Recording-2025-02-27 Chem

Naming Ionic Compounds

• 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

Example of Naming Ionic Compounds

• 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.

Covalent Bonding

• 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).

Types of Molecules

• Diatomic molecules: Consisting of two atoms.

  • E.g., H2 (homonuclear), HCl (heteronuclear).

• Polyatomic molecules: Consist more than two atoms.

  • E.g., H2O (water).

Laws of Chemical Proportions

• 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 and Molecular Formula

• 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.

Determining the Empirical Formula

1. Given percentage composition (e.g., 30.45% Nitrogen, 69.55% Oxygen).

2. Assume 100g sample for simplicity (30.45g N, 69.55g O).

3. 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.

4. Divide by smallest number to find subscripts:

   • N = 2.173 / 2.173 = 1; O = 4.347 / 2.173 = 2 → Empirical formula: NO2.

Calculating Molecular Formula from Empirical Formula

1. Calculate empirical formula mass: NO2 → 14 + (2 * 16) = 46 g/mol.

2. Given molar mass of compound (92 g/mol).

3. Divide molar mass by empirical formula mass: 92 / 46 = 2.

4. Multiply empirical formula by this ratio: 2 * (NO2) → N2O4.

Summary and Common Names of Compounds

• 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.