Chapter 4: Chemical Bonding – Study Notes for General Chemistry

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Chemical Bonding

Introduction to Chemical Bonds

Chemical bonding is the process by which atoms combine to form compounds. The nature of the bond formed depends on the types of atoms involved and how electrons are distributed between them. Understanding chemical bonds is fundamental to predicting the properties and behaviors of substances.

Chemical bond: The force that holds two or more atoms together in a molecule or compound.
Three main types of chemical bonds: ionic, covalent, and metallic.

Types of Chemical Bonds

The three primary types of chemical bonds differ in the way electrons are distributed and the types of elements involved.

Type	Elements Involved	Electron Distribution	Example (Microscopic View)	Example (Macroscopic View)
Ionic	Metals and nonmetals	Transferred	NaCl lattice	Table salt
Covalent	Nonmetals and metalloids	Shared	H2O molecules	Water
Metallic	Metals	Delocalized	Cu lattice	Copper wire

Ionic Bonds

Electrostatic Attractions in Ionic Compounds

Ionic bonds form between cations (positively charged ions) and anions (negatively charged ions) due to electrostatic attraction. The strength of this attraction is described by the electrostatic potential energy equation:

Q1 and Q2: Charges of the ions
d: Distance between the centers of the ions
Energy is directly proportional to the product of the charges and inversely proportional to the distance between them.

Formation and Stability of Ionic Bonds

An ionic bond is formed when a metal transfers electrons to a nonmetal, resulting in the formation of cations and anions.
The resulting electrostatic attraction leads to the formation of a stable ionic compound.
The greater the charge and the smaller the distance between ions, the stronger (more negative) the potential energy and the more stable the compound.

Lattice Energy

Lattice energy is the energy released when one mole of an ionic crystalline compound forms from its free ions in the gas phase. It is a measure of the strength of the ionic bonds in a solid.

Compound	Lattice Energy (kJ/mol)
LiF	1036
LiCl	853
NaF	923
NaCl	787
NaBr	747
KCl	715
KBr	682
MgO	3795
MgCl2	2526

Additional info: Lattice energy increases with higher ionic charges and decreases with larger ionic radii.

Covalent Bonds

Formation and Properties of Covalent Bonds

A covalent bond is formed when two atoms share one or more pairs of electrons. This type of bonding typically occurs between nonmetals.

Bond length: The distance between the nuclei of two bonded atoms.
Bond energy: The energy required to break one mole of covalent bonds in the gas phase.

Bond length and bond energy are inversely related: shorter bonds are generally stronger and require more energy to break.

Bond Polarity and Electronegativity

Bond polarity arises from the unequal sharing of electrons between atoms with different electronegativities (EN).

Electronegativity (EN): A measure of an atom's ability to attract electrons in a chemical bond (Pauling scale, 0.7–4.0, no units).
If ΔEN < 0.5: Nonpolar covalent bond
If 0.5 ≤ ΔEN < 2.0: Polar covalent bond
If ΔEN ≥ 2.0: Ionic bond

Electronegativity increases across a period (left to right) and decreases down a group (top to bottom). Elements with the highest EN: F, O, N, Cl.

Naming and Writing Chemical Formulas

Molecular Compounds

Molecular compounds are formed between nonmetals. Their names and formulas follow specific rules:

The first element is named first; the second element is named with an -ide suffix.
Prefixes (mono-, di-, tri-, etc.) indicate the number of each atom present.
"Mono-" is usually omitted for the first element.

Example: CO2 is carbon dioxide; N2O5 is dinitrogen pentoxide.

Ionic Compounds

Ionic compounds are named by stating the cation first (element name), followed by the anion (element name with -ide ending).

For transition metals, use Roman numerals to indicate the charge (e.g., iron(III) chloride).
Polyatomic ions retain their common names (e.g., sulfate, nitrate).

Example: NaCl is sodium chloride; FeCl3 is iron(III) chloride.

Lewis Structures

Drawing Lewis Structures

Lewis structures represent the arrangement of valence electrons among atoms in a molecule.

Count the total number of valence electrons.
Arrange the atoms and connect them with single bonds.
Complete the octet for each atom (except hydrogen).
Place remaining electrons as lone pairs.
Check that the total number of electrons matches the sum from step 1.

Multiple bonds (double or triple) may be needed to satisfy the octet rule.

Resonance Structures

Some molecules can be represented by two or more valid Lewis structures, called resonance structures. The actual structure is a resonance hybrid, which is an average of the contributing structures.

Resonance increases stability by delocalizing electrons.
Examples: O3, NO3-, CO32-

Formal Charge

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

FC = (number of valence electrons in free atom) – (number of lone pair electrons) – (1/2 number of bonding electrons)
The best structure has formal charges closest to zero, with negative charges on the most electronegative atoms.

Exceptions to the Octet Rule

Electron-Deficient Molecules

Some elements (e.g., H, Be, B, Al) form molecules with fewer than 8 electrons around the central atom.
Example: BeH2 (4 electrons), BF3 (6 electrons).

Odd-Electron Molecules (Radicals)

Molecules with an odd number of electrons (e.g., NO, NO2) are called radicals.
These are highly reactive and often dimerize to achieve stability.

Expanded Octet

Elements in period 3 and beyond can have more than 8 electrons (expanded octet) due to available d orbitals.
Examples: SF6, PCl5, ICl4-

Summary Table: Types of Chemical Bonds

Bond Type	Electron Behavior	Typical Elements	Example
Ionic	Transferred	Metal + Nonmetal	NaCl
Covalent	Shared	Nonmetal + Nonmetal	H2O
Metallic	Delocalized	Metal + Metal	Cu