Lecture 23 © slg CHM 151 TOPICS : 1. Molecular Shapes 2. Polyatomic ion Shapes 3. Introduction to Bond Polarity.

Lecture 23 © slg CHM 151

TOPICS:

1. Molecular Shapes2. Polyatomic ion Shapes3. Introduction to Bond Polarity

Molecular and Polyatomic Ion Shapes

Once a Lewis structure is drawn, the three -dimensional geometry of the species can easilybe determined by utilizing the “valence shell electronpair repulsion theory” called “VSEPR”:

“VSEPR” theory is based on the tendency of negativelycharged regions to repel each other and align as farapart as possible, resulting in predictable shapes forany covalently bonded species.

To utilize “VSEPR”, the number of regions of electron density around the central atom in the species iscounted.

Count as “one region”:

• Single Bonds

• Unshared Pairs

• Multiple bonds between same two atoms

Examples of “four regions”:

A AA

“three regions”:

A A

“two regions”:

A A

Basic Shapes predicted by VSEPR:

Two regions: Three Regions:

A

120o

trigonal planar

A

180o

linear

Bond Angles

Geometry

Four Regions: Five Regions:

Six Regions:

A

trigonal-bipyramidal

90o

120o

A

90o

octahedral

A

tetrahedral

109.5o

trigonal tetrahedral

trigonal-bipyramidal octahedral

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 showincomplete octets (less than 8 e’s)

Metals, metalloids, and halogens do not as a rule formmultiple bonds. Compounds containing these elementswill often show an incomplete octet around the centralatom.

Type One: Two Regions Examples: BeCl2, CO2, NO2

+, HCN

Be Cl 2 BeCl ClBe 22Cl 14

16 e's/2= 8 prs

BeCl Cl (octet violator)

Number of regions around CENTRAL ATOM: 2

BeCl Cl shape : LINEARbond angles: 180o

CO 2 CO O C 42O 12

16 e's/2= 8 prs

CO O


shape : LINEARbond angles: 180o

CO O

CO O

NO 2+

NO O N 52O 12+1 -1e

16 e's/2= 8 prs

NO O



NO O

NO O+

HCN CH N H 1 C 4 N 5

10 e's/2= 5 prs

CH N



CH N

CH N

NO 3-

NO O

N 53O 181- 1

24 e's/2=12 prs

O

NO O

O

NO O

O-

(three re s onance s truc tures )

Type Two: Three Regions NO3

-, NO2-, CH2O


shape : TRIGONAL PLANARbond angles: 120o

NO O

O-

NO 2-

NO O

N 52O 121- 1

18 e's/2=9 prs

NO O NO O

-

(two re s onance s truc tures )



NO O

-

NO O

-

Black orbitalindicates pair ofunshared e’s

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

CH 2O CH H

C 4O 62H 2

12 e's/2=6 prs



O

CH H

O

CH H

O

CH H

O

Type Three: Four RegionsCH2Cl2, NH3, H2O, NH4

+


shape : TETRAHEDRALbond angles: 109.5o

CH Cl

H

Cl

NH 3 NH H

N 53H 3

8e's/2=4 prs


shape : TETRAHEDRALbond angles: < 109.5o

NH H

HNH H

H

H

Note: molecular geometry, trigonal pyramid

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 somewhatmore “repulsive” than bonded pairs.

When grouped with bonded pairs to tiny atoms like H,they tend to distort the bond angles, pushing thebonded pairs closer together.

The bond angles in ammonia are closer to 107o.

H2O OH H

O 62H 2

8e's/2=4 prs


shape : TETRAHEDRALbond angles: < 109.5o

OH

OH H

H

(~105o)

Note: molecular geometry: BENT

NH 4+

NH H

N 54H 41+ -1

8e's/2=4 prs



NH H

H

H

H

H +

GROUP WORK:

Do Lewis structure and assign shape and bond angles:

