Chemical Bonding I: Types, Rules, and Structures

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Chemical Bonding I

Types of Bonding

Chemical bonding describes how atoms combine to form compounds. The nature of the bond depends on the types of atoms involved and their interactions.

Ionic Bonding: Occurs between metals and nonmetals; electrons are transferred from one atom to another.
Covalent Bonding: Occurs between nonmetals; electrons are shared between atoms.
Metallic Bonding: Occurs between metals; electrons are pooled and delocalized throughout the structure.

Types of Atoms	Type of Bond	Characteristic of Bond
Metal and nonmetal	Ionic	Electrons transferred
Nonmetal and nonmetal	Covalent	Electrons shared
Metal and metal	Metallic	Electrons pooled

Octet Rule and Noble Gas Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons, similar to the electron configuration of noble gases (ns2np6).

Elements gain or lose electrons to attain the electron configuration of the nearest noble gas.
H, Li, Be, B attain a configuration like He (1s2).
Transition metals often follow the 18-electron rule.

Example:

Na (1s22s22p63s1) loses 1 electron to form Na+ (1s22s22p6), achieving a noble gas configuration.
O (1s22s22p4) gains 2 electrons to form O2- (1s22s22p6).

Ionic Bonding

Ionic bonds form when electrons are transferred from a metal to a nonmetal, resulting in the formation of cations and anions. The electrostatic attraction between these oppositely charged ions forms the ionic compound.

Example: Formation of NaCl from Na and Cl atoms.
Electron transfer alone does not explain the highly exothermic nature of ionic compound formation; lattice energy is also involved.

Lattice Energy

Lattice energy is the energy associated with the formation of one mole of crystalline salt from gaseous cations and anions.

Lattice energy explains the stability of ionic compounds.

Trends in Lattice Energy

Lattice energy depends on ion size and charge:

Larger ions = weaker attraction = smaller lattice energy.
Larger charge = stronger attraction = larger lattice energy.
Ion charge is generally more important than ion size.

Equation: Example:

LiBr has the highest-magnitude lattice energy among KBr, NaBr, LiBr, RbBr, due to the small size of Li+.

Covalent Bonding

Covalent bonds involve the sharing of electrons between nonmetal atoms. Several electrostatic interactions occur:

Attractions between electrons and nuclei
Repulsions between electrons
Repulsions between nuclei

Polar Covalent Bonds

In polar covalent bonds, electrons are shared unequally due to differences in electronegativity.

Fluorine pulls harder on shared electrons than hydrogen in HF, resulting in greater electron density near fluorine.

Electronegativity

Electronegativity is the ability of an atom in a molecule to attract electrons to itself. It is a combined effect of electron affinity and ionization energy.

On the periodic table, electronegativity increases from left to right across a period and from bottom to top within a group.
Fluorine is the most electronegative element.

Bond Dipole and Dipole Moment

When two atoms share electrons unequally, a bond dipole results. The dipole moment () is calculated as:

Measured in debyes (D).

Trends in Bond Polarity

The greater the difference in electronegativity, the more polar the bond and the larger the dipole moment.

Compound	Bond Length (Å)	Electronegativity Difference	Dipole Moment (D)
HF	0.92	1.9	1.82
HCl	1.27	0.9	1.08
HBr	1.41	0.7	0.82
HI	1.61	0.4	0.44

Lewis Structures

Lewis symbols represent the chemical symbol of an element or ion plus a dot for each valence electron. Lewis structures are used to depict bonding in molecules.

Bonding pairs: Electrons shared between atoms.
Lone pairs (nonbonding pairs): Electrons not shared, belonging to a single atom.

Double and Triple Bonds

Atoms can share more than one pair of electrons, forming double or triple bonds to satisfy the octet rule.

Example: O2 has a double bond; N2 has a triple bond.

Predicting Lewis Structures

Steps to predict Lewis structures:

All atoms tend to follow the octet rule (bonding electrons count for both atoms).
Total number of valence electrons must not change.
Add or subtract electrons for ions with charges.
Elements with lower electronegativity are often central atoms.

Common bonding patterns:

H: 1 bond (always terminal)
C: 4 bonds, 0 lone pairs
N: 3 bonds, 1 lone pair
O: 2 bonds, 2 lone pairs
Halogen: 1 bond, 3 lone pairs (if terminal)
B: 3 bonds, 0 lone pairs

Formal Charge

Formal charge helps determine the most stable Lewis structure. It is calculated as:

The sum of formal charges must equal the net charge of the molecule. The best Lewis structure has the fewest charges, with negative charges on the most electronegative atoms.

Exceptions to the Octet Rule

Some molecules and ions do not follow the octet rule:

Odd number of electrons (e.g., NO)
Less than an octet (e.g., BF3)
More than eight electrons (expanded octet, e.g., PCl5, SO42-)

Resonance Structures

Some molecules cannot be accurately depicted by a single Lewis structure. Resonance structures are used to represent delocalized electrons, as in ozone (O3) and carbonate (CO32-).

Actual structure is a hybrid of all resonance forms.
Delocalization stabilizes the molecule.

Practice and Examples

Draw Lewis structures for CO32- and SO42-.
Identify which molecules are exceptions to the octet rule (e.g., I3-).
Determine the most stable structure of hydrogen cyanide (HCN) based on formal charges.

Summary Table: Bond Types and Properties

Bond Type	Electron Behavior	Example
Ionic	Transferred	NaCl
Covalent	Shared	H2O
Metallic	Pooled	Na (metal)

Additional info:

Expanded explanations of resonance and formal charge were added for clarity.
Practice problems and examples are included to reinforce concepts.