Chapter 3: Molecules and Compounds – Study Notes

Chapter 3: Molecules and Compounds

Introduction

This chapter explores the fundamental concepts of molecules and compounds, including their formation, types of chemical bonds, formulas, and methods for representing and naming compounds. Understanding these concepts is essential for studying chemical reactions and the properties of substances.

Rocket Fuel: Chemical Reactions and Properties

Combustion of Methane

• Combustion Reaction: Methane reacts with oxygen to produce carbon dioxide and water.

• Balanced Equation:

$\mathrm{CH_4 + 2\ O_2 \rightarrow CO_2 + 2\ H_2O}$

• Reactants: 1 molecule of methane (CH4), 2 molecules of oxygen (O2).

• Products: 1 molecule of carbon dioxide (CO2), 2 molecules of water (H2O).

Properties of Substances Involved

Property	Hydrogen	Oxygen	Water
Boiling Point	-253 °C	-183 °C	100 °C
State at Room Temperature	Gas	Gas	Liquid
Flammability	Explosive	Necessary for combustion	Used to extinguish flame

Mixtures vs. Compounds

• Mixture: Hydrogen and oxygen gases can mix in any ratio; no fixed composition.

• Compound: Water (H2O) has a fixed ratio of 2 hydrogen atoms to 1 oxygen atom.

• Key Point: Compounds have definite, fixed ratios of elements; mixtures do not.

Chemical Bonds

Types of Chemical Bonds

• Ionic Bond: Attraction between a cation (positively charged ion) and an anion (negatively charged ion).

• Covalent Bond: Sharing of a pair of electrons between two atoms, with one electron contributed by each atom.

• Bonding Continuum: Ionic and covalent bonds represent extremes; many bonds have characteristics of both.

Ionic Bonding in Solids

• In solids, ions of opposite charge arrange in a three-dimensional lattice structure.

• This arrangement maximizes attractive forces and minimizes repulsive forces, resulting in a stable crystalline solid.

• Example: Sodium chloride (NaCl) forms a cubic lattice of Na+ and Cl- ions.

Covalent Bonds and Electron Sharing

• Atoms share electrons to form a bonding pair, which is energetically stable.

• Electron density is concentrated between the nuclei, holding the atoms together.

• Potential Energy: The lowest potential energy corresponds to the most stable bond length.

Molecular Formulae

Types of Chemical Formulas

• Empirical Formula: Simplest whole-number ratio of atoms in a compound (e.g., HO for hydrogen peroxide).

• Molecular Formula: Actual number of atoms of each element in a molecule (e.g., H2O2 for hydrogen peroxide).

• Structural Formula: Shows the arrangement of atoms and the bonds between them.

Examples of Formulas

• Hexene: C6H12 (empirical formula: CH2)

• Pentene: C5H10 (empirical formula: CH2)

• Butene: C4H8 (empirical formula: CH2)

• Methane: CH4 (empirical and molecular formulas are the same)

• Glucose: C6H12O6 (over 1,500 possible compounds with this formula)

Structural Formula with Wedge Bonds

• D-Glucose: The structural formula uses wedge and dash notation to indicate 3D arrangement of atoms.

• Wedges represent bonds coming out of the plane; dashes represent bonds going behind the plane.

Molecular Models of Structure

• Different models are used for different purposes:

  • Molecular Formula: Shows the types and numbers of atoms.

  • Structural Formula: Shows connectivity of atoms.

  • Ball-and-Stick Model: Shows 3D arrangement and bond angles.

  • Space-Filling Model: Shows relative sizes of atoms and how they fill space.

• Application: 3D models are important in drug discovery and molecular visualization.

Comparison Table: Different Views for Different Needs

Name of Compound	Empirical Formula	Molecular Formula	Structural Formula	Ball-and-Stick Model	Space-Filling Model
Benzene	CH	C6H6	Cyclic ring with alternating double bonds	Ball-and-stick representation	Space-filling representation
Acetylene	CH	C2H2	H–C≡C–H	Ball-and-stick representation	Space-filling representation
Glucose	CH2O	C6H12O6	Ring structure with hydroxyl groups	Ball-and-stick representation	Space-filling representation
Ammonia	NH3	NH3	Trigonal pyramidal structure	Ball-and-stick representation	Space-filling representation

Compounds

Molecular Compounds

• Composed of two or more nonmetals bonded covalently.

