Chapter 10: Chemical Bonding and Molecular Structure

Types of Chemical Bonds

Introduction to Chemical Bonding

Chemical bonds form because they lower the potential energy between the charged particles that compose atoms. The interactions between electrons and nuclei of different atoms, as described by Coulomb's Law, lead to bond formation. However, electrons and nuclei also repel each other, and the net effect of these attractions and repulsions determines whether a bond forms.

• Ionic Bonds: Formed between metals and nonmetals; involve electron transfer.

• Covalent Bonds: Formed between nonmetals; involve electron sharing.

• Metallic Bonds: Formed between metals; involve pooling of electrons.

Example: Sodium (Na, a metal) and chlorine (Cl, a nonmetal) form NaCl via ionic bonding.

Ionic Bonding

In ionic bonding, metals with low ionization energy lose electrons to form cations, while nonmetals with high electron affinity gain electrons to form anions. The resulting electrostatic attraction lowers the overall potential energy, forming a stable ionic compound.

• Key Point: Ionic bonds result from electron transfer and the attraction between oppositely charged ions.

• Example: NaCl, MgO, CaF2

Covalent Bonding

In covalent bonding, atoms share electrons, which interact with the nuclei of both bonding atoms, lowering their potential energy.

• Key Point: Covalent bonds involve electron sharing between nonmetals.

• Example: H2O, Cl2, H2

Valence Electrons and Lewis Structures

Representing Valence Electrons

Valence electrons are responsible for the chemical properties and bonding of main group elements. The most stable configuration is when atoms have eight valence electrons (octet rule), except for helium, which is stable with two (duet).

• Key Point: Lewis dot structures represent valence electrons as dots around the element symbol.

• Example: Oxygen: :O:

Lattice Energy

Definition and Trends

Lattice energy is the energy associated with forming a crystalline lattice of alternating cations and anions from gaseous ions. It is a major contributor to the exothermic nature of ionic compound formation.

• Key Point: Lattice energy increases with higher ionic charge and smaller ionic radius.

• Formula: $E_{lattice} \propto \frac{Q_1 Q_2}{r}$

• Example: CaO has higher lattice energy than KCl due to higher charges.

Properties of Ionic Compounds:

• High melting points (due to strong lattice energy)

• Poor electrical conductivity as solids (ions are fixed)

• Good electrical conductivity in solution (ions are mobile)

Covalent Bonding and Lewis Structures

Single, Double, and Triple Bonds

Covalent bonds can involve one, two, or three shared pairs of electrons, resulting in single, double, or triple bonds, respectively.

• Single Bond: H2O, Cl2, H2

• Double Bond: O2

• Triple Bond: N2

Lewis Structure Guidelines

• Hydrogen atoms are always terminal.

• More electronegative elements are placed at terminal positions.

• Calculate total valence electrons (add for anions, subtract for cations).

• Distribute electrons to give octets (or duets for H and He).

• Form double or triple bonds if necessary to satisfy octet rule.

Electronegativity and Bond Polarity

Electronegativity Trends

Electronegativity is the ability of an atom to attract electrons in a chemical bond. It increases across a period and decreases down a group, similar to ionization energy and electron affinity.

• Key Point: The smaller the atom, the higher the electronegativity.

• Example: F is the most electronegative element.

Bond Polarity

• Polar Covalent Bond: Uneven sharing of electrons (e.g., H–F).

• Nonpolar Covalent Bond: Even sharing of electrons (e.g., H2).

• Ionic Bond: Complete transfer of electrons (e.g., NaCl).

Bond Type by Electronegativity Difference:

Electronegativity Difference	Bond Type
0	Pure Covalent
0.1 – 0.4	Nonpolar Covalent
0.5 – 1.9	Polar Covalent
≥ 2.0	Ionic

Dipole Moment

The dipole moment ($\mu$) measures bond polarity and is calculated as:

$\mu = q \times r$

• Measured in Debyes (D).

• Greater electron sharing and larger atoms result in larger dipole moments.

Example: HCl has a significant dipole moment due to the difference in electronegativity between H and Cl.

*Additional info: The notes cover the core concepts of chemical bonding, including ionic, covalent, and metallic bonds, lattice energy, Lewis structures, electronegativity, and bond polarity, which are all central to General Chemistry Chapter 10.*