Question

Write a chemical equation that shows how limestone neutralizes acid rain.

Short answer

CaCO₃ + H₂SO₄ → CaSO₄ + CO₂ + H₂O

Step-by-step solution

Understand the Problem

Limestone is primarily composed of calcium carbonate (CaCO₃). Acid rain typically contains sulfuric acid (H₂SO₄) along with other acids such as nitric acid (HNO₃). To show how limestone neutralizes acid rain, we need to write a chemical equation demonstrating the reaction between calcium carbonate and one of these acids.

Write the Reaction Equation

Calcium carbonate reacts with sulfuric acid in the following way: \[ \text{CaCO}_{3 (s)} + \text{H}_{2}\text{SO}_{4 (aq)} \rightarrow \text{CaSO}_{4 (s)} + \text{H}_{2}\text{CO}_{3 (aq)} \]Here, solid calcium carbonate reacts with aqueous sulfuric acid to produce solid calcium sulfate and aqueous carbonic acid.

Simplify the Products

The carbonic acid (H₂CO₃) in the solution breaks down further into carbon dioxide and water due to its instability: \[ \text{H}_{2}\text{CO}_{3 (aq)} \rightarrow \text{CO}_{2 (g)} + \text{H}_{2}\text{O}_{(l)} \]Thus, the complete chemical reaction occurring when limestone neutralizes acid rain is:\[ \text{CaCO}_{3(s)} + \text{H}_{2}\text{SO}_{4(aq)} \rightarrow \text{CaSO}_{4(s)} + \text{CO}_{2(g)} + \text{H}_{2}\text{O}_{(l)} \]

Key concepts

Calcium Carbonate Reaction

Calcium carbonate (\( \text{CaCO}_3 \)) is a common, naturally occurring compound we often encounter in various forms, such as limestone and marble. Its ability to neutralize acids makes it valuable in many contexts. When it interacts with an acid like sulfuric acid, it neutralizes the acidity.
In a typical reaction with sulfuric acid (\( \text{H}_2\text{SO}_4 \)), calcium carbonate acts as a base. It reacts to form calcium sulfate (\( \text{CaSO}_4 \)), water (\( \text{H}_2\text{O} \)), and carbon dioxide (\( \text{CO}_2 \)). This is a fundamental way to neutralize harmful acids in the environment. This chemical reaction can be represented as:

• \[ \text{CaCO}_{3(s)} + \text{H}_{2}\text{SO}_{4(aq)} \rightarrow \text{CaSO}_{4(s)} + \text{CO}_{2(g)} + \text{H}_{2}\text{O}_{(l)} \]

This equation shows that the solid limestone dissolves to form calcium sulfate, a solid, while releasing carbon dioxide gas and forming water. Such reactions help reduce the harmful effects of acid rain on natural environments and related structures.

Sulfuric Acid

Sulfuric acid (\( \text{H}_2\text{SO}_4 \)) is a strong, corrosive acid that's prevalent in acid rain. It forms mainly from the chemical reaction of sulfur dioxide (\( \text{SO}_2 \)) in the atmosphere when fossil fuels are burned.
Once released into the environment, it poses significant harm due to its ability to lower the pH of water bodies and soil, which affects ecosystems adversely. When it comes into contact with materials like limestone or marble, it reacts with the calcium carbonate present.
This interactive property of sulfuric acid is crucial in understanding how limestone can help mitigate environmental damage. Sulfuric acid reacts with calcium carbonate to produce calcium sulfate, which is less harmful, along with \( \text{CO}_2 \) and \( \text{H}_2\text{O} \), effectively neutralizing its acidity.

Chemical Equations

Chemical equations are a concise way to depict a chemical reaction using symbols and formulas. They tell us about the substances involved, their amounts, and how they transform during the reaction.
In the limestone and sulfuric acid interaction, the chemical equation is:

• \[ \text{CaCO}_{3(s)} + \text{H}_{2}\text{SO}_{4(aq)} \rightarrow \text{CaSO}_{4(s)} + \text{CO}_{2(g)} + \text{H}_{2}\text{O}_{(l)} \]

The left side shows the reactants, calcium carbonate and sulfuric acid, and the right side lists the products, calcium sulfate, carbon dioxide, and water.
The arrow indicates the direction of the reaction. Chemical equations allow us to understand at a glance what happens in a chemical process. Breakdown of each part, such as the states of compounds—solid \((s)\), aqueous \((aq)\), gas \((g)\), and liquid \((l)\)—gives us insights into what we can expect from this interaction, both in laboratory settings and real-world scenarios.