Module 7

Page 1

7.1

Molecular geometry provides a visual representation of the shape of a molecule and the spatial arrangements or positions of ligand atoms attached to a central atom (the atom present in the least amount) in a molecule.

An understanding of the shape of a molecule helps to explain the physical and chemical properties of the substances.

The Molecular Geometry of a molecule can be predicted from the Lewis structure of the molecule.

The Valence Shell Electron Repulsion (VSEPR) Theory provides the ground rules for predicting the spatial arrangements of atoms in a molecule

VSEPR uses an approach of distributing the ligands (the atoms attached to the central atom) around a central atom in such a manner that the valence shell electron pairs are kept at maximum distances; consequently, minimizing electron pair repulsion in the covalent bonds that have been established.

To understand the bonding that occurs in the covalent bonds between Be and F, we need to first look at the electron configurations of both beryllium and fluorine.

Be 1s2 2s2

F 1s2 2s2 2p5

An electron in the 2s orbital of beryllium can be promoted to a 2p orbital of beryllium.

Energy is required to promote an electron from the 2s orbital to a 2p orbital, but the energy required is more than compensated for by the energy released in forming a stable covalent bond.

The 2p atomic orbital that contains the promoted electron linearly combines with the 2s atomic orbital that contains one electron forming two degenerate (equal in energy) sp hybridized atomic orbitals.

In order to understand the bonding that occurs in the covalent bonds between B and F, we should review the electron configurations of both boron and fluorine. The electron configuration of boron and fluorine are:

B 1s2 2s2 2p1

F 1s2 2s2 2p5

The bonding in boron trifluoride begins with the promotion of an electron in the 2s orbital of boron to the 2p orbital of boron.

The 2p atomic orbital containing the promoted electron linearly combines with the other 2p atomic orbital of boron that contains an electron and the 2s atomic orbital that contains an electron forming three degenerate (equal in energy) sp2 hybridized atomic orbitals. Each sp2 atomic orbital has one electron.

Each B-F covalent bond is formed from the linear combination of a 2p atomic orbital of F and a 2sp2 hybridized atomic orbital of boron to form the bonding and anti-bonding molecular orbitals of the boron/fluorine bond in boron trifluoride. As discussed earlier, two molecular orbitals are produced because molecular orbitals are conserved when atomic orbitals linearly combine to form bonding molecular orbitals.

To understand the bonding that occurs in the covalent bonds between S and O, Let’s revisit the electron configurations of both sulfur and oxygen. The electron configurations of S and O are;

S 1s2 2s2 2p63s2 3p4 or [Ne]3s2 3p4

O 1s2 2s2 2p4

Since sulfur is filling the 3p orbital, there are low-lying d orbitals at energy level 3 that are available to accommodate a promoted electron. Therefore, an electron in the 3p atomic orbital of sulfur can be promoted to a 3d orbital of sulfur.

The 3d atomic orbital that contains the promoted electron can linearly combine with the 3s atomic orbital containing two electrons and a 3p atomic orbital containing one electron to form three degenerate (equal in energy) spd hybridized atomic orbitals.

Each double covalent bond between sulfur and oxygen is comprised of a σ and a π bond. The σ bond is formed from the mixing of a 3spd atomic orbital of sulfur and a 2sp2 hybridized atomic orbital from oxygen. The π bond is formed from the linear combination of a 3p atomic orbital of sulfur and a 2p atomic orbital from oxygen. Therefore, in this case, the double covalent bond is given the notations σ (3spd + 2sp2) and π (3p + 2p) . The lone pair of electrons would reside in a nonbonding 3spd molecular orbital.

Page 2

7.2

To understand the bonding that occurs in the covalent bonds between C and H, we should revisit the electron configurations of both carbon and hydrogen. The electron configurations of carbon and hydrogen are:

C 1s2 2s2 2p2

H 1s1

An electron in the 2s orbital of carbon can be promoted to a 2p orbital of carbon.

The 2p atomic orbital that contains the promoted electron linearly combines with the remaining two 2p atomic orbitals of carbon (each containing an electron) and the 2s atomic orbital that contains one electron to form four degenerate (equal in energy) sp3 hybridized atomic orbitals. The resulting hybridized orbitals are a little lower in energy than the 2p orbitals from which they were formed and a little higher in energy from the 2s orbital from which they were formed. Each sp3 atomic orbital has one electron.

