Question

If \(\mathrm{CO}_{2}\) is passed in excess in to lime water, the milkines first formed disappears due to: (a) Formation of water soluble calcium bicarbonate. (b) The solution getting heated by exothermic reation. (c) Reversal of the original reaction. (d) Formation of volatile calcium derivative.

Short answer

(a) Formation of water soluble calcium bicarbonate.

Step-by-step solution

Understanding the Problem

We need to find out why the milkiness disappears when excess \(CO_2\) is added to lime water. This problem involves understanding chemical reactions between \(CO_2\) and lime water.

Identifying Initial Reaction with CO2

When \( CO_2 \) is passed into lime water, calcium hydroxide (Ca(OH)_2), a reaction occurs producing calcium carbonate (CaCO_3), which is insoluble in water and causes the solution to appear milky: \[ Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O \]

Reaction with Excess CO2

With the continued addition of \( CO_2 \), the initially formed calcium carbonate (CaCO_3) further reacts to form calcium bicarbonate (Ca(HCO_3)_2), which is soluble in water and therefore, the milkiness disappears: \[ CaCO_3 + CO_2 + H_2O \rightarrow Ca(HCO_3)_2 \]

Confirming the Answer

This dissolution of CaCO_3 in excess \( CO_2 \) and water, forming a water-soluble compound (calcium bicarbonate), confirms that option (a) 'Formation of water soluble calcium bicarbonate' is correct.

Key concepts

Calcium Carbonate

Calcium carbonate is a chemical compound with the formula \( \mathrm{CaCO}_3 \). It is a common substance found in rocks as the minerals calcite and aragonite. This compound is also the main component of pearls and the shells of marine organisms. In the context of lime water, calcium carbonate plays a crucial role when lime water reacts with carbon dioxide. Initially, when \( \mathrm{CO}_2 \) is added to lime water, calcium hydroxide \( \mathrm{Ca(OH)_2} \) reacts to form calcium carbonate, which is a white, insoluble solid. This insolubility is what causes the solution to turn milky. However, this milkiness is temporary under the condition of prolonged \( \mathrm{CO}_2 \) exposure. The transformation of calcium carbonate illustrates the dynamic nature of chemical interactions, and how the environment, like the amount of carbon dioxide, can influence the reaction outcomes.

Calcium Bicarbonate

Calcium bicarbonate \( \mathrm{Ca(HCO_3)_2} \) is a soluble compound formed when calcium carbonate reacts further with carbon dioxide and water. This is indeed what occurs following the initial formation of calcium carbonate in the lime water experiment. The increase in carbon dioxide levels in the solution causes the calcium carbonate to transition into calcium bicarbonate, which dissolves in water thereby removing the milkiness of the solution. The formation of calcium bicarbonate is a reversible reaction, which is important in natural processes such as the formation of stalactites and stalagmites in caves. Because it is soluble, calcium bicarbonate can shift between states depending on temperature and carbon dioxide concentration. This solubility is a key factor that distinguishes it from its counterpart, calcium carbonate, showcasing the intricate balance and dependency within chemical reactions.

Carbon Dioxide Reaction

The reaction of carbon dioxide with lime water is an important demonstration of acid-base chemistry. When \( \mathrm{CO}_2 \) is bubbled into lime water (which is a common term for a diluted solution of calcium hydroxide \( \mathrm{Ca(OH)_2} \)), several reactions take place. Initially, carbon dioxide reacts with calcium hydroxide to form

• Calcium carbonate \( \mathrm{CaCO_3} \)
• water \( \mathrm{H_2O} \)

This reaction gives the solution a milky appearance. However, when excess \( \mathrm{CO}_2 \) is present, the situation changes.

• Calcium carbonate further reacts with \( \mathrm{CO}_2 \) and \( \mathrm{H_2O} \) to form calcium bicarbonate \( \mathrm{Ca(HCO_3)_2} \), which is soluble in water, thus clarifying the solution.

This illustrative experiment helps in understanding how chemical equilibria works, and how even stable compounds like calcium carbonate can undergo further reactions under certain conditions. This dynamic process and its environmental implications, like CO2's role in ocean acidity and biological calcification processes, are crucial in science as well as real-world environmental scenarios.