Chapter 5: Molecules and Compounds – Structured Study Notes

Molecules and Compounds

Sugar and Salt: Properties of Compounds vs. Elements

Compounds are substances formed from two or more elements chemically combined in fixed ratios. Their properties are often very different from the properties of the individual elements that compose them. For example, ordinary table sugar (sucrose) is a compound made of carbon, hydrogen, and oxygen atoms, while table salt (sodium chloride) is formed from sodium and chlorine.

• Compound: A substance composed of two or more elements in fixed, definite proportions.

• Example: Sucrose (C12H22O11) and sodium chloride (NaCl).

• Contrast: The properties of sucrose and sodium chloride are very different from those of their constituent elements (carbon, hydrogen, oxygen, sodium, chlorine).

Compounds Display Constant Composition

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

• Law of Constant Composition: All samples of a given compound have the same proportions of their constituent elements.

• Mixture: A combination of elements or compounds in variable proportions.

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

Constant Composition: Mass Ratios

The law of constant composition can be illustrated by decomposing compounds and measuring the mass ratios of their constituent elements. For example, decomposing water yields a fixed mass ratio of oxygen to hydrogen.

• Water: Decomposition of 18.0 g of water yields 16.0 g of oxygen and 2.0 g of hydrogen.

• Mass Ratio Formula: $\text{Mass ratio} = \frac{16.0\ \text{g O}}{2.0\ \text{g H}} = 8.0$ or $8.0:1$

• Ammonia: Decomposition of 17.0 g of ammonia yields 14.0 g of nitrogen and 3.0 g of hydrogen.

• Mass Ratio Formula: $\text{Mass ratio} = \frac{14.0\ \text{g N}}{3.0\ \text{g H}} = 4.7$ or $4.7:1$

Chemical Formulas and Representation

Chemical Formulas: Indicating Elements and Ratios

Chemical formulas are used to represent compounds, indicating the elements present and the relative number of atoms of each. Subscripts show the number of atoms; a subscript of 1 is omitted.

• Example: H2O (water) indicates a 2:1 ratio of hydrogen to oxygen.

• Other Examples: NaCl (table salt, 1:1 ratio), CO2 (carbon dioxide, 1:2 ratio), C12H22O11 (sucrose, 12:22:11 ratio).

• Subscripts: Represent the relative numbers of each type of atom in a compound.

Order of Listing Elements in Chemical Formulas

Formulas list the most metallic element first. If no metal is present, the more metal-like element is listed first. Metals are on the left side of the periodic table; among nonmetals, those to the left are more metal-like.

• Example: NaCl (not ClNa), NO2 (not O2N).

• Table 5.1: Order of listing nonmetal elements in a chemical formula.

Types of Chemical Formulas

There are several ways to represent compounds:

• Empirical Formula: Gives the relative number of atoms of each element.

• Molecular Formula: Gives the actual number of atoms of each element in a molecule.

• Structural Formula: Uses lines to represent chemical bonds and shows how atoms are connected.

• Molecular Models: Ball-and-stick and space-filling models represent the three-dimensional structure of molecules.

Polyatomic Ions in Chemical Formulas

Some chemical formulas contain groups of atoms that act as a unit, called polyatomic ions. When several groups are present, their formula is set off in parentheses with a subscript indicating the number of groups.

• Example: Mg(NO3)2 contains one magnesium atom and two nitrate groups.

• Calculation: Multiply the subscript outside the parentheses by the subscript for each atom inside.

Molecular View and Classification of Elements and Compounds

Classification of Pure Substances

Pure substances can be classified as elements or compounds. Elements may be atomic or molecular, while compounds may be molecular or ionic.

• Atomic Elements: Exist in nature as single atoms (e.g., mercury).

• Molecular Elements: Exist as diatomic molecules (e.g., chlorine, oxygen).

• Molecular Compounds: Composed of molecules formed from two or more nonmetals (e.g., CO2).

• Ionic Compounds: Composed of cations and anions in a three-dimensional array (e.g., NaCl).

Writing and Naming Chemical Compounds

Writing Formulas for Ionic Compounds

Ionic compounds always contain positive and negative ions. The sum of the charges must equal zero. The formula is written by balancing the charges of the cations and anions.

• Steps:

  1. Write the symbol and charge for each ion.

  2. Make the magnitude of the charge on each ion the subscript for the other ion.

  3. Reduce subscripts to the smallest whole-number ratio.

  4. Check that the sum of the charges is zero.

• Example: Magnesium and oxygen form MgO.

Nomenclature: Naming Compounds

Chemists use systematic rules to name compounds. The name can be determined from the formula, and vice versa. Common names exist for some compounds, but systematic names follow specific rules.

• Example: H2O is commonly called water, but its systematic name is dihydrogen monoxide.

Naming Ionic Compounds

Ionic compounds are named based on the type of metal present:

• Type I: Metal with invariant charge (e.g., sodium, magnesium).

• Type II: Metal with variable charge (e.g., iron, copper).

• Binary Ionic Compounds: Name consists of the cation name followed by the base name of the anion with the ending -ide.

• Example: NaCl is sodium chloride.

• Type II Naming: The charge of the metal is specified with a Roman numeral in parentheses.

• Example: FeCl3 is iron(III) chloride.

Naming Ionic Compounds with Polyatomic Ions

Polyatomic ions are groups of atoms with an overall charge. Ionic compounds containing polyatomic ions are named using the name of the polyatomic ion.

• Oxyanions: Anions containing oxygen are named systematically based on the number of oxygen atoms (e.g., nitrate, nitrite, sulfate, sulfite).

• Prefixes: hypo- (less than), per- (more than) are used for series with more than two ions.

• Example: Sodium hypochlorite is the active ingredient in bleach.

Naming Molecular Compounds

Molecular compounds are formed from two or more nonmetals. The first element is the more metal-like one. Prefixes indicate the number of atoms present (mono-, di-, tri-, tetra-, etc.).

• Example: PI3 is phosphorus triiodide.

• Example: N2H4 is dinitrogen tetrahydride.

Naming Acids

Acids are molecular compounds that form ions when dissolved in water. They are composed of hydrogen and one or more nonmetals. Acids are categorized as binary acids (hydrogen and a nonmetal) or oxyacids (hydrogen, a nonmetal, and oxygen).

• Binary Acids: Names have the form hydro- + base name of nonmetal + -ic + acid.

• Oxyacids: Names depend on the oxyanion present. If the oxyanion ends in -ate, the acid name ends in -ic; if it ends in -ite, the acid name ends in -ous.

• Example: HCl(aq) is hydrochloric acid; H2SO4(aq) is sulfuric acid.

Formula Mass: The Mass of a Molecule or Formula Unit

Calculating Formula Mass

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

• Formula: $\text{Formula mass} = \sum \text{atomic masses of all atoms in the formula}$

• Example: The formula mass of sodium sulfite (Na2SO3) is calculated by adding the atomic masses of Na, S, and O.

Summary of Key Chemical Principles and Skills

• Compounds display constant composition.

• Chemical formulas indicate the elements and their ratios.

• Systematic nomenclature allows naming of ionic, molecular compounds, and acids.

• Formula mass is calculated as the sum of atomic masses in the chemical formula.

• Skills include verifying constant composition, writing formulas, classifying elements and compounds, naming compounds, and calculating formula mass.