Classes of Compounds and Nomenclature

Compounds fall into two main categories based on how their atoms are bonded:

• Ionic compounds form when metals transfer electrons to nonmetals, creating positively charged cations and negatively charged anions held together by electrostatic attraction. They typically have high melting points and conduct electricity when dissolved in water. Examples: NaCl, CaO, MgBr₂.
• Molecular (covalent) compounds form when nonmetals share electrons. They tend to have lower melting points and do not conduct electricity in solution. Examples: H₂O, CO₂, CH₄.

A quick rule of thumb: if the compound contains a metal and a nonmetal, it is likely ionic. If it contains only nonmetals, it is likely molecular. Exceptions exist (e.g., ammonium salts like NH₄Cl are ionic despite containing only nonmetals), but this heuristic covers most cases.

Binary ionic compounds contain exactly two elements: a metal cation and a nonmetal anion. The naming rules are straightforward:

1. Name the cation first (simply the element name): sodium, calcium, aluminum
2. Name the anion second, changing the element's ending to -ide: chlorine → chloride, oxygen → oxide, sulfur → sulfide

Examples:

• NaCl = sodium chloride
• CaO = calcium oxide
• Al₂O₃ = aluminum oxide
• MgBr₂ = magnesium bromide

Notice that the formula subscripts are not included in the name for ionic compounds — they are implied by the charges. This is a key difference from molecular compound naming.

To write a formula from a compound's name, you need to know the charges on the ions and then balance them so the compound is electrically neutral.

The process:

1. Identify the cation and anion from the name
2. Determine each ion's charge (from periodic table position or memorization)
3. Find the simplest ratio that balances the total positive and negative charges

For example, “calcium chloride”: Ca²⁺ and Cl^-. To balance 2+ with the 1− charges, you need two chloride ions: CaCl₂. Another example: “aluminum oxide”: Al³⁺ and O^2-. The lowest common multiple of 3 and 2 is 6, requiring 2 Al and 3 O: Al₂O₃.

A quick shortcut: cross the charge numbers as subscripts (reduce to lowest ratio). Ca²⁺ + Cl^- → CaCl₂. Al³⁺ + O^2- → Al₂O₃.

Many ionic compounds contain polyatomic ions — charged groups of covalently bonded atoms that behave as a single unit. When naming these compounds, treat the polyatomic ion as a single entity and use its established name.

Examples:

• NaNO₃ = sodium nitrate
• Ca(OH)₂ = calcium hydroxide
• NH₄Cl = ammonium chloride
• (NH₄)₂SO₄ = ammonium sulfate

When a formula requires more than one polyatomic ion, place parentheses around the ion and add the subscript outside: Ca(NO₃)₂ means one Ca²⁺ and two NO₃^- ions. The parentheses are essential — without them, the subscript applies only to the last atom.

Binary molecular compounds (two nonmetals) use Greek prefixes to indicate the number of each atom, unlike ionic compounds which rely on charge balance.

The prefixes:

• 1 = mono-, 2 = di-, 3 = tri-, 4 = tetra-, 5 = penta-, 6 = hexa-, 7 = hepta-, 8 = octa-, 9 = nona-, 10 = deca-

Rules:

1. Name the first element with its prefix (drop mono- for the first element)
2. Name the second element with its prefix and -ide ending
3. Drop the trailing vowel of a prefix when it precedes a vowel in the element name (mono + oxide = monoxide, not monooxide)

Examples: N₂O₄ = dinitrogen tetroxide; PCl₅ = phosphorus pentachloride; CO = carbon monoxide; CO₂ = carbon dioxide.

Acids are compounds that produce H⁺ (H₃O⁺) ions in water. Their naming depends on whether the anion contains oxygen:

Binary acids (no oxygen, formula H + nonmetal):

• Use the pattern: hydro- + root of element + -ic acid
• HCl = hydrochloric acid; HBr = hydrobromic acid; H₂S = hydrosulfuric acid

Oxyacids (contain oxygen, formula H + polyatomic ion with O):

• If the anion ends in -ate, the acid ends in -ic acid: SO₄^2- (sulfate) → H₂SO₄ = sulfuric acid
• If the anion ends in -ite, the acid ends in -ous acid: SO₃^2- (sulfite) → H₂SO₃ = sulfurous acid

The key pattern: -ate → -ic acid, -ite → -ous acid. This rule applies to all oxyacids (nitric/nitrous, sulfuric/sulfurous, phosphoric/phosphorous, etc.).

A hydrate is an ionic compound that includes a specific number of water molecules within its crystal structure. The water is written after a centered dot in the formula, and named with a Greek prefix + “hydrate.”

Examples:

• CuSO₄·5H₂O = copper(II) sulfate pentahydrate (blue crystals)
• Na₂CO₃·10H₂O = sodium carbonate decahydrate (washing soda)
• CoCl₂·6H₂O = cobalt(II) chloride hexahydrate

When a hydrate is heated, the water is driven off and the compound becomes anhydrous (without water). This often causes a visible color change — for example, hydrated cobalt(II) chloride is pink, but the anhydrous form is blue. This color change makes it useful as a humidity indicator.

The periodic table is a powerful tool for predicting the charges that main-group elements carry when they form ions:

• Group 1 metals form 1+ cations (Li⁺, Na⁺, K⁺)
• Group 2 metals form 2+ cations (Mg²⁺, Ca²⁺, Ba²⁺)
• Group 13 aluminum forms 3+ (Al³⁺)
• Group 17 nonmetals form 1− anions (F^-, Cl^-, Br^-)
• Group 16 nonmetals form 2− anions (O^2-, S^2-)
• Group 15 nitrogen forms 3− (N^3-)

The pattern: metals lose electrons to reach the nearest noble gas configuration, while nonmetals gain electrons to reach the nearest noble gas configuration. The charge equals the number of electrons lost or gained.

