Chapter 7 Covalent Bonds and Molecular Architecture

Chapter 7 Covalent Bonds and Molecular Architecture

Octet rule: Main group elements tend to undergo reactions that leaves them with either 2 or 8 electrons in their outer or valence shell achieved by sharing electrons.

Exceptions to the rule includes all elements that have d levels close in energy. However, much of their chemistry can still be explained by the octet rule.

It is important to realize that this is a model that can also help us understand molecular structure.

Main Group Elements

Covalent bonding and Lewis structures

If H would form a covalent bond, how many bonds would it form?

Octet rule: Elements tend to undergo reactions to form covalent bonds that leaves them with either 2 or 8 electrons in their outer or valence shell.

Valence electrons are the electrons with the highest principle quantum number.

H2 has an electronic environment similar to He

One way chemists have of indicating a sharing of two electrons is to use a line to connect the two atoms, H2 is drawn H-H or H:H. These are called Lewis structures.

atomic orbitals

molecular orbitals

If Cl would form a covalent bond with H, how many bonds would it form?

chlorine is 1 electron short of a filled 3p level

hydrogen is 1 electron short of a filled 1s level

How many valence electrons does each atom posses?

(valence electrons are the electrons with the highest principle quantum number)

How many electrons does H need to share to have an outer shell that resembles an inert gas? Which inert gas?

How many does Cl need to share to have an outer shell that resembles an inert gas? Which inert gas?

H Cl

How many valence electrons does C posses?

If C forms covalent bonds, how many bonds would it form to satisfy the octet rule?

If H would form a covalent bond with C, how many bonds would it form?

H

H-C-H

H

4

4

Lewis structure

How many valence electrons does N posses?

If N forms covalent bonds with H, how many bonds would it form

to satisfy the octet rule?

If H would form a covalent bond with N, how many bonds would it form?

H-N-H

H

5

3

1

How many valence electrons does O possess?

If O forms covalent bonds with H, how many bonds would it form to satisfy the octet rule?

If H would form a covalent bond with O, how many bonds would it form?

H ׀ :O:׀ H

6

2

1

If B would form a covalent bond, how many bonds would it form?

If H would form a covalent bond with B, how many bonds would it form?

How many valence electrons does each atom posses?

How many does B need to share to have an outer shell that resembles an inert gas? Which inert gas?

Boron is a hopeless mess; much of its chemistry is very different from other elements in that the only way it can obtain an octet of electrons is by forming charged complexes; similarly with aluminum

H-B-H ׀ H

H -1׀ Na+ H-B-H ׀ H

Geometry of molecules

Suppose two groups were attached at a single point in space, and suppose these groups repelled each other, what geometric arrangement would these group chose to minimize their repulsion ( how would they arrange themselves to minimize repulsion)?

Suppose two groups were attached at a single point in space, and suppose these groups repelled each other, what geometric arrangement would these group chose to minimize their repulsion ( how would they arrange themselves to minimize repulsion)?

Suppose three groups were attached at a single point in space, and suppose these groups repelled each other, what geometricarrangement would these group chose to minimize their repulsion?

Suppose three groups were attached at a single point in space,

and suppose these groups repelled each other, what geometric

arrangement would these group chose to minimize their repulsion?

.

Suppose four groups were attached at a single point in space, and suppose these groups repelled each other, what geometric arrangement would these group chose to minimize their repulsion?

Suppose four groups were attached at a single point

in space, and suppose these groups repelled each other,

what geometric arrangement would these group chose

to minimize their repulsion?

.

Suppose four groups were attached at a single point

in space, and suppose these groups repelled each other,

what geometric arrangement would these group chose

to minimize their repulsion?

.

Suppose five groups were attached at a single point in space,









Suppose six groups were attached at a single point in space,









http://www.jcrystal.com/steffenweber/POLYHEDRA/p_10.html

octahedral geometry

Let’s define a group as either an atom or a pair of valence electrons not involved in bonding, ignore bonding electrons and electrons in inner shells, draw the Lewis structures, and predict the geometry of the following molecules

H2O

What is the central atom?

How many valence electrons around O?

How many groups around the central atom?

. 6 :O: .

O

4

What is the geometry of the molecule?

The geometry of a molecule is determined only by location of the nuclei. The electrons can not be located because of the uncertainty principle

CCl4; CH4


How many valence electrons around C?

What is the Lewis structure?



C

4

4

tetrahedral

NH3


How many valence electrons around N?

What is the Lewis structure?



N

5

4

pyramidal

Draw Lewis structures and predict the shape of the following compounds:

1. SiCl4

2. CH5N

3. CH2O

4. C2H2Cl2

5. C3H4

PCl5


SF6


Remember that we can only locate the position of a heavy atom; the position of electrons is not determined.

P

S

We have partially explained the geometry observed when atoms combine to form molecules. The geometry of the molecule is determined locally by the central atom. How do we identify central atoms?

Central atoms are determined by the number of bonds needed to complete the octet.

H, Halogens are seldom central atoms

B, C, N O in the first row

Al, Si, P, S in the second row …

Shapes of molecules

1. linear

2. trigonal planar

3. bent

4. tetrahedral

5. trigonal pyramidal

6. trigonal bipyramidal

7. seesaw

8. T shaped

9. octahedral

10. square pyramidal

11. square planar

http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html

What experimental evidense is there that CH4 is tetrahedral and not square planar?

Consider CH2Cl2:

H H

H

H

Cl

Cl

Cl

ClCC

Are there two compounds with the formula CH2Cl2? No

Are these the same?

Consider CHBrClF

Are these the same?

Consider CHBrClF

Are these the same?

Consider CHBrClF

You recall that we were able to explain atomic structure using s, p , d, f orbitals. Can we explain the structure of molecules using these same orbitals?

