing of electrons between two atoms in a molecule
Covalent Bonding and the Octet Rule:
Atoms will share electrons so that each atom
Will fill its valence shell
Bond Length (Bond Distance)
The distance in picometers between the nucleii of two
Bound atoms in a molecule
Bond Energy
The strength of the bond between atoms in a molecule.
The amount of energy (J) required to break the bond.
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Fig 8-3
Pg 330
The interaction energy of a pair of hydrogen atoms varies with
internuclear separation.
Polar Covalent Bond
The unequal sharing of electrons between atoms in a molecule
A δ+ indicates that the atom that doesn’t hog the electron density (less
electronegative) has a partial positive charge. A δ- indicates the
atom that hogs the electrons in the molecule (is more electronegative)
Polar molecules contain dipole moments
Electronegativity
The amount to which an atom attracts electron density
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Covalent Bond - The sharing of
valence electrons between atoms
LEWIS STRUCTURES
1.
Sum the valence electrons from all atoms
– Use the periodic table for reference
– Add an electron for each indicated negative charge,
subtract an electron for each indicated positive charge
2. Write the symbols for the atoms to show which atoms are
attached to which, and connect them with a single bond
– You may need some additional evidence to decide
bonding interactions
– If a central atom has various groups bonded to it, it is
usually listed first: CO32-, SF4
– Often atoms are written in the order of their
connections: HCN
3. Complete the octets of the atoms bonded to
the central atom (H only has two)
4. Place any leftover electrons on the central
atom (even if it results in more than an octet)
5. If there are not enough electrons to give the
central atom an octet, try multiple bonds
–
use one or more of the unshared pairs of electrons
on the atoms bonded to the central atom to form
double or triple bonds
PCl3
Bonding Pairs
Lone Pairs
(a.k.a. nonbonding electrons)
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Let’s Try…
NO+
CH2O
NH3
ClONO2+
PO43-
NH4+
CO
SCNSO42-
Formal Charges
Sometimes, while drawing Lewis structures, you don't have observational data to use.
Formal charges can be used to find out what to choose.
Formal charge = (# of Valence shell e- of an atom) - (# Bond pair e-) - (# of unshared e-)
For example, lets take the first incorrect drawing of sulfuric acid.
Formal charge on S = 6 - 4 - 0 = 2
Formal charge on H = 1 - 1 - 0 = 0
Formal charge on O = 6 - 2 - 4 = 0 (on the ones bonded to H)
Formal charge on O = 6 - 1 - 6 = -1 (on the isolated ones)
We can disregard the ones with 0 formal charge.
The ones that do have formal charge are the sulfur and the isolated oxygens.
Since the 2 oxygens are -1 and sulfur is 2, another bond goes from each oxygen
to the sulfur to cancel the formal charge. (2 + 2(-1) = 0)
When several Lewis structures are possible, those with the smallest formal charges
are the most stable and are preferred.
Resonance
Structures
Sometimes, a single Lewis
structure does not adequately
represent the true structure
of a molecule.
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Resonance Structures
Sometimes, a single Lewis structure does not adequately represent the true structure of
a molecule.
Consider the carbonate ion, CO32carbon (C) has four valence electrons x 1 carbon = 4 eoxygen (O) has six valence electrons x 3 oxygens = 18 eThe ion has an overall negative two charge so we add 2 e- to give a
total of 24 e- to be placed in the Lewis structure.
Carbon is the central atom, the three oxygens are bound to it and electrons
are added to fulfill the octets of the outer atoms.
All the available electrons have been used but carbon is electron deficient - it only has
six electrons around it. So………………………
we share a non-bonding electron pair on an oxygen with the carbon to create a
double bond and thereby fulfill carbon's octet.
becomes
or
becomes
or
becomes
Three equivalent Lewis structures (formal charges are included) can be drawn for
the carbonate ion. The true structure of the carbonate ion is an average of the
three resonance structures connected by a double headed arrow and enclosed in
brackets.
This can also be represented by a resonance hybrid where a dashed line is used to
indicate the delocalized electrons.
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Orbitals and Bonding Theories
Two common approaches to rationalizing
chemical bonding
• Valence bond theory (VB)
Developed by Linus Pauling
• Molecular Orbital theory (MO)
Developed by Robert Mulliken
What are their differences?
