Molecules and Compounds: Chapter 5 Study Notes

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Chapter 5: Molecules and Compounds

Introduction

This chapter explores the nature of molecules and compounds, focusing on how their properties differ from those of the elements that compose them. It introduces key concepts such as chemical formulas, types of compounds, and the law of constant composition.

Properties of Compounds vs. Elements

Properties of Sugar (Sucrose) vs. Its Elements

Ordinary table sugar is a compound called sucrose. A sucrose molecule contains carbon (C), hydrogen (H), and oxygen (O) atoms. The properties of sucrose are very different from those of its constituent elements:

Carbon: In the form of graphite, carbon is a black, solid element.
Hydrogen: A colorless, flammable gas.
Oxygen: A colorless gas essential for respiration.
Sucrose: A white, crystalline solid with a sweet taste.

Example: Sucrose is used as table sugar, while its elements have very different physical and chemical properties.

Properties of Salt (Sodium Chloride) vs. Its Elements

Elemental Sodium (Na): A highly reactive metal that oxidizes rapidly in air.
Elemental Chlorine (Cl): A yellow, poisonous gas with a pungent odor.
Sodium Chloride (NaCl): A stable, edible compound known as table salt.

Key Point: The properties of a compound are generally very different from the properties of the elements that compose it.

Elements, Compounds, and Mixtures

Classification of Substances

Elements: Pure substances consisting of only one type of atom.
Compounds: Substances composed of two or more elements in fixed, definite proportions.
Mixtures: Combinations of elements and/or compounds in variable proportions.

Example: Water (H2O) is a compound, while air is a mixture.

Law of Constant Composition

Definition and Application

The law of constant composition states that all samples of a given compound have the same proportions of their constituent elements.

Formally stated by Joseph Proust (1754–1826).
For example, water always contains hydrogen and oxygen in a mass ratio of 2:16 (or 1:8).

Equation:

Example: Decomposing 18.0 g of water yields 16.0 g of oxygen and 2.0 g of hydrogen, giving a mass ratio of 8.0:1.

Chemical Formulas

Representation of Compounds

Chemical formulas indicate the elements present in a compound and the relative number of atoms of each.

Symbols represent elements; subscripts indicate the number of atoms.
A subscript of 1 is omitted by convention.

Examples:

Water: H2O (2 hydrogen, 1 oxygen)
Table salt: NaCl (1 sodium, 1 chlorine)
Carbon dioxide: CO2 (1 carbon, 2 oxygen)
Sucrose: C12H22O11 (12 carbon, 22 hydrogen, 11 oxygen)

Key Point: Changing a subscript in a chemical formula creates a different compound (e.g., CO vs. CO2).

Order of Elements in Chemical Formulas

Most metallic element is listed first.
Among nonmetals, the more metal-like element (to the left or lower in the periodic table) is listed first.

Example: NO2 (not ON2), SO2 (not O2S).

Polyatomic Ions in Chemical Formulas

Some chemical formulas contain groups of atoms that act as a unit, called polyatomic ions. Parentheses and subscripts indicate the number of such groups.

Example: Mg(NO3)2 contains 1 magnesium, 2 nitrate groups (each with 1 nitrogen and 3 oxygen).

Mg: 1
N: 2 × 1 = 2
O: 2 × 3 = 6

Types of Chemical Formulas

Empirical formula: Simplest whole-number ratio of atoms (e.g., HO for hydrogen peroxide).
Molecular formula: Actual number of atoms in a molecule (e.g., H2O2 for hydrogen peroxide).
Structural formula: Shows how atoms are connected using lines for bonds.
Molecular models: Ball-and-stick and space-filling models represent 3D structure.

Example: Methane (CH4): molecular formula shows 1 carbon and 4 hydrogen; structural formula shows each hydrogen bonded to carbon.

Molecular View of Elements and Compounds

Basic Units

Elements: May be atomic (single atoms) or molecular (diatomic molecules, e.g., Cl2).
Compounds: May be molecular (composed of molecules) or ionic (composed of ions in a lattice).

Example: Mercury is an atomic element; chlorine is a molecular element (Cl2).

