Question

When the following solutions are mixed together, what precipitate (if any) will form? a. \(\mathrm{FeSO}_{4}(a q)+\mathrm{KCl}(a q)\) b. \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}(a q)+\mathrm{Ba}(\mathrm{OH})_{2}(a q)\) c. \(\mathrm{CaCl}_{2}(a q)+\mathrm{Na}_{2} \mathrm{SO}_{4}(a q)\) d. \(\mathrm{K}_{2} \mathrm{S}(a q)+\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}(a q)\)

Short answer

In summary, when the given solutions are mixed, the following precipitates will form: a. No precipitate b. Al(OH)\(_{3}\) c. CaSO\(_{4}\) d. NiS

Step-by-step solution

Determine the possible products

For each reaction, we'll be mixing two ionic compounds in aqueous solutions which will undergo a double displacement reaction. Their cations and anions will be exchanged.

Reaction a: \(\mathrm{FeSO}_{4}(a q)+\mathrm{KCl}(a q)\)

Cations: \(\mathrm{Fe}^{2+}, \mathrm{K}^+\) Anions: \(\mathrm{SO}_4^{2-}, \mathrm{Cl}^-\) Possible products: FeCl\(_{2}\) and K\(_{2}\)SO\(_{4}\)

Reaction b: \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}(a q)+\mathrm{Ba}(\mathrm{OH})_{2}(a q)\)

Cations: \(\mathrm{Al}^{3+}, \mathrm{Ba}^{2+}\) Anions: \(\mathrm{NO}_3^-, \mathrm{OH}^-\) Possible products: Al(OH)\(_{3}\) and Ba(NO\(_{3}\))\(_{2}\)

Reaction c: \(\mathrm{CaCl}_{2}(a q)+\mathrm{Na}_{2} \mathrm{SO}_{4}(a q)\)

Cations: \(\mathrm{Ca}^{2+}, \mathrm{Na}^+\) Anions: \(\mathrm{Cl}^-, \mathrm{SO}_4^{2-}\) Possible products: CaSO\(_{4}\) and NaCl

Reaction d: \(\mathrm{K}_{2} \mathrm{S}(a q)+\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}(a q)\)

Cations: \(\mathrm{K}^+, \mathrm{Ni}^{2+}\) Anions: \(\mathrm{S}^{2-}, \mathrm{NO}_3^-\) Possible products: KNO\(_{3}\) and NiS Now, we will check the solubility of these possible products.

Determine the solubility of the products

We'll use the solubility rules to determine if any of the products formed are insoluble and result in the formation of a precipitate. 1. FeCl\(_{2}\) and K\(_{2}\)SO\(_{4}\): Chlorides and sulfates are generally soluble. No precipitate will form in this case. 2. Al(OH)\(_{3}\) and Ba(NO\(_{3}\))\(_{2}\): Hydroxides are generally insoluble except for Group I elements and some Group II elements (Ca, Sr, Ba). Al(OH)\(_{3}\) is insoluble and will form a precipitate. Ba(NO\(_{3}\))\(_{2}\) is soluble. 3. CaSO\(_{4}\) and NaCl: Sulfates are generally soluble, but CaSO\(_{4}\) is an exception and is moderately soluble. In this situation, it will form a precipitate. NaCl is soluble. 4. KNO\(_{3}\) and NiS: Nitrates are soluble. KNO\(_{3}\) is soluble and Sulfides are generally insoluble except for Group I, II, and NH\(_{4}^+\). NiS is insoluble and will form a precipitate.

Summary of precipitates formed

The following are the precipitates formed in each reaction: a. No precipitate b. Al(OH)\(_{3}\) c. CaSO\(_{4}\) d. NiS After mixing the given solutions, the precipitates formed are Al(OH)\(_{3}\), CaSO\(_{4}\), and NiS.

Key concepts

Solubility Rules

When it comes to determining whether a precipitate will form in a reaction, understanding solubility rules is crucial. Solubility rules are guidelines that help predict the solubility of various ionic compounds in water. These rules are generally simplified descriptions of the solubility behavior of ionic compounds. For instance, *nitrates* (\( NO_3^- \)) and *alkali metal ions* like potassium (K\(^+\)) are generally soluble. This means compounds containing these ions will dissolve in water without forming precipitates.

• *Chlorides*, bromides, and iodides are soluble, except those of silver, lead, and mercury.
• *Sulfates* (\( SO_4^{2-} \)) are usually soluble, with exceptions like barium sulfate, lead sulfate, and calcium sulfate, which tend to be insoluble or moderately soluble.
• *Hydroxides* are mostly insoluble except those of sodium, potassium, and calcium to some extent.

These rules are a helpful shortcut for predicting the behavior of compounds in solution and are essential for conducting successful precipitation reactions.

Double Displacement Reactions

Double displacement reactions are a type of chemical reaction where the ions of two compounds exchange partners to form two new compounds. These reactions usually take place in aqueous solutions where ions can move freely. They are also known as *metathesis reactions*.
In a typical double displacement reaction between two ionic compounds, the cations from one reactant combine with the anions of another. For example, when mixing FeSO\(_4\) and KCl in water, the possible products are FeCl\(_2\) and K\(_2\)SO\(_4\), formed by exchanging the cation-anion pairs.
This type of reaction is often favored to form a product that is insoluble, forming a precipitate, or a gas, which escapes from the reaction mixture. Double displacement reactions are frequently used in laboratories to test for the presence of certain ions due to their predictable products and conditions.

Ionic Compounds

Ionic compounds are fascinating and play a significant role in chemistry. They are made up of positively charged ions (*cations*) and negatively charged ions (*anions*). Because they are oppositely charged, they are held together by strong electrostatic forces known as ionic bonds. This structure gives ionic compounds unique properties such as high melting and boiling points and good electrical conductivity when dissolved in water.
When ionic compounds are mixed in an aqueous solution, they dissociate into their respective ions, sharing these ions in the solution and taking part in chemical reactions.

• Example: FeSO\(_4\) dissociates into Fe\(^{2+}\) and SO\(^{2-}_4\).
• KCl breaks into K\(^+\) and Cl\(^-\).

Understanding the behavior of these dissociated ions helps predict the outcome of reactions, especially for determining whether a precipitate forms in reactions like those seen in double displacement reactions.

Aqueous Solutions

Aqueous solutions are an essential concept in chemistry. They refer to solutions where water acts as the solvent. Most chemical reactions, especially those involving ionic compounds, occur in aqueous solutions because the water's polar nature helps dissolve ionic compounds, splitting them into individual ions.
In such solutions, ions move freely, making aqueous solutions excellent conductors of electricity. This characteristic arises because water stabilizes ions by surrounding them and decreasing the force of attraction between oppositely charged ions, which aids in their dissolution.
When chemical reactions occur in these solutions, such as double displacement reactions, the ions recombine into new compounds. If these new compounds are insoluble in water, they precipitate out as a solid. Studying reactions in aqueous solutions allows chemists to explore many reaction pathways and understand the crucial role of solvents in chemical processes.