• Basic unit is the molecule.

• Examples: Water (H2O), carbon dioxide (CO2), propane (C3H8).

Ionic Compounds

• Composed of cations (usually metals) and anions (usually nonmetals) held together by ionic bonds.

• Basic unit is the formula unit, the smallest electrically neutral group of ions.

• Example: Sodium chloride (NaCl) consists of Na+ and Cl- in a 1:1 ratio.

Naming Ionic Compounds

General Rules

• Cations: Positively charged ions, usually metals.

• Anions: Negatively charged ions, usually nonmetals.

• Metals form cations; nonmetals (halides, oxygen, sulfur) form anions.

• Polyatomic ions contain more than one type of atom (e.g., NH4+, SO42-).

• Oxyanions contain oxygen (e.g., NO3-, SO42-).

Writing Formulas

• Cation is always listed before the anion.

• Polyatomic ions are written as a unit; use parentheses if more than one is present (e.g., Ba(OH)2).

• A compound is ionic if it contains a metal from Group I or II, or a polyatomic ion.

Naming Cations

• Monatomic cations have the same name as the element (e.g., sodium ion, Na+).

• Cations with multiple charges use Roman numerals (e.g., iron(II), Fe2+; iron(III), Fe3+).

• Older names use -ous (lower charge) and -ic (higher charge): ferrous (Fe2+), ferric (Fe3+).

• Cations of nonmetals end in -ium (e.g., ammonium, NH4+).

Naming Anions

• Monatomic anions end in -ide (e.g., chloride, Cl-).

• Some polyatomic anions also end in -ide (e.g., hydroxide, OH-).

• Oxyanions end in -ate (more oxygen) or -ite (less oxygen): nitrate (NO3-), nitrite (NO2-).

Common Anions Table

Formula	Name
Cl-	Chloride ion
O2-	Oxide ion
OH-	Hydroxide ion
NO3-	Nitrate ion
SO42-	Sulfate ion
PO43-	Phosphate ion
CO32-	Carbonate ion

Naming Acids

• If the anion ends in -ide, the acid is named with the prefix hydro- and the suffix -ic acid (e.g., HCl: hydrochloric acid).

• If the anion ends in -ate, the acid is named with the suffix -ic acid (e.g., HNO3: nitric acid).

• If the anion ends in -ite, the acid is named with the suffix -ous acid (e.g., HNO2: nitrous acid).

Naming Molecular (Binary) Compounds

• Formed between two nonmetals.

• The element farther left (or lower in the same group) is named first.

• The second element is named with an -ide ending.

• Prefixes indicate the number of atoms (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-).

• "Mono-" is never used for the first element.

• When a prefix ends in "a" or "o" and the element name begins with a vowel, the final vowel is dropped.

Calculating Empirical Formulae

• Measure the mass of each element in the compound.

• Express as a percentage of the total mass.

• Assume a 100 g sample for ease of calculation.

• Calculate moles of each element.

• Divide by the smallest number of moles to get ratios.

• Multiply to obtain whole numbers if necessary.

Calculating Formula Mass and Molar Mass

Formula Mass

• The sum of the atomic masses of all atoms in a molecule or formula unit.

• Example: For H2O:

$\text{Formula mass} = (2 \times 1.01\ \text{amu}) + (1 \times 16.00\ \text{amu}) = 18.02\ \text{amu}$

Molar Mass

• The mass in grams of one mole of a substance; numerically equal to the formula mass in amu, but with units of g/mol.

• Example: For H2O:

$\text{Molar mass} = 2 \times 1.01\ \text{g/mol} + 16.00\ \text{g/mol} = 18.02\ \text{g/mol}$

Using Molar Mass and Avogadro's Number

• Molar mass allows conversion between mass and moles.

• Avogadro's number ($6.022 \times 10^{23}$) allows conversion between moles and number of molecules.

• Conversion Example:

$\text{grams CO}_2 \rightarrow \text{moles CO}_2 \rightarrow \text{molecules CO}_2$

$\text{grams} \times \frac{1\ \text{mol}}{\text{molar mass}} \times \frac{6.022 \times 10^{23}\ \text{molecules}}{1\ \text{mol}}$

Additional info: Some tables and images were interpreted and expanded for clarity. All key chemical concepts and examples have been included for comprehensive exam preparation.