Download Valence Bond Theory

How does bonding take place with orbitals described
by Quantum Theory?
Learning Goals
 Students will be able to:
1) understand how bonding occurs with respect to
the Quantum Mechanical Model
2) understand the concept of Valence Bond Theory
Success Criteria
 Student will show their understanding by:
1) explaining how half-filled orbitals from two
different atoms overlap to form covalent bonds
2) explaining how hybrid orbitals (sp, sp2, sp3 ,sp3d,
sp3d2) are needed to explain bonding with atoms
such as carbon.
3) explaining how double and triple bonds are
formed.
Pre-requisite Knowledge
 Student will already know:
1) shapes of orbitals (Quantum Mechanical Model)
2) Energy Diagrams
Valence Bond Theory
 Linus Pauling was a long-time
friend of Gilbert Lewis. He
dedicated a famous textbook, The
Nature of the Chemical Bond, to Lewis.
 Pauling is one of only four two-time
winners of a Nobel Prize—his in
two different fields, Chemistry
(1954) and Peace (1962).
 The orange is part of a personal
experiment in which he believed
that high doses of vitamin C could
rid the body of free radicals (ions in
the body) that would damage
tissues and lead to aging. His quest
for longer life was fairly successful he lived to 93.
Quantum Mechanical Model
 We have seen how Quantum
Mechanical Theory has
changed how we think of
electron orbitals.
 The Bohr-Rutherford model
that we learned in Grade 9 and
10 has been replaced.
 Orbitals are now regions where
the electron has a probability of
occurring – more like a cloud
where the electron can exist.
 Now we have to move from
physics to chemistry!
s, p, d, and f orbitals
Simpler Bohr-Rutherford
Diagram
Bonding & the Quantum Mechanical Model
 We need to talk about how




elements actually bond.
The Bohr-Rutherford model helped
us explain bonding and the
combining capacities of elements
(for example Ca wants to give away
2 electrons and Cl wants to gain 1
electron)
Now we are interested in a more
accurate picture of bonding
How are electrons shared between
elements (and these odd shaped
orbitals)?
How can we use this knowledge to
infer the shapes of molecules
 According to valence bond theory, a
covalent bond is formed when two orbitals
overlap (share the same space) to produce a
new combined orbital containing two
electrons of opposite spin.
 This arrangement results in a decrease in the energy
of atoms forming the bonds.
 Great Videos:
http://www.youtube.com/watch?v=vVEqx9gSl3Y
Valence Bond Theory I: Intro to Valence Bond
Theory – Ben’s Chem Videos (start here)
 http://www.youtube.com/watch?v=mTHW9W-2-N8
sp3, sp2, and sp Hybridization – Ben’s Chem
Videos
Figure 1: The formation of a single covalent bond in a
hydrogen molecule by the overlap of
two 1s orbitals of individual atoms. The two shared
electrons spend most of their time between the two
hydrogen nuclei. This represents a new, lower-energy
state of the two atoms.
 Notice that the new combined orbital contains a pair
of electrons of opposite spin, just like a filled orbital.
 Any two half-filled orbitals can overlap in the same
way. The total number of electrons in the bonding
orbital must be two.
Figure 2: Consider hydrogen fluoride, a hydrogen atom
has only one occupied orbital, the 1s orbital. The
hydrogen 1s orbital is believed to overlap with the halffilled 2p orbital of the fluorine atom to form a covalent
bond.
 This approach can also be used for larger molecules.
 An oxygen atom has two half-filled p orbitals.
 It is reasonable to propose that the 1s orbitals of the
two hydrogen atoms overlap with the two half-filled
2p orbitals of the oxygen atom to produce a stable,
lower-energy state.
 Figure 3:
The two covalent bonds are created by two
sets of combined s-p orbitals. The measured angle be the
H-O-H bond is about 105° - this is explained by
additional manipulation of atomic orbitals and a
consideration of the repulsions between pairs of
electrons.
 Overall, when atoms bond, they arrange themselves in
space to achieve the maximum overlap of their halffilled orbitals. Maximum overlap produces a bonding
orbital of lowest energy.
Problems with Lewis Bonding Theory
Problem 1
 Could not explain the four equal bonds represented
by the four pairs of electrons in a carbon compound
like methane, CH4(g)
Problem 2
 Could not explain the existence of double and triple
bonds.
Solution to Problem 1
Explaining equal bonds in methane
 Pauling and others created the idea of electron
promotion from an s to an empty p orbital
 Experimental evidence indicated that the electron
orbitals were equivalent in shape and energy
 The four bonds for carbon in molecules such as
methane are explained by hybridization to four
identical sp3 atomic orbitals

