Chemical Formulas and Molecular Structures

Reading, Interpreting, and Writing Chemical Formulas

Chemical formulas represent the types and numbers of atoms in a compound. Understanding how to read and write these formulas is foundational in chemistry.

• Empirical Formula: Shows the simplest whole-number ratio of atoms in a compound.

• Molecular Formula: Indicates the actual number of each type of atom in a molecule.

• Structural Formula: Depicts the arrangement of atoms within the molecule.

• Example: The molecular formula for water is H2O, indicating two hydrogen atoms and one oxygen atom.

Determining Number of Atoms in a Formula

• Count the subscript for each element to determine the number of atoms present.

• Example: In C6H12O6, there are 6 carbon, 12 hydrogen, and 6 oxygen atoms.

Isomers and Types of Compounds

Isomers are compounds with the same molecular formula but different structures or arrangements of atoms.

• Structural Isomers: Differ in the connectivity of their atoms.

• Stereoisomers: Atoms are connected in the same order but differ in spatial arrangement.

• Types of Compounds: Ionic, covalent (molecular), and metallic compounds.

Ionic and Covalent Compounds

Identifying Ionic vs. Covalent Compounds

Ionic compounds are formed from metals and nonmetals, while covalent compounds are formed between nonmetals.

• Ionic Compounds: Consist of positive (cation) and negative (anion) ions held together by electrostatic forces.

• Covalent Compounds: Atoms share electrons to achieve stable electron configurations.

• Example: NaCl (sodium chloride) is ionic; H2O (water) is covalent.

Bonding Differences Based on Electronic Structure

• Ionic bonding involves transfer of electrons.

• Covalent bonding involves sharing of electrons.

• Bonding type affects properties such as melting point, solubility, and electrical conductivity.

General Properties of Ionic and Covalent Compounds

• Ionic: High melting/boiling points, conduct electricity when molten or dissolved, often soluble in water.

• Covalent: Lower melting/boiling points, poor electrical conductivity, variable solubility.

Nomenclature

Naming Ionic and Covalent Compounds

• Ionic Nomenclature: Name the cation first, then the anion. For transition metals, use Roman numerals to indicate charge.

• Covalent Nomenclature: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom.

• Example: CO2 is carbon dioxide; FeCl3 is iron(III) chloride.

Polyatomic Ions and Roman Numerals

• Polyatomic ions are charged groups of covalently bonded atoms (e.g., SO42−, NO3−).

• Roman numerals indicate the oxidation state of transition metals (e.g., Cu+ is copper(I), Cu2+ is copper(II)).

Chemical Bonding

Covalent and Ionic Bonding

Chemical bonds form when atoms share or transfer electrons to achieve stable electron configurations.

• Ionic Bonding: Transfer of electrons from metal to nonmetal.

• Covalent Bonding: Sharing of electrons between nonmetals.

• Bonding Interactions: In ionic bonds, electrons are transferred; in covalent bonds, electrons are shared.

Electronegativity and Bond Polarity

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

• Difference in electronegativity determines bond type:

  • Large difference: Ionic bond

  • Small difference: Polar covalent bond

  • No difference: Nonpolar covalent bond

• Bond Polarity: Polar bonds have unequal sharing of electrons; nonpolar bonds have equal sharing.

• Notation: Bond polarity can be shown with an arrow (→) pointing toward the more electronegative atom or with delta notation (δ+, δ−).

Lewis Structures and Molecular Geometry

Drawing Lewis Structures

Lewis structures represent the arrangement of valence electrons in molecules and ions.

• Follow the octet rule (8 electrons around most atoms) or duet rule (2 electrons for hydrogen).

• Count total valence electrons, arrange atoms, and distribute electrons to satisfy octet/duet rules.

• Some elements can have expanded octets (e.g., P, S, Cl).

Valid and Invalid Lewis Structures

• Valid structures obey the octet/duet rule and have the correct number of valence electrons.

• Some elements (e.g., B, Be) may have incomplete octets.

VSEPR Theory and Molecular Shape

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Electron pairs (bonded and lone pairs) arrange themselves to minimize repulsion.

• Common Shapes:

  • Linear: 2 electron groups

  • Trigonal planar: 3 electron groups

  • Tetrahedral: 4 electron groups

  • Trigonal bipyramidal: 5 electron groups

  • Octahedral: 6 electron groups

• Lone pairs affect bond angles and molecular shape.

Assigning Molecular Polarity

• Polarity depends on both bond polarity and molecular shape.

• Symmetrical molecules (e.g., CO2) are often nonpolar even if bonds are polar.

• Asymmetrical molecules (e.g., H2O) are polar.

Summary Table: Bond Types and Properties

Key Equations and Concepts

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons.

• Electronegativity Difference:

• Lewis Structure Electron Count:

Additional info: Some content was inferred and expanded for clarity and completeness, such as the summary table and detailed explanations of VSEPR theory and bond polarity.