Compounds and the Mole

Compounds and the Mole

Chapter Outline

  • Section 3.1 Chemical Formulas

  • Section 3.2 Naming Binary Covalent Compounds

  • Section 3.3 Formulas for Ionic Compounds

  • Section 3.4 Naming Ionic Compounds

  • Section 3.5 Naming Acids

  • Section 3.6 Nomenclature Review

  • Section 3.7 The Mole

  • Section 3.8 Molar Mass

  • Section 3.9 Percent Composition

  • Section 3.10 Empirical Formulas

  • Section 3.11 Molecular Formulas

  • Section 3.12 Combustion Analysis

Section 3.1 Chemical Formulas

  • Chemical Nomenclature: Systematic approach to naming chemical compounds.

  • Chemical Formulas: Combinations of elemental symbols and subscripts used to represent compounds.

    • Example: Fe₂O₃ represents rust.

    • Example: O₂ represents elemental oxygen.

  • Formula Units: Chemical formulas that indicate the ratio of atoms in a compound.

    • Example: Formula unit of Ba₃(PO₄)₂ contains 3 Ba atoms, 2 P atoms, and 8 O atoms.

Molecular vs Ionic Compounds

  • Molecular Compounds: Exist as individual molecules made up of bonded atoms (nonmetals), connected via covalent bonds.

    • Example: A molecular compound's formula represents one molecule.

  • Ionic Compounds: Form a continuous 3-D lattice of atoms (metals and nonmetals) held together by ionic bonds. They do not exist as discrete molecules, and the formula unit reflects the ratio of atoms.

Example of Classifying Compounds

  • Classify the following:

    • a. CaCl₂ - ionic

    • b. CCl₄ - covalent (molecular)

    • c. H₂O - covalent (molecular)

    • d. Fe₂O₃ - ionic

Diatomic and Allotropic Elements

  • Some elements exist only as molecules composed of two or more atoms, e.g., Cl₂, P₄, S₈.

  • Allotropes: Different molecular forms of an element, e.g., oxygen (O₂, O₃) and carbon (diamond, graphite).

Section 3.2 Naming Binary Covalent Compounds

  • Prefixes for Naming: Indicate the number of atoms of each element present in covalent compounds.

    • Table of Prefixes (e.g., mono-, di-, tri-, tetra-, etc.)

  • To Name Compounds:

    1. Name the first element with a prefix (if more than one atom).

    2. Name the second element using a prefix and modify to end with -ide.

  • Example: N₂O - dinitrogen monoxide.

Examples of Naming Compounds

  • a. SiCl₄ - silicon tetrachloride

  • b. SO₃ - sulfur trioxide

Special Cases: Binary Compounds of Hydrogen

  • Non-acid binaries are named specifically; e.g., NH₃ is ammonia, PH₃ is phosphine.

Section 3.3 Formulas for Ionic Compounds

  • Ions: Charged atoms/groups that form when electrons are gained/lost.

    • Cations: Positive ions

    • Anions: Negative ions

  • Binary Ionic Compounds: Form when metals react with nonmetals.

    • Metal atoms transfer electrons to nonmetals; metals form cations while nonmetals form anions.

Naming Ionic Compounds

  • Names consist of cation name followed by anion name.

  • Charges should balance to maintain neutrality in the compound (total positive charge = total negative charge).

Examples of Determining Formulas

  • Example: Zn and N

    • Zn: 2+, N: 3-

    • Least common multiple: 6

    • Formula: Zn₃N₂.

    • Example calculations provided for specific compound pairs.

Polyatomic Ions

  • Groups of atoms which have gained/lost electrons and have a net charge (e.g., CO₃²⁻ for carbonate.)

    • Naming conventions and examples provided.

Section 3.4 Naming Ionic Compounds

  • Names of ionic compounds consist of the cation and anion.

    • Polyatomic Cations are commonly used, such as NH₄⁺ (ammonium).

    • Monatomic Metal Cations with constant charge use basic elemental naming.

    • Variable Charge Cations are indicated with Roman numerals (e.g., Fe²⁺ as iron(II)).

  • Anions such as phosphide (P³⁻) follow naming trends for elements/groups.

Section 3.5 Naming Acids

  • Acids release H⁺ ions in water and are categorized into binary acids and oxyacids.

    • Binary Acids: Name starting with hydro-, followed by root and -ic (e.g. HCl = hydrochloric acid)

    • Oxyacids: Named based on the accompanying oxyanion (e.g. ClO₃⁻ to HClO₃ becomes chloric acid).

Section 3.6 Nomenclature Review

  • Overview of various naming conventions and checks for identifying chemical types (ionic vs covalent).

Section 3.7 The Mole

  • Definition: The mole (mol) is the counting unit in chemistry, with Avogadro's number as $N_A = 6.022 imes 10^{23}$ units/mole.

  • Conversion applications explored, such as from moles to particles and vice versa.

Section 3.8 Molar Mass

  • Molar Mass: Defined as the mass of one mole of a substance and related to atomic mass in grams.

    • Calculation of molar masses for elements and basic compounds illustrated.

Example Calculations

  • Example of calculating molar mass for acetic acid: Total mass determined as 60.052 g/mol.

Section 3.9 Percent Composition

  • Percent Composition: Use formula and molar masses to determine the percentage by mass of each element in a compound.

    • Calculation methods and examples provided (e.g., for NaCl).

Section 3.10 Empirical Formulas

  • Empirical formulas express simplest ratio of atoms.

    • Case studies showing conversion of experimental mass data to empirical formulas (e.g., hydrogen peroxide to HO).

Section 3.11 Molecular Formulas

  • Molecular formulas tell the actual number of atoms in a molecule.

    • Difference between empirical and molecular formulas illustrated, including methods for conversion.

Example to Determine Molecular Formula

  • Presented scenario of styrene derived from empirical data and mass data calculation leading to C₈H₈.