 Great Video: Molecular Shape and
Orbital Hybridization
https://www.youtube.com/watch?v=d1E1
8tBTlBg&feature=related (10:09)
 The two 2s electrons and the
two 2p electrons form four sp3
hybrid orbitals with one
bonding electron in each
 This explains the bonding
capacity of four for carbon
 These orbitals are hybridized
only when bonding occurs to
form a molecule; they do not
exist in an isolated atom.
16
Figure 4: (a) An s electron is
promoted to an empty p orbital
in a carbon atom. (b) The four
orbitals are combined to produce
four hybrid sp3 orbitals. (c) Each
sp3 orbital is equivalent in
energy and shape. Electron
repulsion requires that the
orbitals are as far apart as
possible – pointing to the corners
of a regular tetrahedron.
 Pauling suggested that there were a whole series of
hybridizations that could occur
Solution to Problem 2 - Explaining
Double and Triple Covalent Bonds
 Experimental evidence determined that
three substances can be formed when
two carbons atoms bonded with
hydrogen – C2H6(g), C2H4(g), and
C2H2(g).
 Lewis suggested that between the
carbon atoms there must a sharing of
one, two, and three electron pairs in
order to obtain a stable octet around
the carbon atoms.
 How is it possible that electrons in what we would
predict as being sp3 hybrid orbitals could overlap
not once, but twice or three times with just one other
atom?
 According to the valence bond theory, two kinds of
orbital overlap are possible:
 (1) The end-to-end overlap of s orbitals, p orbitals,
hybrid orbitals, or some pair of these orbitals. This
type of overlap produces a sigma (σ) bond. Think of
sigma bonds as the usual single covalent bonds that
you are used to drawing in structural diagrams.
Figure 5: Sigma bonds form with the
overlap of (a) s orbitals (b) p orbitals
and (c) hybrid orbitals.
 (2) Two orbitals can overlap side by side to form a pi
(π) bond. Pi bonds are the second and third lines in
the structural diagrams for double and triple
covalent bonds.
Figure 6:
P orbitals form
with the side-byside overlap of
orbitals.
Double Bonds
 We have already seen that the orbitals of a carbon
atom can be hybridized to form four sp3 hybrid
orbitals.
 To explain double bonds – The key new idea is a
partial hybridization of the available orbitals
leaving one or two p orbitals with single unpaired
electrons.
Double Bond Example –
Ethene (C2H4(g))
 Suppose that after promoting
an electron in carbon’s 2s
orbital to a 2p orbital, we
form three sp2 hybrid orbitals
leaving one p orbital with a
single electron.
 Still have four orbitals to form
bonds but three of these are
hybrids and one is a
“normal” p orbital.
 Figure 7: Instead of mixing all four orbitals, valence bond
theory suggests that only three are mixed to form sp2 hybrid
orbitals and a unhybridized p orbital for a carbon atom
Figure 8: For this carbon atom, the sp2 hybrids are planar at
120° to each other and the p orbital is at right angles to
the plane of the hybrid orbitals.
 In a molecule of
ethene, the three
hybrid orbitals are
used to form
sigma bonds
between the
carbon atoms and
to the hydrogen
atoms.
Figure 9:
(a) The sigma bond for a ethene molecule
use the sp2 hybrid orbitals.
 The half-filled p
orbitals on each
carbon are
believed to
overlap
sideways to
form a pi bond.
Figure 9:
(b) The two half-filled p orbitals of the
adjacent carbon atoms overlap sideways.
 Notice that the
pi bond is a
region of
electron density
appearing
above and
below the
sigma bond
directly joining
the two carbon
atoms.
Figure 9:
(b) The two half-filled p orbitals of the
adjacent carbon atoms overlap sideways.
 A pi bond is a
combined
orbital
containing a
pair or
electrons of
opposite spin.
Figure 9:
(b) The two half-filled p orbitals of the
adjacent carbon atoms overlap sideways.
 The additional shared
pair of electrons in the
pi bond provides
greater attraction to the
two carbon nuclei,
which explains why the
double covalent bond is
shorter and stronger
than a single bond.
Figure 9:
(c) The complete bonding orbitals
for a ethene molecule.
The Triple Bond
Learning Checkpoint
 IN SUMMARY - we have a theory to explain how
covalent bonding occurs but more intriguingly, we
have a theory that can explain the SHAPES OF
MOLECULES (the concept to be explored next class)
 REVIEW: Read p. 239 and explaining how a triple
bond forms in ethyne (using the valence bond theory)
 Assigned Questions
 p. 232 Practice UC # 1, 2, 3, 4
 p. 235 Practice UC # 8, 9, 10, 11
 p. 239 Practice UC # 18, 19, 21
36