Each C-H covalent bond is formed from the linear combination of a 1s atomic orbital of H with a 2sp3 hybridized atomic orbital of carbon to form a bonding and an anti-bonding molecular orbital that describe the carbon/hydrogen bond in methane.

Each bond between carbon and hydrogen is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 2sp3 hybridized atomic orbital of carbon and a 1s atomic orbital from hydrogen. In this case, the single covalent bond is given the notation σ(2sp3 + 1s)

To understand the bonding that occurs in the covalent bonds between N and H, we should revisit the electron configurations of both nitrogen and hydrogen. The electron configurations of nitrogen and hydrogen are:

N 1s2 2s2 2p3

H 1s1

In the case of nitrogen, there is no need to promote an electron in the 2s orbital to a 2p orbital. The only energy requirement would be the need to hybridize the three p orbitals containing each containing one electron and the lower energy 2s orbital containing two electrons to form four degenerate 2sp3 hybridized orbitals.

The four 2sp3 hybridized orbitals are a little lower in energy than the 2p orbitals from which they were formed and a little higher in energy from the 2s orbital from which they were formed. Three 2sp3 atomic orbitals have one electron each and one 2sp3 atomic orbital has two electrons.

Three N-H covalent bonds are formed from the linear combination of a 1s atomic orbital of H with a 2sp3 hybridized atomic orbital of nitrogen to form a bonding and an antibonding molecular orbital that describe the nitrogen/hydrogen bond in ammonia.

Each bond between nitrogen and hydrogen is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 2sp3 hybridized atomic orbital of nitrogen and a 1s atomic orbital from hydrogen. In this case, the single covalent bond is given the notation σ(2sp3 + 1s)

Finally, two electrons would reside in a nonbonding 2sp3 molecular orbital of nitrogen.

To understand the bonding that occurs in the covalent bonds between O and H, we should revisit the electron configurations of both oxygen and hydrogen:

O 1s2 2s2 2p4

H 1s1

In the case of oxygen, there is no need to promote an electron in the 2s orbital to a 2p orbital. The only energy requirement would be the need to hybridize the three p orbitals two of which contain one electron each and one containing two electrons and the lower energy 2s orbital containing two electrons to form four degenerate 2sp3 hybridized orbitals

The four 2sp3 hybridized orbitals are a little lower in energy than the 2p orbitals from which they were formed and a little higher in energy than the 2s orbital from which they were formed. Two of the 2sp3 atomic orbitals have one electron each and two of the 2sp3 atomic orbitals have two electrons each.

The two O-H covalent bonds are formed from the linear combination of a 1s atomic orbital of H with a 2sp3 hybridized atomic orbital of oxygen to form a bonding and an anti-bonding molecular orbital that describe the oxygen/hydrogen bond in water.

Each bond between oxygen and hydrogen is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 2sp3 hybridized atomic orbital of oxygen and a 1s atomic orbital from hydrogen. In this case the single covalent bond is given the notation σ(2sp3 + 1s)

Finally, four electrons would reside in two nonbonding 2sp3 molecular orbitals (each 2sp3 molecular orbital of oxygen would contain two electrons).

To understand the bonding that occurs in the covalent bonds between P and F, we should revisit the electron configurations of both phosphorus and fluoride. The electron configuration for phosphorus and fluorine are:

P 1s2 2s2 2p63s2 3p3

F 1s2 2s2 2p5

An electron in the 3s orbital of phosphorus can be promoted to a low-lying 3d orbital of phosphorus.

The 3d atomic orbital containing the promoted electron linearly combines with the three 2p atomic orbitals of phosphorus (each containing an electron) and the 3s atomic orbital that now contains one electron to form five degenerate sp3d hybridized atomic orbitals.

The resulting hybridized 3sp3d atomic orbitals are a little lower in energy than the 3d orbital from which they were formed and a little higher in energy than the 3p and 3s orbitals from which they were formed. Each 3sp3d atomic orbital has one electron.

Each P-F covalent bond is formed from the linear combination of a 2p atomic orbital of F with a 3sp3d hybridized atomic orbital of phosphorus to form a bonding and an antibonding molecular orbital that describes the phosphorus/fluorine bond in phosphorus pentafluoride.

Each bond between phosphorus and fluorine is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 3sp3d hybridized atomic orbital of phosphorus and a 2p atomic orbital from fluorine. In this case, the single covalent bond is given the notationσ(3sp3d + 2p).

