# week 4 notes.docx

7/28/2019 # Week 4 Notes.docx

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PREDICTING THE GEOMETRY OF MOLECULES AND POLYATOMIC IONS

The valence-shell electron-pair repulsion model (VSEPR) can be used to predict the shapes of many

molecules and polyatomic ions. The model, however, provides no information about bond lengths or the

presence of multiple bonds (Lewis structures do that).

Because electrons repel each other electrostatically, the most stable arrangement of electron groups (i.e.,

the one with the lowest energy) is the one that minimizes repulsions. Groups are positioned around the

central atom in a way that produces the molecular structure with the lowest energy.

The molecular geometry can be described through the VSEPR model by using the following procedure:

1. Draw the Lewis electron structure of the molecule or polyatomic ion. Each group around the centralatom is designated as a bonding pair (BP) or a lone pair (LP).

2. Determine the electron group arrangement around the central atom that minimizes repulsions (Figure9.2).

3. Assign an designation, where is the central atom, is a bonded atom, is the number ofbonded atoms, is a nonbonding valence electron group and is the number of nonbonding valence

electron groups. Identify LPLP, LPBP and BPBP interactions and predict deviations from ideal bond

angles (Figure 9.3).

Because lone pairs are not shared by two nuclei, they occupy more space near the central atomthan bonding pairs. Thus, bonding pairs and lone pairs repel each other electrostatically in the

order BPBP < LPBP < LPLP. This causes bond angles to be smaller than expected.

Multiple bonds also occupy more space around the central atom than single bonds. This is becausea multiple bond has a higher electron density than a single bond, so its electrons occupy more

space than those of a single bond. Multiple bonds also cause bond angles to be smaller than

expected.

4. Describe the molecular geometry.


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The VSEPR model can also be used to predict the structure of somewhat more complex molecules with no

single central atom by treating them as linked fragments.

MOLECULAR DIPOLE MOMENTS

Dipole moments are vectors they possess both magnitude and direction. The dipole moment of a

molecule is the vector sum of the dipole moments of the individual bonds in the molecule. If the individual

bonds cancel each other out, there is no net dipole moment and the molecule is non-polar.

HYBRID ATOMIC ORBITALS: A LOCALIZED BONDING APPROACH


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According to quantum mechanics, bonds form between atoms because their atomic orbitals overlap, with

each region of overlap accommodating a maximum of two electrons with opposite spin, in accordance with

the Pauli Exclusion Principle. Maximum overlap occurs between orbitals with the same spatial orientation

and similar energies.

Using atomic orbitals to predict the stability of bonds forms the basis for a description of chemical bonding

known as the valence bond theory, which is built on two assumptions:

1. The strength of a covalent bond is proportional to the amount of overlap between atomic orbitals. Thegreater the overlap, the more stable the bond.

2. An atom can use different combinations of atomic orbitals to maximize the overlap of orbitals used bybonded atoms.

Atomic orbitals close in energy can hybridize. Electrons in filled subshells can be unpaired by a

combination of processes called promotion and hybridization. Hybridization produces hybrid orbitals and

the unpaired electrons in those orbitals can now form bonds with single electrons in other atoms. The

hybrid orbitals are degenerate.

Because both promotion and hybridization require an input of energy, the formation of a set of singly

occupied hybrid atomic orbitals is energetically uphill. The overall process of forming a compound with

hybrid orbitals will energetically favourable only if the amount of energy released by the formation of

covalent bonds is greater than the amount of energy used to form the hybrid orbitals.

FORMATION OF HYBRID ORBITALS


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FORMATION OF HYBRID ORBITALS

HYBRIDIZATION SUMMARY


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MOLECULAR ORBITALS: A DELOCALIZED BONDING APPROACH

The positions and energies of electrons in molecules can be described in terms ofmolecular orbitals (MOs)

spatial distributions of electrons in a molecule that is associated with a particular orbital energy.

Molecular orbitals are not localized on a single atom (unlike atomic orbitals) but extend over the entire

molecule. Consequently, the molecular orbital theory is a delocalized approach to bonding.

In the molecular orbital approach, the overlapping atomic orbitals are described by mathematical

equations called wave functions. Two atomic orbitals (one from each atom) interact to form two new

molecular orbitals, one produced by the sum of the two wave functions and the other produced by taking

their difference.

Adding two atomic orbitals corresponds to constructive interference between two waves, thus

reinforcing their intensity: the internuclear electron probability density is increased.

Conversely, subtracting one atomic orbital from another corresponds to destructive interference

between two waves, which reduces their intensity and causes a decrease in the internuclear electron

probability density, which results in a node where the electron density is zero.


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Because electrons in the sigma orbital interact simultaneously with both nuclei, they have a lower energy

than electrons that interact with only one nucleus. This means that the sigma orbital has a lower energy

than either of the atomic orbitals.

