Chemical Bonding and Molecular Geometry

Compounds and Ionic Bonding

A compound is a substance composed of two or more elements combined in a specific ratio and held together by chemical bonds. Ionic bonding refers to the electrostatic attraction that holds oppositely charged ions together in an ionic compound. The attraction between the cation and anion forms a stable, electrically neutral compound.

• Ionic compounds typically consist of a metal ion and a non-metal ion.

• The chemical formula denotes the constituent elements and their ratio.

• Example: Sodium chloride (NaCl) is formed by the transfer of an electron from sodium (Na) to chlorine (Cl).

Classification of Elements: Metals, Nonmetals, and Metalloids

Elements are categorized based on their physical and chemical properties:

• Metals: Good conductors of heat and electricity, usually solid, ductile, and malleable.

• Nonmetals: Poor conductors, occur in all physical states, and are not ductile or malleable.

• Metalloids: Exhibit intermediate properties, are semiconductors, and have both metallic and nonmetallic characteristics.

Lattice Energy and Stability of Ionic Compounds

A three-dimensional array of oppositely-charged ions forms a lattice. Lattice energy is the energy required to convert a mole of ionic solid to its constituent ions in the gas phase. The magnitude of lattice energy is a measure of an ionic compound’s stability and depends on the charge and distance between ions.

• Lattice energy increases with higher charge and smaller ion size.

• Formula:

Electron Configurations of Ions and Isoelectronic Species

To write the electron configuration of an ion:

• Write the configuration for the atom.

• Add electrons for negative ions; remove electrons for positive ions.

• Species with identical electron configurations to a noble gas are called isoelectronic.

Types of Chemical Bonds: Covalent, Polar Covalent, and Ionic

There are three main types of chemical bonds:

• Pure covalent bond: Electrons are equally shared between neutral atoms.

• Polar covalent bond: Electrons are unequally shared, resulting in partial charges (δ+ and δ–).

• Ionic bond: Complete transfer of electrons, forming oppositely charged ions.

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract electrons in a compound. The difference in electronegativity between atoms determines bond polarity:

• Difference < 0.5: Nonpolar covalent

• Difference 0.5–2.0: Polar covalent

• Difference > 2.0: Ionic

Dipole Moment and Percent Ionic Character

The dipole moment (μ) quantifies bond polarity:

• Formula:

• Measured in debye units (D).

• Percent ionic character compares observed and calculated dipole moments.

• Formula:

Naming and Writing Formulas for Ionic Compounds

Ionic compounds can be formed from:

• Metal and non-metal

• Metal and polyatomic ion

• Polyatomic ion and non-metal

• Polyatomic ion and polyatomic ion

Rules for naming:

• Name the cation (use Roman numerals for variable charge).

• Name the anion, changing the ending to -ide or using the polyatomic ion name.

• Formulas must be electrically neutral; use cross multiplication or sum of charges to determine ratios.

Polyatomic Ions and Oxoanions

Polyatomic ions are covalently bonded groups of atoms with a net charge. Oxoanions contain oxygen and another element:

• Per-...ate: One more O than -ate

• -ite: One less O than -ate

• Hypo...ite: Two fewer O than -ate

Naming Molecular Compounds

Binary molecular compounds consist of two nonmetals. Greek prefixes denote the number of atoms:

• Mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-

• First element: prefix + name (omit mono-)

• Second element: prefix + name ending in -ide

Compounds Containing Hydrogen and Acids

Hydrogen-containing compounds often have nonsystematic names. Binary acids are named by changing hydrogen to hydro- and the anion ending to -ic.

• Example: Hydrogen chloride → Hydrochloric acid

• Oxoacids: -ate → -ic acid, -ite → -ous acid

Hydrates

A hydrate is a compound with a specific number of water molecules in its solid structure. The anhydrous form contains no water.

• Example: CuCl2·2H2O is copper(II) chloride dihydrate.

Lewis Dot Symbols and Valence Electrons

Valence electrons are the outermost electrons involved in bonding. Lewis dot symbols represent valence electrons as dots around the element symbol.

• For main group metals, dots equal electrons lost.

• For nonmetals, unpaired dots indicate possible bonds.

Lewis Structures and Multiple Bonds

Lewis structures show covalent bonding with shared electron pairs as dashes or dots. Lone pairs are shown as pairs of dots. Multiple bonds (double, triple) involve more shared pairs and are shorter and stronger than single bonds.

• Bond length: Distance between nuclei of bonded atoms.

• Bond strength: Energy required to break the bond.

Formal Charge and Resonance

Formal charge helps determine the most plausible Lewis structure. Resonance structures represent molecules where electron distribution cannot be shown by a single Lewis structure.

• Formal charge = Valence electrons – Associated electrons

• Best structure: Formal charges closest to zero and consistent with electronegativity.

Exceptions to the Octet Rule

Some molecules do not follow the octet rule:

• Central atom has fewer than eight electrons (e.g., boron compounds).

• Odd number of electrons (free radicals).

• Central atom has more than eight electrons (expanded octet, third period and beyond).

Coordinate Covalent Bonds (Dative Bonds)

A coordinate covalent bond occurs when both electrons in a bond are donated by one atom. This is common in Lewis acid-base reactions.

Molecular Geometry: VSEPR Model

The Valence-Shell Electron-Pair Repulsion (VSEPR) model predicts molecular shape based on electron domain repulsion. Electron domains include lone pairs, single, double, and triple bonds.

• ABx notation: A is central atom, x is number of surrounding atoms.

• Common geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Electron-Domain Geometry vs. Molecular Geometry

Electron-domain geometry is the arrangement of electron domains; molecular geometry is the arrangement of bonded atoms. Lone pairs affect molecular geometry.

Bond Angles and Deviations from Ideal Geometry

Lone pairs and multiple bonds cause deviations from ideal bond angles due to increased repulsion.

Molecular Polarity

Molecular polarity depends on both bond polarity and molecular geometry. A molecule is polar if the vector sum of bond dipoles does not equal zero.

• CO2: Polar bonds, nonpolar molecule due to linear geometry.

• H2O: Polar bonds, polar molecule due to bent geometry.

• BF3, CCl4: Polar bonds, nonpolar molecules due to symmetrical geometry.

• CHCl3: Polar bonds, polar molecule due to asymmetrical geometry.

Dipole Moments and Structural Isomers

Dipole moments can distinguish between structural isomers, such as cis- and trans-dichloroethylene.

Bonding Forces: Intramolecular vs. Intermolecular

Intramolecular forces (covalent bonds) are strong attractions within molecules. Intermolecular forces (IMFs) are weaker attractions between molecules and include dispersion, dipole-dipole, and hydrogen bonding.

Types of Intermolecular Forces

• Dispersion forces: Weakest, arise from instantaneous dipoles, present in all molecules.

• Dipole-dipole forces: Stronger, between molecules with permanent dipoles.

• Hydrogen bonds: Strongest dipole-dipole, occur when H is bonded to O, N, or F.

Trends in Intermolecular Forces

Stronger IMFs lead to higher melting and boiling points. Molecular weight and surface area also increase these properties. The order of IMF strength is: Ionic > H-bonding > dipole/dipole > dispersion.

• Example: Sodium chloride (ionic) has a higher boiling point than water (hydrogen bonding) and chlorine gas (dispersion).