Chapter 9 – Molecular Geometry and Bonding Theories Homework: 11, 13, 15, 19, 20, 21, 25, 26, 31,...

Chapter 9 – Molecular Geometry and Bonding Theories

Homework:11, 13, 15, 19, 20, 21, 25, 26, 31, 34, 35, 36, 39, 41, 42, 43, 44, 47,

49, 51, 54, 56, 96, 100

9.2 – The VSEPR Model Two balloons

Linear arrangement Three balloons

Trigonal-planar arrangement Four balloons

Tetrahedral arrangement

Electrons in molecules behave like balloons

A single covalent bond forms between atoms when a pair of electrons is between the atoms A bonding pair of electrons defines a

region in which the electrons are most likely to be found between two atoms

This area we find electrons is called an electron domain

A nonbinding pair (or lone pair) defines an electron domain located around one atom

Example

• Four electron domains here

•In general, each nonbinding pair, single bond or multiple bond produces an electron domain around the central atom

Because electron domains are negatively charged, they repel each other.

The best arrangement of a given number of electron domains is the one that minimizes the repulsions between them. This is the basic idea behind the

VSEPR model.

Similar to Balloons? You bet! Two domains makes linear

arrangement Three domains makes trigonal-

planar arrangement Four domains makes tetrahedral

arrangement

pg. 349

The arrangement of electron domains about the central atom is called its electron–domain geometry.

In contrast, the molecular geometry is the arrangement of only the atoms in a molecule or ion So any non-bonding pairs are not a part

of the molecular geometry

The VSEPR model predicts electron-domain geometry From this and knowing how many domains

are due to nonbinding pairs, we can predict the molecular geometry

When all the electron domains in a molecule come from bonds, the molecular geometry is the same as the electron-domain geometry But if one or more domains comes from lone

pairs, we must ignore those domains for molecular shape

pg. 351

Example NH3

Already done this. 4 electron domains around central atom

So electron-domain geometry is tetrahedral We know 1 of those domains comes

from lone pairs So the molecular geometry of NH3 is trigonal

pyramidal Tetrahedral with one less end, see pg. 347

Steps using VSEPR model to predict shape of molecules

1. Draw Lewis structure Count number of electron domains around

central atom

2. Determine electron-domain geometry Use table 9.1, 9.2 or 9.3

3. Use the arrangement of the bonded atoms to determine the molecular geometry

Use table 9.2 or 9.3

Example CO2

1. Draw Lewis Structure

How many electron domains around the central atom are there?

2. What is the electron-domain geometry for this?

Linear

3. What molecular geometry is possible? Linear

Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles

We refine the VSEPR model to predict and explain slight variances from the ideal bond angles Methane (CH4), ammonia (NH3) and

water (H2O) all have tetrahedral electron-domain geometries

But their bond angles are a little different CH4 = 109.5º, NH3= 107º and H2O = 104.5º Differences based around which type of

electron pairs make up the electron domains

Bond angles decrease as the # of nonbonding electron pairs increase. Bonding pair of electrons attracted by

both nuclei of the bonded atoms Lone pair of electrons attracted

primarily by one nucleus

Because lone pair has less nuclear attraction, it’s domain becomes more spread out So electron domain for lone pairs exert

more repulsive force on adjacent electron domains

This compresses (lessens) the bond angles Since H2O had the most lone pairs, it

gets the shortest bond angles

Multiple Bonds an Bond Angles Multiple bonds have a higher

electron-charge density than single bonds Also creates larger electron domains So electron domains for multiple

bonds exert a greater repulsive force on adjacent electron domains than single bonds do

So multiple bonds (double or triple) will decrease the bond angles too

Phosgene (Cl2CO)

Central atom has three electron domains 3 single bonds Trigonal planar geometry

Double bond acts like a lone pair, reducing the Cl-C-Cl bond angle

How Do These all Compare?

In terms of volume occupied by electron pairs In other words, who compresses the

most? Lone pair > triple bonds > double bonds

> single bonds

Molecules with Expanded Valence Shells So far we have assumed the

molecules have no more than an octet of electrons But the most common exception to

the octet rule is a central atom having greater than 8 valence electrons

So we need to deal with molecules with 5 or 6 electron domains

pg. 354

Example Use the VSEPR model to predict the

electron and molecular geometry of ClF3

Step 1: Lewis structure

How many electron domains around central atom? 5

5 electron domains Gives us an electron geometry of trigonal

bipyramidal How many bonding domains?

3 How many non-binding domains?

2 So its molecular geometry is

T-shaped

Shapes of Larger Molecules The VSEPR model can be extended

to more complex molecules than we’ve been dealing with.

Consider acetic acid CH3COOH

Acetic acid has 3 interior atoms Carbon, and each oxygen

We can use VSEPR to look at each central atom individually

9.3 – Molecular Shape and Molecular Polarity Remember that bond polarity

measures how equally the electrons in a bond are shared between the two atoms Higher bond polarity = less equal

sharing Higher electronegativity difference =

higher bond polarity

The dipole moment depends on both the polarities of the bonds and the geometry of the molecule Last chapter we focused just on the

polarity effect on the dipole moment For every bond in the molecule, we

can look at the bond dipole The dipole moment that is due ONLY

to the two atoms in the bond

Example CO2

O=C=O Each C=O bond is polar (O is more

electronegative than C) Since we have two O=C bonds, the

bonds are identical We end up with high electron density

around the O, and low electron density in the middle

Bond dipoles and dipole moments are vectors The overall dipole moment is the sum of the bond

dipoles that make it up But, must consider both the amount of the

dipole, and the direction of the dipole We have two identical C=O bonds, so the

amount of the dipoles are the same But the DIRECTION of the dipoles are opposite This causes the individual bond dipoles to cancel

each other out So the geometry of CO2 indicates that it is a NONPOLAR

molecule, even though it contains polar bonds.

Bond Dipole Activity Bond Dipole Activity

Steps to Determine Molecular Polarity

1. Draw Lewis structure2. Determine molecular geometry3. Look at effects of

electronegativity differences

9.4 – Covalent Bonding and Orbital Overlap The VSEPR gives as a method to predict

the shape of molecules Does not explain WHY the bonds exist

between atoms A mixture of Lewis’ notion of electron-

pair bonds and atomic orbitals leads to a model of chemical bonding This mixture of views is called the valence-

bond theory

In Lewis theory, covalent bonding occurs when atoms share electrons The sharing concentrates electron density

between the two nuclei involved In valence-bond theory, the build-up of

electron density between the nuclei is thought of as occurring when a valence atomic orbital of one atom merges

with a valence atomic orbital of another atom

This merger of orbitals Means that they share a region of

space Called overlap

The overlap of orbitals allows two electrons of opposite spin to share the common space between the nuclei

Forming an atomic bond

See figure 9.14 on pg. 360

Distance There is always an optimum

distance between the two bonded nuclei in a covalent bond Too close = too much repulsion

between the nuclei Too far = not much overlap, not a

strong bond

9.5 – Hybrid Orbitals