Chapter 9: Chemical Bonding I – Lewis Theory (General Chemistry Study Notes)

Lewis Theory: An Overview

Fundamental Concepts

• Valence electrons play a fundamental role in chemical bonding.

• Electron transfer leads to ionic bonds.

• Electron sharing leads to covalent bonds.

• Electrons are transferred or shared to give each atom a noble gas configuration (the octet rule).

• The octet rule generally applies to main group elements.

Lewis Symbols

Representation of Atoms

• A chemical symbol represents the nucleus and the core electrons.

• Dots around the symbol represent valence electrons.

• Example: Carbon (C) has the electron configuration 1s22s22p2, so its Lewis symbol is C with four dots.

• Group numbers help determine the number of valence electrons (e.g., Group 14 elements have 4 valence electrons).

Lewis Structures for Ionic Compounds

Constructing Ionic Lewis Structures

• Determine how many electrons each atom must gain or lose to achieve a noble gas configuration.

• Use multiples of one or both ions to balance the number of electrons.

• Ionic compounds are rarely molecular; their formulas represent charge and mass-balanced ions.

• In Lewis notation, ions are separated using square brackets, e.g., [Ba2+][O2−].

• Example: Magnesium chloride: Mg2+ and two Cl− ions.

Covalent Bonding

Electron Sharing

• In true covalent bonding, electrons are shared between atoms.

• No atom attracts the electrons from another; the sharing is equal (unless the bond is polar).

• Bond pairs are shared electrons; lone pairs are non-bonding electrons.

• Example: H2 molecule: two H atoms share electrons to form a single bond.

Coordinate Covalent Bonds

Special Covalent Bonds

• A coordinate covalent bond forms when both electrons in a bond are donated by one atom (the Lewis base) to another atom (the Lewis acid).

• Example: Ammonia (NH3) donates a lone pair to H+ to form ammonium (NH4+).

• All N–H bonds in NH4+ are equivalent.

• Chloride ion (Cl−) forms by transfer of a bonding pair to chlorine, resulting in an octet and a charge of −1.

Multiple Covalent Bonds

Double and Triple Bonds

• Multiple bonds result from more than one bonding electron pair being shared.

• Each double bond contains 4 electrons (2 per bond).

• Each triple bond contains 6 electrons (3 per bond).

• Example: CO2 has two double bonds; N2 has a triple bond.

Polar Covalent Bonds

Unequal Sharing of Electrons

• Polar covalent bonds share electrons unequally due to differences in electronegativity.

• Results in partial positive (δ+) and partial negative (δ−) charges.

• Example: H–Cl bond: H is δ+, Cl is δ−.

Electronegativity

Trends and Applications

• Electronegativity (EN) is the ability of an atom in a molecule to attract shared electrons.

• EN increases across a period and up a group in the periodic table.

• Difference in EN between atoms determines bond polarity and ionic/covalent character.

• Large ΔEN (≥2.0) indicates ionic bond; smaller ΔEN indicates covalent bond.

Slater's Rules

Calculating Effective Nuclear Charge (Zeff)

• Slater's Rules estimate the shielding effect and Zeff for valence electrons.

• Formula: $Z_{eff} = Z - S$

• S = sum of shielding contributions from other electrons.

• Higher Zeff leads to higher electronegativity.

Electron Type	Shielding Contribution
Same group (ns/np)	0.35
n-1 shell	0.85
n-2 or lower	1.00

Percent Ionic Character

Bond Character Continuum

• Bonds range from purely covalent to purely ionic.

• Percent ionic character increases with electronegativity difference.

• Example: KF is mostly ionic; HCl is more covalent.

Writing Lewis Structures

General Guidelines

• All valence electrons must be shown.

• Electrons are usually paired.

• Each atom typically requires an octet (except H, which needs only 2 electrons).

• Multiple bonds may be needed, especially for C, N, O, S, and P.

