Benzene C6H6 This is a very important example in organic chemistry C C C C C C

Benzene C6H6 This is a very important example in organic chemistry C C C C C C

Benzene C6H6

The distances between adjacent carbon atoms in benzene are all the same and are equal to 1.40 Å. This distance is between the length of a typical carbon-carbon single bond (1.54 Å) and that of a carbon-carbon double bond (1.33 Å).

Examples: Draw the resonance structures for ozone (O3) and dinitrogen tetraoxide (N2O4).

Summary of rules for drawing resonance structures.

Summary of rules for drawing resonance structures. Rearrange electrons, but not the positions of the atoms.

Summary of rules for drawing resonance structures. Rearrange electrons, but not the positions of the atoms. 2. The same atoms must be bonded to one another in all the resonance structures. (Sometimes this is not true – but in the cases where it is not, the relative weight of the resonance structure is usually small).

3. Consider only those structures in which most or all of the atoms obey the octet rule.

3. Consider only those structures in which most or all of the atoms obey the octet rule. 4. The choice between “reasonable” and “unreasonable” structures is often based on our chemical intuition.

Formal Charge Sometimes when dealing with different resonance structures with unequal energies, we can select the more likely structure on the basis of the formal charge.

Formal Charge Sometimes when dealing with different resonance structures with unequal energies, we can select the more likely structure on the basis of the formal charge. Formal charge: The formal charge of an atom in a structure = number of valence e- – (number of un-shared valence e- + ½ number of bonding e-)

Key Comment: The sum of the formal charges for a structure must equal the charge on the overall structure.

Example: ammonia, NH3 The Lewis structure is:

Example: ammonia, NH3 The Lewis structure is:

Example: ammonia, NH3 The Lewis structure is: formal charge on N = 5 – (2 + ½ x 6) = 0

Example: ammonia, NH3 The Lewis structure is: formal charge on N = 5 – (2 + ½ x 6) = 0 formal charge on each H = 1 – (0 + ½ x 2) = 0

Example: ammonia, NH3 The Lewis structure is: formal charge on N = 5 – (2 + ½ x 6) = 0 formal charge on each H = 1 – (0 + ½ x 2) = 0 Note: sum of formal charges = charge on molecule

Example: ammonium ion, NH4+ The Lewis structure is:

Example: ammonium ion, NH4+ The Lewis structure is: +1

Example: ammonium ion, NH4+ The Lewis structure is: +1 formal charge on N = 5 – (0 + ½ x 8) = 1

Example: ammonium ion, NH4+ The Lewis structure is: +1 formal charge on N = 5 – (0 + ½ x 8) = 1 formal charge on each H = 1 – (0 + ½ x 2) = 0

Example: ammonium ion, NH4+ The Lewis structure is: +1 formal charge on N = 5 – (0 + ½ x 8) = 1 formal charge on each H = 1 – (0 + ½ x 2) = 0 Note: sum of formal charges = charge on ion

Example: hydrogen cyanide, HCN Lewis structure is:

Example: hydrogen cyanide, HCN Lewis structure is:

Example: hydrogen cyanide, HCN Lewis structure is: formal charge on N = 5 – (2 + ½ x 6) = 0

Example: hydrogen cyanide, HCN Lewis structure is: formal charge on N = 5 – (2 + ½ x 6) = 0 formal charge on C = 4 – (0 + ½ x 8) = 0

Example: hydrogen cyanide, HCN Lewis structure is: formal charge on N = 5 – (2 + ½ x 6) = 0 formal charge on C = 4 – (0 + ½ x 8) = 0 formal charge on H = 1 – (0 + ½ x 2) = 0

Example: hydrogen cyanide, HCN Lewis structure is: formal charge on N = 5 – (2 + ½ x 6) = 0 formal charge on C = 4 – (0 + ½ x 8) = 0 formal charge on H = 1 – (0 + ½ x 2) = 0 Note: sum of formal charges = charge on molecule

Key Comment: The formal charge gives an indication of the extent to which electrons have been gained or lost by atoms during bond formation.

Key Comment: The formal charge gives an indication of the extent to which electrons have been gained or lost by atoms during bond formation. Typically, a structure with low values (close to zero) of the formal charges will be energetically favored over a structure with more positive or more negative formal charges.

Example: cyanate ion, NCO- The Lewis structure:

Example: cyanate ion, NCO- The Lewis structure: Valence electron count = 5 + 4 + 6 + 1 = 16

Example: cyanate ion, NCO- The Lewis structure: Valence electron count = 5 + 4 + 6 + 1 = 16

Example: cyanate ion, NCO- The Lewis structure: Valence electron count = 5 + 4 + 6 + 1 = 16 1

Example: cyanate ion, NCO- The Lewis structure: Valence electron count = 5 + 4 + 6 + 1 = 16 1 -2 0 +1

Example: cyanate ion, NCO- The Lewis structure: Valence electron count = 5 + 4 + 6 + 1 = 16 1 2 -2 0 +1

Example: cyanate ion, NCO- The Lewis structure: Valence electron count = 5 + 4 + 6 + 1 = 16 1 2 -2 0 +1 -1 0 0

Example: cyanate ion, NCO- The Lewis structure: Valence electron count = 5 + 4 + 6 + 1 = 16 1 2 3 -2 0 +1 -1 0 0

Example: cyanate ion, NCO- The Lewis structure: Valence electron count = 5 + 4 + 6 + 1 = 16 1 2 3 -2 0 +1 -1 0 0 0 0 -1

Which of the structures is the best representation?

Which of the structures is the best representation?

Which of the structures is the best representation Which of the structures is the best representation? This can be eliminated because of the larger formal charges.

Which of the structures is the best representation Which of the structures is the best representation? This can be eliminated because of the larger formal charges. Both of the structures and have low formal charges.

Which of the structures is the best representation Which of the structures is the best representation? This can be eliminated because of the larger formal charges. Both of the structures and have low formal charges.

Which of the structures is the best representation Which of the structures is the best representation? This can be eliminated because of the larger formal charges. Both of the structures and have low formal charges. This is preferred, since the negative charge is on the more electronegative atom.

Electronegativity: The ability of an atom in a molecule to attract electrons towards itself.

Shapes of Molecules

VSEPR Theory

The shapes of molecules are very important because many molecular physical and chemical properties depend upon the three-dimensional arrangements of their atoms.

Initial chemical event in vision is a shape change for the molecule 11-cis-retinal, which is covalently bonded to a large protein molecule – the combination is called rhodopsin.

Key point: The three-dimensional shapes of molecules can be predicted if we assume that electron pairs in the valence shells of atoms stay as far apart as possible.

VSEPR Model

VSEPR Model VSEPR stands for valence shell electron pair repulsion.

VSEPR Model VSEPR stands for valence shell electron pair repulsion. The basic idea of this theory is that electrons in the valence shell, the electrons usually involved in bonding, repel each other.

In a covalent bond, the pair of electrons (called a bond pair) is responsible for holding the two atoms together.

In a covalent bond, the pair of electrons (called a bond pair) is responsible for holding the two atoms together. However, as a result of electrostatic repulsion, the bond pairs formed between the central atom and the surrounding atoms tend to remain as far apart as possible.

In a covalent bond, the pair of electrons (called a bond pair) is responsible for holding the two atoms together. However, as a result of electrostatic repulsion, the bond pairs formed between the central atom and the surrounding atoms tend to remain as far apart as possible. It is this electrostatic repulsion that ultimately determines the distribution of the bond pairs in space, and hence the overall shape of the molecule.