Chapter 5: Molecules and Compounds – Structured Study Notes

Molecules and Compounds

Introduction to Compounds

Compounds are substances formed from two or more elements chemically combined in fixed proportions. Their properties are often very different from the properties of the individual elements that compose them.

• Compound Example: Sucrose (table sugar) contains carbon, hydrogen, and oxygen atoms, but its properties differ greatly from those of graphite (carbon), hydrogen gas, or oxygen gas.

• Compound Example: Sodium chloride (table salt) is formed from sodium (a reactive metal) and chlorine (a poisonous gas), yet salt is safe and stable.

Elements, Compounds, and Mixtures

Most substances in nature are compounds, not elements. Free atoms are rare. Compounds differ from mixtures in that their constituent elements combine in fixed, definite proportions, while mixtures combine in variable proportions.

• Compound: Elements combine in fixed ratios (e.g., water always has a 2:1 ratio of hydrogen to oxygen).

• Mixture: Elements combine in variable ratios (e.g., a balloon filled with hydrogen and oxygen gas).

The Law of Constant Composition

Joseph Proust's law states that all samples of a given compound have the same proportions of their constituent elements. This is demonstrated by decomposing compounds and measuring the mass ratios of their elements.

• Water: Decomposition yields a constant mass ratio of oxygen to hydrogen (8.0:1).

• Ammonia: Decomposition yields a constant mass ratio of nitrogen to hydrogen (4.7:1).

Formula: Mass ratio = $rac{\text{mass of element 1}}{\text{mass of element 2}}$

Chemical Formulas

Representing Compounds

Chemical formulas indicate the elements present in a compound and the relative number of atoms of each. Subscripts show the number of atoms; a subscript of 1 is omitted.

• Water: $\mathrm{H_2O}$ (2 hydrogen, 1 oxygen)

• Table Salt: $\mathrm{NaCl}$ (1 sodium, 1 chlorine)

• Carbon Dioxide: $\mathrm{CO_2}$ (1 carbon, 2 oxygen)

• Sucrose: $\mathrm{C_{12}H_{22}O_{11}}$ (12 carbon, 22 hydrogen, 11 oxygen)

Changing Subscripts

Changing a subscript in a chemical formula results in a completely different compound. For example, $\mathrm{CO}$ (carbon monoxide) is toxic, while $\mathrm{CO_2}$ (carbon dioxide) is relatively harmless.

Order of Elements in Formulas

Formulas list the most metallic element first. In compounds without metals, the more metal-like element (to the left or lower in the periodic table) is listed first.

• Example: $\mathrm{NaCl}$, not $\mathrm{ClNa}$

• Example: $\mathrm{NO_2}$, not $\mathrm{O_2N}$

Polyatomic Ions

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

• Example: $\mathrm{Mg(NO_3)_2}$ contains 1 magnesium, 2 nitrate groups (each with 1 nitrogen and 3 oxygen).

Types of Chemical Formulas

There are three main types of chemical formulas:

• Empirical Formula: Simplest whole-number ratio (e.g., $\mathrm{H_2O}$ for hydrogen peroxide).

• Molecular Formula: Actual number of atoms (e.g., $\mathrm{H_2O_2}$ for hydrogen peroxide).

• Structural Formula: Shows how atoms are connected (lines represent bonds).

Comparison of Formulas and Models

Molecular, structural, ball-and-stick, and space-filling models help visualize the composition and geometry of molecules.

• Methane ($\mathrm{CH_4}$): One carbon atom bonded to four hydrogen atoms.

Connecting Macroscopic and Molecular Worlds

Structural formulas and models bridge the gap between what we see and the atomic/molecular world.

Molecular View of Elements and Compounds

Classification of Elements

Elements can be atomic (single atoms as basic units) or molecular (diatomic molecules as basic units).

• Atomic Elements: Most metals, e.g., mercury.

• Molecular Elements: Diatomic molecules, e.g., chlorine ($\mathrm{Cl_2}$).

Elements That Exist as Diatomic Molecules

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

Classification of Compounds

Compounds can be molecular (formed from nonmetals, basic unit is a molecule) or ionic (formed from metals and nonmetals, basic unit is a formula unit).

• Molecular Compound Example: Dry ice ($\mathrm{CO_2}$)

• Ionic Compound Example: Table salt ($\mathrm{NaCl}$)

Writing Formulas for Ionic Compounds

Steps for Writing Ionic Compound Formulas

Ionic compounds are formed from metals (cations) and nonmetals (anions). The sum of charges must be zero.

