When polyatomic ions are included, the number of ionic compounds increases significantly. Indeed, most ionic compounds contain polyatomic ions. Well-known examples are sodium hydroxide (NaOH) with OH - as the polyatomic anion, calcium carbonate (CaCO 3 ), and ammonium nitrate (NH 4 NO 3 ), which contains two polyatomic ions: NH + and NO 3 - . A list of the more important polyatomic ions is given in the following table, which can be used for reference while learning the charges of polyatomic ions. A great many of them are oxyanions (polyatomic ions that contain oxygen). Table \(\PageIndex{1}\) Ions Charge Name Formula -3.0 Phosphate PO43- NaN Arsenate AsO43- -2.0 Carbonate CO32- NaN Peroxide O22- NaN Sulfate SO42- NaN Sulfite SO32- NaN Chromate CrO42- NaN Dichromate Cr2O72- NaN Hydrogen phosphate HPO42- -1.0 Hydrogen carbonate (bicarbonate) HCO3- NaN Superoxide O2- NaN Hydrogen sulfate HSO4- NaN Dihydrogen phosphate H2PO4- NaN Hydroxide OH- NaN Nitrate NO3- NaN Nitrite NO2- NaN Acetate C2H3O2- or CH3COO- NaN Cyanide CN- NaN Permanganate MnO4- NaN Perchlorate ClO4- NaN Chlorate ClO3- NaN Chlorite ClO2- NaN Hypochlorite ClO- 1.0 Ammonium NH4+ NaN Hydronium H3O+ The properties of compounds containing polyatomic ions are very similar to those of binary ionic compounds. The ions are arranged in a regular lattice and held together by coulombic forces of attraction. The resulting crystalline solids usually have high melting points (1500 °F for CaCO 3 ) and all conduct electricity when molten. Most are soluble in water and form conducting solutions in which the ions can move around as independent entities. In general polyatomic ions are colorless, unless, like CrO 4 2 – or MnO 4 – , they contain a transition-metal atom. The more negatively charged polyatomic ions, like their monatomic counterparts, show a distinct tendency to react with water, producing hydroxide ions; for example, \[\ce{PO_{4}^{3-} + H_{2}O \rightarrow HPO_{4}^{2-} + OH^{-}} \nonumber \] It is important to realize that compounds containing polyatomic ions must be electrically neutral . In a crystal of calcium sulfate, for instance, there must be equal numbers of Ca 2 + and SO 4 2 – ions in order for the charges to balance. The formula is thus CaSO 4 . In the case of sodium sulfate, by contrast, the Na + ion has only a single charge. In this case we need two Na + ions for each SO 4 2 – ion in order to achieve electroneutrality. The formula is thus Na 2 SO 4 . Structurally, polyatomic ions are similar to the ionic solids we saw earlier. An example of a simple ionic compound, NaCl, is seen below, alongside a more complex ionic solid, AgClO 3 - . Notice how both are tightly packed and form a repeating pattern, which lends both compounds strength and brittleness. In the Silver Chlorate (AgClO 3 - ), however, polyatomic ions are present where the Cl - ions are present in the Sodium Chloride (NaCl). On the left is the lattice structure of the ionic solid NaCl, with Na represented by the green spheres and Cl represented by the purple spheres. On the right is the lattice structure of AgClO 3 , with the silver spheres representing Ag and the red and green molecules representing the polyatomic ion ClO 3 Example \(\PageIndex{1}\) : Ionic Formula What is the formula of the ionic compound calcium phosphate? Solution It is necessary to have the correct ratio of calcium ions, Ca 2 + , and phosphate ions, PO 4 3 – , in order to achieve electroneutrality. The required ratio is the inverse of the ratio of the charges on ions. Since the charges are in the ratio of 2:3, the numbers must be in the ratio of 3:2. In other words the solid salt must contain three calcium ions for every two phosphate ions: The formula for calcium phosphate is thus Ca 3 (PO 4 ) 2 .