Chemical Bonding and Molecular Structures

Ionic Compounds and Lattice Energy

Ionic compounds are formed when atoms transfer electrons, resulting in the formation of positively charged cations and negatively charged anions. These ions are held together by strong electrostatic forces known as ionic bonds.

• Ions in Ionic Compounds: Ions are held together in a repeating three-dimensional arrangement called a crystal lattice.

• Lattice Energy: The energy required to separate one mole of an ionic solid into its gaseous ions. It is a measure of the strength of the ionic bonds in the crystal lattice.

• Crystal Lattice: A regular, repeating arrangement of ions in an ionic solid.

• Energy Changes: When two neutral atoms approach, potential energy decreases as they attract, reaching a minimum at the bond length. If they get too close, repulsion increases the energy.

Equation for Lattice Energy (Born-Haber Cycle):

Ionic vs. Covalent Bonding

• Ionic Bonding: Involves the transfer of electrons from a metal to a nonmetal, forming ions.

• Covalent Bonding: Involves the sharing of electrons between two nonmetals.

• Comparison: Ionic compounds tend to have high melting points and conduct electricity when molten; covalent compounds have lower melting points and do not conduct electricity.

Naming and Writing Formulas

• Binary Ionic Compounds: Consist of two elements: a metal and a nonmetal. Name the metal first, then the nonmetal with an '-ide' ending.

• Transition Metals: Use Roman numerals to indicate the metal's charge (e.g., Iron(III) chloride: FeCl3).

• Polyatomic Ions: Compounds containing ions like NO3-, SO42-, etc.

• Acids: Binary acids (e.g., HCl) are named 'hydro-' + root + '-ic acid'; oxyacids are named based on the polyatomic ion.

• Binary Molecular Compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom (e.g., CO2: carbon dioxide).

The Rule of Eight (Octet Rule)

Atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell, resembling the electron configuration of noble gases.

Valence and Core Electrons

• Valence Electrons: Electrons in the outermost shell, involved in bonding.

• Core Electrons: Electrons in inner shells, not usually involved in bonding.

Avogadro’s Number and Mole Calculations

• Avogadro’s Number: particles per mole.

• Conversions: To convert grams to moles, divide by the molar mass; to convert moles to grams, multiply by the molar mass.

Formula:

Polar and Nonpolar Covalent Bonds

• Polar Covalent Bond: Electrons are shared unequally due to a difference in electronegativity.

• Nonpolar Covalent Bond: Electrons are shared equally between atoms of similar electronegativity.

Electronegativity and Trends

• Electronegativity: The ability of an atom to attract shared electrons in a bond.

• Trend: Increases across a period (left to right) and decreases down a group (top to bottom) on the periodic table.

Resonance Structures and Formal Charges

• Resonance Structures: Different valid Lewis structures for the same molecule, showing delocalized electrons.

• Formal Charge: Used to determine the most stable Lewis structure. Calculated as:

• Choosing Lewis Structures: The structure with the lowest formal charges and negative charges on the most electronegative atoms is preferred.

Expanded Octets

• Explanation: Atoms in period 3 or higher can have more than eight electrons due to available d orbitals (e.g., SF6).

Bond Order, Bond Length, and Bond Strength

• Bond Order: The number of chemical bonds between a pair of atoms (single = 1, double = 2, triple = 3).

• Bond Length: Inversely related to bond order; higher bond order means shorter bond length.

• Bond Strength: Directly related to bond order; higher bond order means stronger bonds.

Lewis Structures

• Writing Lewis Structures: Show all valence electrons as dots or lines, indicating bonds and lone pairs.

Molecular Shapes and Bond Angles

The shape of a molecule is determined by the arrangement of atoms and electron pairs around the central atom, described by VSEPR (Valence Shell Electron Pair Repulsion) theory.

Predicting Molecular Shape

• Draw the Lewis structure.

• Count the number of bonding pairs and lone pairs on the central atom (steric number).

• Use VSEPR theory to predict the shape.

Molecular Polarity

• Polarity: Determined by the difference in electronegativity and the symmetry of the molecule.

• Polar Molecule: Has a net dipole moment (e.g., H2O).

• Nonpolar Molecule: Symmetrical, dipoles cancel (e.g., CO2).

Steric Number

• Steric Number: The total number of atoms bonded to the central atom plus the number of lone pairs on the central atom.

• Used to determine molecular geometry.

Additional info: This summary expands on the learning objectives by providing definitions, explanations, and examples for each key concept in chemical bonding and molecular structure, as outlined in a typical general chemistry curriculum.