Textbook Notes (280,000)
CA (170,000)
UTSC (20,000)
Chemistry (300)
Chapter 2

CHMB31 Chapter 2


Department
Chemistry
Course Code
CHMB31H3
Professor
Alen Hadzovic
Chapter
2

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Chapter 2: Molecular Structure and Bonding
Lewis Structures
Covalent Bond: formed when two neighboring atoms share an electron pair.
Single Bond: a shared electron pair, denoted A-B.
Double Bond: two shared electron pairs, denoted A=B.
Triple Bond: three shared electron pairs, denoted AB.
Lone Pair: an unshared pair of valence electrons on an atom.
Atoms share electron pairs until they have acquired an octet of valence
electrons.
oOctet Rule: each atom shares electrons with neighboring atoms to
achieve a total of eight valence electrons.
oLewis Structure: diagram that shows pattern of bonds and lone pairs in
a molecule.
Decide on the number of electrons that are to be included in the
structure by adding together the numbers of all the valence
electrons provided by the atoms.
Write the chemical symbols of the atoms in the arrangement
that shows which atoms are bonded together.
Distribute the electrons in pairs so that there is one pair of
electrons forming a single bond between each pair of atoms
bonded together and then supply electron pairs until each atom
has an octet.
Resonance between Lewis structures lowers the calculated energy of the
molecule and distributes the bonding character of electrons over the
molecule.
oResonance: the actual structure of the molecule is taken to be a
superposition or average of all the feasible Lewis structures
corresponding to a given arrangement.
oResonance Hybrid: the blended structure of two or more Lewis
structures.
Lewis structures with similar energies provide the greatest resonance
stabilization.
oResonance averages the bond characteristics over the molecule.
oThe energy of a resonance hybrid structure is lower than that of any
single contributing structure.
oAll the structures of the same energy contribute equally to the overall
structure.
oThe greater the energy difference between two Lewis structures, the
smaller the contribution of the higher energy structure.
Valence Shell Electron Pair Repulsion (VSEPR): regions of enhanced electron
density (bonding pairs or lone pairs) take up positions as far apart as possible
and the shape of the molecule is identified by referring to the locations of the
atoms in the resulting structure.
oWrite down Lewis structure for the molecule or ion and identify the
central atom.
oCount the number of atoms and lone pairs carried by that atom
because each atom and each lone pair counts as one region of high
electron density.
oNote which locations correspond to atoms and identify the shape of the
molecule.
Stereochemically Inert: lone pairs that do not influence the molecular
geometry and are usually in the non-directional s orbitals.

