Question

Give the empirical formulas and names of the compounds formed from the following pairs of ions: (a) \(\mathrm{Rb}^{+}\) and \(\mathrm{I}^{-},\) (b) \(\mathrm{Cs}^{+}\) and \(\mathrm{SO}_{4}^{2-}\) (c) \(\mathrm{Sr}^{2+}\) and \(\mathrm{N}^{3-},(\mathrm{d}) \mathrm{Al}^{3+}\) and \(\mathrm{S}^{2-}\).

Short answer

The empirical formulas and names of the compounds formed from the ions pair are (a) \(\mathrm{RbI}\) - Rubidium Iodide, (b) \(\mathrm{Cs}_{2}\mathrm{SO}_{4}\) - Cesium Sulfate, (c) \(\mathrm{Sr}_{3}\mathrm{N}_{2}\) - Strontium Nitride, and (d) \(\mathrm{Al}_{2}\mathrm{S}_{3}\) - Aluminum Sulfide.

Step-by-step solution

Determine the ratios for each ion pair

We know that when ions combine to form a compound, they do so in a way that makes the total charge of the compound neutral. This means the positive and negative charges must balance each other out. For (a), the ions are \(\mathrm{Rb}^{+}\) and \(\mathrm{I}^{-}\), both with a charge of 1, so they can combine in a 1:1 ratio. For (b), the ions are \(\mathrm{Cs}^{+}\) and \(\mathrm{SO}_{4}^{2-}\); Cs has a charge of +1 and SO4 a charge of -2, so they will combine in a 2:1 ratio. For (c), the ions are \(\mathrm{Sr}^{2+}\) and \(\mathrm{N}^{3-}\); Sr has a charge of +2 and N a charge of -3, so they will combine in a 3:2 ratio. For (d), the ions are \(\mathrm{Al}^{3+}\) and \(\mathrm{S}^{2-}\); Al has a charge of +3 and S a charge of -2, so they will combine in a 2:3 ratio.

Write out the empirical formulas based on the ratios

Now that we have the ratios, we can write out the empirical formulas. For (a), \(\mathrm{Rb}^{+}\) and \(\mathrm{I}^{-}\) will combine to give \(\mathrm{RbI}\). For (b), \(\mathrm{Cs}^{+}\) and \(\mathrm{SO}_{4}^{2-}\) combine to give \(\mathrm{Cs}_{2}\mathrm{SO}_{4}\). For (c), \(\mathrm{Sr}^{2+}\) and \(\mathrm{N}^{3-}\) combine to give \(\mathrm{Sr}_{3}\mathrm{N}_{2}\). For (d), \(\mathrm{Al}^{3+}\) and \(\mathrm{S}^{2-}\) combine to give \(\mathrm{Al}_{2}\mathrm{S}_{3}\).

Name the compounds

We now use the convention for naming simple ionic compounds, which is 'cation + anion'. For anions, we also replace the ending of the element's name with 'ide'. For polyatomic ions like \(\mathrm{SO}_{4}^{2-}\), we use their common names. So, compound (a) is named Rubidium Iodide, compound (b) is named Cesium Sulfate, compound (c) is named Strontium Nitride, and compound (d) is named Aluminum Sulfide.

Key concepts

Ionic Compounds

Ionic compounds are formed through the electrostatic attraction between positively and negatively charged ions. These ions usually result from the loss or gain of electrons to achieve stable electron configurations. Ionic bonds occur between metals and non-metals due to their differing electronegativities, where metals lose electrons while non-metals gain them.

Ions in ionic compounds come together to balance out each other's charges and create a neutral compound. For example, rubidium iodide (\(\mathrm{Rb}^{+}\) and \(\mathrm{I}^{-}\)) combines in a 1:1 ratio because both ions have a single positive and negative charge, respectively. This results in a neutral compound. Ionic compounds typically form crystalline structures at room temperature and have high melting and boiling points.

Understanding the fundamental nature of ionic compounds helps us to predict the ratios in which ions will combine, based on their charges.

Charge Balance

An essential aspect of writing the empirical formulas for ionic compounds is maintaining charge balance. The concept of charge balance requires that the total positive charge from cations (positive ions) equals the total negative charge from anions (negative ions). This way, the compound remains electrically neutral.

For example, when cesium (\(\mathrm{Cs}^{+}\)) combines with sulfate (\(\mathrm{SO}_{4}^{2-}\)), the charge (positive and negative) determines how the ions join together. Since sulfate has a -2 charge, two cesium ions, each with a +1 charge, must be used to balance the charge, resulting in the formula for cesium sulfate (\(\mathrm{Cs}_{2}\mathrm{SO}_{4}\)). Charge balance ensures the compound is stable and serves as a key principle in determining the empirical formulas of ionic compounds.

Compound Naming

Naming ionic compounds involves straightforward rules. We begin by naming the cation (metal), followed by the anion (non-metal) with an "-ide" suffix or its given polyatomic ion name. For instance, with rubidium iodide (\(\mathrm{RbI}\)), the cation is rubidium, and the anion is iodide. Naming conventions are vital as they provide a universal language for chemists to identify and communicate the chemical constituents easily.

• If the anion is a single element, modify its ending to "-ide" (e.g., iodide from iodine).
• For polyatomic ions like sulfate (\(\mathrm{SO}_{4}^{2-}\)), use their standard names.

By following these naming conventions, one can systematically name compounds such as strontium nitride (\(\mathrm{Sr}_{3}\mathrm{N}_{2}\)) or aluminum sulfide (\(\mathrm{Al}_{2}\mathrm{S}_{3}\)), revealing their ionic nature and composition.

Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded together, acting as a single charged unit. These ions tend to have special names and consistent charges. Knowing these standard names makes it easier to write and name complex ionic compounds.

For example, sulfate (\(\mathrm{SO}_{4}^{2-}\)) is a common polyatomic ion. Unlike monatomic ions (consisting of a single atom), polyatomic ions do not alter their naming pattern when forming compounds. When cesium (\(\mathrm{Cs}^{+}\)) joins sulfate, the result is cesium sulfate (\(\mathrm{Cs}_{2}\mathrm{SO}_{4}\)).

• Polyatomic ions often end in "-ate" or "-ite" depending on the oxidation state of oxygen present in the ion.
• Convey a single charge for the entire group, not per atom.

Understanding polyatomic ions is crucial for deciphering chemical formulas and forming accurate names for many ionic compounds.