1. Write the symbol and charge for each ion.

2. Balance charges by adjusting subscripts.

3. Reduce subscripts to smallest whole numbers.

4. Check for charge neutrality.

Example: Magnesium and chlorine form $\mathrm{MgCl_2}$.

Example: Magnesium and oxygen form $\mathrm{MgO}$.

Polyatomic Ions in Ionic Compounds

Polyatomic ions are groups of atoms with a charge. When more than one polyatomic ion is present, use parentheses and a subscript.

• Example: Calcium and nitrate form $\mathrm{Ca(NO_3)_2}$.

Common Polyatomic Ions

Name	Formula
Acetate	C2H3O2-
Carbonate	CO32-
Bicarbonate	HCO3-
Hydroxide	OH-
Nitrate	NO3-
Phosphate	PO43-
Ammonium	NH4+

Naming Compounds

Naming Ionic Compounds

Identify the compound as ionic (metal + nonmetal). There are two types:

• Type I: Metal forms only one type of ion (main group metals).

• Type II: Metal forms more than one type of ion (transition metals).

Type I Binary Ionic Compounds

Name = [name of cation] + [base name of anion + -ide]

• Example: $\mathrm{NaCl}$ is sodium chloride.

• Example: $\mathrm{Al_2O_3}$ is aluminum oxide.

Type II Binary Ionic Compounds

For metals with variable charge, specify the charge with a Roman numeral.

• Example: $\mathrm{FeCl_3}$ is iron(III) chloride.

• Example: $\mathrm{CuO}$ is copper(II) oxide.

Naming Ionic Compounds with Polyatomic Ions

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

• Example: $\mathrm{KNO_3}$ is potassium nitrate.

• Example: $\mathrm{Fe(OH)_2}$ is iron(II) hydroxide.

• Example: $\mathrm{NH_4NO_3}$ is ammonium nitrate.

Naming Oxyanions

Oxyanions are polyatomic ions containing oxygen. The naming depends on the number of oxygen atoms:

• -ate: More oxygen atoms (e.g., nitrate $\mathrm{NO_3^-}$)

• -ite: Fewer oxygen atoms (e.g., nitrite $\mathrm{NO_2^-}$)

• per- ... -ate: Most oxygen atoms (e.g., perchlorate $\mathrm{ClO_4^-}$)

• hypo- ... -ite: Least oxygen atoms (e.g., hypochlorite $\mathrm{ClO^-}$)

Polyatomic Ions in Everyday Chemistry

• Sodium hypochlorite: Active ingredient in bleach.

• Sodium bicarbonate: Baking soda, source of CO2 in baking.

• Calcium carbonate: Ingredient in antacids.

• Sodium nitrite: Preservative in meats.

Naming Molecular Compounds

Molecular compounds are formed from two or more nonmetals. Use prefixes to indicate the number of atoms.

• Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, etc.

• Example: $\mathrm{N_2O}$ is dinitrogen monoxide.

• Example: $\mathrm{CO_2}$ is carbon dioxide (mono- omitted for first element).

Naming Acids

Acids are molecular compounds that form ions in water. They are classified as binary acids (hydrogen + nonmetal) or oxyacids (hydrogen + nonmetal + oxygen).

Naming Binary Acids

Name = hydro + [base name of nonmetal + -ic] + acid

• Example: $\mathrm{HCl(aq)}$ is hydrochloric acid.

• Example: $\mathrm{HBr(aq)}$ is hydrobromic acid.

• Example: $\mathrm{H_2S(aq)}$ is hydrosulfuric acid.

Naming Oxyacids

Oxyacids contain oxyanions. The name depends on the ending of the oxyanion:

• -ate: [base name of oxyanion + -ic] + acid (e.g., $\mathrm{HNO_3}$ is nitric acid)

• -ite: [base name of oxyanion + -ous] + acid (e.g., $\mathrm{HNO_2}$ is nitrous acid)

Formula Mass

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

• Example: For $\mathrm{MgCl_2}$: $24.31\ \text{amu} + 2 \times 35.45\ \text{amu} = 95.21\ \text{amu}$

Summary Table: Learning Objectives

• Restate and apply the law of constant composition.

• Write chemical formulas and determine atom counts.

• Classify elements and compounds.

• Write and name formulas for ionic and molecular compounds, including acids.

• Calculate formula mass.