Chapter 7 Student Notes

Covalent Bonding and Electron-Dot (Lewis) Structures

Introduction to Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms of a molecule and the lone pairs of electrons that may exist. They are essential for understanding molecular structure, predicting reactivity, and rationalizing chemical properties.

• Covalent bonds involve the sharing of electron pairs between atoms.

• Lewis structures show all valence electrons as dots or lines (bonds) around atoms.

• Hydrogen (H) can only form one bond and never has lone pairs in neutral molecules.

Steps for Drawing Lewis Structures

1. Sum the valence electrons for all atoms, adjusting for charges (add for anions, subtract for cations).

2. Determine the central atom (usually the least electronegative, never hydrogen).

3. Connect atoms with single bonds (lines).

4. Distribute remaining electrons as lone pairs to complete octets (or duets for H).

5. Form multiple bonds if necessary to satisfy the octet rule.

6. Check formal charges and minimize them for the most stable structure.

Example: For CO2, sum valence electrons (C: 4, O: 6x2 = 12; total = 16), connect C to O with double bonds, and assign lone pairs to O to complete octets.

Octet Rule and Exceptions

The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons. However, there are important exceptions:

• Hydrogen (H): Only 2 electrons (duet rule).

• Boron (B), Beryllium (Be), Aluminum (Al): Often stable with fewer than 8 electrons.

• Expanded octet: Elements in period 3 or below (e.g., P, S, Cl) can have more than 8 electrons due to available d orbitals.

Example: SF6 has 12 electrons around sulfur (expanded octet).

Formal Charge

Formal charge helps determine the most stable Lewis structure by assigning charges to atoms based on electron ownership.

• Formula:

$\text{Formal Charge} = \text{(Valence electrons)} - \text{(Nonbonding electrons)} - \frac{1}{2}\text{(Bonding electrons)}$

• The sum of all formal charges in a molecule or ion must equal the overall charge.

• Structures with the smallest (or zero) formal charges are generally most stable.

Example: In O3, the central O has a formal charge of +1, and one terminal O has -1.

Resonance Structures

Some molecules cannot be represented by a single Lewis structure. Resonance structures are multiple valid Lewis structures for the same molecule, differing only in the placement of electrons.

• Resonance structures are connected by double-headed arrows (↔).

• The actual structure is a resonance hybrid, an average of all resonance forms.

• Resonance stabilizes molecules by delocalizing electrons.

• Best resonance forms minimize formal charges and place negative charges on more electronegative atoms.

Example: The nitrate ion (NO3-) has three equivalent resonance structures.

Examples of Lewis Structures and Resonance

Species	Lewis Structure	Notes
NO3-	Three resonance forms with one N=O double bond and two N–O single bonds, negative charge delocalized.	All O atoms equivalent in the hybrid.
CO32-	Three resonance forms, each with one C=O and two C–O bonds.	Bond order = 1.33
SO2	Two resonance forms, S=O double bond alternates position.	Actual structure is a hybrid.
O3	Two resonance forms, central O with +1 formal charge, terminal O with -1.	Delocalized pi electrons.

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract electrons in a bond. The difference in electronegativity between atoms determines bond polarity:

• Nonpolar covalent: ΔEN < 0.5 (electrons shared equally)

• Polar covalent: 0.5 ≤ ΔEN < 2.0 (electrons shared unequally)

• Ionic: ΔEN ≥ 2.0 (electrons transferred)

Example: H–Cl is polar covalent; Na–Cl is ionic.

Bond Energy and Bond Order

Bond energy is the energy required to break one mole of a bond in the gas phase. Bond order is the number of shared electron pairs between two atoms.

• Higher bond order = stronger, shorter bonds.

• Bond energy can be used to estimate reaction enthalpy changes.

Formula for estimating reaction enthalpy:

$\Delta H_{\text{rxn}} = \sum \text{(Bond energies of bonds broken)} - \sum \text{(Bond energies of bonds formed)}$

Example: For CH4 combustion, sum energies for all C–H and O=O bonds broken and subtract energies for C=O and O–H bonds formed.

Practice Problems and Applications

• Draw Lewis structures for a variety of molecules and ions, including organic and inorganic examples.

• Assign formal charges and identify resonance structures as needed.

• Estimate bond orders and reaction enthalpies using bond energies.

Summary Table: Common Lewis Structures and Resonance

Compound/Ion	Key Features
Acetic acid (CH3COOH)	Carboxyl group, resonance in carboxylate ion
Nitrate (NO3-)	Three resonance forms, delocalized charge
Phosphate (PO43-)	Four equivalent resonance forms
SO3	Three resonance forms, expanded octet on S
CO2	Two double bonds, linear geometry
CH4	Tetrahedral, single bonds only

Additional info:

• Lewis structures are foundational for understanding molecular geometry, polarity, and reactivity.

• Practice drawing structures for ions and molecules, paying attention to octet rule exceptions and resonance.

• Use formal charge to select the most stable resonance form.