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Topic 12 : Advanced Bonding Concepts. LECTURE SLIDES VSEPR shape theory Bond Polarity Molecular Polarity. Kotz & Treichel, 9.6- 9.8. Molecular and Polyatomic Ion Shapes. Once a Lewis structure is drawn , the three - dimensional geometry of the species can easily

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topic 12 advanced bonding concepts
Topic 12 : Advanced Bonding Concepts
  • LECTURE SLIDES
  • VSEPR shape theory
  • Bond Polarity
  • Molecular Polarity

Kotz & Treichel, 9.6- 9.8

molecular and polyatomic ion shapes
Molecular and Polyatomic Ion Shapes

Once a Lewis structure is drawn, the three -

dimensional geometry of the species can easily

be determined by utilizing the “valence shell electron

pair repulsion theory” called “VSEPR”:

“VSEPR” theory is based on the tendency of negatively

charged regions to repel each other and align as far

apart as possible, resulting in predictable shapes for

any covalently bonded species.

slide3

To utilize “VSEPR”, the number of regions of electron

  • density around the central atom in the species is
  • counted.
  • Count as “one region”:
  • Single Bonds
  • Unshared Pairs
  • Multiple bonds between same two atoms
slide4

Examples of “four regions”:

“two regions”:

“three regions”:

slide5

Basic Shapes predicted by VSEPR:

Two regions:

Three Regions:

Bond Angles

Geometry

slide6

Five Regions:

Four Regions:

Six Regions:

slide8

Before we begin, some guidelines about forming

double and triple bonds in Lewis structures:

C, N, O, S form double and triple bonds and never show incomplete octets (less than 8 e’s)

Metals, metalloids, and halogensdo not as a rule form multiple bonds. Compounds containing these elements will often show an incomplete octet around the central

atom.

slide9

Type One: Two Regions

Examples: BeCl2, CO2, NO2+

slide15

Black orbital

indicates pair of

unshared e’s

NOTE: “molecular geometry” (bonds only): BENT

group work 12 1 geometry 2 3 regions
Group Work 12.1: Geometry, 2,3 Regions

1. Do a Lewis Structure for HCN and CH2O.

2. Draw each molecule “to shape”

3. Describe geometry and bond angles for each

slide17

Type Three: Four Regions

CH2Cl2, NH3, H2O, NH4+

slide20

As is turns out, unshared pairs of electrons around

the central atom are not held in place between two

atoms as bonded pairs are.

They tend to occupy more space and to be somewhat

more “repulsive” than bonded pairs.

When grouped with bonded pairs to tiny atoms like H,

they tend to distort the bond angles, pushing the

bonded pairs closer together.

The bond angles in ammonia are closer to 107o.

slide23

GROUP WORK 12.2: Geometry, 3, 4 Regions

Do Lewis structure and assign shape and bond angles:

CO32-, SiCl4

slide24

Type Four: Five Regions

PF5 , ClF3 ,IF2-, SF4

slide25

Bond angles,

each “axial” F,

90o from trigonal

plane

Bond angles

in “equatorial”

position 120o

slide34

GROUP WORK 12.3: Geometry, 5, 6 Regions

Do Lewis structure and assign shape and bond angles:

ICl4+, XeOF2, ICl4-

Note: O in XeOF2 is equatorial, experimental

evidence

slide35

To see relevance of “shape work”, let’s turn next

to bond and molecular polarity. To help examine

this topic we turn back to the property of

“electronegativity”:

electronegativity
ELECTRONEGATIVITY

The trends in ionization energies and electron affinities

can be thought of as summarized in a single property

called “electronegativity” (en or X).

Electronegativityis a unit-less set of assigned values

on a scale of 0 --> 4 describing the ability of an atom to

attract electrons to itself.

The values reaches a maximum at fluorine, with an X =4.

Nonmetals have thelargest values,metals the lowest.

Noble gases have no assigned Xvalue.

slide37

Most active non-metals

Most active metals

slide39

We have classified bonds “ionic” and “covalent”,

depending on whether electron pairs are shared or

electrons are completely transferred from one atom

to another.

In actuality, there is no sharp dividing line between the

two types but rather a continuum:

Evenly shared

electrons

Unevenly shared

electrons

Transferred

electrons

To determine where a bond lies in this “continuum”,

it is useful to consider the difference in electronegativity

( X) between the two atoms making up the bond:

slide40

When the difference (X) is close to zero, sharing is fairly

even and electrons are not much closer to one atom than

the other. The bonds are considered “non-polar.”

When the difference is above zero to about 1.7, the

electrons are closer to the more electronegative atom

and partial charge buildup, polarization, develops.

When a metal or polyatomic cation is present and the (X) is 1.7or higher,ionic bonding becomes the more likely bond type and valence electrons aretransferredto the more electronegative atom.

slide42

So, we need to consider a third more specialized

type of bond, “the polar covalent bond:”

This type of bond will be the important factor to

be considered when we look at molecular polarity,

which arises from molecular shape and bond

polarity.

The polar molecular in turn will exhibit different

solubilities and boiling points than non polar

molecules.

slide43

Let us consider the bond between H and Cl in a

molecule of hydrogen chloride (only hydrochloric

acid when in water!):

Orbital between H and Cl

E pair closer to Cl,

more electronegative

slide44

The electron cloud from the pair of shared

electrons is more dense closer to the chlorine,

and much less dense closer to the hydrogen.

The bond has become “polarized”: it has developed a region (or “pole”) of partial positive charge buildup and a region (or “pole”) of partial negative buildup.

slide45

Major portion of

electron density

slide46

“partially

negative”

“partially

positive”

Arrow to indicate polar

bond, pointing to more

(-) atom

slide47

The molecule has only one bond, and it is polar.

This makes the entire molecule a “dipole”, one which

has a positive and negative pole and will align in an

electrical or magnetic field:

All diatomic

molecules with

polarized bonding

between the two

atoms are DIPOLES.

molecular polarity larger molecules
MOLECULAR POLARITY, LARGER MOLECULES
  • All diatomic molecules with polar bond(s) are dipoles,
  • but the situation is not so simple for larger molecules.
  • There are two factors to consider:
  • Are the bonds polar?
  • Are they arranged so that the center of positive
  • charge and the center of negative charge do not
  • “coincide”?
slide50

BOND POLARITY

MOLECULAR

POLARITY

slide51

Center of +,-

charges coincide,

center of molecule

slide57

In conclusion, to be a dipole, a polar molecule (or

polyatomic ion), the presence of polar bonds is

required.

However, in addition, the polar bonds must be

arranged so that they are not canceling.

Molecular shape must be such that the center of the

negative charge buildup does not coincide with the

center of positive charge buildup.

slide58

Determine the polarity of each of the below:

Group Work

12.4

  • Draw to shape
  • Check  en
  • Draw arrow, if
  • dipole
importance of polarity
Importance of Polarity

As it turns out, it is the difference in polarity which

determines, for the same sized species, whether

it is soluble in water and whether it is a gas at

room temperature or a volatile liquid which

evaporates quickly or a high boiling liquid which

does not evaporate at all.

The attractions between molecules which causes

these differences all arise from increasing polarity or

its complete lack...

slide60

End

Topic 12