4.2.1 Resonance Structures

• Some atoms or elements have structures that don’t seem to fit with what you would expect their typical Lewis Structure to be

• This can be explained by the delocalization of electrons

• Delocalized electrons = electrons in a molecule, ion, or solid metal that are not permanently associated with one atom or covalent bond

Example:

Nitrate (V) Ion

• A molecule with 1 double bond and 2 single bonds

• It has 3 possible Lewis Structures where the double bond moves around and is with each of the three oxygens

• Since there are different possibilities, these Lewis Structures are also called Resonance Structures

• So you would expect each of these Resonance Structures to have explicitly 1 double bond and 2 single bonds, right?

• That’s not the case, however, because studies of the electron density and bond length show that all 3 bonds are equal in length

• In fact, the electron density is spread evenly between the three oxygen atoms

  • The actual bond length for all of them is somewhere between a single and a double bond

  • The actual structure is something between all the resonance structures and is called a resonance hybrid

Resonance Structures of the Nitrate (V) Ion

Steps to Determine a Lewis Structure:

1. Count the # of Valence Electrons

2. Consider the Charge

   • Add more electrons for negative charges

   • Subtract electrons for positive charges

N + 3O + 1

5 + (3 × 6) +1

= 24 electrons

3. Draw a Skeleton

   • Put single bonds between atoms first

   • Generally, the least electronegative atom goes in the center

   • On the ends, put Hydrogens (because Hydrogen can only have 2 valence electrons and can only bond once) and Halogens

4. Subtract Skeletal Electrons from Valence Electrons

   • Use the remaining electrons to create Lone Pairs

   • OR additional bonds (double/triple) as needed

   • Remember the overall goal is to satisfy the Octet Rule

• 3 structures are possible for Nitrate (V) Ion, with 1 double and 2 single bonds

• The negative charge is distributed throughout the ion and is depicted with the negative sign outside of the resonance brackets

• Electron pairs rapidly oscillate between different positions, never really staying a single or a double bond for long

• Criteria for forming resonance hybrid structures: molecules must have a double bond that is capable of migrating from one part of a molecule to another

• In other words, when there are adjacent atoms with equal electronegativity and lone pairs of electrons that can move to another position in order for the double bonds to be in other positions

  • Ex. Carbonate Ion, Benzene, Ozone, and the Carboxylate Anion

4.2.2 Shapes of Molecules

• VSEPR (Valence Shell Electron Pair Repulsion) Theory = A theory that predicts molecular shape and the angles between bonds based on the concepts:

  1. All electron pairs and all lone pairs arrange themselves as far apart in space as possible

  2. Lone Pairs repel more strongly than bonding pairs

  3. Multiple bonds (double/triple) behave as single bonds

• Domains = The regions of negative cloud charge

• Steric Number (SN) = # of atoms + # of Lone Pairs around the central atom (same concept as Domains)

• Can also be denoted as:

  • A = Central Atom

  • B = Bonded Pair

  • E = Lone Pair

• ex. SN = 2 → AB₂

• ex. SN = 3 → AB₃

• ex. SN = 3 (but one of the domains is a lone pair) → AB₂E₁

Steric Number = 2

• If SN = 2, then the angle between bonds is 180°

• SN = 2 → AB₂

• Molecular Geometry = “Linear”

  • ex. BeCl₂, CO₂, HC≡CH

Steric Number = 3

• If SN = 3, then the angle between the bonds is 120°

• SN =3 → AB₃

• Molecular Shape = “Trigonal Planar”

  • ex. BF₃ and CH₂CH₂ and CH₂O

AB₂E₁

• If one of the electron domains is a lone pair, then the bond angle is slightly less than 120° since lone pairs repulse more, pushing against the other two bonding pairs closer together

  • ex. SO₂

• Molecular Geometry = “Bent Linear”

Steric Number = 4

• If SN = 4, then the angle between bonds is 109.5°

  • E.g. CH₄, NH₄⁺

• SN = 4 → AB₄

• Molecular Geometry = “Tetrahedral”

AB₃E₁

• If one of the electron domains is a lone pair, the bond angle is slightly less than 109.5° due to increased lone pair repulsion

