Understanding Polyatomic Anions and Electronegativity in Chemistry

Example: Sulfur Oxyanions Consider the relationship between sulfate and sulfite:

• $\text{SO}_4^{2-}$ is sulfate. Here, the sulfur atom is in a $+6$ oxidation state.
• $\text{SO}_3^{2-}$ is sulfite. Here, the sulfur atom is in a $+4$ oxidation state.

Exam Tip: When you see a name ending in "-ate," think "more oxygen." When you see "-ite," think "less oxygen." This helps you quickly deduce the chemical formula if you know the base ion.

Extended Oxyanion Families

Some elements belong to an "extended family" of oxyanions, where the number of oxygen atoms can vary significantly. These families use specific prefixes to denote the extremes of the oxygen count:

1. per- ... -ate: The highest number of oxygen atoms (highest oxidation state).
2. -ate: The standard high oxygen count.
3. -ite: A lower oxygen count.
4. hypo- ... -ite: The fewest oxygen atoms (lowest oxidation state).

Example: The Chlorine Family The chlorine atom can form several distinct oxyanions:

• $\text{ClO}_4^-$: perchlorate (highest oxidation state, $\text{Cl} = +7$)
• $\text{ClO}_3^-$: chlorate
• $\text{ClO}_2^-$: chlorite
• $\text{ClO}^-$: hypochlorite (lowest oxidation state, $\text{Cl} = +1$)

From Anions to Oxyacids and Hydroxides

The naming rules for polyatomic anions extend directly into the naming of acids and hydroxides.

Oxyacids

Oxyacids are acids that contain hydrogen, oxygen, and one other element. To name them, you omit the "hydrogen" part of the name and modify the anion's suffix:

• If the anion ends in -ate, the acid ends in -ic (e.g., sulfate $\rightarrow$ sulfuric acid).
• If the anion ends in -ite, the acid ends in -ous (e.g., sulfite $\rightarrow$ sulfurous acid).

Hydroxides

Hydroxides are compounds containing the $\text{OH}^-$ ion. When naming ionic compounds involving transition metals (which can have multiple oxidation states), we use the Stock System. This involves naming the metal, followed by its oxidation state in Roman numerals, and then the word "hydroxide".

Examples:

• $\text{NaOH}$: sodium hydroxide
• $\text{Fe(OH)}_2$: iron(II) hydroxide
• $\text{Fe(OH)}_3$: iron(III) hydroxide

Electronegativity and Bonding

To understand why these atoms bond the way they do, we must look at electronegativity. Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond.

The difference in electronegativity between two atoms determines the type of bond formed:

• Ionic Bonds: Occur when there is a large difference in electronegativity, typically between a metal and a nonmetal. Electrons are transferred from the metal (forming a cation) to the nonmetal (forming an anion).
• Covalent Bonds: Occur when atoms share electrons, typically between two nonmetals. This is the type of bonding found within polyatomic ions themselves.

In a periodic table context, the most electronegative elements are generally found toward the top and right, while the most metallic (least electronegative) elements are toward the bottom and left.

Practical Application: Formula Writing

When writing formulas for ionic compounds containing polyatomic ions, you must ensure the total positive charge from the cations balances the total negative charge from the anions to reach a net charge of zero.

Crucial Rule: The Use of Parentheses If you need more than one of a polyatomic ion to balance the charge, you must use parentheses to indicate the quantity.

Worked Example: Write the formula for a compound made of Magnesium ($\text{Mg}^{2+}$) and Nitrate ($\text{NO}_3^-$).

1. The total positive charge is $+2$.
2. The total negative charge from one nitrate is $-1$.
3. To balance the charges, we need two nitrate ions: $\text{Mg}^{2+} + 2(\text{NO}_3^-) \rightarrow \text{Mg}(\text{NO}_3)_2$

Without the parentheses, $\text{MgNO}_3$ would incorrectly imply there is only one oxygen atom, which changes the identity of the ion entirely.