To understand the bonding that occurs in the covalent bonds between S and F, we need to revisit the electron configuration of both sulfur and fluorine. The electron configurations of sulfur and fluorine are:

S 1s2 2s2 2p63s2 3p4

F 1s2 2s2 2p5

An electron in the 3p orbital of sulfur can be promoted to a low-lying 3d orbital of sulfur.

The 3d atomic orbital that contains the promoted electron linearly combines with the three 2p atomic orbitals of phosphorus (each containing an electron) and the 3s atomic orbital that contains two electrons to form five degenerate 3sp3d hybridized atomic orbitals. The resulting hybridized 3sp3d orbitals are a little lower in energy than the 3d orbital from which they were formed and a little higher in energy than the 3p and 3s orbitals from which they were formed.

The four S-F covalent bonds in sulfur tetrafluoride are formed from the linear combination of 2p atomic orbitals of F with 3sp3d hybridized atomic orbitals of sulfur to form bonding and anti-bonding molecular orbitals.

Each bond between sulfur and fluorine is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 3sp3d hybridized atomic orbital of sulfur and a 2p atomic orbital from fluorine. In this case, the single covalent bond is given the notation σ(3sp3d + 2p)

One 3sp3d hybridized atomic orbital has two electrons that do not participate in the formation of an S-F bond. These two electrons reside in a non-bonding 3sp3d molecular orbital.

The geometry of the trigonal bipyramidal molecule has two kinds of spatial arrangements of the five bonds. One spatial arrangement is described as the axial position, and the other is described as the equatorial position.

Page 3

7.3

To understand the bonding that occurs in the covalent bonds between Cl and F, we should revisit electron configurations of both chlorine and fluoride. The electron configurations of chlorine and fluorine are:

Cl 1s2 2s2 2p63s2 3p5

F 1s2 2s2 2p5

An electron in the 3p orbital of chlorine can be promoted to a low-lying 3d orbital of chlorine.

The 3d atomic orbital that contains the promoted electron can now mix with the three 2p atomic orbitals of chlorine (two contain one electron each and the third contains two electrons) and the 3s atomic orbital that now contains two electrons to form five degenerate sp3d hybridized atomic orbitals. The resulting hybridized 3sp3d orbitals are a little lower in energy than the 3d orbital from which they were formed and a little higher in energy from the 3p and 3s orbitals from which they were formed. Three sp3d atomic orbitals have one electron and two hybridized sp3d have two electrons.

The three Cl-F covalent bonds in chlorine trifluoride are formed from the linear combination of 2p atomic orbitals of F with 3sp3d hybridized atomic orbitals of chlorine to form bonding and anti-bonding molecular orbitals.

Each bond between chlorine and fluorine is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 3sp3d hybridized atomic orbital of chlorine and a 2p atomic orbital from fluorine. In this case, the single covalent bond is given the notation σ(3sp3d + 2p)

Two 3sp3d hybridized atomic orbitals have two electrons each that do not participate in the formation of a Cl-F bond. These two electrons reside in two non-bonding 3sp3d molecular orbitals.

To understand the bonding that occurs in the covalent bonds between Xe and F, we should revisit the electron configurations of both xenon and fluoride. The electron configurations of xenon and fluorine are:

Xe 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6

F 1s2 2s2 2p5

An electron in the 5p orbital of xenon can be promoted to a low-lying 5d orbital of xenon.

The 5d atomic orbital that contains the promoted electron linearly combines with the three 5p atomic orbitals of xenon (one contains one electron and the other two contain two electrons) and the 5s atomic orbital that contains two electrons to form five degenerate 5sp3d hybridized atomic orbitals. The resulting hybridized 5sp3d orbitals are a little lower in energy than the 5d orbital from which they were formed and a little higher in energy from the 5p and 5s orbitals from which they were formed. Two 5sp3d atomic orbitals have one electron and three hybridized 5sp3d have two electrons.

The two Xe-F covalent bonds in xenon difluoride are formed from the linear combination of 2p atomic orbitals of F with 5sp3d hybridized atomic orbitals of xenon to form bonding and anti-bonding molecular orbitals.

Each bond between xenon and fluorine is a single covalent bond (referred to as a sigma, σ bond).

The σ bond is formed from the linear combination of a 5sp3d hybridized atomic orbital of xenon and a 2p atomic orbital from fluorine. In this case, the single covalent bond is given the notation σ (5sp3d + 2p)

Three 5sp3d hybridized atomic orbitals have two electrons each that do not participate in the formation of a Xe-F bond. These three electrons reside in three non-bonding 5sp3d molecular orbitals.

