# Representing Molecules | General Chemistry 1

The representation of molecules is studied in this chapter: Lewis structures, formal charges, resonance hybrids, electronegativity and polarity, dipole moment and percent ionic character, expanded valence shells

## The Octet Rule

Atoms tend to form molecules in such a way that they reach an octet in the valence shell and reach a noble gas configuration. The octet rule applies to almost all compounds made up of second period elements ⇒ this rule is particularly important in the study of organic compounds (compounds containing mainly C, N, and O atoms)

Carbon: Z = 6 ⇒ 6 electrons ⇒ 1s2s2p2 ⇒ 4 electrons 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

## 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

5 bonds ⇒ 10 electrons

... then lone pairs

5 bonds + 2 lone pairs ⇒ 14 e-

## 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 + $\frac{1}{2}$ 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 + 6 unshared e- ⇒ 7 valence e- in bound atom
Formal charge = 6 - 7 = - 1

## 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

## Electronegativity & Polarity

Electronegativity:

The ability of an atom in a molecule to draw shared electrons toward itself. Electrons in a bond are pulled toward the atom with the highest electronegativity. Electronegativity increases across a period and towards the top of a group ⇒ same trend as electron affinity

Distinction between covalent, polar covalent and ionic:

The difference in electronegativity between atoms is an indicator of the bond polarity: the greater the difference in electronegativity, the more polar the bond

• Purely covalent or nonpolar bond: a bond between atoms whose electronegativities differ by less than 0.5
• Polar covalent bond: a bond between atoms whose electronegativities differ by the range of 0.5 to 2.0
• Ionic bond: a bond between atoms whose electronegativities differ by 2.0 or more

H2 ⇒ no difference of electronegativity ⇒ nonpolar bond
HCl (electronegativity: H = 2.1, Cl = 3.0) ⇒ difference of electronegativity = 0.9 ⇒ polar covalent bond
NaCl (electronegativity: H = 0.9, Cl = 3.0) ⇒ difference of electronegativity = 2.1 ⇒ ionic bond

## Dipole Moment & Percent Ionic Character

Dipole moment (μ):

A quantitative measure of the polarity of a bond. It occurs whenever there is a separation of positive and negative charges. It is represented as a crossed arrow above the Lewis structure pointing along the bond from δ+ to δ-

Dipole moments are usually expressed in debye unit (1 D = 3.336 x 10-30 C.m). The dipole moment μ can be calculated (in C.m):

$\mathrm{\mu }$ = Q x r

Q = charge (in coulomb C)
r = distance between the charges (in m)

Percent ionic character:

A quantitative way to describe the nature of a bond and quantify its polarity. The measured dipole moment is compared to the one predicted by assuming that the bonded atoms have discrete charges

% ionic character =  x 100

## Exceptions to the Octet Rule

Exceptions to the octet rule:

• Incomplete octet (B and Be are prime candidates for this)
• Odd number of electrons (free radical)
• More than 8 electrons on the central atom ⇒ expanded valence shell

A species formed by homolytic cleavage that contains one or more unpaired electrons. Many radicals are highly reactive: the unpaired electron tends to form a covalent bond with an unpaired electron on another molecule

CH3 has one unpaired electron on the carbon, NO has one unpaired electron on the nitrogen ⇒ these are free radicals

Expanded valence shell

Atoms with a valence shell n $\ge$ 3 can have more than 8 valence electrons around them. This is because they also have d orbitals which can be used in bonding to form an expanded octet

Phosphorus is in the third period of the periodic table and therefore has 3s, 3p and 3d orbitals. The 3d orbitals of phosphorus in PCl5 are used and form an expanded octet ⇒ P has 10 electrons in its outer shell