Transition metals are the exception — they can form multiple different charges (Fe²⁺ and Fe³⁺, Cu⁺ and Cu²⁺), which is why Roman numerals are needed in their names.

Every ionic compound is electrically neutral — the total positive charge must equal the total negative charge. This charge balance principle determines the formula subscripts.

Step-by-step method:

1. Write the symbols and charges of the cation and anion
2. Find the least common multiple (LCM) of the charge magnitudes
3. Divide the LCM by each ion's charge to get the subscripts
4. Write the formula: cation first, anion second, subscripts as needed

Example: iron(III) oxide. Fe³⁺ and O^2-. LCM of 3 and 2 = 6. Subscript for Fe = 6/3 = 2. Subscript for O = 6/2 = 3. Formula: Fe₂O₃.

For polyatomic ions, enclose in parentheses when the subscript is greater than 1: Ca²⁺ and NO₃^- → Ca(NO₃)₂.

Transition metals (and some main-group metals like Pb and Sn) can form ions with different charges. To specify which charge is present, a Roman numeral in parentheses follows the metal name.

Examples:

• FeCl₂ = iron(II) chloride (Fe²⁺)
• FeCl₃ = iron(III) chloride (Fe³⁺)
• CuO = copper(II) oxide (Cu²⁺)
• Cu₂O = copper(I) oxide (Cu⁺)

To determine the Roman numeral from a formula: calculate the charge on the metal ion by working backward from the known anion charges. In FeCl₃, each Cl is 1−, total anion charge is 3−, so Fe must be 3+. The name is iron(III) chloride.

Main-group metals with only one possible charge (like Na⁺, Ca²⁺, Al³⁺) do not need Roman numerals.

A relatively small set of polyatomic ions appears repeatedly in chemistry. Memorizing them is essential for naming and writing formulas:

• Positive: NH₄⁺ (ammonium)
• −1 charge: OH^- (hydroxide), NO₃^- (nitrate), NO₂^- (nitrite), ClO₃^- (chlorate), C₂H₃O₂^- (acetate), MnO₄^- (permanganate), HCO₃^- (bicarbonate), CN^- (cyanide)
• −2 charge: SO₄^2- (sulfate), SO₃^2- (sulfite), CO₃^2- (carbonate), CrO₄^2- (chromate), Cr₂O₇^2- (dichromate), HPO₄^2- (hydrogen phosphate)
• −3 charge: PO₄^3- (phosphate)

Notice a helpful naming pattern: many polyatomic ions come in pairs. The -ate form has more oxygen atoms; the -ite form has one fewer oxygen. Nitrate (NO₃^-) vs. nitrite (NO₂^-). Sulfate (SO₄^2-) vs. sulfite (SO₃^2-). The charge stays the same between pairs.

An older (but still widely used) naming system uses the suffixes -ous and -ic to distinguish between lower and higher oxidation states of metals:

• -ous = lower charge: ferrous (Fe²⁺), cuprous (Cu⁺), stannous (Sn²⁺)
• -ic = higher charge: ferric (Fe³⁺), cupric (Cu²⁺), stannic (Sn⁴⁺)

These older names often use Latin roots for the metal: iron → ferr-, copper → cupr-, tin → stann-, lead → plumb-.

The same -ous/-ic pattern appears in acid naming: the -ite polyatomic ion gives an -ous acid, and the -ate ion gives an -ic acid. Sulfite → sulfurous acid; sulfate → sulfuric acid.

While the modern IUPAC system using Roman numerals is preferred, you will encounter -ous/-ic names in older references and on many reagent bottles.

Being able to classify a compound quickly determines which naming rules to apply:

• Ionic compounds: Contain a metal + nonmetal (or polyatomic ion). Named with cation first, anion second (-ide ending or polyatomic ion name). Use Roman numerals for transition metals. Example: CuSO₄ = copper(II) sulfate.
• Molecular compounds: Contain only nonmetals (excluding acids). Named with Greek prefixes. Example: N₂O₅ = dinitrogen pentoxide.
• Acids: Produce H⁺ in water. Binary acids: hydro___ic acid. Oxyacids: ___ic acid (-ate anion) or ___ous acid (-ite anion). Example: HNO₃ = nitric acid.

When you see a formula, the first step is always classification. Check: Does it contain a metal? → Ionic. Only nonmetals and no H at the front? → Molecular. H at the front + nonmetal or polyatomic anion? → Acid. This decision tree guides you to the correct naming conventions.

Naming compounds correctly requires a systematic decision path:

1. Classify the compound: Is it ionic (metal + nonmetal or contains a polyatomic ion) or molecular (nonmetal + nonmetal)?
2. If ionic: Does the metal have a variable charge? If yes, use Roman numerals. Name the cation first, then the anion (change ending to -ide for monatomic anions).
3. If ionic with polyatomic ions: Use the polyatomic ion name as-is. Watch for -ate vs. -ite endings.
4. If molecular: Use Greek prefixes for both elements (omit mono- on the first element). Change the second element’s ending to -ide.
5. If an acid: Binary acid (no oxygen) → hydro-___-ic acid. Oxyacid from -ate ion → ___-ic acid. Oxyacid from -ite ion → ___-ous acid.

Common mistakes: omitting Roman numerals for transition metals that require them, using prefixes in ionic compound names, confusing -ate/-ite polyatomic ion pairs, writing the formula of an ionic compound without balancing charges to zero, and naming hydrates without specifying the number of water molecules using Greek prefixes.