Empirical formulae

1. The empirical formula gives the simplest whole-number ratio of atoms or ions, which can differ from the actual numbers of atoms in a molecule.
2. To find the empirical formula from a molecular formula, divide every subscript by their highest common factor.
   1. Ethane, $C_2{H}_6$, has a ratio of $2:6$, which simplifies to $1:3$, so its empirical formula is ${CH}_3$.
   2. Glucose, $C_6{H}_{12}{O}_6$, simplifies by dividing every subscript by 6, giving the empirical formula ${CH}_{2}O$.
   3. Hydrogen peroxide, $H_2{O}_2$, has the empirical formula $HO$.
3. When the subscripts share no common factor, the empirical formula and the molecular formula are the same, as in $H_{2}O$, ${CO}_2$ and ${CH}_4$.
4. Different compounds can share an empirical formula, so the empirical formula alone cannot identify a molecular substance.
   1. Ethene, $C_2{H}_4$, and propene, $C_3{H}_6$, both have the empirical formula ${CH}_2$.
5. A giant ionic lattice contains an enormous number of ions, so the formula of an ionic compound always shows the simplest ratio of its ions.
   1. Sodium chloride is written $NaCl$ because its lattice contains sodium ions and chloride ions in the ratio $1:1$.
6. To deduce an empirical formula from a model or diagram, count each type of atom or ion in the part shown, write the numbers as a ratio, and simplify it.

Formulae of ionic compounds

1. An ionic compound has no overall charge, so the total positive charge of its ions must cancel the total negative charge.
2. The charge on a simple ion can be predicted from the position of the element in the Periodic Table.
   1. Group I metals form $1+$ ions, such as ${Na}^{+}$ and $K^{+}$.
   2. Group II metals form $2+$ ions, such as ${Mg}^{2+}$ and ${Ca}^{2+}$.
   3. Aluminium, in Group III, forms ${Al}^{3+}$ ions.
   4. Group VI non-metals form $2-$ ions, such as the oxide ion, $O^{2-}$.
   5. Group VII non-metals form $1-$ ions, such as the chloride ion, ${Cl}^{-}$.
3. For a transition element, the Roman numeral in the name gives the charge, so iron(III) oxide contains ${Fe}^{3+}$ ions and copper(II) sulfate contains ${Cu}^{2+}$ ions.
4. A compound ion is a group of atoms that stays together and carries one overall charge, and its formula never changes from one compound to another.
   1. The common negative compound ions are hydroxide, ${OH}^{-}$, nitrate, ${NO}_3^{-}$, carbonate, ${CO}_3^{2-}$, and sulfate, ${SO}_4^{2-}$.
   2. The common positive compound ion is ammonium, ${NH}_4^{+}$.
5. To write a formula from the charges, find the smallest number of each ion that makes the charges add up to zero.
6. Write the positive ion first, show the number of each ion as a subscript, and leave the charges out of the finished formula.
7. Put a compound ion in brackets when more than one of it is needed, as in ${Ca(OH)}_2$ and ${Mg(NO}_3{)}_2$.
8. From a model or diagram of an ionic lattice, count each type of ion, simplify the ratio, and then check that the charges balance.

Deducing symbol equations

1. Many questions describe a reaction in words and expect you to turn the description into a balanced symbol equation with state symbols.
2. Pick out every reactant and product named in the description, then write the correct formula for each one.
3. Use the clues in the wording to choose each state symbol.
   1. A solution, or anything described as dissolved in water, is $(aq)$.
   2. A precipitate, a powder, a metal, crystals or any other solid is $(s)$.
   3. Bubbles, fizzing or a gas given off show that the product is $(g)$.
   4. Water formed in a reaction is $(l)$ unless it is given off as steam.
4. Balance the equation using coefficients only, then check the atoms on both sides.

Ionic equations

1. When ionic compounds dissolve, their ions separate and move freely through the water.
2. In many reactions between solutions, some of these ions take no part and are present unchanged at the end.
3. An ionic equation leaves out these spectator ions and shows only the particles that actually change.
4. Write an ionic equation by following these steps:
   1. Write the balanced symbol equation with state symbols.
   2. Split every aqueous ionic compound into its separate ions, keeping the coefficients.
   3. Leave solids, liquids, gases and molecular substances as complete formulae.
   4. Cross out any ion that appears unchanged on both sides.
   5. Write out what remains, keeping the state symbols.
5. A correct ionic equation is balanced for atoms and also for charge, so the total charge on the left equals the total charge on the right.
6. One ionic equation can describe many different reactions.
   1. Mixing any solution containing ${Ag}^{+}$ ions with any solution containing ${Cl}^{-}$ ions forms the same white precipitate of silver chloride.