Ionic compounds have no overall charge

1. Metal atoms lose electrons to form positive ions, and non-metal atoms gain electrons to form negative ions.
2. The charge on a simple ion can be predicted from the group number of the element.
   1. Group 1 metals form $1+$ ions, Group 2 metals form $2+$ ions and aluminium, in Group 3, forms ${Al}^{3+}$.
   2. Group 6 non-metals form $2-$ ions and Group 7 non-metals form $1-$ ions, such as $O^{2-}$ and ${Cl}^{-}$.
3. A negative ion made from one element takes the ending -ide, so chlorine forms chloride ions and oxygen forms oxide ions.
4. Transition metals can form ions with different charges, shown by a Roman numeral in the name, so iron(II) is ${Fe}^{2+}$ and iron(III) is ${Fe}^{3+}$.
5. In the compound, the total positive charge must equal the total negative charge.

Polyatomic ions stay together with one overall charge

1. Some ions contain several atoms covalently bonded together, and the charge belongs to the whole group.
2. The common negative ones are hydroxide, ${OH}^{-}$, nitrate, ${NO}_3^{-}$, carbonate, ${CO}_3^{2-}$, and sulfate, ${SO}_4^{2-}$.
   1. An ion name ending in -ate shows that the ion contains oxygen.
3. The ammonium ion, ${NH}_4^{+}$, is a common positive polyatomic ion.
4. When a formula needs more than one polyatomic ion, the ion goes inside brackets with the number after them, as in ${Ca(OH)}_2$.
   1. The 2 after the bracket doubles every atom inside it, so ${Ca(OH)}_2$ contains one calcium, two oxygen and two hydrogen atoms.
5. The atoms inside a polyatomic ion never change, so sulfate is always ${SO}_4$ in a formula.

Writing ionic formulae so the charges cancel

1. Write down the formula and charge of each ion.
2. Find the smallest numbers of each ion that give equal and opposite total charges.
   1. For ${Mg}^{2+}$ and ${Cl}^{-}$, one magnesium ion balances two chloride ions, giving ${MgCl}_2$.
3. Write the positive ion first and show the number of each ion as a subscript, leaving out any 1.
4. Leave the charges out of the final formula, because the compound has no overall charge.
5. When the charges are the same size, one of each ion is needed, so ${Mg}^{2+}$ and $O^{2-}$ give $MgO$.
6. Ammonium sulfate needs two ${NH}_4^{+}$ ions for each ${SO}_4^{2-}$ ion, giving $({NH}_4{)}_2{SO}_4$.

State symbols show solid, liquid, gas or aqueous

1. A state symbol is written in brackets straight after each formula in an equation.
2. The four state symbols are $(s)$ for solid, $(l)$ for liquid, $(g)$ for gas and $(aq)$ for aqueous.
   1. Aqueous means dissolved in water, so a solution of sodium chloride is $NaCl(aq)$.
3. Pure water at room temperature is $H_{2}O(l)$, but steam is $H_{2}O(g)$.
4. Acids and alkalis in the laboratory are solutions, so they are written as $HCl(aq)$ or $NaOH(aq)$.
5. A precipitate forms as a solid when two solutions are mixed, so it takes the symbol $(s)$.

Choosing state symbols from the reaction conditions

1. Metals and solid compounds added to a reaction, such as magnesium ribbon or marble chips, are $(s)$.
2. Gases given off as bubbles, such as hydrogen and carbon dioxide, are $(g)$.
3. Magnesium reacting with hydrochloric acid is written as: $$Mg(s)+2HCl(aq)\to {MgCl}_2(aq)+H_2(g)$$
4. Marble chips reacting with hydrochloric acid are written as: $${CaCO}_3(s)+2HCl(aq)\to {CaCl}_2(aq)+H_{2}O(l)+{CO}_2(g)$$
5. Mixing silver nitrate and sodium chloride solutions gives a white precipitate of silver chloride: $${AgNO}_3(aq)+NaCl(aq)\to AgCl(s)+{NaNO}_3(aq)$$
   1. The sodium nitrate stays dissolved, so it is $(aq)$.