Bonding and Molecular Geometry

Chapter 10

MOLECULAR GEOMETRY AND CHEMICAL BONDING THEORY

INTRODUCTION

A) Lewis structures do not indicate the three dimensional shape of a molecule. They do not show the arrangement space of the atoms, what we call the molecular geometry or molecular structure.

B) Molecules have definite shapes and the shape of a molecule controls some of its chemical and physical properties.

II. Valence Shell Electron Pair Repulsion Theory - VSEPR - predicts the shapes of a number of molecules and polyatomic ions.

A) Assumptions of VSEPR Theory

1) Electron pairs in the valence shell of an atom tend to orient themselves so that the total energy is minimized. This means that: the electrons will approach the nucleus as close as possible yet take positions as far away from each other as possible to minimize _______________ .

2) Because lone pairs of electrons are spread out more broadly than bond pairs, repulsions are greatest between two lone pairs, intermediate between a lone pair and a bond pair, and weakest between two bonding pairs of electrons.3) Repulsive forces decrease rapidly with increasing interpair angle - greatest at 90o, much weaker at 120o, and very weak at 180o.

B) What are the ideal arrangements of electron pairs to minimize repulsions?

1) We need to identify the number of regions of high electron density, i.e. the number of stereoactive sets of electrons. One stereoactive set can be a single, double or triple bond.

2) If only single bonds, the number of stereoactive sets is equal to the bond pairs plus the lone pairs of electrons around the central atom.

3) If the number of stereoactive sets is 2, the arrangement of the pairs is:



III. PREDICTION OF GEOMETRIESA) Start with a Lewis dot structure - focus attention on the central atom.

B) Count the regions of high electron density around the central atom. Remember that one, two and three pairs of electrons between the same two atoms count as _________________ .

Each lone pair counts as one region, and each can be called a stereoactive set since each affects the geometry of the molecule.

C) Arrange the pairs of electrons (stereoactive sets) in such a way as to minimize repulsions especially at 90o.

D) VSEPR Notation - looks very much like the formula for a compound with a few changes.

1) The central atom is called A.

2) All the outer atoms are designated with an X.

3) Any lone pair electrons are designated with an E.

D) EXAMPLES

1) Two regions of high electron density

AX2 molecules

a) Look at the example of the BeF2(g) molecule.

The Lewis Structure is:

Look at the number of regions of high electron density around the center atom. There are two regions of high electron density.

2 pairs as far apart as possible are at 180o from each other. This gives a linear molecule.

Is the molecule polar?

No. Even though the bonds are polar, the symmetry of the molecule is such that the polar bonds "cancel" each other out and the center of positive charge coincides with the center of negative charge.

b) CO2 The Lewis Structure is:

Focus on the C atom. How many regions of high electron density? ___________

What is the geometry of the molecule? ______

Is the molecule polar?_________ Why?

c) HCN - The Lewis Structure is:

What is the geometry of the molecule?

Is it polar?

2) Three regions of high electron density

AX3 molecules

a) Look at the example of BF3 molecules.


Look at the number of regions of high electron density around the center atom. There are three regions of high electron density.3 pairs as far apart as possible are at 120o from each other. This gives a trigonal planar molecule.Is this molecule polar? __________________

AX2E molecules.

Example is GeF2 A "quick and dirty" way to find the number of electrons of halide compounds around the central atom is to take the number of electrons contributed to the bonding by the central atom and add ONE more electron for each halogen atom in the molecule.

Ge is in column 4A so we have 4e from it and we obtain 1e from each of the 2 F's which gives a total of 6e.

This gives a total of three pairs of electrons. There are two bonding pairs and one nonbonding pair at about 120o from each other.

This gives an angular geometry.

Is this molecule polar?

3) Four regions of high electron density

AX4 molecules

a) Look at the example of CH4

molecules.


Look at the number of regions of high electron density around the center atom. There are four regions of high electron density. The name of the geometry is TETRAHEDRAL and the angle is approximately 109.5o.

Is the molecule POLAR? _________

AX3E molecules

b) 3 bond pairs and 1 lone pair: EXAMPLE IS NH3. The Lewis structure is:

From where did we obtain the 8 electrons around the N? The name of the geometry is TRIGONAL PYRAMIDAL, OR trigonal pyramid.

We do not call it tetrahedral, because we do not include the position of the lone pair in the name of the geometry.

The geometry of the molecule name only includes the position of the atoms. The lone pair certainly affects the geometry of the molecule, however, its position is not included in the name.

Is the molecule POLAR? ________________

c) AX2E2

Two bonding regions and 2 lone pairs: EXAMPLE IS H2O. The Lewis structure is:

From where did we obtain the 8 electrons around the O? The name of the geometry is ANGULAR OR BENT.

Why don’t we call it tetrahedral?

Is the molecule POLAR?

Remember from Chapter 9, the polarity of a molecule can be indicated with partial signs followed by a + sign or a minus sign. The structure of the water molecule can be written:

4) FIVE regions of high electron density

AX5 molecules

a) Look at the example of PF5 molecules.

There is no Lewis Structure, but the dot structure is:

There are five regions of high electron density around the center atom.

Why is this possible?

The name of the geometry is TRIGONAL BIPYRAMIDAL. How do we get that name?

This is the first time we see a geometry in which there are two different angles; a set of 120o angles and a set of 90o angles.How do we get the five pairs around the P?