To understand the bonding that occurs in the covalent bonds between S and F, we should revisit the electron configurations of both sulfur and fluorine. The electron configurations of sulfur and fluorine are:

S 1s2 2s2 2p63s2 3p4

F 1s2 2s2 2p5

An electron in the 3s orbital of sulfur can be promoted to a low-lying 3d orbital of sulfur. Also, an electron in the 3s orbital of sulfur is promoted to a 3d orbital of sulfur.

The two 3d atomic orbitals containing the promoted electrons linearly combine with the three 2p atomic orbitals of sulfur (each containing an electron) and the 3s atomic orbital that now contains one electron to form six degenerate 3sp3d2 hybridized atomic orbitals. The resulting six hybridized 3sp3d2 orbitals are a little lower in energy than the 3d orbital from which they were formed and a little higher in energy from the 3p and 3s orbitals from which they were formed. Each 3sp3d2 atomic orbital has one electron.

Sulfur heptafluoride is an octahedral molecule with two

F-S-F bond angles (90o and 180o)

Each of the six S-F covalent bonds in sulfur hexafluoride is formed from the linear combination of a 2p atomic orbital of F with a 3sp3d2 hybridized atomic orbital of sulfur to form bonding and anti-bonding molecular orbitals.

Each bond between sulfur and fluorine is a single covalent bond (referred to as a sigma, σ bond).

The σ bond is formed from the linear combination of a 3sp3d2 hybridized atomic orbital of sulfur and a 2p atomic orbital from fluorine. In this case, the single covalent bond is given the notation σ (3sp3d2+ 2p)

Sulfur hexafluoride is an octahedral molecule

with F-S-F bond angles equal to

90o and 180o

To understand the bonding that occurs in the covalent bonds between Br and F, we should revisit the electron configurations of both bromine and fluorine. The electron configurations of bromine and fluorine are:

Br [Ar] 3d10 4s2 4p5

F 1s2 2s2 2p5

An electron in two of the 4p atomic orbitals of bromine is promoted to two low-lying 4d orbitals of bromine.

The two 4d atomic orbitals containing the promoted electrons linearly combine with the three 4p atomic orbitals of bromine (each containing an electron) and the 4s atomic orbital that contains two electrons to form six degenerate 4sp3d2 hybridized atomic orbitals. The resulting six hybridized 4sp3d2 orbitals are a little lower in energy than the 4d orbital from which they were formed and a little higher in energy than the 3p and 3s orbitals from which they were formed. Five of the 4sp3d2 atomic orbitals have one electron and one 4sp3d2 has two electrons.

The five Br-F covalent bonds in bromine pentafluoride are formed from the linear combination of a 2p atomic orbital of F with a 4sp3d2 hybridized atomic orbital of bromine to form bonding and anti-bonding molecular orbitals.

Each bond between bromine and fluorine is a single covalent bond (referred to as a sigma, σ bond).

The σ bond is formed from the linear combination of a 4sp3d2 hybridized atomic orbital of bromine and a 2p atomic orbital from fluorine. In this case, the single covalent bond is given the notation σ (4sp3d2+ 2p).

Bromine pentafluoride is a square pyramidal (a special case of octahedral)

a molecule with bond angles equal to 90o and 180o.

To understand the bonding that occurs in the covalent bonds between Br and F, we should revisit the electron configurations of both xenon and fluorine. The electron configurations of xenon and fluorine are:

Xe 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6

F 1s2 2s2 2p5

Two electrons in two 5p orbitals of xenon are promoted to two low-lying 5d orbitals of xenon.

The two 5d atomic orbitals containing the promoted electrons linearly combine with the three 4p atomic orbitals of xenon (two containing one electron and one containing two electrons) and the 5s atomic orbital that contains two electrons to form six degenerate 5sp3d2 hybridized atomic orbitals. The resulting six hybridized 5sp3d2 orbitals are a little lower in energy than the 5d orbital from which they were formed and a little higher in energy than the 5p and 5s orbitals from which they were formed. Four of the 5sp3d2 atomic orbitals have one electron and two 5sp3d2 have two electrons.

The four Xe-F covalent bonds in xenon tetrafluoride are formed from the linear combination of a 2p atomic orbital of F with a 5sp3d2 hybridized atomic orbital of xenon to form bonding and anti-bonding molecular orbitals.

Each bond between bromine and fluorine is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 5sp3d2 hybridized atomic orbital of xenon and a 2p atomic orbital of fluorine. In this case, the single covalent bond is given the notation σ (5sp3d2+ 2p)

A square planar (special case of octahedral) molecule

with F-Xe-F bond angles equal to 90o and 180o.