CO32-, SiCl4



CO O

O2-

CO32- 4 + 18 + 2 e 's = 24 e 's , 12 pairs



SiCl Cl

Cl

Cl

Si Cl4 4 + 28 e's = 32 e's, 16 pairs

Type Four: Five RegionsPF5 , ClF3 , IF2

-, SF4

PF 5 PF

F

P 55F 35

40e's/2=20 prs

F

FF


shape : Trigonal Bipyramidalbond angles: 120o, 90o

PF

F

F

FF

F PF

F

F

F

Bond anglesin triangle,120o

Bond angles,each “axial” F,90o from trigonalplane

ClF 3 Cl F

Cl 73F 21

28e's/2=14 prs


shape : Trigonal Bipyramidalbond angles: 90o

F

F

Cl F

F

F

Note: “T-shaped”; unshared pairs always trigonal planar

IF 2-

I F

I 72F 141- 1

22e's/2=11 prs



F

I

F

F

-

-

Note: “linear” molecular geometry

Note: “Seesaw” molecular geometry

SF 4 S F

S 64F 28

34e's/2=17 prs



F

S

F

F

F

F

F

F

Type Five: Six RegionsSF6, IF5, XeF4

SF 6 SF

S 66F 42

48e's/2=24prs

F

F

F

F F


shape : OCTAHEDRALbond angles: 90o

SF

F F

F

F

F

IF 5 IF

I 75F 35

42e's/2=21 prs

F

F

F F



IF

F F

FF

Note: molecular geometry “square pyramidal”

Xe F4 XeF

Xe 84F 28

36e's/2=18 prs



F

F F

XeF

F F

F

Note: “Square planar”

GROUP WORK:

Do Lewis structure and assign shape and bond angles:

ICl4+, XeOF2, ICl4

-

ICl4+



I

Cl

Cl

Cl

Cl

7 + 28 -1 e's = 34 e's, 17 prs

+



Xe

F

F

XeO F2 8 + 6 + 14 e's = 28 e's, 14 prs

O



ICl

Cl Cl

Cl

ICl4- 7 + 28 +1 e's = 36 e's, 18 prs

-

To see relevance of “shape work”, let’s turn nextto bond and molecular polarity. To help examinethis topic we turn back to the property of“electronegativity”:

Unit 5, Lecture 24, next week!!

ELECTRONEGATIVITY

The trends in ionization energies and electron affinities can be thought of as summarized in a single property called “electronegativity” (en or X).

Electronegativity is 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 the largest values, metals the lowest.Noble gases have no assigned X value.

ELECTRONEGATIVITY

INCREASES

>2

.8-1.4

1.5-1.9

Most active metals

Most active non-metals

The electronegativity values are quite useful inevaluating bond type and what we will term“bond polarity,” which arises when electrons areshared unevenly.

In summation:

Metals: larger size, lower ionization potential, lower electron affinity, and lower electronegativity; tend to form positive ions

Non-Metals: smaller size, higher ionization potential,higher electron affinity, higher electronegativity; function as anions in ionic compounds.

We have classified bonds “ionic” and “covalent”,depending on whether electron pairs are shared orelectrons are completely transferred from one atom to another.

In actuality, there is no sharp dividing line between thetwo types but rather a continuum:

Evenly shared electrons

Unevenly shared electrons

Transferredelectrons

To determine where a bond lies in this “continuum”, it is useful to consider the difference in electronegativitybetween the two atoms making up the bond:

When the difference is less than 0.5, sharing is fairlyeven and electrons are not much closer to one atom than the other.

When the difference is between 0.5 and about 1.5, theelectrons are closer to the more electronegative atom and partial charge buildup, polarization, develops.

When electronegativity difference is greater than1.5 or so, ionic bonding becomes the more likelytype and valence electrons are transferred to themore electronegative atom.

So, we need to consider a third more specializedtype of bond, “the polar covalent bond:”

This type of bond will be the important factor tobe considered when we look at molecular polarity, which arises from molecular shape and bondpolarity.

The polar molecular in turn will exhibit different solubilities and boiling points than non polar molecules. The relationship between these iswhat we will next examine.

Lecture 23 © slg CHM 151 TOPICS : 1. Molecular Shapes 2. Polyatomic ion Shapes 3. Introduction to Bond Polarity.

Documents

o slide

bent slide

hcn slide

bond polarity slide

bond angles geometry

trigonal planar slide

square planar slide

trigonal plane slide