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

Study Guide - Smart Notes

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Bond Dissociation Energy and Covalent Bond Strength

Understanding Bond Dissociation Energy (D)

Bond dissociation energy (D) is a measure of the strength of a covalent bond, representing the energy required to break the bond between two atoms in a molecule. This concept is essential for comparing the stability and reactivity of different chemical bonds.

Bond Dissociation Energy (D): The energy needed to break a specific bond in a molecule in the gas phase.
Bond Order: Refers to the number of shared electron pairs between two atoms. Higher bond order generally means stronger bonds.
Example: Triple bonds (bond order = 3) have higher bond dissociation energies than double (bond order = 2) or single bonds (bond order = 1).

Types of Covalent Bonds: Polar, Nonpolar, and Ionic

Comparing Bond Types

Covalent bonds can be classified as polar or nonpolar based on the difference in electronegativity between the bonded atoms. Ionic bonds result from the complete transfer of electrons between atoms.

Polar Covalent Bonds: Electrons are shared unequally due to differences in electronegativity.
Nonpolar Covalent Bonds: Electrons are shared equally between atoms with similar electronegativity.
Ionic Bonds: Electrons are transferred from one atom to another, resulting in oppositely charged ions.
Bonding Types: There is a continuum of bond types, with varying degrees of electron sharing and transfer.
Example: HCl (polar covalent), Cl2 (nonpolar covalent), NaCl (ionic).

Electronegativity and Bond Polarity

Understanding Electronegativity

Electronegativity is the tendency of an atom to attract electrons in a chemical bond. Differences in electronegativity determine the polarity of bonds and molecules.

Electronegativity Scale: Commonly measured using the Pauling scale.
Bond Polarity: The greater the difference in electronegativity, the more polar the bond.
Dipole Moment: A measure of bond polarity, represented as a vector in molecular drawings.
Example: Water (H2O) has a significant dipole moment due to the difference in electronegativity between H and O.

Characteristics of Ionic and Covalent Compounds

General Properties

Ionic and covalent compounds exhibit distinct physical and chemical properties, largely due to the nature of their 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: Ionic compounds have strong electrostatic forces; covalent compounds may have weaker forces such as hydrogen bonding, dipole-dipole, or London dispersion forces.

Lewis Structures and Bonding Principles

Drawing Lewis Structures

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

Bonding Pairs: Shared pairs of electrons between atoms.
Lone Pairs: Non-bonding pairs of electrons localized on a single atom.
Octet Rule: Atoms tend to form bonds to achieve eight electrons in their valence shell.
Types of Bonds: 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

Application of Lewis Structures

Lewis formulas are used to represent molecules and polyatomic ions, showing all valence electrons and bonds.

Polyatomic Ions: Charged species composed of two or more atoms covalently bonded.
Example: The sulfate ion (SO42-) requires careful placement of charges and bonds in its Lewis structure.

Exceptions to the Octet Rule

Limitations and Special Cases

Not all elements follow the octet rule. Exceptions include molecules with odd numbers of electrons, atoms with expanded octets, and electron-deficient species.

Expanded Octets: Elements in period 3 and beyond can have more than eight electrons (e.g., SF6).
Odd-Electron Molecules: Molecules like NO have an odd number of electrons and cannot satisfy the octet rule for all atoms.
Electron-Deficient Molecules: Some compounds, such as BF3, have fewer than eight electrons around the central atom.

Resonance Structures

Understanding Resonance

Resonance structures are alternative Lewis structures for a molecule or ion that differ only in the placement of electrons. Resonance helps explain the delocalization of electrons and the stability of certain compounds.

Resonance Structures: Multiple valid Lewis structures for the same molecule, connected by double-headed arrows.
Bond Lengths: Resonance leads to bond lengths that are intermediate between single and double bonds.
Example: The nitrate ion (NO3-) has three resonance structures.

Formal Charge and Resonance Evaluation

Assigning Formal Charges

Formal charge is used to evaluate the stability and plausibility of Lewis structures and resonance forms.

Formal Charge Formula:

Best Resonance Structure: The most stable resonance structure has the lowest formal charges and places negative charges on the most electronegative atoms.
Example: In the carbonate ion (CO32-), formal charges help identify the most reasonable resonance forms.

Summary Table: Bond Types and Properties

Comparison of Bond Types

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

Additional info: Some explanations and examples have been expanded for clarity and completeness.