Question

lime \((\mathrm{CaO})\) is one of the most abundantly produced chemicals in the world. Write balanced equations for these reactions: (a) The preparation of lime from natural sources (b) The use of slaked lime to remove \(\mathrm{SO}_{2}\) from flue gases (c) The reaction of lime with arsenic acid \(\left(\mathrm{H}_{3} \mathrm{~A} \mathrm{~s} \mathrm{O}_{4}\right)\) to manufacture the insecticide calcium arsenate (d) The regeneration of \(\mathrm{NaOH}\) in the paper industry by reaction of lime with aqueous sodium carbonate

Short answer

(a) \text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2, (b) \text{Ca(OH)}_2 + \text{SO}_2 \rightarrow \text{CaSO}_3 + \text{H}_2\text{O}, (c) 3\text{CaO} + 2\text{H}_3\text{AsO}_4 \rightarrow \text{Ca_3(AsO_4)_2} + 3\text{H}_2\text{O}, (d) \text{CaO} + \text{Na}_2\text{CO}_3 + \text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{CaCO}_3.

Step-by-step solution

Preparation of Lime

Lime (\text{CaO}) is prepared by heating calcium carbonate (\text{CaCO}_3) from natural sources such as limestone. The balanced equation for this reaction is:\[ \text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2 \]

Use of Slaked Lime to Remove \text{SO}_2 from Flue Gases

Slaked lime (\text{Ca(OH)}_2) reacts with sulfur dioxide (\text{SO}_2) present in flue gases to form calcium sulfite (\text{CaSO}_3) and water (\text{H}_2\text{O}). The balanced equation for this reaction is:\[ \text{Ca(OH)}_2 + \text{SO}_2 \rightarrow \text{CaSO}_3 + \text{H}_2\text{O} \]

Reaction of Lime with Arsenic Acid

Lime (\text{CaO}) reacts with arsenic acid (\text{H}_3\text{AsO}_4) to form calcium arsenate (\text{Ca_3(AsO_4)_2}), which is used as an insecticide. The balanced equation for this reaction is:\[ 3\text{CaO} + 2\text{H}_3\text{AsO}_4 \rightarrow \text{Ca_3(AsO_4)_2} + 3\text{H}_2\text{O} \]

Regeneration of \text{NaOH} in the Paper Industry

Lime (\text{CaO}) reacts with aqueous sodium carbonate (\text{Na}_2\text{CO}_3) to regenerate sodium hydroxide (\text{NaOH}) in the paper industry. The balanced equation for this reaction is:\[ \text{CaO} + \text{Na}_2\text{CO}_3 + \text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{CaCO}_3 \]

Key concepts

Lime Preparation

Lime, also known as calcium oxide (CaO), is a fundamental chemical in various industries. Its preparation primarily involves heating natural limestone, which is composed largely of calcium carbonate (CaCO₃). When limestone is subjected to high temperatures in a kiln, it decomposes into lime (CaO) and carbon dioxide (CO₂). This process can be summarized by the chemical equation:
\[ \text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2 \]
This reaction is fundamental in industry and is known as calcination. The high temperature required for this process often necessitates energy-intensive methods, making it costly. However, the significance of lime in various applications justifies these energy expenditures.

Understanding the production of lime is crucial, as it's a building block for many other chemical processes.

Sulfur Dioxide Removal

Sulfur dioxide (SO₂) is a harmful pollutant commonly found in flue gases emitted from industrial settings like power plants. To mitigate its environmental impact, industries use slaked lime (Ca(OH)₂) for its removal.
When slaked lime reacts with sulfur dioxide, it results in the formation of calcium sulfite (CaSO₃) and water (H₂O). The balanced chemical equation for this process is:
\[ \text{Ca(OH)}_2 + \text{SO}_2 \rightarrow \text{CaSO}_3 + \text{H}_2\text{O} \]
This reaction not only helps in reducing the SO₂ emissions but also produces a useful byproduct, calcium sulfite, which can be further processed. Removing sulfur dioxide effectively reduces acid rain and air-quality issues.

Studying this process is critical for those wanting to understand environmental protection mechanisms associated with industrial activities.

Calcium Arsenate Production

Calcium arsenate (Ca₃(AsO₄)₂) is an insecticide formed through the reaction of lime (CaO) with arsenic acid (H₃AsO₄). This reaction is significant in the agricultural industry due to its pest control benefits. The balanced equation for this chemical reaction is:
\[ 3\text{CaO} + 2\text{H}_3\text{AsO}_4 \rightarrow \text{Ca}_3(\text{AsO}_4)_2 + 3\text{H}_2\text{O} \]
When lime reacts with arsenic acid, it produces calcium arsenate and water. The importance of this reaction lies in producing effective insecticides that help in protecting crops from various pests. However, safety measures should be taken, given the toxicity associated with arsenic compounds.

Grasping this process is important for students interested in agriculture and chemical production industries.

Sodium Hydroxide Regeneration

In the paper industry, sodium hydroxide (NaOH) is a valuable chemical used for various processes, including paper pulping. It can be regenerated using lime (CaO) and aqueous sodium carbonate (Na₂CO₃). The balanced chemical equation for this reaction is:
\[ \text{CaO} + \text{Na}_2\text{CO}_3 + \text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{CaCO}_3 \]
This reaction highlights the efficiency of reusing chemicals within industrial processes. When lime reacts with sodium carbonate in the presence of water, it forms sodium hydroxide and calcium carbonate. The regenerated NaOH can then be reused, reducing the need for fresh chemicals and lowering overall costs.

Understanding the regeneration of sodium hydroxide is crucial for those interested in sustainable industrial practices and chemical engineering.