Ionic Bonding Concepts

Ionic Bonding

• Ionic bonds are forces holding charged particles together in compounds.

Valence Electrons

• Valence electrons are outermost shell electrons.

Cations

• Metals form cations.

Anions

• Nonmetals form anions.

Endothermic

• Energy is absorbed during endothermic reactions.

Exothermic

• Energy is released during exothermic reactions.

Ionic Compound Alias

• An ionic compound is also called a salt.

Crystal Lattice Structure

• Ionic compounds have a crystalline structure called a crystal lattice.

Lattice Energy

• Lattice energy is the energy needed to separate the ions in an ionic compound.

Ionic Compound Formation

• Ionic compound formation is exothermic; stable, low-energy systems.

Ionic Compound Characteristics

• High melting and boiling points
• Strongest bonds
• Non-conductors as solids
• Good conductors when melted or dissolved in water (aqueous solutions).

Anion Reaction Equation

• ELEMENT + e- -> ELEMENT-1 + electron affinity

Electron Affinity

• Electron affinity is energy released when an electron joins a gaseous atom.

Cation Reaction Equation

• ELEMENT + ionization energy -> ELEMENT+1 + e-

Ionization Energy

• Ionization energy is energy needed to remove an electron from an atom.

Atomic Radii and Bonding

• Smaller atomic radii lead to stronger bonds.
• Higher total oxidation charges lead to stronger bonds.
• Stronger bonds have more negative lattice energy

Oxidation Charge

• Oxidation charge is the charge of a monatomic ion (one-atom ion).

Ionic Solid Brittleness

• Ionic solids break due to the alignment of ions.

Pseudo Noble Gas Configuration

• Relatively stable electron arrangements, similar to noble gases.

Elements with Pseudo Noble Gas Configurations

• Copper, Silver, Zinc

Stable Bonding Configuration

• Atoms achieve a noble gas configuration when they bond.

Covalent Bonding

• Covalent bonds form through electron sharing.

Diatomic Elements

• H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂

Bonds and Shared Atoms

• Single bond: 2 shared atoms
• Double bond: 4 shared atoms
• Triple bond: 6 shared atoms

Sigma Bond Position

• Sigma bonds are located directly between two bonding atoms.

Sigma Bond Formation Explanation

• Electron orbitals overlap, centering the shared electrons.

Pi Bond Location

• Pi bonds are positioned above and below the bonding axis.

Pi Bond Pre-Conditions

• Occur when there's one sigma bond plus one or two pi bonds.

Molecular Geometry - 2,0

• Linear (180°) sp hybridised

Molecular Geometry - 2,1

• Bent (120°) sp² hybridised

Molecular Geometry - 3,0

• Trigonal planar (120°) sp² hybridised

Molecular Geometry - 3,1

• Trigonal pyramidal (109.5°) sp³ hybridised

Molecular Geometry - 4,0

• Tetrahedral (109.5°) sp³ hybridised

Molecular Geometry - 2,2

• Bent (109.5°) sp³ hybridised

Molecular Geometry - 5,0

• Trigonal bipyramidal

Molecular Geometry - 6,0

• Octahedral

Dissociation Energy

• Dissociation energy represents the energy needed to break a bond.

Lewis Structures

• Lewis structures illustrate molecular structure using symbols and dots to show bonds and lone pairs.

Resonance

• Resonance occurs when multiple Lewis structures represent a molecule, especially with multiple double and single bonds to the same atom.

Hybridization

• Hybridization is the mixing of atomic orbitals to form new, identical hybrid orbitals.

Electronegativity

• Electronegativity is the ability of an element to attract electrons in a chemical bond.

Electron Affinity (diff.)

• Electron affinity is the tendency of an atom to accept electrons.

Electronegativity vs. Electron Affinity

• Electronegativity's about attracting electrons; electron affinity's about accepting them.

Exceptions to Octet Rule

• BH₃, NO₂, ClO₂, PCl₅, XeF₃⁺

Expanded Octets

• Elements in Period 3 or lower can have expanded octets (more than 8 valence electrons).

Van der Waals Forces

• Includes dispersion, dipole-dipole, and hydrogen bonding intermolecular forces.

Dispersion Forces

• Weak intermolecular forces between nonpolar molecules (London Dispersion Forces).

Dipole-Dipole Forces

• Stronger intermolecular forces between polar molecules.

Dispersion Force Scope

• Exist between all molecules, but only present in nonpolar ones.

Hydrogen Bonding

• Hydrogen bonding involves hydrogen atoms in one dipole and fluorine, oxygen, or nitrogen from another.

Atomic Charged Regions

• Atomic charged regions referred to as poles or dipoles.

Bond Charge Indication

• Used to show opposing charges, delta negative (δ-) and delta positive (δ+).

Electronegativity Placement

• Lower electronegative element receives positive delta and higher the negative delta.

Bond Length and Strength Relationship

• Shorter bonds are generally stronger.

Bond Number and Length Relationship

• Increasing bond numbers decrease bond length.

Electronegativity Difference - 0

• Zero electronegativity difference indicates nonpolar covalent bonds.

Electronegativity Difference (0-1.7)

• Electronegativity difference between 0 and 1.7 indicates polar covalent bonds.

Electronegativity Difference - 1.7

• Electronegativity difference of 1.7 or greater suggests ionic bonds.

Single Bond Characteristics

• Longest length, weakest strength, least dissociation energy

Double Bond Characteristics

• Shorter length, stronger strength, greater dissociation energy

Triple Bond Characteristics

• Shortest length, strongest strength, greatest dissociation energy

Ionic vs. Covalent Physical Properties

• Ionic: High melting/boiling points, good conductors in aqueous solutions, hard/brittle.
• Covalent: Low melting/boiling points, easily molded (malleable/ductile), poor conductors.

Malleable/Ductile

• Malleable: Can be hammered into sheets.
• Ductile: Can be drawn into wires.

Property Differences Explanation

• Property differences result from different attractive forces.

Elements with Variable Charges/Roman Numerals

• 10 elements may need a charge/roman numeral indicating oxidation.

Non-Metal Elements Needing Roman Numerals

• 5 nonmetals require Roman numerals indicating oxidation.

Specific Non-Metal Elements Needing Roman Numerals

• Gallium, Indium, Thallium, Lead, Tin

Metal Elements with Specific Charges

• Zinc (+2); Silver (+1); Cadmium (+2)

Elements Requiring Prefix System

• Antimony and Bismuth always use the prefix system to show charge.