Write the name from the formula or the formula from the name for each hydrated ionic compound. a. cobalt(II) phosphate octahydrate b. \(\mathrm{BeCl}_{2} \cdot 2 \mathrm{H}_{2} \mathrm{O}\) c. chromium(III) phosphate trihydrate d. \(\mathrm{LiNO}_{2} \cdot \mathrm{H}_{2} \mathrm{O}\)

Understanding the names of ionic compounds is critical for effective communication in chemistry. To begin with, an ionic compound is formed when positive and negative ions attract each other, creating a neutral compound. The cation, usually a metal, is named first followed by the anion, which can be a nonmetal or a polyatomic ion.

When naming ionic compounds with transition metals like cobalt(II), the Roman numeral indicates the charge on the cation, essential for determining the correct chemical formula. For polyatomic ions, such as phosphate (\(\mathrm{PO}_4^{3-}\)), the name stays the same, regardless of the charge.

Hydration numbers are denoted by prefixes such as ‘mono’ for one, ‘di’ for two, 'tri' for three, up to ‘octa’ for eight, followed by the word 'hydrate'. So, 'cobalt(II) phosphate octahydrate' tells us the type of cobalt ion, the polyatomic anion involved, and the number of water molecules associated with the compound.

Writing the chemical formula for an ionic compound involves a clear understanding of the charges on the ions and how they balance. Take the formula (\(\mathrm{BeCl}_2 \cdot 2\mathrm{H}_2\mathrm{O}\)) for example. Here, beryllium has a charge of +2 and chlorine has a charge of -1. Since the compound must be neutral, one Be^2+ ion pairs with two Cl^- ions. The hydrated part, 'dihydrate', indicates two water molecules, represented by '2 H_2 O'.

In chemical formula writing, it's important to consider subscripts to indicate the number of each ion and to use parentheses to denote the multiple of polyatomic ions, as seen in the compound (\(\mathrm{CrPO}_4\cdot 3 \mathrm{H}_2\mathrm{O}\)) for chromium(III) phosphate trihydrate.

Polyatomic ions are charged entities composed of two or more atoms covalently bonded together. They behave as a single unit in the formation of an ionic compound. Common polyatomic ions include (\(\mathrm{NH}_4^+\)) (ammonium), (\(\mathrm{NO}_3^-\)) (nitrate), and the previously mentioned (\(\mathrm{PO}_4^{3-}\)) (phosphate).

When chemicals like chromium(III) phosphate trihydrate and cobalt(II) phosphate octahydrate are formed, their formulas reflect the need to balance the overall charge by incorporating the appropriate number of polyatomic ions. The use of parentheses in the formulas around the polyatomic ion, as in (\(\mathrm{Co}_3(\mathrm{PO}_4)_2\)), is crucial when multiple polyatomic ions are required to balance the charge.

Balancing ionic charges is the cornerstone of writing correct formulas for ionic compounds. The total positive charge must equal the total negative charge to achieve a neutral compound. For instance, in lithium nitrite, the lithium ion has a charge of +1, and the nitrite ion has a charge of -1. They combine in a one-to-one ratio to form a neutral compound.

In cases involving transition metals or polyatomic ions, charge balancing can become more complex. With cobalt(II) phosphate, a 3 to 2 ratio was necessary because cobalt was +2 and phosphate was -3. Understanding how to apply the 'crisscross' method to deduce the required subscript numbers or how to adjust the number of ions is essential, ensuring the total charges from both ions cancel each other out perfectly.