Ionic compounds typically consist of a metal and a nonmetal.
Example: NaCl (Sodium Chloride)
Sodium (Na) is the metal; Chlorine (Cl) is the nonmetal.
When naming ionic compounds:
If the metal cation has only one possible charge, name the metal first, then add ‘-ide’ to the nonmetal name.
E.g., NaCl becomes Sodium Chloride.
If the metal can have more than one charge (like Iron, Fe), specify the charge in Roman numerals:
Iron (Fe) can be Fe2+ or Fe3+.
Naming conventions:
FeCl2 = Iron(II) Chloride
FeCl3 = Iron(III) Chloride
Sodium Bromide:
Sodium (Na) has one possible charge (+1).
Resulting Name: Sodium Bromide.
Calcium Oxide:
Calcium (Ca) has a charge of +2 (only one possible charge).
Resulting Name: Calcium Oxide (CaO).
Aluminum Oxide (Al2O3):
Aluminum is +3, while Oxygen is -2.
Empirical formula acknowledges charge balance.
In covalent bonds, electrons are shared between atoms rather than transferred.
Example: Hydrogen (H) molecules (H2).
Differentiation between Ionic and Covalent Bonds:
Ionic bonds involve the transfer of electrons (e.g., NaCl).
Covalent bonds involve sharing of electrons (e.g., H2).
Diatomic molecules: Consisting of two atoms.
E.g., H2 (homonuclear), HCl (heteronuclear).
Polyatomic molecules: Consist more than two atoms.
E.g., H2O (water).
Law of Definite Proportions: Different samples of a compound will always contain the same elements in the same ratio.
Law of Multiple Proportions: When two elements form multiple compounds, the mass ratios of one element that combine with a fixed mass of the other can be expressed in small whole numbers.
Empirical Formula: Simplest whole number ratio of elements in a compound.
Molecular Formula: Represents the actual number of atoms of each element in a molecule.
Given percentage composition (e.g., 30.45% Nitrogen, 69.55% Oxygen).
Assume 100g sample for simplicity (30.45g N, 69.55g O).
Convert grams to moles:
Moles of N = mass (30.45g) / molar mass (14g/mol) → 2.173 moles.
Moles of O = mass (69.55g) / molar mass (16g/mol) → 4.347 moles.
Divide by smallest number to find subscripts:
N = 2.173 / 2.173 = 1; O = 4.347 / 2.173 = 2 → Empirical formula: NO2.
Calculate empirical formula mass: NO2 → 14 + (2 * 16) = 46 g/mol.
Given molar mass of compound (92 g/mol).
Divide molar mass by empirical formula mass: 92 / 46 = 2.
Multiply empirical formula by this ratio: 2 * (NO2) → N2O4.
Some compounds have specific common names that do not conform to IUPAC naming rules:
H2O = Water, NH3 = Ammonia, CO2 = Carbon Dioxide.
Hydrates: Compounds that include water in their structure (e.g., CuSO4•5H2O).
Anhydrous: Compounds without water in their structure; often result from heating a hydrate.