Ionic and Molecular Compounds

Simple Ions

Simple ions are atoms that have gained or lost electrons to achieve a stable electron configuration. They are classified as either cations or anions based on their charge.

• Cation: A positively charged ion formed when an atom loses electrons. Example: Na+.

• Anion: A negatively charged ion formed when an atom gains electrons. Example: Cl-.

• Formation of Non-metal Ions: Non-metals typically gain electrons to form anions.

Biologically Important Ions

• Examples: Ca2+ (calcium ion), CO32- (carbonate ion), PO43- (phosphate ion).

• Naming Ionic Compounds: The cation is named first, followed by the anion. For example, NaCl is sodium chloride.

• Ion Formation: Atoms form ions to achieve noble gas electron configurations.

Physical and Chemical Properties of Ionic Compounds

• Strong Attraction: Ionic compounds are held together by electrostatic forces between oppositely charged ions.

• Electrical Conductivity: Ionic compounds conduct electricity when molten or dissolved in water, but not in solid state.

• High Melting and Boiling Points: Due to strong ionic bonds.

Ionic Compounds with Simple Ions

Naming and Determining Formulas

• Naming: Name the cation first, then the anion. For example, CaBr2 is calcium bromide.

• Formula Determination: Balance the charges to ensure the compound is neutral. Example: Calcium ion (Ca2+) and bromide ion (Br-) combine to form CaBr2.

Ionic Compounds with Polyatomic Ions

Names and Formulas

• Polyatomic Ion: A charged group of covalently bonded atoms acting as a single ion. Example: NO3- (nitrate).

• Naming: Name the cation, then the polyatomic anion. Example: NaNO3 is sodium nitrate.

• Formula Writing: Use parentheses if more than one polyatomic ion is needed. Example: Ca(NO3)2.

Covalent Bonding and Molecular Compounds

Definition and Formation

Covalent bonds are formed when two non-metal atoms share electrons to achieve stability. Molecular compounds are composed of molecules held together by covalent bonds.

• Formation: Atoms share electrons to complete their valence shells.

• Example: H2O (water) is formed by sharing electrons between hydrogen and oxygen.

Naming and Formula Conversion

• Prefixes: Used to indicate the number of atoms. Example: mono-, di-, tri-, tetra-, etc.

• Example: CO2 is carbon dioxide; SO3 is sulfur trioxide.

Lewis Structures

Drawing Lewis Dot Structures

Lewis structures represent the valence electrons of atoms and show how they are shared or transferred in compounds.

• Rules: Place dots around the element symbol to represent valence electrons. Pair electrons to form bonds.

• Electron Counting: Sum the valence electrons for all atoms in the molecule or ion.

• Difference: Ionic compounds transfer electrons; molecular compounds share electrons.

Resonance Structures

Sharing of Electrons in Resonance

Resonance structures occur when more than one valid Lewis structure can be drawn for a molecule, indicating delocalized electrons.

• Rule: Only the position of electrons changes, not the arrangement of atoms.

• Example: NO3- (nitrate ion) has three resonance structures.

VSEPR Theory and Molecular Shapes

Predicting Electron Pair Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shape of molecules based on repulsion between electron pairs.

• Lone Pair Effect: Lone pairs occupy more space and can distort molecular shapes.

• Electron Geometry vs. Molecular Structure: Electron geometry considers all electron pairs; molecular structure considers only bonded atoms.

• Example: H2O has a bent shape due to two lone pairs on oxygen.

Molecular Geometry and Polarity

Predicting Polarity

Polarity depends on the distribution of electrons and the shape of the molecule.

• Polar Molecule: Has an uneven distribution of charge. Example: H2O.

• Non-Polar Molecule: Has an even distribution of charge. Example: CO2.

Intermolecular Forces and Properties of Compounds

Relationship Between Phases and Intermolecular Forces

Intermolecular forces determine the physical properties of compounds, such as melting and boiling points, and their phase (solid, liquid, gas).

• Temperature Changes: Stronger intermolecular forces require more energy (higher temperature) to change phase.

Types of Intermolecular Forces

• Dispersion Forces: Weak forces present in all molecules, caused by temporary dipoles.

• Dipole-Dipole Interactions: Occur between polar molecules due to permanent dipoles.

• Hydrogen Bonding: Strongest type, occurs when hydrogen is bonded to N, O, or F.

Summary Table: Types of Intermolecular Forces

Key Equations

• Electrostatic Force (Coulomb's Law):

• Lewis Structure Electron Counting:

• Polarity Determination:

Additional info: Academic context and examples have been added to clarify concepts and provide a self-contained study guide.