Chapter10 Study Notes: Chemical Bonding I – The Lewis Model

Chemical Bonding I: The Lewis Model

Lewis Dot Symbols

Lewis Dot Symbols (Electron Dot Diagrams) are diagrams that represent the valence electrons of an atom or ion. These symbols are essential for visualizing how atoms bond and interact in chemical compounds.

• Valence Electrons: Electrons in the outermost shell of an atom, involved in chemical bonding.

• Main Group Elements: Number of valence electrons equals the group number (for Groups 1A–8A).

• Transition Metals: Number of valence electrons varies and is less predictable.

• Lewis Dot Symbol: The element symbol surrounded by dots representing valence electrons.

Example: Which element will possess the most valence electrons?

• Answer: Group 8A elements (e.g., Ne, Ar) have the most valence electrons (8).

Drawing Lewis Dot Symbols

1. Identify the element and its group number.

2. Place one valence electron as a dot on each side of the element symbol before pairing.

3. For cations, remove electrons; for anions, add electrons.

Example: Draw the Lewis Dot Symbol for Te.

• Te is in Group 6A, so it has 6 valence electrons.

Chemical Bonds

Chemical bonds are the attractive forces that hold atoms or ions together in a chemical compound. There are three primary types of chemical bonds: ionic, covalent, and metallic.

Ionic Bonding

• Ionic Bond: Formed by the transfer of electrons from a metal to a nonmetal, resulting in oppositely charged ions.

• Electrostatic Attraction: The force that holds the cation and anion together.

Example: NaCl is an ionic compound formed from Na+ and Cl-.

Covalent Bonding

• Covalent Bond: Formed by the sharing of valence electrons between nonmetals.

• Molecular Compounds: Compounds containing covalent bonds.

Example: H2O is a covalent compound.

Metallic Bonding

• Metallic Bond: The attraction between free-flowing electrons and positively charged ions in a metal.

• Responsible for properties such as conductivity, malleability, and luster.

Example: Copper (Cu) exhibits metallic bonding.

Electronegativity and Dipole Moment

Electronegativity (EN) is a measure of an atom's ability to attract electrons in a bond. Differences in electronegativity between atoms lead to bond polarity and dipole moments.

• Electronegativity Trend: Increases across a period and decreases down a group.

• Dipole Moment: Occurs when a bond has a significant difference in electronegativity, resulting in partial charges.

Example: Calculate the difference in EN between C (2.5) and F (4.0): $\Delta EN = 4.0 - 2.5 = 1.5$

Bond Classification by Electronegativity Difference

Bond Type	Electronegativity Difference	Example
Nonpolar Covalent	0–0.4	Cl–Cl
Polar Covalent	0.4–2.0	H–Cl
Ionic	>2.0	Na–Cl

The Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons, resembling the electron configuration of noble gases.

• Valence Electrons: Electrons available for bonding.

• Shared Electrons: Electrons shared between atoms in a covalent bond.

Example: In NH3, nitrogen has 5 valence electrons and shares 3 electrons with hydrogen to complete its octet.

Incomplete and Expanded Octets

• Incomplete Octet: Some elements (e.g., B, Be) are stable with fewer than 8 electrons.

• Expanded Octet: Elements in period 3 or higher can have more than 8 electrons (e.g., P, S).

Formal Charge

Formal charge is a bookkeeping tool used to determine the distribution of electrons in a molecule. It helps identify the most stable Lewis structure.

• Formula: $\text{Formal Charge} = \text{Valence Electrons} - \text{Nonbonding Electrons} - \frac{1}{2} \times \text{Bonding Electrons}$

• Bonding Electrons: Shared between atoms.

• Nonbonding Electrons: Lone pairs on the atom.

Example: Calculate the formal charge of N in NH3. $\text{Formal Charge}_N = 5 - 2 - \frac{1}{2} \times 6 = 0$

Lewis Dot Structures for Neutral Compounds

Lewis Dot Structures show the arrangement of valence electrons in molecules. The most stable structure minimizes formal charges and satisfies the octet rule.

1. Count total valence electrons.

2. Arrange atoms (central atom is usually the least electronegative).

3. Connect atoms with single bonds.

4. Distribute remaining electrons to complete octets.

5. Assign formal charges and adjust structure if needed.

Example: Draw the Lewis Dot Structure for CH2O (formaldehyde).

