Chapter 5: Molecules and Compounds – Study Notes for Introductory Chemistry

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Chapter 5: Molecules and Compounds

Introduction to Molecules and Compounds

This chapter explores the nature of molecules and compounds, focusing on how their properties differ from those of their constituent elements. Understanding the formation, composition, and nomenclature of compounds is fundamental in introductory chemistry.

Properties of Compounds vs. Elements

Properties of Sugar (Sucrose, C12H22O11)

Sucrose is an ordinary table sugar, a compound made of carbon, hydrogen, and oxygen atoms.
The properties of sucrose are very different from those of its component elements:
- Carbon (as graphite): black, solid, conducts electricity.
- Hydrogen: colorless, flammable gas.
- Oxygen: colorless, supports combustion.
When combined in fixed ratios, these elements form sucrose, which is sweet, soluble, and non-conductive.
Example: Sucrose (C12H22O11) is used as a sweetener in food.

Properties of Salt (NaCl) vs. Sodium and Chlorine

Elemental sodium: highly reactive metal, oxidizes rapidly in air.
Elemental chlorine: yellow, poisonous gas with a pungent odor.
Sodium chloride (NaCl): safe, edible table salt, very different from its elements.
Key Point: The properties of a compound are generally very different from those of the elements that compose it.
Example: NaCl is used in food preservation and seasoning.

Elements, Compounds, and Mixtures

Classification of Substances

Most substances encountered daily are compounds, not elements.
Atoms are rare in nature; most matter exists as compounds or mixtures.
Compound: Elements combine in fixed, definite proportions.
Mixture: Elements or compounds combine in variable proportions.
Example: Water (H2O) is a compound; air is a mixture.

Law of Constant Composition

Definition and Examples

Law of Constant Composition: All samples of a given compound have the same proportions of constituent elements.
Formally stated by Joseph Proust (1754–1826).
Example (Water): Decomposing 18.0 g of water yields 16.0 g oxygen and 2.0 g hydrogen.
Mass ratio:
Example (Ammonia): Decomposing 17.0 g of ammonia yields 14.0 g nitrogen and 3.0 g hydrogen.
Mass ratio:

Chemical Formulas

Representing Compounds

Chemical formula: Indicates the elements present and the relative number of atoms of each.
Example: H2O (water) has 2 hydrogen atoms and 1 oxygen atom.
Subscripts indicate the number of atoms; a subscript of 1 is omitted.
Examples:
- NaCl: sodium and chlorine in a 1:1 ratio.
- CO2: carbon and oxygen in a 1:2 ratio.
- C12H22O11: sucrose, 12:22:11 ratio.
Changing a subscript changes the compound entirely (e.g., CO vs. CO2).

Order of Elements in Formulas

Most metallic element is listed first (e.g., NaCl, not ClNa).
Among nonmetals, the more metal-like element (to the left/bottom in the periodic table) is listed first.
Example: NO2 not O2N; SO2 not O2S.

Polyatomic Ions in Formulas

Groups of atoms that act as a unit and carry a charge are called polyatomic ions.
When multiple groups are present, use parentheses and a subscript (e.g., Mg(NO3)2).
To determine the number of each atom, multiply the subscript outside the parentheses by the subscript inside.
Example: Mg(NO3)2 contains 1 Mg, 2 N, and 6 O atoms.

Types of Chemical Formulas and Models

Empirical, Molecular, and Structural Formulas

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

Molecular View of Elements and Compounds

Atomic and Molecular Elements

Atomic elements: Exist in nature as single atoms (e.g., mercury).
Molecular elements: Exist as molecules, usually diatomic (e.g., Cl2, O2).
Table: Elements That Occur as Diatomic Molecules

Element	Formula
Hydrogen	H2
Nitrogen	N2
Oxygen	O2
Fluorine	F2
Chlorine	Cl2
Bromine	Br2
Iodine	I2

Molecular and Ionic Compounds

Molecular compounds: Composed of molecules formed from two or more nonmetals (e.g., CO2).
Ionic compounds: Composed of cations (usually metals) and anions (usually nonmetals), arranged in a 3D array (e.g., NaCl).
Table: Classification of Substances

Type	Basic Unit	Example
Atomic Element	Atom	Hg
Molecular Element	Molecule	O2
Molecular Compound	Molecule	CO2
Ionic Compound	Formula Unit	NaCl

Writing Formulas for Ionic Compounds

Rules and Examples

Formed when a metal bonds to a nonmetal.
Metals form positive ions (cations), nonmetals form negative ions (anions).
The sum of the charges must be zero (charge-neutral compound).
Example: Na+ and Cl- combine to form NaCl.
Use the crossover technique: the magnitude of each ion's charge becomes the subscript for the other ion.
Reduce subscripts to the smallest whole-number ratio.
Example: Mg2+ and Cl- combine to form MgCl2.

