Bonding, Electronegativity, and Resonance: Key Concepts in General Chemistry

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

Bonding and Bond Dissociation Energy

Bond Dissociation Energy (D)

Bond dissociation energy is a measure of the strength of a covalent bond, representing the energy required to break the bond between two atoms in a molecule. It is a key concept for comparing the stability and strength of different chemical bonds.

Definition: The energy needed to break one mole of a specific bond in a gaseous molecule.
Bond Strength: Higher bond dissociation energy indicates a stronger bond.
Bond Order: The term "bond order" refers to the number of shared electron pairs between two atoms. Higher bond order generally means stronger bonds.
Example: Triple bonds (e.g., in N2) have higher bond dissociation energies than double or single bonds.

Types of Covalent Bonds

Polar, Nonpolar, and Ionic Bonds

Covalent bonds can be classified based on the distribution of electrons between atoms. Understanding these differences is essential for predicting molecular properties.

Polar Covalent Bonds: Electrons are shared unequally due to differences in electronegativity. Example: H2O.
Nonpolar Covalent Bonds: Electrons are shared equally. Example: Cl2.
Ionic Bonds: Electrons are transferred from one atom to another, resulting in charged ions. Example: NaCl.
Bond Polarity: Determined by the difference in electronegativity between atoms.
Continuous Range: Bond types exist on a spectrum, with varying degrees of polarity and electron sharing.

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is the tendency of an atom to attract electrons in a chemical bond. It is a fundamental property that influences bond polarity and molecular behavior.

Definition: Quantitative measure of an atom's ability to attract shared electrons.
Periodic Trends: Electronegativity increases across a period and decreases down a group.
Bond Polarity: The greater the difference in electronegativity, the more polar the bond.
Dipole Moment: A vector quantity representing the separation of charge in a molecule. Can be illustrated in molecular drawings.
Example: In HCl, Cl is more electronegative than H, resulting in a polar bond.

Characteristics of Ionic and Covalent Compounds

General Properties

Ionic and covalent compounds exhibit distinct physical and chemical properties due to differences in bonding.

Ionic Compounds: High melting and boiling points, conduct electricity when dissolved in water, form crystalline solids.
Covalent Compounds: Lower melting and boiling points, poor electrical conductivity, can exist as gases, liquids, or solids.
Intermolecular Forces: Forces between molecules, such as hydrogen bonding, dipole-dipole interactions, and London dispersion forces, affect physical properties.
Strength Comparison: Ionic bonds are generally stronger than covalent bonds due to electrostatic attraction.

Lewis Structures and Bonding Principles

Lewis Dot Structures

Lewis structures are diagrams that represent the bonding between atoms and the distribution of valence electrons in a molecule.

Bonding Pairs: Shared pairs of electrons between atoms.
Lone Pairs: Non-bonding pairs of electrons localized on an atom.
Octet Rule: Atoms tend to form bonds to achieve eight electrons in their valence shell.
Bond Types: Single, double, and triple bonds represent one, two, or three shared pairs of electrons, respectively.
Coordinate Covalent Bonds: Both electrons in the bond originate from the same atom.
Example: CO2 has two double bonds between C and O.

Writing Lewis Formulas for Molecules and Polyatomic Ions

Lewis Structures for Polyatomic Ions

Polyatomic ions are charged species composed of two or more atoms covalently bonded. Writing Lewis structures for these ions requires accounting for the overall charge.

Procedure: Count total valence electrons, adjust for charge, distribute electrons to satisfy the octet rule.
Example: The sulfate ion (SO42−) has 32 valence electrons (including the charge).
Note: Familiarity with writing Lewis structures for ions is essential for understanding molecular geometry and reactivity.

Exceptions to the Octet Rule

Limitations and Exceptions

Not all molecules and ions follow the octet rule. Several important exceptions exist, especially for elements in period 3 and beyond, and for molecules with odd numbers of electrons.

Expanded Octet: Elements in period 3 or higher (e.g., P, S, Cl) can have more than eight electrons in their valence shell.
Odd-Electron Molecules: Molecules with an odd number of valence electrons (e.g., NO) cannot satisfy the octet rule for all atoms.
Incomplete Octet: Some elements (e.g., B, Be) may have fewer than eight electrons.
Example: PF5 has 10 electrons around phosphorus (expanded octet).

Resonance Structures

Resonance and Bond Lengths

Resonance structures are alternative Lewis structures for a molecule or ion that differ only in the arrangement of electrons. They help explain delocalized bonding and observed bond lengths.

Definition: Resonance occurs when more than one valid Lewis structure can be drawn for a molecule.
Bond Lengths: Actual bond lengths are intermediate between those predicted by individual resonance structures.
Example: Ozone (O3) has two resonance structures, resulting in equal bond lengths.
Resonance Hybrid: The true structure is a weighted average (hybrid) of all resonance forms.

Evaluating Resonance Structures

Criteria for Resonance Importance

Not all resonance structures contribute equally to the resonance hybrid. Chemists use specific criteria to evaluate their relative importance.

Formal Charge: Structures with formal charges closest to zero are favored.
Electronegativity: Negative formal charges are more stable on more electronegative atoms.
Octet Rule: Structures in which all atoms satisfy the octet rule are preferred.
Example: In the nitrate ion (NO3−), three resonance structures distribute the negative charge over the oxygen atoms.

Assigning Formal Charges

Formal Charge Calculation

Formal charge is a bookkeeping tool used to determine the distribution of electrons in a molecule. It is essential for evaluating resonance structures and predicting molecular stability.

Formula:

Application: Assign formal charges to each atom in a molecule to identify the most stable resonance structure.
Example: In CO2, both oxygen atoms have a formal charge of 0 in the most stable structure.

Summary Table: Bond Types and Properties

The following table summarizes the main types of chemical bonds and their key properties.

Bond Type	Electron Sharing	Polarity	Example
Ionic	Transfer	High	NaCl
Polar Covalent	Unequal Sharing	Moderate	H2O
Nonpolar Covalent	Equal Sharing	Low	Cl2

Additional info: Academic context and examples have been expanded for clarity and completeness.