Chemical Formulas, Compounds, and the Law of Constant Composition

Chapter 5: Chemical Formulas and Compounds

Objectives

• Apply the law of constant composition to chemical compounds.

• Interpret and write chemical formulas.

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

• Classify each element as atomic or molecular.

• Identify whether compounds are ionic or molecular.

• Write formulas for ionic compounds, including those with polyatomic ions.

How Elements and Compounds Differ

Properties of Elements vs. Compounds

Elements and compounds have distinct properties. When elements combine to form compounds, the resulting substances often have properties very different from the original elements.

• Elemental Sodium (Na): A highly reactive and toxic metal that dulls instantly upon exposure to air.

• Elemental Chlorine (Cl2): A yellow, poisonous gas with a pungent odor.

• Sodium Chloride (NaCl): A stable, edible compound (table salt) formed from sodium and chlorine, with properties unlike either element.

Example: Table salt (NaCl) is safe to eat, even though its constituent elements are dangerous in their pure forms.

Compounds and Mixtures

Natural Occurrence and Composition

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

• Compounds: Elements combine in fixed ratios (law of constant composition).

• Mixtures: Elements or compounds combine in variable proportions (no fixed ratio).

Example: In a mixture of hydrogen and oxygen gases, the ratio of hydrogen to oxygen can vary. In water (H2O), the ratio of hydrogen to oxygen is always 2:1.

The Law of Constant Composition

Definition and Historical Context

The law of constant composition (also called the law of definite proportions) states that all samples of a given compound have the same proportions of their constituent elements by mass. This law was first formulated by Joseph Proust.

• For any pure compound, the ratio of the masses of the elements is always the same, regardless of the sample size or source.

Examples of Constant Composition

• Water (H2O): In an 18.0 g sample of water, there are 16.0 g of oxygen and 2.0 g of hydrogen. The mass ratio is:

$\text{O:H mass ratio} = \frac{16.0\,g\,O}{2.0\,g\,H} = 8.0:1$

• This ratio holds for any sample of pure water.

• Ammonia (NH3): In a 17.0 g sample, there are 14.0 g of nitrogen and 3.0 g of hydrogen. The mass ratio is:

$\text{N:H mass ratio} = \frac{14.0\,g\,N}{3.0\,g\,H} \approx 4.7:1$

Chemical Formulas

Representing Compounds

Chemical formulas indicate the elements present in a compound and the relative number of atoms of each element. The subscript to the right of an element symbol shows the number of atoms of that element in one unit of the compound.

• H2O: 2 hydrogen atoms and 1 oxygen atom.

• CO2: 1 carbon atom and 2 oxygen atoms.

• NaCl: 1 sodium atom and 1 chlorine atom.

• C12H22O11 (sucrose): 12 carbon, 22 hydrogen, and 11 oxygen atoms.

Important: Changing a subscript changes the compound (e.g., CO vs. CO2).

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 (less electronegative) element is listed first, often following periodic table trends (e.g., NO, CO2).

Polyatomic Ions in Formulas

Polyatomic ions are groups of atoms that act as a single charged unit. When more than one polyatomic ion is present, parentheses are used with a subscript outside to indicate the number of groups (e.g., Mg(NO3)2).

• Mg(NO3)2: 1 magnesium ion (Mg2+) and 2 nitrate ions (NO3-).

• Atom count: Mg: 1, N: 2, O: 6

Types of Chemical Formulas

• Empirical formula: Simplest whole-number ratio of atoms (e.g., CH2O for glucose).

• Molecular formula: Actual number of atoms of each element (e.g., C6H12O6 for glucose).

• Structural formula: Shows how atoms are connected using lines for bonds.

Models: Ball-and-stick and space-filling models visually represent molecular structure.

Basic Units of Substances

Elements

• Atomic elements: Exist as single atoms (e.g., Hg, Na).

• Molecular elements: Exist as molecules with two or more atoms of the same element (e.g., O2, N2).

Compounds

• Molecular compounds: Composed of two or more nonmetals; basic unit is a molecule (e.g., CO2, H2O).

• Ionic compounds: Composed of metals and nonmetals; basic unit is a formula unit (e.g., NaCl, MgO). These do not contain individual molecules but consist of a lattice of alternating cations and anions.

Writing Formulas for Ionic Compounds

Rules for Ionic Compounds

• Ionic compounds form between metals (cations) and nonmetals (anions).

• The total positive charge must balance the total negative charge (the sum of the charges must be zero).

• Write the symbol and charge for the metal, then the nonmetal. Use subscripts to balance charges.

• 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 with Polyatomic Ions

• When a compound contains polyatomic ions, treat the entire ion as a unit.

• Use parentheses if more than one polyatomic ion is needed.

• Balance the total charges to zero.

Examples:

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

Key Principle: The sum of the charges on the ions in an ionic compound must equal zero.

Summary Table: Types of Substances and Their Basic Units

Type	Basic Unit	Example
Atomic Element	Single atom	Hg (mercury)
Molecular Element	Molecule	O2 (oxygen gas)
Molecular Compound	Molecule	CO2 (carbon dioxide)
Ionic Compound	Formula unit	NaCl (sodium chloride)

Additional info: Polyatomic ions commonly encountered in introductory chemistry include sulfate (SO42-), nitrate (NO3-), carbonate (CO32-), phosphate (PO43-), and hydroxide (OH-).