  • ex. NH₃

• Molecular Geometry = “Trigonal Pyramidal”

AB₂E₂

• If 2 electron domains are lone pairs, bond angle also less than 109.5°

  • ex. H₂O

• Molecular Geometry = “Bent”

Summary

4.2.3 Predicting Shapes & Bond Angles

1. Draw Lewis Structure

   • Determine the number of bonding (B) and Lone Pairs (E) around the central atom (A)

2. Apply VSEPR Rules

   • Deduce shape and bond angle

4.2.4 Molecular Polarity

Bond Polarity ≠ Molecular Polarity

• Previously, you learned that bond polarity was determined by the difference in electronegative felt between two bonded atoms

• However, now you can determine if a molecule is polar or not

  • Consider:

    1. The polarity of each bond in the molecule

    2. How the bonds are arranged in the molecule

• Note: Some molecules have polar bonds, yet are overall not molecularly polar since the polar bonds in the molecule are arranged in a way that the individual bond dipole moments cancel each other out

  • ex. CH₃Cl vs. CCl₄

• CH₃Cl

  • Has 4 polar covalent bonds that don’t cancel each other out

  • This means the molecule is polar overall

  • The overall dipole moment is pointing towards the electronegative chlorine atom

• CCl₄

  • Also has 4 polar covalent bonds, BUT the individual bond dipole moments cancel each other out

  • So CCl₄ is a nonpolar molecule

4.2.5 Giant Covalent Structures

Giant Covalent Structures

Covalent Lattices

• Covalent Bonds = Bonds between nonmetals in which electrons are shared between atoms

• Giant Covalent Substances = Sometimes, a substance can't bond like a regular molecule. Instead, the bonds between atoms continue forever, forming a big lattice. There are no separate molecules in this situation, and all the nearby atoms are connected by covalent bonds.

  • ex. C

• Allotrope = Different atomic or molecular arrangements of the same element in the same physical state

• Graphite, diamond, buckminsterfullerene and graphene are allotropes of carbon

Giant Covalent Structures Examples

Diamond

• Diamond is a giant lattice of carbon atoms

• Each carbon is covalently bonded to 4 others in a tetrahedral geometry with a bond angle of 109.5°

• This results in a giant lattice with strong bonds in all directions and causes diamond to be the hardest known substance

  

Graphite

• Each carbon atom is bonded to 3 others in a layered structure of hexagons with a bond angle of 120°

• The spare electron is delocalized and moves around in the space between the layers

• All atoms in the same layer are held together by strong covalent bonds while the different layers are held together by weak intermolecular forces

Buckminsterfullerene

• Contains 60 carbon atoms

  • Each atom is bonded to 3 others by single covalent bonds

• The fourth electron is delocalized so the electrons can migrate throughout the structure

  • This allows for the structure to be a semi-conductor

• Has the same shape as a soccer ball, so it is nicknamed the football molecule

Graphene

• Some substances infinitely covalent bond only in two dimensons, forming only layers

  • ex. Graphene

• Graphene is a single layer of carbon atoms bonded in a repeating hexagonal pattern

  • It is so thin, 1 million times thinner than paper, that Graphene is actually considered 2D

Properties of Giant Covalent Structures

• As always, different structures and bonding types have different effects on the physical properties of substances (ie. melting/boiling points, electrical conductivity, and solubility)

Covalent Bonding & Giant Covalent Lattice Structures

• Giant Covalent Lattices:

  • Very High melting and boiling points

  • Large # of covalent bonds

  • A lot of energy is needed to break the lattice

  • Can be hard or soft

    • Hard (difficult to break their 3D network of strong covalent bonds)

      • Diamond

      • Silicon (IV) Oxide

    • Soft (forces between carbon layers are weak)

      • Graphite

    • (Graphene is strong, flexible, and transparent, making it a very useful material)

  • Insoluble in water (Most)

  • Do NOT conduct electricity (Most)

    • The some that do have delocalized electrons:

      • Graphite

      • Graphene

      • Buskminsterfullerene (semi-conductor)