Skeletal Structure

Identifying Atom Positions

• Identify central and terminal atoms.

• Hydrogen atoms are always terminal and can only accommodate two electrons.

• Central atoms are generally those with the lowest electronegativity (except H and O).

• Carbon atoms are always central in organic compounds.

• Inorganic structures are usually compact and symmetrical; organic structures may not be.

Strategy for Writing Lewis Structures

Stepwise Approach

1. Count the total number of valence electrons.

2. Draw a skeletal structure.

3. Place two electrons in each bond.

4. Identify terminal atoms and complete their octets.

5. Subtract electrons used from the total; place remaining electrons on the central atom.

6. If atoms lack octets, form multiple bonds as needed.

Formal Charge

Calculating and Using Formal Charge

• Formal charge (FC) helps determine the most plausible Lewis structure.

• Formula: $FC = \text{valence electrons} - \text{lone pair electrons} - \frac{1}{2} \times \text{bonding electrons}$

• Structures with the lowest formal charges are preferred.

• Negative formal charges are usually placed on the most electronegative atoms.

Resonance Structures

Delocalization of Electrons

• Some molecules cannot be represented by a single Lewis structure; instead, multiple resonance structures are drawn.

• The real structure is a hybrid of all resonance forms.

• Example: Ozone (O3) has two resonance structures; bond order is 1.5.

• Bond order formula: $\text{Bond Order} = \frac{\text{Number of bonds}}{\text{Number of atom pairs}}$

Exceptions to the Octet Rule

Types of Exceptions

• Odd-electron species: Molecules with an odd number of electrons (radicals).

• Incomplete octets: Some elements (e.g., B, Be) are electron-deficient and do not achieve an octet.

• Expanded octets: Elements in Period 3 and below can have more than eight electrons (e.g., SF6, PCl5).

Bond Order and Bond Length

Relationship Between Bonding and Structure

• Bond order: Number of shared electron pairs between two atoms.

• Single bond: order = 1; double bond: order = 2; triple bond: order = 3.

• Higher bond order = shorter and stronger bond.

• Bond length: Distance between nuclei of bonded atoms.

Bond	Bond Length (pm)
C–C	154
C=C	134
C≡C	120
N–N	145
N=N	125
N≡N	110
O–O	148
O=O	121

Bond Energies

Energy Required to Break Bonds

• Bond energy: Energy required to break one mole of a bond in gaseous molecules.

• Shorter bonds are stronger and require more energy to break.

• Bond energies can be used to estimate enthalpy changes in reactions.

• Formula: $\Delta H_{rxn} = \sum \text{Bond energies of bonds broken} - \sum \text{Bond energies of bonds formed}$

Bond	Bond Energy (kJ/mol)
H–H	436
C–H	413
C–C	348
C=C	614
C≡C	839
N–N	163
N=N	418
N≡N	941
O–O	146
O=O	498

Practice Problems and Examples

Sample Calculations and Applications

• Calculate the number of valence electrons for a molecule (e.g., acetone, C3H6O: 24 electrons).

• Draw Lewis structures for molecules and ions (e.g., HCN, glycine).

• Assign formal charges and select the most plausible Lewis structure.

• Draw resonance structures and calculate average bond order (e.g., NO3−: bond order = 1.33).

• Estimate enthalpy changes using bond energies.

Summary Table: Key Lewis Structure Rules

Rule	Description
Octet Rule	Atoms (except H) tend to have 8 electrons in their valence shell
Formal Charge	Lowest formal charge preferred; negative charge on most electronegative atom
Resonance	Delocalization of electrons; real structure is a hybrid
Exceptions	Odd electrons, incomplete octets, expanded octets

Additional info:

• These notes are based on lecture slides and annotated class notes for a General Chemistry course (CHEM 1050), focusing on Lewis Theory and chemical bonding.

• Tables and values are inferred from standard textbook data (Tro, "Chemistry: A Molecular Approach").