Naming Chemical Compounds: Ionic, Molecular, Acids and Hydrates

Choosing the naming system from the formula

Look at what the formula contains before you try to name it. Ionic compounds are held together by charges, so their names report the ions and leave the counting to charge balance. Molecular compounds have no ionic charges to balance, so their names must state the number of each atom with Greek prefixes. Acids have their own names because they are named for how they behave in water.

Read each formula through the checklist below, in order, and stop at the first match.

• Formula followed by a dot and water, such as ·5H₂O: a hydrate. Name the compound, then add the hydrate term.
• Formula that begins with H and is written as an aqueous solution, (aq): an acid.
• A metal, or the ammonium ion NH₄⁺, combined with nonmetals: an ionic compound.
• Only nonmetals, with no ammonium: a molecular compound named with prefixes.

Ionic names and Stock numbers for variable charges

Name the cation first, then the anion. A monatomic cation keeps the element name (sodium, calcium), and a monatomic anion takes the root of the element name plus -ide (chloride, oxide, nitride). Ionic names never use prefixes such as di- or tri- to count the ions; the subscripts follow from the charges, so stating them would be redundant.

Many metals, especially transition metals and tin and lead, form more than one cation. For these, the Stock system puts the metal's oxidation number, which for a monatomic cation equals its charge, in Roman numerals in parentheses: iron(II) is Fe²⁺ and iron(III) is Fe³⁺. Metals with only one common charge do not get a numeral: group 1 metals (+1), group 2 metals (+2), aluminum (+3), zinc (+2) and silver (+1).

Worked example: name Co₃(PO₄)₂. Phosphate is PO₄³⁻, so two phosphates carry 2 × (−3) = −6. The three cobalt ions must total +6, so each is +6 ÷ 3 = +2. The name is cobalt(II) phosphate. Note that the Roman numeral is the charge of one ion, not the subscript.

Worked example in reverse: write the formula of tin(IV) sulfide. Tin(IV) is Sn⁴⁺ and sulfide is S²⁻. The smallest whole-number ratio that balances charge is one Sn⁴⁺ for every two S²⁻, since +4 + 2(−2) = 0. The formula is SnS₂.

Polyatomic ions and the -ate/-ite pattern

A polyatomic ion is a group of covalently bonded atoms with an overall charge. It keeps its name and formula as a unit inside an ionic compound. When a formula needs more than one copy of a polyatomic ion, enclose it in parentheses before the subscript: calcium nitrate is Ca(NO₃)₂, not CaNO₃₂.

Many oxyanions come in families. The -ate ion has one more oxygen than the -ite ion with the same charge: sulfate is SO₄²⁻ and sulfite is SO₃²⁻. For the chlorine family, per- adds one more oxygen than -ate, and hypo- removes one more than -ite. Adding H⁺ to an anion makes its charge one unit less negative and adds hydrogen to the name. The list below covers the ions most general-chemistry courses expect you to know.

• Ammonium NH₄⁺, hydroxide OH⁻, cyanide CN⁻, peroxide O₂²⁻, acetate CH₃COO⁻.
• Nitrate NO₃⁻, nitrite NO₂⁻; sulfate SO₄²⁻, sulfite SO₃²⁻, hydrogen sulfate HSO₄⁻.
• Carbonate CO₃²⁻, hydrogen carbonate HCO₃⁻; phosphate PO₄³⁻, hydrogen phosphate HPO₄²⁻, dihydrogen phosphate H₂PO₄⁻.
• Perchlorate ClO₄⁻, chlorate ClO₃⁻, chlorite ClO₂⁻, hypochlorite ClO⁻ (bromine and iodine oxyanions are named the same way).
• Permanganate MnO₄⁻, chromate CrO₄²⁻, dichromate Cr₂O₇²⁻.

Prefixes for binary molecular compounds

For a compound of two nonmetals, the element farther left in the periodic table is generally named first; if both are in the same group, the lower one comes first. The second element takes the -ide ending. Traditional names for some compounds of oxygen with chlorine, bromine and iodine break this pattern by naming the halogen first, as in dichlorine monoxide, Cl₂O; the 2005 IUPAC recommendations remove this exception and name the same compound oxygen dichloride, OCl₂.

