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

The octet rule:

The 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 reach 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 covalent 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 in the CH3NO2 molecule?

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 e- + 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 e- + 6 unshared e- ⇒ 7 valence e- in bound atom
Formal charge = 6 - 7 = - 1

 

 

Hybridization and Geometry

Hybridization:

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-orbital to form 2 sp hybrid orbitals
 

Geometry:

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
 


The different resonance forms of a substance are not all equal: the form with more bonds and less charges has a higher contribution to the resonance hybrid

 

Selecting the best resonance structure:

  • Lower formal charges (positive or negative) are preferable to higher charges
  • Formal charges on adjacent atoms are not desirable
  • A more negative formal charge should reside on a more electronegative atom

 

Difference between isomers and resonance structures:

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

Lewis Structures

Lewis structure:

A representation of the arrangement of atoms and the position of all valence electrons in a molecule or polyatomic ion. Shared electron pairs are represented by lines between 2 atoms, and lone pairs are represented by pairs of dots on individual atoms. We always try to satisfy the octet rule (or duet rule for hydrogen) when writing Lewis structures

  • Dot: one non-bonding electron
  • Pair of dots: lone electron pair (lone pair)
  • Line: two shared electrons (bond)
     
Lewis structure of NH3: 

N: 5 valence electrons ⇒ needs 3 shared electrons ⇒ 3 covalent bonds​​​
H: 1 valence electron ⇒ needs 1 shared electron ⇒ 1 covalent bond


Lewis structure of CO2:

C: 4 valence electrons ⇒ needs 4 shared electrons ⇒ 4 covalent bonds
O: 6 valence electrons ⇒ needs 2 shared electrons ⇒ 2 covalent bonds

 

How to write Lewis structures:
 

  1. Count the total number of valence electrons. Add or subtract electrons as necessary if you have a negative or a positive charge
  2. Determine how many covalent bonds / lone pairs the molecule will have by dividing the number of valence electrons by 2. Use the molecular formula to draw the skeletal structure
  3. Distribute the remaining valence electrons to satisfy the octet rule, completing the octets of the more electronegative atoms first. Include double or triple bonds if necessary
  4. Check the number of valence electrons of the drawn molecule
  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- 

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
  • The 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