Question

All but two of the following solutions yield a precipitate when the solution is also made \(2.00 \mathrm{M}\) in \(\mathrm{NH}_{3}\). Those two are (a) \(\mathrm{MgCl}_{2}(\mathrm{aq}) ;\) (b) \(\mathrm{FeCl}_{3}(\mathrm{aq})\); (c) \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}(\mathrm{aq}) ;(\mathrm{d}) \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})\); (e) \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}(\mathrm{aq})\).

Short answer

The two solutions that will not form a precipitate when the solution is made 2.00M in \(NH_3\) are: (c) \((NH_4)_2SO_4\) and (d) \(Cu(NO_3)_2\).

Step-by-step solution

Identify the possible precipitates

Using knowledge of solubility rules, the possible salts that can form a precipitate were identified. These are: 1. \(MgCl_2\)2. \(FeCl_3\)3. \(Cu(NO_3)_2\)4. \(Al_2(SO_4)_3\)

Identify the soluble salts in ammonia

The rules of solubility show that most nitrate (NO3) salts and most ammonium (NH4) salts are soluble. Therefore, \(NH_4)_2SO_4\) and \(Cu(NO_3)_2\) will not form a precipitate when the solution is made 2.00M in \(NH_3\)

Key concepts

Precipitate Formation

Precipitate formation occurs when two soluble substances react in solution to form an insoluble substance, called a precipitate. It's like mixing two clear liquids and ending up with a cloudy or solid substance, which then falls to the bottom of the container. In chemistry, this usually happens when ions in the solutions react to form a compound that is no longer soluble in the given solution.

When predicting whether a precipitate will form, chemists rely on a set of guidelines known as solubility rules. These rules help determine if a certain combination of ions will remain dissolved or will form a solid. For instance, when a solution of magnesium chloride ( MgCl_2 ) or iron(III) chloride ( FeCl_3 ) is mixed with ammonia, certain ion combinations may result in new compounds that are insoluble in water, such as hydroxides or other compounds.

• To identify potential precipitates, you often must consider the solubility of the resulting compounds.
• A compound forms a precipitate if it is considered insoluble or only slightly soluble in the solution.
• Common types of precipitates include hydroxides, sulfates, and carbonates.

Nitrate Solubility

Nitrate ions ( NO_3^- ) are renowned for their high solubility in water. This is an important rule in chemistry: almost all nitrate salts are soluble, making them very unlikely to precipitate under normal circumstances. The reason for this high solubility stems from the strong interactions between nitrate ions and water molecules.

For example, copper(II) nitrate ( Cu(NO_3)_2 ) remains soluble even when the solution is made 2.00M in ammonia. This behavior fits within the solubility rules, which state that adding ammonia won't affect the solubility of nitrate salts significantly.

• This inherent solubility is due to the effective interaction with water, which keeps the nitrate ions dispersed throughout the solution.
• Knowing that nitrate salts are generally soluble helps predict that they will not form a precipitate.
• This makes separating or isolating specific compounds easier in processes where solubility manipulation is required.

Ammonium Solubility

Ammonium ions ( NH_4^+ ) share a similar fate with nitrates concerning solubility in water. Most ammonium salts are soluble, making them predictable non-precipitates in many chemical reactions. The ammonia molecule is naturally soluble, and when it forms NH_4^+ ions in water, these ions typically remain in solution.

Consider ammonium sulfate ((NH_4)_2SO_4) . This compound remains in solution without forming a precipitate under high ammonia conditions due to the strong solubility of NH_4^+ . This behavior is crucial in chemistry, as it allows NH_4^+ ions to remain in solution for various reactions and processes.

• Solubility rules indicate most ammonium compounds won't precipitate, providing a basis for predictability in chemical processes.
• This characteristic is useful in reactions requiring dissolved ions without unwanted solid formation.
• Knowing when and why ammonium salts stay in solution helps chemists plan and execute reaction mechanisms effectively.