Greek prefixes give the number of each atom: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9) and deca- (10). Mono- is never used on the first element. Most textbooks drop the final a or o of a prefix before oxide, giving monoxide and pentoxide; the 2005 IUPAC recommendations keep the vowel except in monoxide, so you will also see names such as dinitrogen pentaoxide.

Worked example: name P₄S₁₀. Both elements are nonmetals, so use prefixes. Phosphorus is farther left, so it comes first: tetraphosphorus. Sulfur gets deca- and the -ide ending: decasulfide. The name is tetraphosphorus decasulfide. Going the other way, dinitrogen trioxide is N₂O₃. Some molecular compounds are known almost only by common names, such as water, ammonia and methane.

Acid names from their anions

An acid's name is derived from the anion that remains after the hydrogen ions are removed. If the anion ends in -ide, the acid is named hydro- + root + -ic acid: HBr(aq) is hydrobromic acid. Without water present, HBr is a molecular gas named hydrogen bromide; the acid name describes the aqueous solution.

For oxyacids, change -ate to -ic and -ite to -ous, and keep any per- or hypo- prefix. Oxyacid names never take the hydro- prefix. Nitrate gives nitric acid, HNO₃; nitrite gives nitrous acid, HNO₂; sulfate gives sulfuric acid, H₂SO₄; phosphate gives phosphoric acid, H₃PO₄.

Worked example: name HIO₃ and write the formula of hypobromous acid. Removing H⁺ from HIO₃ leaves IO₃⁻, iodate, by analogy with chlorate. The -ate ending becomes -ic, so the acid is iodic acid. For hypobromous acid, the -ous ending and hypo- prefix point to hypobromite, BrO⁻. One H⁺ balances one negative charge, so the formula is HBrO.

Hydrates: names, formulas and percent water

A hydrate is a solid, usually an ionic compound, with water molecules built into its crystal in a fixed ratio. The formula shows the water after a centered dot, and the name adds a prefix to the word hydrate: CoCl₂·6H₂O is cobalt(II) chloride hexahydrate. Heating a hydrate can drive off the water and leave the anhydrous compound.

Worked example: percent water in CoCl₂·6H₂O. Use IUPAC standard atomic weights rounded to two decimal places: Co 58.93, Cl 35.45, H 1.01 and O 16.00 g/mol. The molar mass of CoCl₂ is 58.93 + 2(35.45) = 129.83 g/mol. Water is 2(1.01) + 16.00 = 18.02 g/mol, so six waters are 6 × 18.02 = 108.12 g/mol. The hydrate's molar mass is 129.83 + 108.12 = 237.95 g/mol, and the mass percent of water is (108.12 ÷ 237.95) × 100% = 45.44%.

The same ratio predicts what heating should leave. If 3.500 g of the hydrate loses all of its water, the anhydrous CoCl₂ remaining is 3.500 g × (129.83 ÷ 237.95) = 1.910 g, and the water driven off is 3.500 g − 1.910 g = 1.590 g. Four significant figures are justified because the sample mass and every molar mass carry at least four.

Naming errors that change the formula

A wrong name is a wrong substance, so treat every name as a formula you can check. After naming an ionic compound, rebuild the formula from your name and confirm that the positive and negative charges sum to zero. After naming a molecular compound, confirm that each prefix matches a subscript.

The following errors are the ones to look for first.

• Using prefixes in ionic names, such as calling MgCl₂ magnesium dichloride. The correct name is magnesium chloride.
• Reading the Roman numeral as a subscript. Iron(III) oxide is Fe₂O₃, not Fe₃O.
• Omitting the numeral for a variable-charge metal, or adding one for sodium or calcium.
• Confusing sulfide (S²⁻), sulfite (SO₃²⁻) and sulfate (SO₄²⁻).
• Adding hydro- to an oxyacid. Hydrosulfuric acid is H₂S(aq); H₂SO₄ is sulfuric acid.
• Missing that mercury(I) exists as the Hg₂²⁺ unit, so mercury(I) chloride is Hg₂Cl₂.