2.5.A.1 Lewis diagrams can be constructed according to an established set of principles:

1. Chemical Bonds:

2. Valence Electrons:

• Valence electrons are the electrons in the outermost shell of an atom.
• These electrons are involved in chemical bonding.
• The number of valence electrons determines an atom’s reactivity and its ability to form bonds.
• Example: Sodium (Na) has 1 valence electron, and chlorine (Cl) has 7 valence electrons.

i. Octet Rule:

• Atoms tend to gain, lose, or share electrons to have 8 electrons in their outermost shell.
• Achieving 8 valence electrons results in a stable electron configuration.
• The Octet Rule applies to most elements, especially those in the second period and beyond.

Examples:

• Ionic Bonding:
  • Sodium (Na) loses its 1 valence electron, and chlorine (Cl) gains it.
  • Both now have 8 electrons in their valence shells, satisfying the Octet Rule.
• Covalent Bonding:
  • Two oxygen (O) atoms, each with 6 valence electrons, share 2 pairs of electrons to form O₂, achieving 8 electrons in their valence shells.

3. Drawing Lewis Structures:

Step 1: Count Valence Electrons

• Determine how many valence electrons are present in each atom of the molecule.
  • For atoms in Groups 1-18, the number of valence electrons corresponds to the group number (e.g., Oxygen in Group 16 has 6 valence electrons).
• Sum the valence electrons for all atoms in the molecule.
  • Example: In H₂O, H has 1 valence electron (×2 for two hydrogen atoms), and O has 6 valence electrons. Total = 2 + 6 = 8 valence electrons.

Step 2: Determine the Skeleton Structure

• Identify the central atom (usually the least electronegative atom, except hydrogen).
• Place the atoms around the central atom and connect them with single bonds.
  • Example: For H₂O, O is the central atom, and H atoms are attached to it.

Step 3: Distribute Electrons as Lone Pairs

• Place remaining electrons as lone pairs around the atoms, starting with the outer atoms.
• Ensure that atoms (other than hydrogen) achieve a full octet (8 electrons in their valence shell).
  • Example: In H₂O, after forming two O-H bonds, place the remaining 4 electrons as lone pairs on oxygen.

Step 4: Check the Octet Rule

• Verify that all atoms (except hydrogen) have a full octet (8 electrons in their valence shell).
  • Hydrogen only needs 2 electrons to fill its shell.

Step 5: Form Multiple Bonds (if needed)

• If any atom (except hydrogen) does not have a full octet, convert lone pairs into additional bonds to form double or triple bonds.
• Example: In CO₂, after forming single bonds between carbon and oxygen, convert lone pairs on oxygen into double bonds to ensure carbon achieves an octet.

Step 6: Final Check

• Ensure that the total number of electrons in the structure matches the number of valence electrons you started with.
• Check that all atoms (except hydrogen) follow the octet rule, and hydrogen follows the duet rule (2 electrons).

4. Resonance Structures:

i. Multiple Valid Lewis Structures: Some molecules can be represented by more than one valid Lewis structure.

ii. Electron Distribution: The different structures differ only in the arrangement of electrons, not atoms.

iii. Resonance Hybrid: The true structure is a hybrid (average) of all possible resonance structures, meaning electrons are delocalized across the molecule.

iv. Resonance Arrows: The different structures are connected with a double-headed arrow (↔) to show that the molecule oscillates between these forms.

v. Electron Delocalization: Electrons are spread over several atoms rather than being confined to one bond or location, which stabilizes the molecule.

vi. Molecular Stability: Resonance increases stability by reducing the overall energy through electron distribution.

vii. Examples:

1. • Ozone (O₃): Two resonance structures where the double bond shifts between oxygen atoms, creating a hybrid with partially double bonds.
   • Carbonate Ion (CO₃²⁻): Three resonance structures where the double bond shifts between the three oxygen atoms, leading to equal bond character.

viii. Importance of Resonance:

• • • Stability: Resonance makes molecules more stable by spreading electron density.
    • Electron Delocalization: Shows how electrons are distributed over several atoms, especially important in molecules like benzene (C₆H₆).
    • Explains Molecular Behavior: Resonance helps explain molecular properties, reactivity, and stability.

5. Exceptions to the Octet Rule:

6. VSEPR Theory: Valence Shell Electron Pair Repulsion Theory

• Purpose: VSEPR theory predicts the shapes of molecules based on the repulsion between electron pairs (bonding and lone pairs) in the valence shell of atoms.
• Electron Pair Repulsion: Electron pairs (whether in bonds or lone pairs) repel each other because they carry negative charge. This repulsion determines the molecular geometry.
• Electron Domains: Each region where electrons are located (bonding pairs or lone pairs) is called an electron domain. These domains arrange themselves to minimize repulsion.
• Electron Pair Arrangement: The arrangement of electron pairs around a central atom dictates the overall shape of the molecule.

i. Steps in VSEPR Theory:

Step 1: Determine the Lewis Structure

• What to Do: Draw the Lewis structure of the molecule, ensuring you follow the rules for bonding and placing electrons around atoms.
• Why It’s Important: The Lewis structure shows the number of bonds and lone pairs around the central atom, which are the key factors in predicting the molecular shape.
  • Example: For CO₂, the Lewis structure is:
    • Carbon (C) is in the center, with 2 double bonds to two oxygen (O) atoms, and no lone pairs on carbon. Each oxygen has two lone pairs.

Step 2: Count Electron Domains

• What to Do: Count the total number of electron domains (bonding and lone pairs) around the central atom.
  • Bonding domains: Count the number of single, double, or triple bonds as one electron domain per bond.
  • Lone pairs: Each lone pair on the central atom also counts as one electron domain.
• Why It’s Important: The number of electron domains determines the electron pair geometry, which is key to predicting the molecular shape.
  • Example: For CO₂ (carbon dioxide), there are 2 bonding pairs (C=O double bonds) and no lone pairs on the central carbon atom. So, there are 2 electron domains around carbon.

Step 3: Predict Geometry

• What to Do: Use the number of electron domains to predict the molecular geometry.

  • Electron domains arrange themselves to minimize repulsion (based on the principle of VSEPR).
  • The number of electron domains determines the ideal geometry (shape) around the central atom.

• Why It’s Important: The molecular shape (geometry) depends on how the electron domains are spaced around the central atom. This minimizes repulsion and determines the angles between bonds.

iii. Common Molecular Shapes:

iv. Lone Pairs and Molecular Shape:

• Lone pairs take up more space than bonding pairs, affecting the molecular shape by pushing bonding pairs closer together.
• Example: In NH₃ (ammonia), the shape is trigonal pyramidal due to one lone pair on nitrogen, even though nitrogen has 4 electron domains.