1. Structural Formulas of OrganIc Compounds: Isomers
Section 1.6
Structural Formulas of OrganIc Compounds: Isomers

2. Levels of Organic Structure
We think about the structures of organic molecules at various levels
The composition of a molecule is captured by its molecular formula
The constitution (connectivity) of a molecule is captured by a structural formula showing how atoms are connected
The configuration of a molecule is captured by a structural drawing showing the positions of atoms in space
C4H10O
Compounds can have identical structures on a broad level but differ on a more specific level; such compounds are called isomers.

3. Constitutional Isomers
Constitutional isomers have the same composition (molecular formula) but differ in how their atoms are connected
Ethanol and diethyl ether both have the molecular formula C2H6O, but they have different connectivity
Constitutional isomers have different physical and chemical properties!

4. Condensed and Bond-line Formulas
Organic chemists make use of structural shortcuts to make structure drawings easier to parse
Condensed formulas are textual and list atoms in order of their connectivity
Bond-line formulas omit labels for carbon atoms and omit hydrogens where they are implied (assume the octet rule is satisfied)
becomes (CH3)OCH2CH3
Carbons are implied here
becomes

5. Expanding a Bond-line Formula
To expand a bond-line formula into a full Lewis structure, add in all implied carbons, hydrogens, and lone pairs
Draw carbon atoms at each vertex and at the end of each line

6. Expanding a Bond-line Formula
To expand a bond-line formula into a full Lewis structure, add in all implied carbons, hydrogens, and lone pairs
Draw carbon atoms at each vertex and at the end of each line
Add implicit hydrogens to the carbons such that the octet rule is satisfied (consistent with formal charges as well)

7. Expanding a Bond-line Formula
To expand a bond-line formula into a full Lewis structure, add in all implied carbons, hydrogens, and lone pairs
Draw carbon atoms at each vertex and at the end of each line
Add implicit hydrogens to the carbons such that the octet rule is satisfied (consistent with formal charges as well)
If missing, add implicit lone pairs on heteroatoms and/or carbons

8. Resonance and Curved Arrows
Section 1.7
Resonance and Curved Arrows

9. Resonance
The assumption of electron localization built into Lewis structures is not always true!
Molecules containing delocalized electrons have multiple Lewis-structural representations called resonance structures
For example, protonated formaldehyde has two important resonance forms
The true resonance hybrid contains a partial double bond between C and O and partial positive charge on C and O. A B
The compound does not convert between A and B; it is in reality a weighted sum of structures A and B!

10. Effects of Resonance
A Lewis structure of ozone suggests that the two O–O bonds are not equivalent
However, experiment confirms that they are equivalent—they have the same length and the outer oxygens have the same partial negative charge!
Resonance structures reveal that the bond orders of both O–O bonds are between 1 and 2 and that both outer oxygens share negative charge
The resonance hybrid
Resonance represents the delocalization and stabilization of charge. Molecules with resonance are stabilized!

11. Rules of Resonance When can resonance be considered?
Rule 1. The connectivity and positions of the atoms must remain the same in all resonance structures
Rule 2. Each contributing structure must have the same number of electrons and the same net charge
A and B are not resonance structures because they differ in connectivity.
They are constitutional isomers.
A, C, and D are resonance structures. All have a net charge of zero, although the formal charges of atoms differ between them.

12. Rules of Resonance When can resonance be considered?
Rule 3. Each contributing structure must have the same number of unpaired electrons
Rule 4. Atoms of second-row elements must not violate the octet rule
F contains two unpaired electrons while A, C, and D contain none. F is not an important resonance form of A.
G and H are valid resonance forms, but I is not because within it, nitrogen violates the octet rule.

13. Rules of Resonance Which resonance form is most important?
Rule 5. Resonance structures containing more covalent bonds are typically more important
Rule 6. Resonance structures with minimal separation of opposite charges are most important
Structure J with a double bond is the major contributor and is more important than structure K with a single bond (see also Rule 6).
Structure L lacks formal charges and is more important than structure M, which contains formal charges.

14. Rules of Resonance Which resonance form is most important?
Rule 7. Resonance structures with negative charge on the most electronegative atom and positive charge on the most electropositive atom are most important
In the major contributor N, the more electronegative oxygen atom bears the negative charge.

15. The Importance of Resonance
Recognizing resonance and drawing valid resonance structures are important skills for at least two reasons:
Resonance indicates that molecules are stabilized due to the delocalization of charge.
Resonance forms containing formal charges can reveal hidden points of reactivity in organic molecules.
Carbonate anion is fairly stable despite its overall charge of –2 due to delocalization of the charge over three oxygen atoms.
Minor contributor K reveals that carbon in C=O can act as a Lewis acid!

16. Sulfur and Phosphorus-contaInIng OrganIc Compounds and the octet Rule
Section 1.8
Sulfur and Phosphorus-contaInIng OrganIc Compounds and the octet Rule

17. Expanded Octets of Third-row Atoms
Atoms of elements in the third row can bear more than eight electrons in covalent compounds due to hypervalency
Compounds of phosphorus and sulfur are the most important examples of this phenomenon
Important examples include PCl5, phosphates, sulfates, sulfonates, and SF4.

18. Examples of Expanded Octets
Trimethylphosphine oxide has two important resonance forms; form G is more important (see Rule 1)
Phosphates in molecules such as adenosine triphosphate (ATP) contain expanded octets at phosphorus

19. Section 1.9
Molecular GeometrIes

20. Valence Shell Electron Pair Repulsion
VSEPR theory: In covalent compounds, valence electron pairs strive to be as far away from one another as possible to minimize electron-electron repulsion
This effect is one basis of the geometries of molecules. For example, methane’s four bonding pairs point to the vertices of a tetrahedron
We can distinguish between the electron-group arrangement of bonding pairs and lone pairs and the molecular geometry defined by the bonds and atoms only

21. Electron Groups
Electron groups are collections of electrons situated close in space. Two types:
Lone (unshared, non-bonding) electron pairs
Bonding electrons – multiple bonds are treated as a single group
The number of electron groups around an atom dictates the geometry at that atom
Unshared electron pairs are considered “larger” and “more repulsive” than bonded pairs

22. A Survey of VSEPR Geometries

23. Molecular Dıpole Moments
Section 1.10
Molecular Dıpole Moments

24. Applying Geometry and Polarization
With knowledge of the geometry of a molecule and the polarization of its bonds, we can deduce the molecular dipole moment
Individual bond dipoles are added in a vector sense to produce the overall molecular dipole vector
Compounds containing polarized bonds may still be nonpolar overall!

25. Nonpolar and Polar Molecules
Carbon tetrachloride (CCl4) is a nonpolar molecule. The four C–Cl bond dipoles add to a net molecular dipole of zero
Methylene chloride (CH2Cl2) is a polar molecule. The C–H and C–Cl bond dipoles add to give a net molecular dipole

26. Identifying Molecular Dipoles
Symmetric molecules and those with weakly polarized or unpolarized bonds are nonpolar; all others are polar
The polarization of a molecule determines its chemical behavior (reactivity) and physical properties
Structure Determines Properties!