An ionic compound is formed by the strong electrostatic attraction between positively and negatively charged particles called ions. These compounds typically form when a metal (which loses electrons) reacts with a nonmetal (which gains electrons). The resulting positive ions (cations) and negative ions (anions) are held together by ionic bonding in a highly ordered, three-dimensional structure called a crystal lattice.
The Zero Net Charge Rule
The overall electrical charge of any stable ionic compound is zero, meaning the compound is electrically neutral. This principle, known as the Zero Net Charge Rule, is a fundamental requirement for stability. The total positive charge contributed by all cations must perfectly cancel out the total negative charge contributed by all anions. This neutrality is maintained because electrons are transferred, not created or destroyed, during the compound’s formation. The precise ratio of cations to anions within the crystal lattice ensures this perfect charge cancellation.
Predicting Individual Ion Charges
The charge an individual ion acquires is directly related to its position on the periodic table and the desire of atoms to achieve a stable electron configuration. Atoms tend to gain or lose electrons so that their outermost electron shell is full. This systematic relationship means the charge on most main-group ions can be predicted simply by knowing the element’s group number.
Cations (Positive Ions)
Metals, found on the left side of the table, have few valence electrons and readily lose them to form positive cations. Elements in Group 1 (Alkali Metals) lose one electron to form ions with a +1 charge. Group 2 elements (Alkaline Earth Metals) lose two electrons, resulting in a +2 charge.
Anions (Negative Ions)
Conversely, nonmetals on the right side of the table gain electrons to form negative anions. Group 17 elements (Halogens) gain one electron to form a \(-1\) charge. Group 16 elements gain two electrons for a \(-2\) charge.
Combining Ions to Achieve Neutrality
Once the individual charges of the cation and anion are known, the correct ratio in which they must combine is determined to satisfy the Zero Net Charge Rule. This requires finding the smallest whole number ratio that makes the total positive charge equal to the total negative charge. The number of ions needed for each element is represented using subscripts in the chemical formula.
Consider the formation of the compound between Aluminum ions (\(\text{Al}^{3+}\)) and Oxygen ions (\(\text{O}^{2-}\)). The lowest common multiple of the charges (3 and 2) is 6, meaning the compound must have a total charge of \(6+\) and \(6-\). This requires two aluminum ions (\(2 \times 3+\)) and three oxygen ions (\(3 \times 2-\)), resulting in the neutral chemical formula \(\text{Al}_2\text{O}_3\). A simple technique to quickly determine these subscripts is to use the magnitude of the charge from one ion as the subscript for the other ion.
Understanding Polyatomic Ions
Some ionic compounds contain polyatomic ions, which are groups of two or more atoms covalently bonded together that collectively carry a single electrical charge. These charged groups act as a single unit when forming an ionic compound, such as the sulfate ion (\(\text{SO}_4^{2-}\)) or the ammonium ion (\(\text{NH}_4^{+}\)). When combining with other ions, the rule of zero net charge still applies, and the entire charge of the polyatomic unit must be factored into the balance. If more than one polyatomic ion is needed for neutrality, the group is enclosed in parentheses before applying the subscript, as seen in \(\text{Ca}(\text{NO}_3)_2\).