Polyatomic Ions in Ionic Compounds

Recognize common polyatomic ions (e.g., NO3-, SO42-, OH-).
When writing formulas, use parentheses if more than one polyatomic ion is present (e.g., Ca(NO3)2).
Do not reduce subscripts inside parentheses.
Table: Common Polyatomic Ions

Name	Formula
Nitrate	NO3-
Sulfate	SO42-
Hydroxide	OH-
Carbonate	CO32-
Phosphate	PO43-
Ammonium	NH4+

Naming Ionic Compounds

Type I and Type II Ionic Compounds

Type I: Metal forms only one type of cation (main group metals).
Type II: Metal forms more than one type of cation (transition metals, Sn, Pb).
Type I Naming: Name of cation (metal) + base name of anion (nonmetal) + -ide.
Example: NaCl = sodium chloride; Al2O3 = aluminum oxide.
Type II Naming: Name of cation (metal) + (Roman numeral for charge) + base name of anion + -ide.
Example: FeCl3 = iron(III) chloride; PbO2 = lead(IV) oxide.

Naming Ionic Compounds with Polyatomic Ions

Use the name of the polyatomic ion in place of the anion or cation name.
Example: KNO3 = potassium nitrate; Fe(OH)2 = iron(II) hydroxide; NH4NO3 = ammonium nitrate.

Naming Oxyanions

Oxyanions are polyatomic ions containing oxygen.
If two ions in a series, the one with more oxygen atoms ends in -ate, the one with fewer ends in -ite.
Examples: NO3- = nitrate, NO2- = nitrite; SO42- = sulfate, SO32- = sulfite.
If more than two ions, use prefixes hypo- (less than) and per- (more than).
Example: ClO- = hypochlorite, ClO4- = perchlorate.

Naming Molecular Compounds

Rules and Prefixes

Molecular compounds are usually formed from two or more nonmetals.
Use prefixes to indicate the number of atoms:
- mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9), deca- (10)
First element: prefix + element name (omit mono- if only one atom).
Second element: prefix + base name + -ide.
Examples: CO2 = carbon dioxide; N2O = dinitrogen monoxide.

Naming Acids

Binary Acids

Contain hydrogen and a nonmetal.
Name: hydro + base name of nonmetal + -ic + acid.
Examples: HCl(aq) = hydrochloric acid; HBr(aq) = hydrobromic acid.

Oxyacids

Contain hydrogen, a nonmetal, and oxygen (as part of a polyatomic ion).
If the oxyanion ends in -ate: base name of oxyanion + -ic + acid.
If the oxyanion ends in -ite: base name of oxyanion + -ous + acid.
Examples: HNO3(aq) = nitric acid; HNO2(aq) = nitrous acid; H2SO4(aq) = sulfuric acid; H2SO3(aq) = sulfurous acid.

Table: Names of Some Common Oxyacids and Their Oxyanions

Acid Name	Oxyanion Name	Formula
Nitric acid	Nitrate	NO3-
Nitrous acid	Nitrite	NO2-
Sulfuric acid	Sulfate	SO42-
Sulfurous acid	Sulfite	SO32-
Chloric acid	Chlorate	ClO3-
Perchloric acid	Perchlorate	ClO4-

Polyatomic Ions in Everyday Chemistry

Sodium hypochlorite (NaClO): Active ingredient in household bleach.
Sodium bicarbonate (NaHCO3): Used in baking soda.
Calcium carbonate (CaCO3): Active ingredient in antacids.
Sodium nitrite (NaNO2): Used to preserve packaged meats.

Summary and Review

Compounds display constant composition; their elements combine in fixed, definite proportions.
Chemical formulas indicate the elements present and the relative number of atoms.
Chemical nomenclature allows systematic naming of ionic compounds, molecular compounds, and acids.