Lewis Structures, the Octet Rule, and Molecular Geometry

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Lewis Structures

Drawing Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms and the distribution of valence electrons in a molecule or ion. They are essential for understanding molecular structure and reactivity.

Step 1: Count the valence electrons. The sum gives the total number of electrons that must be used in the Lewis structure.
Step 2: Arrange the atoms. Place atoms next to each other, using bonding patterns or placing atoms around a central atom.
Step 3: Place bonds between each pair of atoms. Draw single bonds to connect atoms.
Step 4: Add remaining electrons as lone pairs. Fill octets with lone pairs, starting with atoms on the periphery.

Bonding, Lone Pairs, and the Octet Rule

The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons, achieving a stable electron configuration similar to noble gases.

Hydrogen: Forms 1 bond, no lone pairs (2 electrons total).
Carbon: Forms 4 bonds, no lone pairs.
Nitrogen: Forms 3 bonds, 1 lone pair.
Oxygen: Forms 2 bonds, 2 lone pairs.
Halogens (e.g., F, Cl): Form 1 bond, 3 lone pairs.

Formula:

Number of bonds + Number of lone pairs = 4 (for atoms obeying the octet rule)

Exceptions to the Octet Rule

While the octet rule is useful, there are important exceptions:

Boron (B) and Aluminum (Al): Often have only three valence electrons and typically form 3 bonds.
Hydrogen (H): Only wants 2 electrons (duet rule).
Elements in period 3 or higher: Can have expanded octets (more than 8 electrons). For example, sulfur in SF6 has 12 valence electrons.

Indicating Charge in a Lewis Structure

When drawing Lewis structures for ions, the charge must be included:

Draw brackets around the structure and add the charge outside the brackets.
Adjusting valence electrons:
- Add one electron for every negative charge.
- Subtract one electron for every positive charge.

Example: Lewis Structure for NH2-

N: 5 valence e-
H: 2 × 1 valence e-
Negative charge: +1 e-
Total: 8 valence e-

Structure: [NH2]- with 4 electrons in lone pairs.

Example: Lewis Structure for H3O+

O: 6 valence e-
H: 3 × 1 valence e-
Positive charge: -1 e-
Total: 8 valence e-

Structure: [H3O]+ with 2 electrons in lone pairs.

Double and Triple Bonds

Formation of Multiple Bonds

Double and triple bonds form when the number of valence electrons is not enough to complete the octets of all atoms with single bonds.

Double bond: Two pairs of electrons are shared between two atoms.
Triple bond: Three pairs of electrons are shared between two atoms.
Hydrogen and halogens do not form double or triple bonds.

How to Draw Multiple Bonds

Convert one lone pair of an outer atom to one bonding pair between the outer atom and the central atom, or between two atoms needing octets.
Double bond: Contains four electrons in two two-electron bonds.
Triple bond: Contains six electrons in three two-electron bonds.

Example: Lewis Structure for C2H4 (Ethene)

Count valence electrons: 2 × 4 (C) + 4 × 1 (H) = 12 e-
Arrange atoms: H–C–C–H, with two H on each C.
Place bonds and add lone pairs. If octets are not complete, convert lone pairs to double bonds.
Final structure: H2C=CH2 (each C has a complete octet).

Example: Lewis Structure for Cyanide, CN-

C: 4 valence e-
N: 5 valence e-
Negative charge: +1 e-
Total: 10 valence e-
Arrange atoms: [C≡N]- (triple bond between C and N, with lone pairs to complete octets).

Resonance Structures

Definition and Purpose

Resonance structures are different possible Lewis structures for a molecule or ion where the arrangement of electrons (not atoms) varies. The true structure is a resonance hybrid, an average of all resonance forms.

Resonance is indicated by a double-headed arrow between structures.
Atoms remain in the same positions; only electrons move.
All resonance structures must have complete octets where possible.

Example: Resonance in Ozone (O3)

Count valence electrons: 3 × 6 = 18 e-
Draw possible structures with double bonds in different positions.
Resonance hybrid is the average of these structures.

VSEPR Theory and Molecular Geometry

Valence Shell Electron-Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the three-dimensional orientation of electron groups around a central atom. Electron groups (atoms or lone pairs) arrange themselves as far apart as possible to minimize repulsion.

Electron group: Any atom or lone pair attached to the central atom.
Bond angle: The angle between adjacent electron groups.

Common Molecular Geometries

Electron Groups	Lone Pairs	Geometry	Bond Angle	Example
2	0	Linear	180°	CO2
3	0	Trigonal planar	120°	BF3
3	1	Bent	120°	SO2
4	0	Tetrahedral	109.5°	CH4
4	1	Trigonal pyramidal	109.5°	NH3
4	2	Bent	109.5°	H2O

Predicting Molecular Shape

Draw the Lewis structure.
Count the number of electron groups (atoms + lone pairs) around the central atom.
Determine electron-group geometry using VSEPR theory.
Determine molecular shape based on the number of atoms and lone pairs.

Example: Shape of H2S

Lewis structure: H–S–H
S has 4 electron groups (2 atoms + 2 lone pairs) → tetrahedral electron geometry.
Shape: Bent, bond angle ≈ 109°.

Example: Shape of NH3

Lewis structure: N with 3 H atoms and 1 lone pair.
4 electron groups → tetrahedral electron geometry.
Shape: Trigonal pyramidal, bond angle ≈ 109°.

Summary Table: Electron Groups and Molecular Shapes

Electron Groups	Lone Pairs	Electron Geometry	Molecular Shape	Bond Angle
2	0	Linear	Linear	180°
3	0	Trigonal planar	Trigonal planar	120°
3	1	Trigonal planar	Bent	120°
4	0	Tetrahedral	Tetrahedral	109.5°
4	1	Tetrahedral	Trigonal pyramidal	109.5°
4	2	Tetrahedral	Bent	109.5°

Key Equations

(for octet rule)

Additional info:

Expanded octets are possible for elements in period 3 and beyond due to available d orbitals.
Resonance structures do not imply rapid switching; the actual molecule is a hybrid.