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Lone pairs repel other pairs more strongly than bonding pairs do.
oLone pair/lone pair > Lone pair/bonding region > Bonding
region/bonding region
oThe greater repelling effect of a lone pair is explained by supposing
that the lone pair is on average closer to the nucleus than a bonding
pair and therefore, repels other electrons more strongly.
oGiven the choice between an axial and an equatorial site for a lone
pair, the lone pair occupies the equatorial site.
oIn a molecule with two adjacent bonding pairs and one or more lone
pairs, the bond angle is decreased relative to that expected when all
pairs are bonding.
oA deficiency of the VSEPR model is that it cannot be used to predict the
ACTUAL BOND ANGLE adopted by the molecule.
Valence Bond Theory
Valence Bond Theory (VB Theory): considers the interaction of atomic orbitals
on separate atoms as they are brought together to form a molecule.
In VB Theory, the wavefunction of an electron pair is formed by superimposing
the wavefunctions for the separated fragments of the molecule.
oWhen the atoms are close, it is not possible to know whether it is
electron 1 that is on A or electron 2.
oOnly paired electrons can contribute to a bond in VB Theory.
oSpin Pairing: VB wavefunction is formed by spin pairing of the electrons
in the two contributing atomic orbitals.
oBond: electron distribution described by the wavefunction.
Molecular Potential Energy Curve: shows the variation of the molecular energy
with internuclear separation.
oIt is calculated by changing the internuclear separation R and
evaluating the energy at each selected separation.
oDeeper the minimum of the curve, the more strongly the atoms are
bonded together.
oThe steepness of the curve shows how rapidly the energy of the
molecule rises as the bond is stretched or compressed.
Electrons in atomic orbitals of the same symmetry but on neighboring atoms
are paired to form σ and π bonds.
oHomonuclear Diatomic Molecules: diatomic molecules in which both
atoms belong to the same element (N2, O2 and H2).
oIf the wavefunction remains unchanged when the bond is rotated
around the internuclear axis, the bond is classified as σ.
oIf the signs of the lobes of the orbital are interchanged when the bond
is rotated through 180° around the internuclear axis, the bond is
classified as π.
Each σ bond in a polyatomic molecule is formed by the spin pairing of
electrons in any neighboring atomic orbitals with cylindrical symmetry about
the relevant internuclear axis.
π bonds are formed by pairing electrons that occupy neighboring atomic
orbitals of the appropriate symmetry.
Two deficiencies of the VB Theory: the failure to account for bond angles and
the valence of carbon because carbon can form 4 bonds but according to this
theory, it can only form 2.
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Promotion: the excitation of an electron to an orbital of higher energy in the
course of bond formation; may occur if the outcome is to achieve more or
stronger bonds and a lower overall energy.
oPromotion is a characteristic feature of carbon and of its congeners in
Group 14.
oBecause the promotion energy is small, the promoted electron leaves a
doubly occupied ns orbital and enters a vacant np orbital, relieving the
electron-electron repulsion it experiences in the ground state.
Hypervalence and octect expansion occur for elements following Period 2.
oHypervalent: species that demand the presence of more than an octet
of electrons around at least one atom (PCl5 and SF6).
oThe rarity of hypervalence in Period 2 may be the geometrical difficulty
of packing more than four atoms around a small central atom and may
have little to do with the availability of d orbitals.
Hybrid orbitals are formed when atomic orbitals on the same atom interfere.
oSpecific hybridization schemes correspond to each local molecular
geometry.
osp^3 Hybrid Orbital (tetrahedral): a hybrid built from 1 s orbital and 3
p orbitals.
oHybrid orbitals have pronounced directional characters, it enhances
the amplitude in the internuclear region due to the constructive
interference between the s and the positive lobes of the p orbitals.
oThe bond strength of the hybrid orbital is greater than for an s or p
orbital alone.
Molecular Orbital Theory
Molecular Orbital Theory (MO Theory): a more sophisticated model of bonding
that can be applied equally successfully to simple and complex molecules.
Molecular orbitals are constructed as linear combinations of atomic orbitals.
oThe square of one-electron wavefunction (ψ^2) gives the probability
distribution for that electron in the molecule.
oAn electron in a molecular orbital is likely to be found where the orbital
has a large amplitude.
oWhen an electron is close to the nucleus of one atom, its wavefunction
closely resembles an atomic orbital of that atom.
oLinear Combination of Atomic Orbitals (LCAO): we can construct a
reasonable first approximation to the molecular orbital by
superimposing atomic orbitals contributed by each atom; we combine
the atomic orbitals of contributing atoms to give molecular orbitals that
extend over the entire molecule.
oψ = cχ (of A) + cχ (of B)
oBasis Set: atomic orbitals χ from which the molecular orbital is built.
oThe coefficients c in the linear combination show the extent to which
each atomic orbital contributes to the molecular orbital.
oThe greater the value of c, the greater the contribution of that atomic
orbital to the molecular orbital.
There is a high probability of finding electrons in atomic orbitals that have
large coefficients in the linear combination.
Each molecular orbital can be occupied by up to two electrons.
One molecular orbital lies below that of the parent atomic energy levels, one
lies higher in energy than they do and the remainder are distributed between
these to extremes.
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