Conversely, electrons in the sigma star orbital interact with only one nucleus at a time. In addition, they are

farther away from the nucleus than they were in the parent atomic orbitals. Consequently, the sigma star

molecular orbital has a higher energy than either of the atomic orbitals.

The (bonding) molecular orbital is stabilized relative to the atomic orbitals, and the (antibonding)

molecular orbital is destabilized by about as much as the bonding orbital is stabilized.

The energy-level diagram is filled with valence electrons according to the Pauli Exclusion Principle and

Hunds Rule: each orbital can accommodate a maximum of two electrons with opposite spins (but

preferable unpaired), and the orbitals are filled in order of increasing energy.


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MOLECULAR ORBITALS FORMED FROM ATOMIC ORBITALS


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The molecular orbitals formed between two atomic orbitals or between two atomic orbitals have

spherical symmetry along the interaction axis (the z-axis) and are called sigma () orbitals.

The remaining orbitals on each of the two atoms, and , do not point directly toward each other.

Instead, they are perpendicular to the internuclear axis. Although these two pairs are equivalent in energy

(degenerate), the orbital in one atom can interact with only the orbital on the other, and the orbitals interact only with each other as well. These interactions are side-to-side, rather than the head-to-

head interactions characteristic of orbitals. The resulting molecular orbitals do not have spherical

symmetry along the internuclear axis and are called pi () orbitals.


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BOND ORDER IN MOLECULAR ORBITAL THEORY

In molecular orbital theory, electrons in antibonding orbitals effectively cancel the stabilization resulting

from electrons in bonding orbitals. Consequently, any system that has equal numbers of bonding and

antibonding electrons will have a bond order of 0, and it is predicted to be unstable and therefore not to

exist in nature.

Molecules with fractional bond orders are predicted to exist, but are to be highly reactive.

FILLING MOLECULAR ORBITAL ENERGY-LEVEL DIAGRAMS FOR HOMONUCLEAR DIATOMIC MOLECULES

Diatomic molecules with unpaired electrons in their molecular orbitals are paramagnetic.

MOLECULAR ORBITALS FOR HETERONUCLEAR DIATOMIC MOLECULES


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POLYATOMIC SYSTEMS WITH MULTIPLE BONDS


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INTERMOLECULAR FORCES

Intermolecular forces are forces required to break apart molecules from one another, as opposed to

intramolecular forces, which are forces required to break apart bonds within molecules. Since

intermolecular forces are much weaker, they determine bulk properties, such as melting and boiling points.

Intermolecular forces are electrostatic in nature, that is, they arise from the interaction between positively

and negatively charged species. Like covalent and ionic bonds, intermolecular interactions are the sum ofboth attractive and repulsive components. Because electrostatic interactions fall off rapidly with increasing

distance between molecules, intermolecular interactions are most important for solids and liquids, where

the molecules are close together.

DIPOLE-DIPOLE INTERACTIONS

If the structure of the molecule is such that the individual bond dipoles do not cancel one another, then the

molecule has a net dipole moment. Such molecules orient themselves in a substance to maximize the

attraction.


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Substances comprised of molecules with higher dipole moments (and correspondingly higher energies of

interaction) have higher boiling points, because of the extra force of attraction.

LONDON DISPERSION FORCES

Temporary fluctuations in the electron distributions within atoms and nonpolar molecules could result in

the formation of short-lived instantaneous dipole moments, which produce attractive forces called London

dispersion forces between otherwise nonpolar substances. The instantaneous dipole moment on one atom

can interact with the electrons in an adjacent atom, pulling them toward the positive end of the

instantaneous dipole and repelling them from the negative end. The net effect is that the first dipole causes

the temporary formation of a dipole, called induced dipole, in the second. Interactions between these

temporary dipoles cause atoms to be attracted to one another.

This effect tends to become more pronounced as atomic and molecular masses increase. The reason for

this trend is that the strength of London dispersion forces is related to the ease with which the electron

distribution in a given atom can be perturbed. The ease of deformation of the electron distribution in an

atom or molecule is called its polarizability ().

Polarizability:

Increases as volume occupied by electrons increases (larger atom or molecule size)


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Increases with number of electrons (heavier elements) Decreases as electrons are more tightly bound (more filled orbitals)

HYDROGEN BONDING

The bonds between hydrogen and highly electronegative elements such as O, N and F are highly polar. The

large difference in electronegativity results in a large partial positive charge on hydrogen and a

correspondingly large partial negative charge on the O, N or F atom. Consequently, HO, HN and HF

bonds have very large bond dipoles that can interact strongly with one another. Because a hydrogen atom

is so small, these dipoles can also approach one another more closely than most other dipoles. Thecombination of large bond dipoles and short dipole-dipole distances results in very strong dipole-dipole

interactions called hydrogen bonds.


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Because of the large bond dipoles, partially positively charged hydrogen atoms form hydrogen bonds with

the lone (non-bonding) pairs of electrons on the partially negatively charged O, N or F atoms.

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