Page 4

7.4

The simplest molecule with a multiple bonds (a double bond) is ethene (ethylene), C2H4. Following the octet rule, the structure for C2H4 would be:

A sigma, σ, and a pi, π, bond comprises the double bond.

In order to understand the nature of the double bond, let’s revisit the electron configuration of carbon. This configuration is 1s2 2s2 2p2.

An electron in the 2s orbital of carbon can be promoted to the 2p orbital.

The 2p atomic orbital that contains the promoted electron linearly combines with the 2s atomic orbital that contains one electron to form three degenerate 2sp2 hybridized atomic orbitals.

A 2sp2 hybridized atomic orbital of carbon linearly combines with a 1s atomic orbital of hydrogen to form a C-H sigma bond. Four of these C-H bonds are formed in the ethylene molecule.

The third 2sp2 hybridized atomic orbital linearly combines with the 2sp2 hybridized atomic orbital of the second carbon to form a 2sp2 + 2sp2 σ bond: The remaining electron in the 2p orbital of carbon linearly combines with the electron in the 2p orbital of the second carbon to form a π bond. These two types of linear combinations of atomic orbitals comprise the C-C double bond.

The simplest molecule with a triple bond is acetylene (ethyne), C2H2. Following the octet rule, the structure for C2H2 would be: H-C=-C-H

A sigma, σ, and two pi, π, bonds comprise the triple bond.

In order to understand the nature of the triple bond, let’s revisit the electron configuration of each of the two carbons that comprise the acetylene molecule. The electron configuration of carbon is 1s2 2s2 2p2.

In forming the triple bond, an electron in the 2s orbital of carbon is promoted to the 2p orbital of carbon.

A 2p atomic orbital of carbon-containing one electron linearly combines with the 2s atomic orbital to form two degenerate 2sp hybridized atomic orbitals.

A 2sp hybridized atomic orbital of carbon linearly combines with a 1s atomic orbital of hydrogen to form a C-H sigma bond. Two of these C-H bonds are formed in the acetylene molecule.

The second 2sp hybridized atomic orbital linearly combines with the 2sp hybridized atomic orbital of the second carbon to form a (2sp + 2sp) σ bond: The remaining two 2p orbitals of carbon linearly combines with the remaining two 2p orbitals of the second carbon to form two π bonds.

MOT describes the electronic structure of molecules in terms of the molecular orbitals (bonding and antibonding) spread over the entire molecule.

The electron configuration of the above diatomic molecule is (σ1s)2 (σ*1s)0

The bond order is given by the following equation:

Bond Order (BO) = ½ (nb - na) where nb = the number of electrons in the bonding molecular orbitals and na = the number of electrons in the antibonding molecular orbitals

For H2, the bond order is ½ (2-0) = 1. This means that there is a single bond between the two hydrogen atoms that comprise the diatomic hydrogen molecule.

If a BO = 0 is obtained, then this implies that the proposed molecule does not exist.

Where K is 1s2

This MOT picture describes the paramagnet property of liquid oxygen (liquid oxygen is attracted to a magnet) because the picture shows diatomic oxygen with two unpaired electrons.

The diagram gives a bond order of 2* for diatomic oxygen; therefore, the two oxygen atoms in diatomic oxygen form a double bond. This also follows the octet rule.

*½ (8-4) = 2

Molecular Polarity

"A molecule can be polar or nonpolar depending upon the nature of the bonds and the shape of the molecule. For a molecule that has different outer atoms, the molecular symmetry will decide the polarity.

If the molecular geometry is such that the dipole moments of each polar bond cancel each other then the molecule is nonpolar.

However, if the molecular geometry is such that the dipole moments of each polar bond don’t cancel each other then the molecule is polar."

Formation of a Chemical Bond

"Chemical bonds hold atoms together in molecules. A chemical bond is a strong electrostatic attraction between two or more atoms. It is formed either by a transfer of electrons or by a sharing of electrons such that each participating atom approaches a near-noble-gas configuration. A bond formed by a transfer of electrons is called an ionic bond. A bond formed by a sharing of electrons is called a covalent bond. A coordinate covalent bond is formed when an electron-rich atom donates a pair of electrons to an electron-deficient atom."

Magnetic Properties

As indicated earlier, liquid oxygen is attracted to a magnet because the MO picture shows that it has unpaired electrons (a condition that makes atoms and molecules paramagnetic).