Writing Formulas for Ionic Compounds

Rules for Ionic Compounds

Formed from metals (cations) and nonmetals (anions).
The sum of charges must be zero (charge-neutral).
Use subscripts to balance charges.

Examples:

Na+ and Cl- → NaCl
Mg2+ and Cl- → MgCl2
Al3+ and O2- → Al2O3

Steps:

Write symbols and charges.
Cross charges to become subscripts.
Reduce subscripts to smallest whole numbers.

Polyatomic Ions in Ionic Compounds

Recognize common polyatomic ions and use parentheses when more than one is present.

Example: Ca2+ and NO3- → Ca(NO3)2

Naming Ionic Compounds

Type I Ionic Compounds (Metal with Constant Charge)

Binary compounds: name of cation (metal) + base name of anion (nonmetal) + -ide
Main group metals usually have constant charge.

Examples:

NaCl: sodium chloride
Al2O3: aluminum oxide

Type II Ionic Compounds (Metal with Variable Charge)

Transition metals may have variable charges.
Specify charge with Roman numeral in parentheses.

Examples:

FeCl3: iron(III) chloride
CuO: copper(II) oxide

Naming Ionic Compounds with Polyatomic Ions

Use the name of the polyatomic ion in the compound name.

Examples:

KNO3: potassium nitrate
Fe(OH)2: iron(II) hydroxide
NH4NO3: ammonium nitrate

Naming Oxyanions

Oxyanions are anions containing oxygen.
For two ions in a series: -ate (more oxygen), -ite (less oxygen).
For more than two: hypo- (least), per- (most).

Examples:

NO3-: nitrate
NO2-: nitrite
ClO4-: perchlorate
ClO3-: chlorate
ClO2-: chlorite
ClO-: hypochlorite

Naming Molecular Compounds

Rules for Naming Binary Molecular Compounds

Composed of two nonmetals.
Prefix + name of first element + prefix + base name of second element + -ide
Prefixes indicate number of atoms: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-
Mono- is usually omitted for the first element.

Examples:

CO2: carbon dioxide
N2O: dinitrogen monoxide

Naming Acids

Types of Acids

Binary acids: Contain hydrogen and a nonmetal.
Oxyacids: Contain hydrogen, a nonmetal, and oxygen (as part of a polyatomic ion).

Naming Binary Acids

hydro + base name of nonmetal + -ic + acid

Examples:

HCl (aq): hydrochloric acid
HBr (aq): hydrobromic acid
H2S (aq): hydrosulfuric acid

Naming Oxyacids

If oxyanion ends in -ate: base name of oxyanion + -ic + acid
If oxyanion ends in -ite: base name of oxyanion + -ous + acid

Examples:

HNO3 (aq): nitric acid (from nitrate)
HNO2 (aq): nitrous acid (from nitrite)
H2SO4 (aq): sulfuric acid (from sulfate)
H2SO3 (aq): sulfurous acid (from sulfite)

Formula Mass

Definition and Calculation

The formula mass of a compound is the sum of the atomic masses of all the atoms in its chemical formula.

Equation:

Example: For MgCl2:

Mg: 1 × 24.31 amu = 24.31 amu
Cl: 2 × 35.45 amu = 70.90 amu
Total: 95.21 amu

Summary of Key Learning Objectives

Restate and apply the law of constant composition.
Write chemical formulas and determine the number of each type of atom.
Classify elements and compounds as atomic, molecular, or ionic.
Write and name formulas for ionic and molecular compounds, including those with polyatomic ions.
Name binary acids and oxyacids.
Calculate formula mass.

Type of Formula	Description	Example
Empirical Formula	Simplest whole-number ratio of atoms	HO (hydrogen peroxide)
Molecular Formula	Actual number of atoms in a molecule	H2O2 (hydrogen peroxide)
Structural Formula	Shows how atoms are connected	H–O–O–H

Common Polyatomic Ion	Formula	Name
Carbonate	CO32-	carbonate ion
Bicarbonate	HCO3-	hydrogen carbonate ion
Nitrate	NO3-	nitrate ion
Sulfate	SO42-	sulfate ion
Hydroxide	OH-	hydroxide ion

Additional info: These notes expand on the brief points and images in the original slides, providing definitions, examples, and formulas for clarity and completeness.