Question

Which one of the following statement is correct? (a) Ferric ions give a deep green precipitate on adding potassium ferrocyanide solution (b) From a mixed precipitate of \(\mathrm{AgCl}\) and \(\mathrm{AgI}\), ammonia solution dissolves only \(\mathrm{AgCl}\) (c) Manganese salts give a violet borax bead test in the reducing flame (d) On boiling a solution having \(\mathrm{K}^{+}, \mathrm{Ca}^{2+}\) and \(\mathrm{HCO}_{3}^{-}\) ions we get a precipitate of \(\mathrm{K}_{2} \mathrm{Ca}\left(\mathrm{CO}_{3}\right)_{2}\)

Short answer

Statement (b) is correct; ammonia dissolves only AgCl from AgCl and AgI.

Step-by-step solution

Analyze Statement (a)

Statement (a) suggests that ferric ions should produce a deep green precipitate when treated with potassium ferrocyanide. However, the reaction between ferric ions and potassium ferrocyanide actually produces a Prussian blue precipitate, not a green one. Therefore, statement (a) is incorrect.

Analyze Statement (b)

Statement (b) claims that ammonia can dissolve only AgCl from a mixed precipitate of AgCl and AgI. Ammonia does dissolve AgCl but not AgI because the latter is less soluble in ammonia. Hence, statement (b) is correct.

Analyze Statement (c)

Statement (c) states that manganese salts give a violet color in the borax bead test under reducing flame conditions. However, manganese salts produce a color ranging from colorless to pink/lavender depending on their state of oxidation, not violet. This makes statement (c) incorrect.

Analyze Statement (d)

Statement (d) suggests that boiling a solution containing K⁺, Ca²⁺, and HCO₃⁻ would result in the formation of K₂Ca(CO₃)₂ as a precipitate. Normally, upon heating, calcium bicarbonate decomposes to form calcium carbonate (CaCO₃), which precipitates, not a double salt as suggested. Therefore, statement (d) is incorrect.

Key concepts

Precipitation Reactions

Precipitation reactions occur when two soluble salts in aqueous solution react to form one or more insoluble products, called precipitates. These reactions are important in chemistry for isolating compounds and identifying substances based on their physical properties.
For example, if you mix solutions containing silver ions (\(\text{Ag}^+\)) and chloride ions (\(\text{Cl}^-\)), you will get a white precipitate of silver chloride (\(\text{AgCl}\)). This reaction is an example of a precipitation reaction.
Precipitation reactions can be used in qualitative analysis to identify the presence of specific ions. The formation of the precipitate often depends on the solubility rules of the ions involved.
To predict whether a precipitation reaction will occur, you can consult solubility tables. These tables indicate whether specific salts tend to dissolve or precipitate in water. Keep in mind that the precipitate's color and solubility may vary with changes in temperature or the presence of other ions. In the case of ferric ions with potassium ferrocyanide, a distinct Prussian blue precipitate forms.

Selective Dissolution Using Ammonia

Selective dissolution using ammonia is a technique where ammonia solution is used to dissolve certain metal salts while leaving others undissolved. This process helps in the separation and identification of ions in a mixture.
An important example is the selective dissolution of silver chloride (\(\text{AgCl}\)) and silver iodide (\(\text{AgI}\)). When you add ammonia to a mixture of these two salts, \(\text{AgCl}\) dissolves to form a complex ion, \(\text{Ag(NH}_3)_2^+\), but \(\text{AgI}\) remains undissolved.
Ammonia acts as a complexing agent because it can form coordinate covalent bonds with silver ions, increasing their solubility in the solution. This selective solubility is useful for laboratory analysis, helping chemists to distinguish between ions that may otherwise appear similar.
The underlying principle of selective dissolution is based on the differences in the formation constants (also known as stability constants) of the possible complexes. Silver iodide is much less soluble in ammonia compared to silver chloride, aiding in straightforward separation of these two compounds.

Borax Bead Test

The borax bead test is a qualitative analysis method used to identify metal ions based on the color they impart to a borax bead when heated in a flame. Borax (\(\text{Na}_2\text{B}_4\text{O}_7\)) is known to form glass-like beads that can change color upon exposure to metal salts.
This test is especially useful for detecting different oxidation states of metals, which can result in different colors. For instance, in the reducing flame, manganese salts can produce a color ranging from colorless to pink or lavender, not violet as sometimes claimed.
The test involves looping a small amount of borax on a metal wire and heating it to form a bead. Then, the bead is dipped in the powdered sample to be tested, and reheated in the flame. The color that the bead displays is indicative of the specific metal ions present.

• The bead color can change under oxidizing or reducing conditions, making it highly versatile.
• This method is often used in experimental inorganic chemistry to quickly identify metal components in a mineral or compound.

Thermal Decomposition of Bicarbonates

Thermal decomposition of bicarbonates refers to the breakdown of bicarbonate ions (\(\text{HCO}_3^-\)) upon heating, resulting in simpler compounds. When bicarbonate solutions are heated, they typically decompose to form carbonate ions, water, and carbon dioxide gas.
A common example is calcium bicarbonate [\(\text{Ca(HCO}_3)_2\)], which decomposes when heated to form calcium carbonate (\(\text{CaCO}_3\)), carbon dioxide, and water:
\[ \text{Ca(HCO}_3)_2 \rightarrow \text{CaCO}_3 + \text{CO}_2 + \text{H}_2\text{O} \]
This reaction is crucial in understanding natural processes like the formation of limestone from groundwater solutions and is an important consideration in hard water formation and treatment.

• Typically, bicarbonates of alkali earth metals like calcium or magnesium show noticeable changes when heated, often resulting in precipitation of their carbonates.
• This reaction has practical implications because it explains why heating hard water results in the formation and deposition of scale—an important factor in plumbing and industrial systems.

In this context, it's clear that the incorrect statement suggests both an improbable product and an unrealistic outcome, making it a fictional claim.