Chemical Compounds, Formulas, and the Law of Constant Composition

Chapter 5: Chemical Compounds and Formulas

Objectives

• Understand and apply the law of constant composition

• Interpret chemical formulas

• Determine the total number of each type of atom in a chemical formula

• Classify each element as atomic or molecular

• Classify compounds as ionic or molecular

• Write formulas for ionic compounds

How Elements and Compounds Differ

Properties of Elements vs. Compounds

Elements and compounds have distinct physical and chemical properties. When elements combine to form compounds, the resulting compound often has properties very different from its constituent elements.

• Elemental Sodium: Highly reactive, toxic metal that dulls instantly upon air exposure.

• Elemental Chlorine: Yellow gas with a pungent odor, highly reactive and poisonous.

• Sodium Chloride (NaCl): A compound formed from sodium and chlorine; it is a stable, edible table salt.

Example: Table salt (NaCl) is safe to eat, while its constituent elements (sodium and chlorine) are hazardous in their pure forms.

Classification of Substances

Elements vs. Compounds

Most substances found in nature are compounds, not elements. Compounds are formed when elements chemically combine in fixed ratios, while mixtures have variable proportions.

• Elements: Pure substances consisting of only one type of atom.

• Compounds: Substances composed of two or more elements chemically combined in fixed ratios.

• Mixtures: Combinations of substances in variable proportions, not chemically bonded.

Example: Water (H2O) always has a fixed ratio of hydrogen to oxygen, while a mixture of hydrogen and oxygen gases can have any ratio.

The Law of Constant Composition

Definition and Application

The law of constant composition (also known as the law of definite proportions) states that all samples of a given compound have the same proportions of their constituent elements by mass.

• Formulated by Joseph Proust.

• Regardless of the source or amount, a compound's elemental composition is always the same.

Example: In water, the mass ratio of oxygen to hydrogen is always 8:1.

$\text{Mass ratio of O to H in water} = \frac{16.0\,\text{g O}}{2.0\,\text{g H}} = 8.0:1$

For ammonia (NH3):

$\text{Mass ratio of N to H in ammonia} = \frac{14.0\,\text{g N}}{3.0\,\text{g H}} = 4.7:1$

Chemical Formulas

Types and Interpretation

Chemical formulas represent the elements present in a compound and the number of atoms of each element.

• Subscript: Indicates the number of atoms of an element in the compound.

• If no subscript is shown, it is assumed to be 1.

Examples:

• Water: H2O (2 hydrogen atoms, 1 oxygen atom)

• Carbon dioxide: CO2 (1 carbon atom, 2 oxygen atoms)

• Sodium chloride: NaCl (1 sodium atom, 1 chlorine atom)

• Sucrose: C12H22O11 (12 carbon, 22 hydrogen, 11 oxygen atoms)

Changing the subscript changes the compound. For example, CO and CO2 are different compounds.

Order of Elements in Formulas

• For compounds with metals, the metal is listed first (e.g., NaCl, not ClNa).

• For compounds with only nonmetals, the more metal-like element (usually to the left in the periodic table) is listed first (e.g., NO2, not O2N).

Polyatomic Ions

Definition and Representation

Polyatomic ions are groups of atoms that act as a single charged unit in chemical reactions and compounds.

• When multiple polyatomic ions are present, parentheses are used with a subscript outside to indicate the number of groups.

Example: Magnesium nitrate: Mg(NO3)2 (1 magnesium ion, 2 nitrate ions)

Types of Chemical Formulas

Empirical, Molecular, and Structural Formulas

• Empirical formula: Simplest whole-number ratio of atoms in a compound.

• Molecular formula: Actual number of atoms of each element in a molecule.

• Structural formula: Shows how atoms are connected by chemical bonds.

Example: Glucose has an empirical formula CH2O, molecular formula C6H12O6, and a structural formula showing the arrangement of atoms.

Basic Units of Substances

Atomic, Molecular, and Formula Units

• Elements: Can exist as single atoms (e.g., Hg) or molecules (e.g., H2).

• Molecular compounds: Basic unit is a molecule (e.g., CO2).

• Ionic compounds: Basic unit is a formula unit (e.g., NaCl), representing the simplest ratio of ions.

Writing Formulas for Ionic Compounds

Rules and Examples

Ionic compounds are formed from metals and nonmetals. The formula must balance the total positive and negative charges.

• Write the symbol and charge for the metal (cation) and nonmetal (anion).

• Balance the charges so the total is zero.

• Use subscripts to indicate the number of each ion needed.

• Reduce subscripts to the smallest whole-number ratio.

Examples:

Cation	Anion	Formula
Li+	Cl-	LiCl
Mg2+	Cl-	MgCl2
K+	O2-	K2O
Mg2+	O2-	MgO

Formulas for Compounds with Polyatomic Ions

When writing formulas for compounds containing polyatomic ions, balance the charges as with simple ions.

Cation	Polyatomic Anion	Formula
Na+	OH- (hydroxide)	NaOH
K+	PO43- (phosphate)	K3PO4
Al3+	CO32- (carbonate)	Al2(CO3)3
Fe3+	NO3- (nitrate)	Fe(NO3)3

Rule: The sum of the charges on the ions must equal zero.

Summary Table: Types of Compounds and Their Basic Units

Type	Basic Unit	Example
Element (atomic)	Atom	Hg
Element (molecular)	Molecule	H2
Molecular compound	Molecule	CO2
Ionic compound	Formula unit	NaCl

Additional info: Polyatomic ions are listed in standard chemistry reference tables (e.g., pg. 151 of textbook). For more practice, see problems 1, 3-4, 10, 25-26, 33-34, 37-38, 41, 47-48, 53, 58.