Lone Pairs

• Lone Pair: A pair of nonbonding electrons localized on an atom.

Example: In H2S, sulfur has 2 lone pairs.

Lewis Dot Structures for Ions

Lewis Dot Structures for ions must account for the extra or missing electrons due to the charge.

• Cations: Remove electrons.

• Anions: Add electrons.

Example: Draw the Lewis Dot Structure for NaCl (Na+ and Cl-).

Lewis Dot Structures: Exceptions

• Free Radicals: Molecules with an odd number of electrons.

• Expanded Octets: Atoms in period 3 or higher can have more than 8 electrons.

Example: Draw the Lewis Dot Structure for NO (nitric oxide), a free radical.

Lewis Dot Structures: Acids

Acids are covalent compounds that release H+ ions in solution. Their Lewis structures show the hydrogen attached to the rest of the molecule.

• Common acids: HBr, HNO3, H2SO4, CH3COOH

Example: Draw the Lewis Dot Structure for HNO3.

Resonance Structures

Resonance structures are multiple valid Lewis structures for a molecule or ion, differing only in the placement of electrons.

• Double-Headed Arrow: Used to show resonance between structures.

• Resonance Hybrid: The true structure is a weighted average of all resonance forms.

Example: Draw all resonance structures for CO32-.

Average Charge

• The average charge on an atom is the sum of its formal charges in all resonance structures divided by the number of structures.

Example: Calculate the average charge of oxygen in PO43-.

Average Bond Order

Bond order is the average number of chemical bonds between a pair of atoms in a molecule. In resonance structures, bond order is calculated by dividing the total number of bonds by the number of bond locations.

• Formula: $\text{Bond Order} = \frac{\text{Total Number of Bonds}}{\text{Number of Bond Locations}}$

Example: Calculate the bond order of S–O bonds in SO3.

Bond Energy

Bond energy is the amount of energy required to break one mole of a bond in a molecule. It is used to estimate the enthalpy change of reactions.

• Enthalpy of Reaction Formula: $\Delta H_{rxn} = \sum \text{Bond Energies of Reactants} - \sum \text{Bond Energies of Products}$

Example: Calculate $\Delta H_{rxn}$ for the formation of NH3 from N2 and H2.

Standard Bond Energies Table

Bond	Bond Energy (kJ/mol)
H–H	432
N≡N	941
S–H	347
F–H	565
F–F	159
C–H	413
C–C	347
C=O	799
C–O	358
O=O	498
O–H	467

Lattice Energy

Lattice energy is the energy change when separated gaseous ions combine to form an ionic solid. It is a measure of the strength of the ionic bond.

• Exothermic Reaction: Lattice formation releases energy.

• Endothermic Reaction: Lattice dissociation absorbs energy.

• Formula: $\text{Lattice Energy} = \frac{Q_1 Q_2}{r}$ where $Q_1$ and $Q_2$ are the charges of the ions, and $r$ is the distance between them.

Example: Which compound possesses the strongest ionic bond: MgF2 or KCl?

Physical Properties Related to Lattice Energy

• Higher lattice energy leads to higher melting and boiling points.

Born-Haber Cycle

The Born-Haber Cycle is a series of steps used to calculate the lattice energy of an ionic compound from its constituent elements.

• Includes enthalpy changes for atomization, ionization, electron affinity, and formation.

Example: Calculate the lattice energy for BaBr2 using the Born-Haber Cycle values.

Step	ΔH (kJ/mol)
Atomization	113
Ionization	503
Electron Affinity	-736
Lattice Energy	-1000

Sigma and Pi Bonds

Sigma (σ) and Pi (π) bonds are types of covalent bonds formed by the overlap of atomic orbitals.

• Sigma Bond (σ): Formed by head-on overlap; strongest type of covalent bond.

• Pi Bond (π): Formed by side-to-side overlap; present in double and triple bonds.

Example: Ethylene (C2H4) contains one sigma and one pi bond between the carbon atoms.

Practice Problems

• Draw Lewis Dot Structures for various ions and molecules.

• Calculate formal charges and bond orders.

• Identify resonance structures and average charges.

• Estimate enthalpy changes using bond energies.

Additional info: These notes expand upon the provided study prep by including definitions, formulas, and examples for each major concept in Lewis bonding, as well as tables for bond classification and standard bond energies.