Chemical Formulas

Chemical formulas are essential in understanding the composition and structure of compounds. They reveal how atoms combine, their ratios, and how to calculate important properties like molar mass, which is crucial for mastering concepts in Honors Chemistry.

1. Empirical formula

   • Represents the simplest whole-number ratio of atoms in a compound.
   • Derived from experimental data, often through combustion analysis.
   • Does not provide information about the actual number of atoms in a molecule.

2. Molecular formula

   • Shows the actual number of each type of atom in a molecule.
   • Can be a multiple of the empirical formula.
   • Useful for determining the molecular weight of a compound.

3. Structural formula

   • Depicts the arrangement of atoms within a molecule.
   • Indicates how atoms are bonded and the geometry of the molecule.
   • Can be represented in various ways, including Lewis structures and condensed formulas.

4. Ionic compound formulas

   • Composed of cations and anions, resulting in a neutral compound.
   • The formula reflects the ratio of ions needed to balance charges.
   • Typically written with the cation first followed by the anion.

5. Covalent compound formulas

   • Consist of nonmetals sharing electrons to form molecules.
   • The formula indicates the number of each type of atom present.
   • Prefixes (mono-, di-, tri-, etc.) are often used to denote the number of atoms.

6. Hydrate formulas

   • Compounds that include water molecules in their structure.
   • Written with a dot separating the compound and the water (e.g., CuSO4·5H2O).
   • The number of water molecules is indicated by a coefficient.

7. Percent composition

   • Represents the percentage by mass of each element in a compound.
   • Calculated using the formula: (mass of element in 1 mole of compound / molar mass of compound) × 100%.
   • Useful for determining the purity of a substance.

8. Molar mass calculation

   • The mass of one mole of a substance, expressed in grams per mole (g/mol).
   • Calculated by summing the atomic masses of all atoms in the molecular formula.
   • Essential for converting between grams and moles in stoichiometry.

9. Balancing chemical equations

   • Ensures that the number of atoms for each element is the same on both sides of the equation.
   • Follows the law of conservation of mass.
   • Involves adjusting coefficients, not subscripts, to balance the equation.

10. Oxidation numbers

    • Indicate the degree of oxidation of an atom in a compound.
    • Help in identifying redox reactions and determining electron transfer.
    • Follow specific rules, such as elements in their elemental form have an oxidation number of zero.

11. Polyatomic ions

    • Charged entities composed of two or more atoms bonded together.
    • Commonly found in ionic compounds and can have positive or negative charges.
    • Often memorized as they frequently appear in chemical formulas.

12. Nomenclature rules

    • Systematic naming of chemical compounds based on their composition and structure.
    • Includes rules for naming ionic, covalent, and acid compounds.
    • Essential for clear communication in chemistry.

13. Formula units

    • The simplest ratio of ions in an ionic compound.
    • Represents the smallest repeating unit in a crystal lattice.
    • Used to describe ionic compounds, as they do not exist as discrete molecules.

14. Stoichiometry calculations

    • Involves using balanced chemical equations to calculate quantities of reactants and products.
    • Relies on mole ratios derived from the coefficients in a balanced equation.
    • Essential for predicting yields and understanding reaction efficiency.

15. Lewis structures

    • Diagrams that represent the bonding between atoms and the lone pairs of electrons.
    • Help visualize molecular geometry and predict reactivity.
    • Useful for understanding resonance structures and formal charges.