Structure and Bonding | Organic Chemistry 1

Electron configuration:

The arrangement of electrons in the atomic orbitals of an atom. An atomic orbital can hold a maximum of 2 electrons

The electronic configuration of an element can easily be determined by using the mnemonic below. The diagonal lines give the order of the subshells. The maximum number of electrons in each subshell is as follows: 2 e^- in the s-subshells, 6 e^- in the p-subshells, 10 e^- in the d-subshells, 14 e^- in the f-subshells
 

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 that they reach an octet in the valence shell and reach a noble gas configuration

Ionic bonds vs. covalent bonds

Bonding is a process of joining 2 atoms that results in a decrease in energy and an increase in stability: the atoms get a complete outer shell of valence electrons. There are 2 main types of 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 are bonds with 2 electrons. These bonds result from the sharing of electrons between 2 atoms (especially those in the middle of the periodic table). The electrons are shared to allow the atoms to reach noble-gas configurations
  Expected 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:

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 + $\frac{1}{2}$ 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:

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

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

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

Resonance structures:

A group of Lewis structures with the same placement of the 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 the most bonds and fewer charges has a higher contribution to the resonance hybrid

Selecting the best resonance structure that contributes the most to the resonance hybrid:

• 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 by the arrangement of their atoms and electrons, while 2 resonance structures differ only by the arrangement of their electrons

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 nonbonding 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 if you have a negative or a positive charge
2. Determine the number of 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 octet of the more electronegative atoms first. Include double or triple bonds if necessary
4. Check the number of valence electrons in 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 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. The 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: it is assumed that a carbon is at the junction of 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 the heteroatoms are drawn as well as the hydrogens that are directly bonded to them