Molecules and Compounds: Structure, Nomenclature, and Composition

Chapter 3: Molecules and Compounds

Introduction

This chapter explores the fundamental concepts of molecules and compounds, including their classification, chemical formulas, nomenclature, and methods for determining composition. Understanding these topics is essential for mastering the language and calculations of chemistry.

Compounds and Their Properties

Compounds Are Not Their Elements!

• Chemical compounds have unique chemical and physical properties distinct from the elements from which they are made.

• Example: Water (H2O) is composed of hydrogen and oxygen, but its properties are very different from those of either element.

Substance	State at 25°C	Color	Properties
Hydrogen	Gas	Colorless	Flammable
Oxygen	Gas	Colorless	Supports combustion
Water	Liquid	Colorless	Extinguishes fire

Kinds of Compounds

• Molecular Compounds: Composed of atoms held together by covalent bonds (shared electrons). Usually consist of nonmetals.

• Ionic Compounds: Composed of positively and negatively charged ions (cations and anions) held together by ionic bonds (electrostatic attraction). Usually a metal and a nonmetal; often form crystalline solids.

Chemical Formulas and Terms

Types of Chemical Formulas

• Molecular formula: Shows the exact number and type of atoms in a molecule (e.g., C6H12O6 for glucose).

• Empirical formula: Shows the simplest whole-number ratio of elements in a compound (e.g., CH2O for glucose).

• Formula unit: The smallest electrically neutral unit in an ionic compound (e.g., NaCl).

Representing Chemical Formulas

• Chemical formulas use elemental symbols to represent atoms.

• Subscripts indicate the number of each atom present (e.g., CH4 for methane).

Classification of Elements and Compounds

Atomic-Level Classification

• Elements can exist as atomic elements (single atoms, e.g., Ne) or molecular elements (diatomic or polyatomic molecules, e.g., O2, S8).

• Compounds are classified as molecular or ionic based on their bonding and composition.

Molecular Elements

• Diatomic elements: Exist as molecules of two atoms (H2, N2, O2, F2, Cl2, Br2, I2).

• Polyatomic elements: Exist as molecules with more than two atoms (e.g., P4, S8).

Naming Compounds

Naming Binary Ionic Compounds

• The cation (positive ion) is named first; the anion (negative ion) is named second with an -ide ending.

• If the metal can form more than one charge, indicate the charge with Roman numerals in parentheses.

• Example: FeCl2 is iron(II) chloride.

Formulas of Binary Ionic Compounds

• Write the ions with their charges.

• Balance the charges so the total is zero; use subscripts as needed.

• Example: Titanium(IV) oxide: Ti4+ and O2− combine to form TiO2.

Polyatomic Ions

• Common polyatomic ions include ammonium (NH4+), acetate (C2H3O2−), carbonate (CO32−), hydroxide (OH−), nitrate (NO3−), phosphate (PO43−), and sulfate (SO42−).

• Names and formulas must be memorized for common polyatomic ions.

Oxyanion Name Conventions

• Oxyanions are anions containing oxygen and another element.

• "-ate" and "-ite" endings indicate more or fewer oxygens, respectively (e.g., nitrate NO3−, nitrite NO2−).

• Prefixes "per-" (most oxygen) and "hypo-" (least oxygen) are used for series with more than two oxyanions.

Hydrated Ionic Compounds

• Hydrates are ionic compounds containing a specific number of water molecules per formula unit.

• Water is indicated by a dot and a prefix (e.g., CuSO4·5H2O is copper(II) sulfate pentahydrate).

Naming Binary Molecular Compounds

• Composed of two nonmetals.

• The first element is named normally; the second gets an "-ide" ending.

• Prefixes (mono-, di-, tri-, etc.) indicate the number of atoms. "Mono-" is omitted for the first element.

• Example: CO is carbon monoxide, CO2 is carbon dioxide.

Naming Acids

• Binary acids (hydrogen + nonmetal): Use "hydro-" prefix, root of nonmetal, and "-ic acid" (e.g., HCl is hydrochloric acid).

• Oxyacids (hydrogen + oxyanion): If the anion ends in "-ate," the acid ends in "-ic acid"; if "-ite," use "-ous acid" (e.g., HNO3 is nitric acid, HNO2 is nitrous acid).

Summary of Inorganic Nomenclature

• Systematic rules exist for naming ionic, molecular, and acid compounds based on their composition and structure.

• Refer to flowcharts or tables for stepwise guidance.

Molar Mass and Composition Calculations

Terminology

• Molecular mass: Mass of one molecule of a molecular compound (in amu).

• Formula mass: Mass of one formula unit of an ionic compound (in amu).

• Molar mass: Mass of one mole of particles (in grams per mole, g/mol).

Calculation of Molar Mass

• Sum the atomic masses of all atoms in the formula.

• Example: For CO2:

  • C: 12.01 g/mol

  • O: 2 × 16.00 g/mol = 32.00 g/mol

  • Total: 44.01 g/mol

Percent Composition

• Percent composition = (mass of element in compound / molar mass of compound) × 100%

• Example: %C in CO2 = (12.01 / 44.01) × 100% = 27.3%

Conversion Factors from Chemical Formulas

• Subscripts in chemical formulas can be used as conversion factors between moles of elements and moles of molecules or formula units.

• Example: 1 mol CO2 contains 2 mol O atoms.

Empirical and Molecular Formulas

Determining Empirical Formulas from Mass Percent

1. Assume 100 g of sample; convert % to grams.

2. Convert grams to moles for each element.

3. Divide by the smallest number of moles to get the simplest ratio.

4. If necessary, multiply to get whole numbers.

Example: A compound with 52.14% C, 13.13% H, and 34.73% O yields empirical formula C2H6O.

Determining the Molecular Formula

• Divide the molar mass (from experiment) by the empirical formula mass to get a whole number (n).

• Molecular formula = (empirical formula) × n

• Example: If empirical formula is CH2O (mass = 30.03 g/mol) and molar mass is 180.18 g/mol, n = 6, so molecular formula is C6H12O6.

Practice Problems and Examples

• Practice naming and writing formulas for various compounds, including binary ionic, binary molecular, acids, and hydrates.

• Calculate molar mass, percent composition, and convert between mass, moles, and number of particles using Avogadro's number ($6.022 \times 10^{23}$).

• Determine empirical and molecular formulas from composition data.

Key Equations

• Molar mass: $\text{Molar mass} = \sum (\text{atomic mass} \times \text{number of atoms})$

• Percent composition: $\%\text{Element} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100\%$

• Number of particles: $\text{Number of particles} = \text{moles} \times 6.022 \times 10^{23}$

Additional info: For a comprehensive list of polyatomic ions and nomenclature rules, refer to standard tables in your textbook or course materials.