Structure and Bonding | Organic Chemistry 1

The structure and bonding in organic molecules are studied in this chapter: atomic structure and electron configurations, ionic and covalent bonds (octet rule), formal charges, hybridization types, resonance forms, Lewis structure and other molecular structures

Atomic Structure


X: Symbol of the element
Z: Atomic number = number of protons
A: Mass number = number of protons + number of neutrons

Number of neutrons = A-Z
Number of electrons = number of protons - charge 


Z is always the same for a specific element

A can be different. XZA and XZA' are two isotopes

Electron Configuration

Electron configuration:

Distribution of electrons of an atom in atomic orbitals

To easily determine the electron configuration of an element, you need to draw a table like the one below and use diagonal lines to determine the order of the subshells. The maximum number of electrons in each subshell is: 2 e- in s subshells, 6 e- in p subshells, 10 e- in d subshells, 14 e- in f subshells

Oxygen  ⇒ Z = 8 and neutral atom ⇒ 8 electrons ⇒ 1s22s22p4
Iron ⇒ Z = 26 and neutral atom ⇒ 26 electrons ⇒ 1s22s22p63s23p64s23d6


Ionic and Covalent Bonds

Octet rule:

General rule governing the bonding process for second-row elements, the main elements in organic chemistry: atoms tend to form molecules in such a way as to reach an octet in the valence shell and attain a noble-gas configuration


Bonding is a process of joining 2 atoms that leads to lowered energy and increased stability: the atoms attain a complete outer shell of valence electrons. There are 2 types of bonding:

Ionic bonds:

Ionic bonds are based on the strong electrostatic attraction between 2 ions of opposite charges. These bonds result from the transfer of electrons from one element to another  in order to follow the octet rule

Covalent bonds:

Covalent bonds are two-electron bonds. These bonds result from the sharing of electrons between two atoms (especially those in the middle of the periodic table). Electrons are shared to allow the atoms to attain noble-gas configurations
Predicted number of bonds around an atom = 8 - number of valence electrons

Carbon: Z = 6 ⇒ 6 e- and neutral ⇒ 1s2 2s2 2p2  ⇒ 4 e in its valence shell
Carbon needs 4 more e- to get the configuration of Neon and will therefore form 4 covalent bonds with other atoms ⇒ carbon atom is tetravalent

Nitrogen: Z = 7⇒ 7 e- and neutral ⇒ 1s2 2s2 2p3 ⇒ 5 e- in its valence shell
Nitrogen needs 3 more e- to get the configuration of Neon and will therefore form 3 covalent bonds with other atoms ⇒ nitrogen is trivalent

Formal Charge

Formal charge:

Charge assigned to individual atoms in a Lewis structure
Formal charge = number of valence electrons in free atom - number of valence electrons in bound atom
Number of valence electrons in bound atom = number of unshared electrons + 12 number of shared electrons


What are the formal charges of CH3NO2?

N: 1s2 2s2 2p3 = [He] 2s2 2p3 ⇒ 5 valence electrons in free atom
4 bonds:  8 shared electrons⇒ 4 valence electrons in bound atom
Formal charge = 5 - 4 = + 1

O: 1s2 2s2 2p4 = [He] 2s2 2p4 ⇒ 6 valence electrons in free atom
2 bonds + 2 lone pairs: 4 shared + 4 unshared e- ⇒ 6 valence e- in bound atom
Formal charge = 6 - 6 = 0

O: 1s2 2s2 2p4 = [He] 2s2 2p4 ⇒ 6 valence electrons in free atom
1 bond + 3 lone pairs: 2 shared + 6 unshared e- ⇒ 7 valence e- in bound atom
Formal charge = 6 - 7 = - 1



Hybridization and Geometry


Combination of 2 or more atomic orbitals to form the same number of hybrid orbitals:

sp3 hybridization: combination of 1 s-orbital and 3 p-orbitals to form 4 sp3 hybrid orbitals
sphybridization: combination of 1 s-orbital and 2 p-orbitals to form 3 sp2 hybrid orbitals
sp hybridization: combination of 1 s-orbital and 1 p-orbitals to form 2 sp hybrid orbitals


The number of electron domains (atoms or lone pairs of electrons) around an atom determines both its geometry and hybridization:

2 electron domains ⇒ linear, angle bond: 180° ⇒ sp hybridization
3 electron domains ⇒ trigonal planar, angle bond: 120° ⇒ sp2 hybridization
4 electron domains ⇒ tetrahedral, angle bond: 109.5° ⇒ sp3 hybridization


Resonance Forms

Resonance structures:

Lewis structures having the same placement of atoms but a different placement of their π or nonbonded electrons ⇒ their single bonds remain the same but the position of their multiple bonds and nonbonded electrons differ

Resonance structures must be valid Lewis structures. Different resonance forms of a substance are not all equal: the form with more bonds and fewer charges has a higher contribution to the resonance hybrid

Difference between isomers and resonance structures:

2 isomers differ in the arrangement of both atoms and electrons while 2 resonance structures differ only in the arrangement of electrons

Shorthand Representations

Shorthand methods are used to abbreviate the structure of organic molecules. The 2 main types of shorthand representations are:

Condensed structure:

  • The main carbon chain is written horizontally. Atoms are drawn next to the atoms to which they are bonded
  • Covalent bonds and lone pairs are omitted
  • Parentheses are used around similar groups bonded to the same atom. If different substituents are bonded to the same atom, vertical lines can be used

Skeletal structure:

  • Carbon atoms are not shown: a carbon is assumed to be at the junction of any 2 lines and at the end of any line
  • Hydrogens around each carbon are not drawn, but we assume that there are enough hydrogens for the carbons to follow the octet rule
  • All heteroatoms are drawn as well as hydrogens directly bonded to them


Lewis Structures

Lewis structure:

Electron dot representation of a molecule

  • Only the valence electrons are drawn
  • Each element in the second-row has no more than 8 electrons
  • Each hydrogen has 2 electrons

Drawing a Lewis Structure:

  1. Count how many valence electrons there are in the molecule 
  2. Arrange the atoms you think are bonded together
  3. Add bonds first then lone pairs
  4. Check the number of valence electrons
  5. Assign formal charges to all atoms


Draw a Lewis structure for methanol CH3OH

Count the valence electrons:
1 C x 4 e- + 4 H x 1 e- + 1 O x 6 e- = 14 valence electrons ⇒ 7 bonds and lone pairs

Arrange the atoms


Add bonds ...

5 bonds ⇒ 10 electrons 

... then lone pairs

5 bonds + 2 lone pairs ⇒ 14 e-