Question

Separate samples of a solution of an unknown ionic compound are treated with dilute \(\mathrm{AgNO}_{3}, \mathrm{~Pb}\left(\mathrm{NO}_{3}\right)_{2},\) and \(\mathrm{BaCl}_{2}\). Precipitates form in all three cases. Which of the following could be the anion of the unknown salt: \(\mathrm{Br}^{-}, \mathrm{CO}_{3}^{2-}, \mathrm{NO}_{3}^{-}\) ?

Short answer

The anion of the unknown salt is \(\mathrm{CO}_{3}^{2-}\), as it is the only anion that forms insoluble products (precipitates) when reacted with all three reagents: \(\mathrm{AgNO}_{3}\), \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\), and \(\mathrm{BaCl}_{2}\).

Step-by-step solution

Write the general equations

To begin, write the general equations for the reactions of each reagent with each possible anion: 1. \(\mathrm{Ag}^{+} + \mathrm{X}^- \rightarrow \mathrm{AgX}\) (with \(\mathrm{AgNO}_{3}\)) 2. \(\mathrm{Pb}^{2+} + \mathrm{X}^{2-} \rightarrow \mathrm{PbX}\) (with \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{₂}\)) 3. \(\mathrm{Ba}^{2+} + 2\mathrm{X}^{-} \rightarrow \mathrm{BaX}_{2}\) (with \(\mathrm{BaCl}_{₂}\)) Next, we need to determine if the products will precipitate.

Check precipitation with reagent

To check whether a formed product will precipitate in each reaction, we can refer to the solubility rules. If a product is insoluble, it should form a precipitate. We will go through each possible anion and check if it forms an insoluble product with the given reagents: - With \(\mathrm{Br}^{-}\): 1. \(\mathrm{AgBr}\): insoluble (precipitate) 2. \(\mathrm{PbBr}_{2}\): insoluble (precipitate) 3. \(\mathrm{BaBr}_{2}\): soluble (no precipitate) - With \(\mathrm{CO}_{3}^{2-}\): 1. \(\mathrm{Ag}_{2}\mathrm{CO}_{3}\): insoluble (precipitate) 2. \(\mathrm{PbCO}_{3}\): insoluble (precipitate) 3. \(\mathrm{BaCO}_{3}\): insoluble (precipitate) - With \(\mathrm{NO}_{3}^{-}\): 1. \(\mathrm{AgNO}_{3}\): soluble (no precipitate) 2. \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\): soluble (no precipitate) 3. \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\): soluble (no precipitate)

Determine the anion of the unknown compound

Based on the precipitation reactions, the only possible anion that forms a precipitate in all three cases is \(\mathrm{CO}_{3}^{2-}\): - With \(\mathrm{Br}^{-}\), \(\mathrm{BaBr}_{2}\) did not form a precipitate. - With \(\mathrm{NO}_{3}^{-}\), all the products were soluble and no precipitate formed. - With \(\mathrm{CO}_{3}^{2-}\), all products were insoluble and precipitates formed in all cases. So, the anion of the unknown salt is \(\mathrm{CO}_{3}^{2-}\).

Key concepts

Precipitation Reactions

Understanding precipitation reactions is crucial for grasping how certain ionic compounds behave in solutions. A precipitation reaction occurs when two soluble salts in aqueous solutions combine and form a product, an insoluble salt, which settles out of the solution as a solid that we call a precipitate.

These reactions are guided by the principle that not all ionic compounds remain soluble in water. When performing experiments or working with chemical equations, you can predict the formation of a precipitate by considering the solubility rules. For the student exploring chemistry exercises, recognizing these reactions requires familiarity with the possible outcomes when different ions interact.

For example, if you mix silver nitrate (AgNO3) with hydrochloric acid (HCl), the silver (Ag+) ions from the AgNO3 and the chloride (Cl-) ions from HCl form silver chloride (AgCl), which is an insoluble salt that appears as a white precipitate. Thus, this is a classic example of a precipitation reaction.

Ionic Compounds

An ionic compound is made up of positively charged ions called cations and negatively charged ions called anions. These ions are held together by ionic bonds – the electrostatic forces of attraction between oppositely charged particles. In an aqueous solution, these compounds often dissociate into their constituent ions.

When learning about ionic compounds, one must differentiate between those that dissolve in water (soluble salts) and those that do not (insoluble salts). It is essential to know that the solubility of ionic compounds varies and is dictated by the nature of their ions. Common cations include metal ions like sodium (Na+) and magnesium (Mg2+), while common anions include chloride (Cl-) and sulfate (SO42-).

For example, sodium chloride (NaCl) is a soluble ionic compound that dissociates entirely into Na+ and Cl- ions in an aqueous solution. Conversely, compounds like lead iodide (PbI2) are poorly soluble in water and are prone to forming a precipitate, especially when the concentration of ions exceeds their solubility limits.

Insoluble Salts

The concept of insoluble salts ties directly into the study of precipitation reactions. Insoluble salts are ionic compounds that have very low solubility in water; they do not dissolve significantly and tend to form a precipitate. The term 'insoluble' is somewhat relative since most so-called insoluble compounds are slightly soluble to some extent.

To predict whether a salt is likely to be insoluble, one can refer to a set of empirical guidelines known as the solubility rules. These rules suggest, for instance, that carbonates (CO32-) and phosphates (PO43-) are generally insoluble except when paired with ammonium (NH4+) or alkali metal cations (such as Na+ or K+).

Referring to the original exercise, we see that silver carbonate (Ag2CO3) and lead carbonate (PbCO3) are examples of insoluble salts that form as precipitates. Understanding these principles helps students to identify the outcome of mixing different ionic compounds in solution and reinforces the importance of solubility rules for predicting and explaining chemical behavior.