Chapter 10 Chemical Bonding

Chemical Bonding: The Lewis Model

Bonding Theories

• Explain how and why atoms attach to form molecules.
• Predict molecular shapes and chemical/physical properties.

Lewis Model

• Simple bonding theory emphasizing valence electrons.
• Valence electrons are represented as dots in Lewis structures.
• Predicts molecular stability and properties.

Why Atoms Bond

• Bonds form to lower potential energy between charged particles.
• Potential energy is calculated considering:
  • Nucleus-to-nucleus repulsions
  • Electron-to-electron repulsions
  • Nucleus-to-electron attractions

Types of Bonds

• Ionic: Electrons transferred between metal and nonmetal.
• Covalent: Electrons shared between nonmetals.
• Metallic: Electrons pooled in metals.

Ionic Bonds

• Metals lose electrons (low ionization energy) to form cations.
• Nonmetals gain electrons (high electron affinity) to form anions.
• Oppositely charged ions attract, forming an ionic bond.

Covalent Bonds

• Nonmetals share valence electrons (high ionization energies).
• Potential energy is lowest when electrons are between nuclei.

Metallic Bonds

• Metals easily lose electrons (low ionization energy).
• Valence electrons are shared as a pool among atoms/ions.
• Cations exist in a "sea" of delocalized electrons.
• Bonding arises from cation attraction to delocalized electrons.

Valence Electrons

• Lewis model focuses on valence electrons because they are most important in bonding.
• Main-group column number indicates the number of valence electrons.
• Transition elements typically have two valence electrons.

Lewis Structures of Atoms

• Valence electrons are represented as dots around the element symbol.
• Pair the first two dots for s orbital electrons.
• Add one dot on each open side for the first three p electrons and then pair the rest.

Lewis Bonding Theory

• Atoms bond to achieve a more stable electron configuration (lower potential energy).
• Atoms transfer or share electrons to obtain an outer shell with eight electrons (octet rule).

Charge and Ion Stability

• Metals lose valence shell electrons to form cations.
• Nonmetals gain valence electrons to form anions.

Lewis Theory and Ionic Bonding

• Lewis symbols show electron transfer from metal to nonmetal.
• Predicts the number of electrons a metal atom should lose/nonmetal atom should gain to attain a stable electron arrangement (octet rule).
• Predicts formulas of ionic compounds.
• Predicts relative strengths of ionic bonds using Coulomb’s law.

Ionic Bonding and Crystal Lattice

• Crystal lattice is an alternating cation and anion structure.
• Held together by electrostatic attraction, maximizing stability.

Lattice Energy

• Lattice energy: energy released when a solid crystal forms from separate gas-state ions.
• Always exothermic.
• Depends on the size of charges and inversely on the distance between ions.

Born-Haber Cycle

• A hypothetical series of reactions representing the formation of an ionic compound from its constituent elements.
• Uses Hess’s law to determine lattice energy by summing enthalpy changes.
• ΔHf°(salt) = ΔHf°(metal atoms, g) + ΔHf°(nonmetal atoms, g) + IE + EA + ΔH{lattice}Where IE is ionization energy, and EA is electron affinity.

Trends in Lattice Energy

• Larger ionic radius = less exothermic lattice energy.
• Increasing magnitude of ionic charge = more exothermic lattice energy.
• Ion charge is generally more important than ion size.

Ionic Bonding Model vs. Reality

• Predicts high melting/boiling points due to strong attractions between ions.
• Predicts brittleness because displacing ions leads to repulsive forces.
• Predicts ionic solids do not conduct electricity because ions are locked in position.
• Predicts liquid ionic compounds and aqueous solutions conduct electricity because ions can move.

Covalent Bonding: Lewis Structures

• Lewis model is useful for covalent bonding.
• Atoms share valence electrons to attain octets (or duets for hydrogen).
• Shared electrons count toward each atom’s octet.
• Shared pair of electrons = covalent bond.

Bonding and Lone Pairs

• Bonding pairs: electrons shared by atoms.
• Lone pairs: electrons not shared but belong to a particular atom.

Single, Double, and Triple Covalent Bonds

• Single: One pair of electrons shared.
• Double: Two pairs of electrons shared.
• Triple: Three pairs of electrons shared.

Covalent Bonding: Model vs. Reality

• Predicts which combinations of atoms form stable molecules.
• Predicts covalent bonds are highly directional.
• Predicts covalently bonded compounds will be found as individual molecules with low melting and boiling points.
• Predicts molecular compounds do not conduct electricity.

Electronegativity and Bond Polarity

• Unequal sharing of electrons results in a polar covalent bond.
• One end of the bond has larger electron density (partial negative charge).
• The other end is electron deficient (partial positive charge).

Bond Polarity Classification

• Covalent: electron transfer is insufficient for the material to display the classic properties of ionic compounds.
• Polar covalent: unequal sharing produces a dipole.

Electronegativity

• The ability of an atom to attract bonding electrons to itself.
• Increases across a period (left to right) and decreases down a group (top to bottom).
• Fluorine is the most electronegative element; Francium is the least.
• Larger electronegativity difference = more polar bond.

Electronegativity Difference and Bond Type

• 0: Pure covalent (equal sharing).
• 0.1-0.4: Nonpolar covalent.
• 0.4-1.9: Polar covalent.
• ≥2.0: Ionic.

Bond Dipole Moments

• Dipole moment (\mu): measure of bond polarity.
• \mu = q \times r where q is the magnitude of the partial charges, and r is the distance between them.
• Measured in Debyes (D).
• More shared electrons and larger atoms generally result in larger dipole moments.

Percent Ionic Character

• Percentage of a bond’s measured dipole moment compared to complete electron transfer.
• Indicates the degree to which the electron is transferred.

Writing Lewis Structures of Molecules

• Write the correct skeletal structure (Hydrogen atoms are always terminal).
• Calculate the total number of valence electrons.
• Distribute electrons to give octets (or duets for hydrogen) to atoms.
• Form double or triple bonds if atoms lack an octet.

Resonance

• Extensions of Lewis model suggest delocalization of electrons.
• Delocalization of charge helps stabilize the molecule.

Resonance Structures

• Lewis structures that differ only in electron positions.
• The actual molecule is a resonance hybrid (combination of resonance forms).

Formal Charge

• Fictitious charge assigned to each atom in a Lewis structure.
• Helps to distinguish among competing Lewis structures.
• FC = # \text{ valence } e^- - [\text{nonbonding } e^- + \frac{1}{2} \text{bonding } e^-]
• Sum of formal charges in a molecule = 0; in an ion, equals the charge.

Evaluating Resonance Structures

• Sum of formal charges in a neutral molecule must be zero.
• Sum of formal charges in an ion must equal the charge of the ion.
• Small (or zero) formal charges are better.
• Negative formal charge should reside on the most electronegative atom.

Exceptions to the Octet Rule

• Odd number electron species (free radicals) are very reactive.
• Incomplete octets (B, Al).
• Expanded octets (elements with empty d orbitals).

Bond Energies

• Bond energy (bond enthalpy): energy to break one mole of a bond in the gas phase.
• Always positive (endothermic).
• \Delta H_{rxn}° can be estimated as the sum of enthalpy changes associated with breaking bonds in reactants and forming bonds in products.

Trends in Bond Energies

• More shared electrons = stronger bond (higher bond energy).
• Shorter bond = stronger bond.

Bond Lengths

• Bond length: the distance between the nuclei of bonded atoms.
• More shared electrons = shorter bond.