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Resonance, Formal Charge, and Molecular Shape: General Chemistry Study Notes

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Recap: Ionic and Covalent Bonding

Ionic Bond Formation

The transfer of electrons is a key aspect of ionic bond formation, but the true driving force is the lattice energy of the resulting ionic compound.

  • Lattice energy is the energy released when ions come together to form a crystalline lattice.

  • Lattice energies and ionic bond strengths are directly proportional.

  • Lattice energies can be estimated using experimental or theoretical methods.

Covalent Bonding

Covalent bonding involves the sharing of electrons between atoms, resulting in measurable bond characteristics.

  • Bond polarity describes the degree to which bonding electrons are transferred or shared unequally.

  • Electronegativity is the tendency of an atom to attract shared electrons.

  • The extent of bond polarity is typically estimated from differences in electronegativity values.

  • Lewis structures are diagrams that show the sharing of electrons between atoms in covalent or polar covalent bonds.

  • The octet rule helps predict the formula and structure of covalent compounds.

Resonance, Formal Charge, and Molecular Shape

Lewis Structures: Stepwise Construction (Example: O3)

Lewis structures are used to represent the arrangement of electrons in molecules. The following steps outline the process for ozone (O3):

  1. Arrange the atoms: Place the atoms in a reasonable order, often with the least electronegative atom in the center.

  2. Count valence electrons: Each oxygen atom has 6 valence electrons. For O3:

  3. Connect atoms with single bonds: Each bond uses 2 electrons.

  4. Distribute remaining electrons: Place electrons as lone pairs to satisfy the octet rule for each atom.

  5. Form multiple bonds if needed: If the central atom lacks an octet, convert a lone pair from a terminal atom into a bonding pair, forming a double bond.

O3 Bonds: Bond Energies and Lengths

Covalent bonds between the same pairs of atoms have characteristic bond dissociation energies and bond lengths. In O3, the bonds are intermediate between single and double bonds.

Bond

Bond Dissociation Energy (kJ/mol)

Bond Length (pm)

Average O–O

180

148

Average O=O

498

121

O3

445

128

Ozone has only one measurable bond dissociation energy and bond length; both bonds are identical and intermediate between single and double bonds.

Resonance

Resonance occurs when more than one valid Lewis structure can be drawn for a molecule. The actual electronic structure is a weighted average of these possibilities, called a resonance hybrid.

  • Resonance is due to electron delocalization.

  • In O3, electrons are shared equally among all three atoms, rather than being localized between two.

  • Resonance structures are depicted with double-headed arrows between them.

Formal Charge

Formal charge helps determine the most significant resonance structure. It is calculated as:

Structures with formal charges closest to zero and negative charges on the most electronegative atoms are preferred.

Criteria for Resonance Structure Importance

  • Each atom should satisfy the octet rule.

  • Like charges on adjacent atoms are unfavorable.

  • Smaller formal charges (positive or negative) are preferable.

  • A more negative formal charge should reside on a more electronegative atom.

Exceptions to the Octet Rule

Fewer Than Eight Valence Electrons

Some compounds, such as those of beryllium and boron, are electron-deficient and do not satisfy the octet rule.

  • Example: BeCl2 and BF3

Odd Number of Electrons (Free Radicals)

Free radicals are chemical species with an unpaired electron, resulting in molecules that cannot satisfy the octet rule for every atom.

  • Example: NO (Nitric oxide)

More Than Eight Valence Electrons (Expanded Octet)

Atoms in period 3 or greater can have more than eight valence electrons, forming expanded octets.

  • Examples: PCl5 (10 electrons), SF6 (12 electrons)

Molecular Shape and the VSEPR Model

VSEPR Model

The Valence Shell Electron Pair Repulsion (VSEPR) model predicts molecular shapes by minimizing electron-electron repulsions around a central atom.

  • A single, double, or triple bond counts as one charge cloud.

  • Lone pairs and unpaired electrons also count as charge clouds.

  • Charge clouds arrange themselves as far apart as possible.

Common Molecular Geometries

Charge Clouds

Bonding Pairs

Lone Pairs

Geometry

Bond Angle

Example

2

2

0

Linear

180°

CO2

3

3

0

Trigonal planar

120°

BF3

3

2

1

Bent

~120°

SO2

4

4

0

Tetrahedral

109.5°

CH4

4

3

1

Trigonal pyramidal

~107°

NH3

4

2

2

Bent

~104.5°

H2O

5

5

0

Trigonal bipyramidal

120°, 90°

PCl5

5

4

1

Seesaw

<120°, <90°

SF4

5

3

2

T-shaped

<90°

ClF3

5

2

3

Linear

180°

I3-

6

6

0

Octahedral

90°

SF6

6

5

1

Square pyramidal

<90°

BrF5

6

4

2

Square planar

90°

XeF4

Shapes of Molecules with Multiple Central Atoms

For molecules with two or more central atoms, state the geometry about each central atom separately. For example, methylamine is tetrahedral about carbon and trigonal pyramidal about nitrogen.

Summary

  • Resonance describes the delocalization of electrons in molecules with multiple valid Lewis structures.

  • Formal charge calculations help identify the most significant resonance contributors.

  • Exceptions to the octet rule include electron-deficient molecules, free radicals, and expanded octets.

  • The VSEPR model predicts molecular shapes based on the arrangement of charge clouds around the central atom.

  • Memorizing common molecular geometries and bond angles is essential for understanding molecular structure and properties.

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