VB
MO
• Closely tied to Lewis’
idea of bonding electrons
• Usually the method of
choice to provide a
qualitative, visual
picture
• Particularly good for
molecules made up of
many atoms
• Good for describing
molecules in their
ground state
• Atomic orbitals of the
atoms in the molecule
combine to form a set of
orbitals that are the
property of the orbital
• Gives a quantitative
picture of a molecule
• MO is essential in
describing the bonding
in the excited states of
molecules.
Valence Bond Theory (VB)
Orbital Overlap Model of Bonding
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The idea that bonds are formed
by overlap of orbitals is the basis
of the valence bond theory
The Main Points for the Valence
Bond Approach to Bonding Are:
• Orbitals overlap to form a bond between
two atoms
• Two electrons can be accommodated in the
overlapping of orbitals
• Because of orbital overlap, the bonding
electrons have a higher probability of being
found within a region of space influenced
by both nuclei.
The overlap between two s orbitals is called:
s
s
A Sigma Bond (σ)
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But sigma bonds can also form between
other types of orbitals such as p and d.
H:
↑
1s
F:
↑↓
2s
↑↓
↑↓
2p
↑
The white electrons represent nonbonding electrons
The red electrons represent bonding electrons
The hybridization of the
s and p orbitals
What about trying to draw an overlap for CH4.
C: ↑↓ ↑ ↑ —
2s
2p
We know that carbon makes 4 bonds
and tetrahedral shape, but it only
has two available electronsfor bonding.
How do we solve this problem???
By Hybridizing the s and p orbitals
Add the one s orbital to the 3 p orbitals and get the
sp3 hybrid orbital
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What about the type of hybrid you would need for
a trigonal-planar electron-pair geometry
Consider BF3:
B: ↑↓ ↑ — —
2s
2p
How many bonds around B do we need
To make a trigonal-planar shape?
3
What orbitals would we hybridize to get three identical orbitals?
One s orbital and 2 p orbitals → sp2 hybridized orbital
Hybridization Involving s, p,
and d Atomic Orbitals
What would be the hybridized orbitals for PF5?
sp3d
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Multiple Bonds
Consider ethylene:
Draw the structure: H
H
C=C
H
H
What are the hybridized
orbitals for the C atoms?
energy
sp2
↑
unhybridized p orbital
C: ↑ ↑ ↑
sp2
What about the double bond?
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Molecular Orbital Theory
It is a different way to view orbitals in molecules
*** MO theory assumes that pure s and p atomic orbitals of
the atoms in the molecule combine to produce orbitals that are
spread out over several atoms or even over the entire molecule.
The new orbitals are called molecular orbitals. ***
Principles of MO theory
• 1st – the total number of molecular orbitals
will be equal to the number of atomic
orbitals
• 2nd – the bonding molecular orbital is lower
in energy than the parent orbitals and the
anitbonding orbital is higher in energy
• 3rd – electrons of the molecule are assigned
to orbitals of successively higher energy.
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First Principle of MO Theory
Two molecular orbitals result from the addition and
subtraction of the overlapping orbitals
Consider Hydrogen
BONDING
Addition of the atomic orbitals → increased probability
that electrons will reside
in the bond region between
the two nuclei
ANTIBONDING
Subtraction of the atommic orbitals → decreased probability
that electrons will reside
between the nuclei, but
a higher probability of
finding the electron in
other regions.
This also represents the second principle of MO theory- antibonding
is always higher in energy
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Third principle states
that the electrons fill
from the bottom on up
to the top.
Consider the bonding of
two lithium atoms
The Fourth Principle of MO theory
Atomic orbitals combine to form molecular orbitals
most effectively when the atomic orbitals are of similar
energy
This becomes important when we move past He2 to Li2
The molecular orbitals
come from 1s ± 1s
and 2s ± 2s
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Bond Order
Bond order = ½ (number of electrons in bonding MOs –
number of electrons in antibonding MOs)
Molecular Orbitals from Atomic
p Orbitals
Three types of interactions can occur for two atoms that
both have s and p orbitals
• Sigma bonding and antibonding are formed by s orbitals
• p orbitals head-to-head interaction
• p orbitals interacting sideways
Sigma bonding and antibonding are formed by s orbitals
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p orbitals head-to-head interaction
p orbitals interacting sideways
So, what would the molecular orbital
look like?
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