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 a Lewis structure:
 

  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- 

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

Resonance Forms

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 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

Electronegativity & Polarity

Electronegativity:

A measure of the ability of an atom in a molecule to attract electrons to itself. Electrons shared in a bond are attracted to the atom with the highest electronegativity. Electronegativity increases across a period and decreases down 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 polarity of the bond: 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 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 by 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):

μ = 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 = μ (observed)μ (calculated) x 100

 

Exceptions to the Octet Rule

Exceptions to the octet rule:

  • An 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

 

Free radical:

A chemical species with an odd number of valence electrons. It is 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 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

Check your knowledge about this Chapter

The octet rule states that 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

A Lewis structure is 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

  1. Count the total number of valence electrons. Add or subtract electrons if you have a negative or 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

The formal charge of an atom is the number of valence electrons in the free atom minus the number of valence electrons in the bound atom. The number of valence electrons in the bound atom is equal to the number of unshared electrons +  number of shared electrons

Resonance structures are a group of Lewis structures with the same placement of the atoms but 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

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

Principles for determining which resonance structure is most stable:

  • 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 is a measure of the ability of an atom in a molecule to attract electrons to itself. Electrons shared in a bond are attracted to the atom with the highest electronegativity

Electronegativity increases across a period and decreases down a group ⇒ same trend as electron affinity

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

  • Purely covalent or non-polar bond: a bond between atoms whose electronegativities differ by less than 0.5
  • Polar covalent bond: a bond between atoms whose electronegativities differ by 0.5 to 2.0
  • Ionic bond: a bond between atoms whose electronegativities differ by 2.0 or more

The dipole moment is a quantitative measure of the polarity of a bond. It occurs whenever there is a separation of positive and negative charges. It is represented by 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):

μ = Q x r

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

Percent ionic character is 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 = μ (observed)μ (calculated) x 100

 

  • An 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 free radical is a chemical species with an odd number of valence electrons. Therefore, it violates the octet rule. It is formed by homolytic cleavage and contains one or more unpaired electrons. Many radicals are highly reactive: the unpaired electron tends to form a covalent bond with an unpaired electron in another molecule

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