2.3 Formulas

An ionic formula represents a neutral ionic compound: the total positive charge from the cations must equal the total negative charge from the anions. Use the ion charges and choose the simplest whole-number ratio of ions that gives overall charge zero.

Use a fixed sequence: (1) write the cation and anion with their charges, (2) choose subscripts so total charge balances, (3) simplify the ratio if a common factor remains, and (4) remove charge labels from the final neutral formula. Do not change a polyatomic ion's internal subscripts.

If more than one polyatomic ion is needed, put the ion in brackets before adding its subscript, for example Ca²⁺ with NO₃⁻ gives Ca(NO₃)₂. Common ions such as NO₃⁻, CO₃²⁻, SO₄²⁻, OH⁻, NH₄⁺ and PO₄³⁻ must be treated as intact charged groups when constructing the formula.

Use a Roman numeral as the oxidation number of a variable-charge metal, not as a formula subscript: iron(III) is Fe³⁺ before balancing. Do not balance an equation by changing these formula subscripts; equation coefficients belong to the neighbouring objective.

A balanced chemical equation conserves each element: the number of atoms of every element is the same on both sides. For an ionic equation, total charge must also be conserved, and state symbols should identify the reacting species and products.

Use a fixed sequence: (1) write the correct formulae and state symbols, (2) count atoms on both sides, (3) change coefficients until every element balances, (4) simplify to the smallest whole-number ratio, and (5) check atoms, charge and states. Never change a formula subscript to balance an equation.

To obtain a net ionic equation, write the complete ionic equation, split only appropriate aqueous strong electrolytes into ions, then cancel ions that are unchanged on both sides. The remaining species are the particles that undergo the chemical change; recheck charge as well as atom counts.

Coefficients change the amounts of substances while subscripts define the substances themselves. Do not cancel a reacting ion as a spectator, assume every dissolved species should be split, or omit state symbols when they are needed to distinguish the ionic form.

An empirical formula shows the simplest whole-number ratio of atoms or ions in a substance. A molecular formula shows the actual number of each atom in one molecule and is a whole-number multiple of the empirical formula when the substance is molecular.

Use the question being asked to choose the representation: empirical formula answers 'what is the simplest composition ratio?', whereas molecular formula answers 'how many atoms are present in one molecule?'. For an ionic compound, the formula unit gives the simplest ratio rather than a discrete molecule.

To relate the two, divide the molecular formula by the empirical formula to obtain a whole-number multiple. The molecular formula may equal the empirical formula when the simplest ratio is already the actual molecular composition; otherwise every empirical subscript is multiplied by the same integer.

Do not call an empirical formula an incomplete molecular formula, and do not change individual subscripts independently. Calculating the ratio from mass or percentage data belongs to the neighbouring calculation objective; this card establishes the meanings and relationship of the two terms.

An anhydrous salt contains no water of crystallisation in its crystal structure. A hydrated salt contains a fixed number of water molecules associated with each formula unit; the water is part of the crystal composition and is written after a dot in the hydrate formula.

Water of crystallisation is not simply liquid solvent or surface moisture. It is held in a definite stoichiometric ratio, so a hydrate can be represented as salt · nH₂O, where n is the number of water molecules per formula unit in that crystal.

Heating a hydrated salt can remove its water of crystallisation and form the anhydrous salt. The salt formula and its recorded mass therefore change together; use the stated hydrate ratio rather than treating the water content as arbitrary.

Do not interpret hydrated as dissolved, wet or chemically bonded into a new molecular compound, and do not assume every salt has water of crystallisation. The quantitative mass-loss calculation belongs to the neighbouring formula-calculation objective.

To calculate an empirical formula from composition data, first record the mass or percentage of each element. Convert each value to amount using its relative atomic mass, divide every amount by the smallest amount, and scale the ratios to the smallest whole numbers that fit the data.

Keep the element order and units consistent, then check that the final integer ratio reproduces the supplied composition within the stated precision. Do not round each raw ratio independently before comparing the ratios; use a common multiplier when a simple fraction is present.

If a relative molecular mass is supplied, calculate the empirical-formula mass and find the whole-number multiple: multiplier = relative molecular mass ÷ empirical-formula mass. Multiply every empirical subscript by that same integer to obtain the molecular formula.

An empirical formula is the simplest ratio, whereas a molecular formula gives the actual atom numbers in one molecule. Do not use the molecular-multiple step for an ionic lattice, alter individual subscripts, or extend the calculation into reacting-mass or percentage-yield problems.