10.2 Lewis structures

Lewis structures are the backbone of understanding chemical bonding. They show how atoms share or transfer electrons to form molecules and compounds. This concept is crucial for grasping how elements combine and interact in various chemical reactions.

By learning to draw and interpret Lewis structures, you'll be able to predict molecular shapes, bond types, and reactivity. This knowledge is essential for understanding more complex topics in chemical bonding and molecular behavior.

Lewis symbols for atoms and ions

Representing valence electrons as dots

• Lewis symbols are diagrams that show the valence electrons of an atom as dots around the chemical symbol of the element
• The number of valence electrons for an atom is equal to the group number on the periodic table
  • Exception: helium has 2 valence electrons
• For ions, the number of valence electrons is adjusted based on the charge
  • Positive ions have fewer valence electrons
  • Negative ions have more valence electrons
• Valence electrons are typically arranged in pairs, with a maximum of 4 pairs (8 electrons) around an atom, following the octet rule

Examples of Lewis symbols

• Carbon (C): 4 valence electrons, represented as ●C●●
• Oxygen (O): 6 valence electrons, represented as ●●O●●
• Sodium ion (Na+): 8 valence electrons, represented as ●●Na+●●
• Chloride ion (Cl-): 8 valence electrons, represented as ●●Cl-●●

Applying the octet rule

Stable electronic configurations

• The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with 8 valence electrons (or 2 for hydrogen)
• Atoms can achieve an octet by forming:
  • Ionic bonds (transferring electrons)
  • Covalent bonds (sharing electrons)
• In ionic compounds:
  • Metals tend to lose electrons to form cations
  • Nonmetals gain electrons to form anions
• In covalent compounds, atoms share electrons to form single, double, or triple bonds

Examples of achieving octets

• Sodium chloride (NaCl): Na loses an electron to form Na+, Cl gains an electron to form Cl-, both achieving octets
• Water (H2O): O shares electrons with two H atoms, forming single covalent bonds and achieving an octet

Constructing Lewis structures

Steps for drawing Lewis structures

1. Determine the total number of valence electrons in the molecule or ion

2. Arrange the atoms and connect them with single bonds

   • The least electronegative atom is usually central

3. Distribute the remaining electrons as lone pairs to achieve an octet for each atom (except hydrogen, which only needs 2 electrons)

4. If there are not enough electrons for each atom to have an octet, form double or triple bonds by moving lone pairs between atoms

5. Ensure the total number of electrons in the structure matches the number of valence electrons in the molecule or ion

Examples of Lewis structures

• Ammonia (NH3): N is the central atom, with 3 single bonds to H atoms and 1 lone pair
• Carbon dioxide (CO2): C is the central atom, with 2 double bonds to O atoms
• Nitrate ion (NO3-): N is the central atom, with 1 double bond, 2 single bonds to O atoms, and 1 negative charge

Exceptions to the octet rule

Odd-electron molecules and expanded valence shells

• Some molecules, such as nitric oxide (NO), have an odd number of electrons and cannot satisfy the octet rule for all atoms
• Elements in the third period and beyond can have expanded valence shells with more than 8 electrons
  • Example: sulfur hexafluoride (SF6)
• Some atoms, particularly boron and beryllium, tend to have fewer than 8 electrons in their valence shell when bonded
• Resonance structures can be used to represent molecules or ions that cannot be accurately depicted by a single Lewis structure

Examples of exceptions

• Nitric oxide (NO): N has 7 electrons and O has 8 electrons in their valence shells
• Sulfur hexafluoride (SF6): S has 12 electrons in its expanded valence shell
• Boron trifluoride (BF3): B has 6 electrons in its valence shell

Formal charges and resonance structures

Calculating formal charges

• Formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure
• To calculate formal charge, use the formula:
  • Formal charge = [number of valence electrons in free atom] - [number of non-bonding electrons] - 1/2 [number of bonding electrons]
• In a molecule or ion with multiple resonance structures, the structure with the lowest formal charges (closest to zero) on each atom is generally the most stable

Identifying stable resonance structures

• Resonance structures with the least separation of formal charges are typically more stable than those with greater charge separation
• When comparing resonance structures, those with the negative formal charge on the most electronegative atom are generally more stable

Examples of formal charges and resonance

• Ozone (O3): Two resonance structures with formal charges of 0, +1, and -1; the structure with the negative charge on the central O is more stable
• Carbonate ion (CO3^2-): Three resonance structures with formal charges of 0, 0, and -1; all structures contribute equally to the overall hybrid structure