Remember the shape of these orbitals:

s: spherically symmetric

p: 3 orbitals each with two lobes touching each other at the nucleus and oriented 90 ° to each other.

d: 5 orbitals with a more complicated structure.

s

p

d

Chemists like to think that the electrons are the glue that hold atoms together. Therefore, can the structure of the molecules we just described, for example, CH4, PCl5, be explained using the hydrogen atom atomic orbitals just shown?

Schroedinger Equation is a differential equation. :

Properties of a differential equation:

1. the equation may have more than one solution.

2. any combination of solutions (sum or difference) is also a solution

2s = 1/4(1/2a3).5(2-r/a)(2.718)r/2a

2p = 1/4(1/2a3).5(r/a)(2.718)r/2acos

…

Linus Pauling: hybridization of atomic orbitals

What were to happen if we combined ½ of a 2s orbital with one of the ½ 2p orbitals mathematically?

- + +a 2s + b 2p

a 2s - b 2p- + +

- +

- +

These hybrid orbitals are directional, pointing 180° away from each other and are called sp hybrid orbitals

2 s 2 p

Combining a 2s orbital with 2 2p orbital can result in 3 sp2 hybrid orbital that point at 120 ° to each other

Combining a 2s orbital with 2 2p orbital results in 3 sp2 hybrid orbital that point at 120 ° to each other; note that one p orbital remains unchanged by these mathematics.

A summary of the types of hybridization necessary to product maximum electron density in the necessary direction as dictated by experimental geometries

sp hybridization: 2 orbitals pointing 180 ° to each other; 2 atomic p orbitals remain unchanged

sp2 hybridization: 3 orbitals pointing 120 ° to each other; 1 atomic p orbitals remains unchanged

sp3 hybridization: 4 orbitals pointing to the corners of a regular tetrahedron; all atomic p orbitals used

dsp2 hybridization: 4 orbitals pointing to the corners of a square; 4 d orbitals, 1 p orbital unchanged

dsp3 hybridization: 5 orbitals pointing to the corners of a trigonal bipramid; 4 d orbitals unchanged

d2sp3 hybridization: 6 orbitals pointing to the corners of a octahedron; 3d orbitals unchanged

sp2

sp3

d2sp3

Shapes of the hybrid orbitals

Draw the Lewis structure of C2H4 so that every carbon has a filled octet and each hydrogen has a He configuration

How many groups around each carbon?

3

What is the geometry at each carbon?

trigonal

sp2 hybrid orbitals

Atomic p orbitals on each C

Draw the Lewis structure of C2H2 so that every carbon has

a filled octet and each hydrogen has a He configuration

How many groups around each carbon?

2

What is the geometry about each carbon?

digonal: 180 °

H-CC-H

Bond lengths

C-C 1.54 *10-10 m

C=C 1.34*10-10 m

CC 1.2* 10-10 m

How do we explain the formation of CH4 using the electronic configuration of C

Valence electrons of C _____ _____ _____2p

____ 2s

hybridize

promote 1 electron: requires investment of energy

add electrons from H: allows the formation of 4 bonds instead of 2 and satisfies the octet rule

How do we explain the structure of molecules such as SF4?

3d ____ _____ _____ ______ ____

3p ___ ____ ____

3s ____

hybridize

3d ____ _____ _____ ______ ____ 3p ___ ____ ____ 3s ____

promote an electron

add electron from F

3d _____ _____ ______ ____ ____ ___ ____ ____ _____ dsp3

3d ____ _____ _____ ______ ____

3p ___ ____ ____ 3s ____

3d ____ _____ _____ ______ ____

3p ___ ____ ____ 3s ____

How do we explain the structure of molecules such as SF6?

3d _____ _____ ______ ____ ___ ____ ____ _____ _____

promote 2 electrons

hybridize

add electron from F

How good are Lewis structures at explaining molecular properties?

For compounds of carbon: excellent very few exceptions

For other elements: very good, some exceptions

Consider the Lewis structures of the diatomic molecules of the elements

H2,

O2

N2

F2

Any exceptions?

How good is this model at explaining molecular properties?

For compounds of carbon: excellent very few exceptions

For other elements: very good, some exceptions

Consider the Lewis structures of the diatomic molecules of the elements

H2,

O2

N2

F2

Any exceptions? O2 is paramagnetic

Molecular orbital Theory

Basic tenets: molecules are formed by combining atomic orbitals on each atom that have the proper orientation.

Whenever two atomic orbitals combine to form a molecular orbital, one combination is obtained by mathemetically adding the two together; this orbital goes down in energy relative to the atomic level. The other, obtained by the mathematical difference between the two orbitals, goes up in energy.

If a level goes down in energy relative to the atomic level, it is referred to as a bonding molecular orbital

If a level goes up in energy relative to the atomic level, it is called an antibonding level

If a level is not affected relative to the atomic level it is called a non-bonding orbital

H2

This model predicts that H2- should have some stability; the total

energy of He2 is the same as two isolated He atoms so nothing keeps the molecule together and it falls apart in two He atoms.

O2

F2

N2

Draw the Lewis structure of ozone: O3

Draw the Lewis structure of ozone

:O: :O::

::O:

:

+

-

Resonance structures: two structures that are identical except for the location of the electrons

Symbol chemists use to denote resonance structures

C

CC

C

CC

H

H

H

H

H

HC

CC

CC

C

H

H

H

H

H

H

Draw the structure of sulfuric acid; H2SO4

Draw the structure of the sulfate ion; SO4-2

Why do covalent bonds form?

Covalent bonds: the sharing of electrons

Electronegativity: generally meant to identify the unevenness in sharing electrons in a covalent bond

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Chapter 7 Covalent Bonds and Molecular Architecture