Question

Describe in each case how you would separate the cations or anions in the following aqueous solutions: (a) \(\mathrm{NaNO}_{3}\) and \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) (b) \(\mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}\) and \(\mathrm{K} \mathrm{NO}_{3},\) (c) \(\mathrm{KBr}\) and \(\mathrm{KNO}_{3},\) (d) \(\mathrm{K}_{3} \mathrm{PO}_{4}\) and \(\mathrm{KNO}_{3},\) (e) \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) and \(\mathrm{NaNO}_{3}\)

Short answer

Use selective precipitation reactions to separate ions based on solubility differences.

Step-by-step solution

Separate Ba2+ from Na+

In solution (a), \(\text{NaNO}_{3}\) contains sodium ions \(\text{Na}^+\) and \({\text{Ba}({\text{NO}_3})_2}\) contains barium ions \({\text{Ba}^{2+}}\). Add a sulfate solution, such as \(\text{H}_2\text{SO}_4\), to the mixture. \({\text{BaSO}_4}\), a white precipitate, will form because it is insoluble in water. The sodium ions \(\text{Na}^{+}\) will remain in solution because sodium sulfate is soluble.

Separate Mg2+ from K+

In solution (b), \ ext{Mg}({\text{NO}_3})_2\ and \(\text{KNO}_3\), add a carbonate solution, such as \(\text{Na}_2\text{CO}_3\). \(\text{MgCO}_3\) is insoluble and will form a precipitate, while \(\text{K}_2\text{CO}_3\) remains soluble, allowing \(\text{K}^+\) ions to stay in solution.

Separate Br- from NO3-

In solution (c), \(\text{KBr}\) and \(\text{KNO}_3\), use silver nitrate \(\text{AgNO}_3\) as a reagent. A white precipitate of \(\text{AgBr}\) forms, while \(\text{NO}_3^-\) ions remain in solution because \(\text{AgNO}_3\) is soluble.

Separate PO4^3- from NO3-

In solution (d), \(\text{K}_3\text{PO}_4\) and \(\text{KNO}_3\), adding calcium chloride \(\text{CaCl}_2\) will cause \(\text{Ca}_3(\text{PO}_4)_2\) to form a precipitate due to its low solubility, while the nitrates remain soluble.

Separate CO3^2- from NO3-

In solution (e), \(\text{Na}_2\text{CO}_3\) and \(\text{NaNO}_3\), add a strong acid such as hydrochloric acid \(\text{HCl}\). This will convert \(\text{CO}_3^{2-}\) to \(\text{CO}_2\) gas, which escapes as bubbles, while \(\text{NO}_3^{-}\) ions remain in solution.

Key concepts

Cation Separation

In chemistry, cation separation is a technique used to distinguish between different metal ions present in a solution. This method is crucial because certain reactions only occur with specific metal ions, making it necessary to isolate one from another. A typical example from our exercise involves separating sodium ions ( Na^+ ) from barium ions ( Ba^{2+} ). By adding a sulfate solution, such as sulfuric acid ( H_2SO_4 ), barium sulfate ( BaSO_4 ), an insoluble compound, precipitates out of the solution. This reaction does not affect sodium ions, as sodium sulfate is soluble in water and remains dissolved. The key steps in cation separation include: - Selecting the appropriate reagent that will react specifically with one cation to form an insoluble compound. - Ensuring the remaining cations stay dissolved and do not form precipitates. This selective precipitation is guided by solubility rules, allowing chemists to effectively separate ions based on their differing chemical behaviors.

Anion Separation

Anion separation focuses on isolating different negatively charged ions in a solution. It's vital for detecting and analyzing various compounds containing similar anions. Observing this concept in our exercise, potassium bromide ( KBr ) is mixed with potassium nitrate ( KNO_3 ). To separate bromide ions ( Br^- ) from nitrate ions ( NO_3^- ), silver nitrate ( AgNO_3 ) is added to the solution. This causes a reaction where silver bromide ( AgBr ), which is insoluble, forms a precipitate, while NO_3^- ions remain in solution because silver nitrate is soluble. Key points to remember for anion separation are: - Choosing a reagent that will selectively precipitate a particular anion. - Ensuring other anions do not react similarly and remain in the solution. Such separation techniques are particularly useful in qualitative analysis, aiding in the identification of specific anions within a mixture.

Precipitation Reactions

The heart of separation techniques in chemistry often revolves around precipitation reactions. These reactions occur when two solutions combine to form an insoluble solid, known as a precipitate. In our exercise, numerous examples illustrate this concept, such as the formation of barium sulfate ( BaSO_4 ) when barium ions react with sulfate ions. Here's what makes precipitation reactions so important: - They provide a visual indication of a chemical reaction through the formation of a solid. - Precipitation allows for the easy separation of components due to the differential solubility of compounds. To effectively use precipitation reactions for separation: - Utilize solubility rules to predict the solubility of different compounds in solution. - Carefully choose reactants that will induce precipitation without unwanted side reactions. These reactions are fundamental in fields such as environmental science for removing unwanted ions from water and in analytical chemistry for identifying chemical substances.

Solubility Rules

Solubility rules are a set of guidelines that help predict whether an ionic compound will dissolve or form a precipitate in water. Understanding these rules allows chemists to anticipate and control the outcome of mixing various solutions. They were instrumental in solving the exercise problems where soluble and insoluble compounds were identified to allow separation of ions. Consider the solubility rules applied in the given examples: - Most sulfates are soluble, except for those of barium, lead, and calcium, allowing barium sulfate to precipitate while others stay dissolved. - Most carbonate compounds are insoluble, except those with alkali metals, leading to magnesium carbonate precipitating while potassium carbonate stays dissolved. These rules help decide on the correct reagents for reactions and can predict the behavior of ions in a mixture. Basic solubility guidelines include:

• Nitrates, acetates, and most alkali metal salts are typically soluble.
• Chlorides, bromides, and iodides are generally soluble, except for those formed with silver, lead, and mercury.
• Carbonates, phosphates, sulfides, and hydroxides are mostly insoluble except when paired with alkali metals or ammonium.

Chemists rely on these rules for designing effective step-by-step separation processes in the lab.