The P contributes 5 electrons since it is in column 5A of the Periodic chart, and each F contributes 1 electron giving a total of 10 electrons or 5 pairs.


AX4E molecules

a) Look at the example of SF4 molecules.There is no Lewis Structure, but the possible dot structures are:

We see four bond pairs and one lone pair.How do we decide which of the two possible structures is the correct one?

Where do the 4 F's and the lone pair go?

We haven't had the problem until now, because all the positions were equivalent.

Now we need the two assumptions of VSEPR theory which we haven't used until now.

We have to look at the number and kind of lone pair - bond pair interactions especially at 90o.

If the lone pair is put in the axial position, there are 3 lone pair - bond pair interactions at 90o.

If the lone pair is put in the equatorial position there are only 2 lone pair - bond pair interactions at 90o.

The lowest energy is associated with the lone pair in the equatorial position so the correct structure is:

What name do we give to this structure? If we rotate it 90 degrees it will be easier to see the name.

We call it a seesaw or an irregular tetrahedron.Is the molecule POLAR?

AX3E2 molecules

a) Look at the example of ClF3 molecules.

There is no "Lewis Structure", but the possible dot structures are:

We see three bond pairs and two lone pairs. How do we decide which of the three possible structures is the correct one? Where do the 3 F's and the two lone pairs go?

We have to look at the number and kind of lone pair - bond pair interactions especially at 90o.Most students want to put them on the axial positions, to put them at 180o from each other. If the lone pairs are put in the axial positions, there are 6 lone pair - bond pair interactions at 90o.

If the lone pairs are put at 90o, one pair on an axial position and one on an equatorial position then there is lone pair - lone pair at 90o, which is very repulsive.

The lowest energy is associated with the two lone pairs in equatorial positions where there are only 4 lone pair - bond pair interactions at 90o and one lone pair - lone pair interaction at 120o which is not important. This is the favored one, so the correct structure is:

What name do we give to this structure?

We call it a T - shaped


AX2E3 molecules

a) Look at the example of XeF2

molecules.

We see two bond pairs and three lone pairs.

If the lone pairs are put at 120o, a pair on each axial position we minimize the energy.



4) SIX regions of high electron density

AX6 molecules

a) Look at the example of SF6 molecules.

There are SIX regions of high electron density. Why is this possible?

The name of the geometry is OCTAHEDRAL. How do we get that name?

AX5E molecules

a) Look at the example of IF5 molecules.

WHY is this the only structure possible?What name do we give to this structure? Is it polar?

AX4E2 molecules

a) Look at the example of XeF4 molecules.

WHY?



IV. UNDERSTANDING COVALENT BONDING

A) Valence Bond Theory is a useful explanation for beginning students in general and organic chemistry. It is only an approximation, and it doesn't explain all bonding.

B) In valence bond theory, we explain bonding through the overlap of orbitals. An orbital on one atom occupies a portion of the same region in space as another orbital on another atom. Two orbitals overlap.

C) One pair of electrons occupies both orbitals.D) For the H2 molecule, the H-H bond axis is

through the center of the two atoms. The overlap is that of two spheres.

There is increased probability of finding the electron along the axis between the nuclei of the two atoms. The overlap region is symmetrical around the bond axis. This is called a sigma bond:

E) For the HCl(g), H contributes 1 electron

from its 1s orbital, spherical in nature and half empty. The Cl has a half empty p orbital. The s and p orbitals can overlap and a bond can be formed.

There is increased probability of finding the bonding electrons along the axis between the two atoms.

F) For the F2 molecule, there is a half empty p

orbital on each F atom so that end to end overlap of the p orbitals is possible giving a sigma bond.

G) We can generalize that where there is a single bond between two atoms, there is a sigma bond.

H) The second kind of bonding is pi ( ) bonding. It occurs when p orbitals overlap side by side.

The increased electron density is concentrated above and below the nuclear axis.

I) bonding occurs in molecules with multiple bonds.

1) This means there can be only one sigma bond between the same 2 atoms. All other bonds between the same 2 atoms are pi bonds.

2) A double bond will contain 1 sigma and 1 pi bond, a triple bond will contain 1 sigma and 2 pi bonds.

V. THE CARBON ATOM AND ITS FOUR BONDS

A) How can we explain the ability of the carbon atom to form four bonds with valence bond theory?

B) From the electron configuration of C - 1s22s22p2 and the formation of bonds takes place with the valence electrons 2s22p2 we have two p orbitals with 1 electron in each apparently available for orbital overlap.

C) It appears that there are 2 orbitals available for bonding. From this we could deduce that CH2 would be a stable molecule. It

is NOT.

D) The stable molecule is CH4. From where do

we obtain 4 orbitals on C with one electron in each? In addition, the molecule is tetrahedral. From where does the carbon atom get 4 orbitals oriented in space at 109.5 degrees? The s orbital is spherical and has no particular orientation and the p orbitals are oriented at 90 degrees to each other.1) Beginning organic chemistry using the valence bond approach rationalizes this bonding as follows:

We then mix these four orbitals 1s + 3p = 4 sp3 orbitals. Each orbital is a hybrid (like corn) with a new name sp3.

This means that the orbital is one part s and three parts p. These four new orbitals point to the corners of a tetrahedron.

Now carbon has the ability to form 4 sigma bonds to other atoms.The type of hybrid orbitals on the central atom is determined by the number of regions of high electron density around the central atom.

Bonding and Molecular Geometry

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