Problem 19 - Chemistry: The Central Science in SI Units, Global Edition

Question

(a) The hydrogen carbonate ion \(\left(\mathrm{HCO}_{3}^{-}\right)\) is amphiprotic. Write a balanced chemical equation showing how it acts as an acid toward water and another equation showing how it acts as a base toward water.

Step-by-Step Solution

Answer

When the hydrogen carbonate ion, HCO₃⁻, acts as an acid, it donates a proton (H⁺) to a water molecule, resulting in the formation of a hydronium ion (H₃O⁺) and the carbonate ion (CO₃²⁻): \(HCO_3^- + H_2O \rightarrow H_3O^+ + CO_3^{2-}\) When acting as a base, HCO₃⁻ accepts a proton (H⁺) from a water molecule, forming a hydroxide ion (OH⁻) and carbonic acid (H₂CO₃): \(HCO_3^- + H_2O \rightarrow OH^- + H_2CO_3\)

Step 1: Write the equation for the hydrogen carbonate ion acting as an acid

When HCO₃⁻ acts as an acid, it donates a proton (H⁺) to a water molecule. This results in the formation of a hydronium ion (H₃O⁺) and the carbonate ion (CO₃²⁻). The balanced chemical equation for this reaction is: \(HCO_3^- + H_2O \rightarrow H_3O^+ + CO_3^{2-}\)

Step 2: Write the equation for the hydrogen carbonate ion acting as a base

When HCO₃⁻ acts as a base, it accepts a proton (H⁺) from a water molecule. This results in the formation of a hydroxide ion (OH⁻) and carbonic acid (H₂CO₃). The balanced chemical equation for this reaction is: \(HCO_3^- + H_2O \rightarrow OH^- + H_2CO_3\) In conclusion, we have demonstrated the amphiprotic behavior of the hydrogen carbonate ion through the two balanced chemical equations presented above.

Key Concepts

Understanding Acid-Base Reactions

Acid-base reactions are a fundamental part of chemistry, crucial for understanding how substances interact with each other in aqueous solutions. In the simplest terms, these reactions involve the transfer of protons

• Acids are substances that donate protons ( ext{H}^+), which means they have a proton to "give" to another substance.
• Bases, on the other hand, are substances that accept protons, essentially "taking" the proton from an acid.

This concept is rooted in Brønsted-Lowry theory, which defines acids and bases in terms of their ability to donate or accept protons. Acid-base reactions often lead to the formation of water and a salt when the acid and base neutralize each other. Thus, understanding these interactions helps to predict the products of chemical reactions and their potential equilibrium states.
A key factor to note in acid-base chemistry is the strength of the participating acids and bases, which can affect the extent and direction of a reaction. Different substances can react in various ways depending on their nature and the conditions under which they meet. This leads to a diverse range of chemical behavior evident in many natural processes.

The Role of Hydrogen Carbonate Ion

Hydrogen carbonate ion ( ext{HCO}_3^-), also known as bicarbonate, is a fascinating example of an amphiprotic substance, which means it can act both as an acid and as a base.

• When acting as an acid, ext{HCO}_3^- donates a proton to water, forming hydronium ion ( ext{H}_3O^+) and carbonate ion ( ext{CO}_3^{2-}).
• Conversely, when acting as a base, ext{HCO}_3^- accepts a proton from water, resulting in the formation of hydroxide ion ( ext{OH}^-) and carbonic acid ( ext{H}_2 ext{CO}_3).

This dual nature is due to the hydrogen carbonate ion's composition, where both a proton donor site and an acceptor site coexist. This amphiprotic characteristic makes it immensely important in biological and environmental systems, like the human blood buffer system, where it helps maintain constant pH levels.

Balancing Chemical Equations

Writing and balancing chemical equations is an essential skill in chemistry for representing chemical reactions accurately. Essentially, chemical equations are a shorthand notation for chemical reactions. They tell us the reactants, the products, and the proportions in which they react or are produced.

• A balanced chemical equation has the same number of each type of atom on both sides of the equation.
• Balancing involves adjusting coefficients—the numbers before molecules—so that atom counts are equal on both sides.

For example, when showing the hydrogen carbonate ion ( ext{HCO}_3^-) acting as an acid, the equation is balanced as:\[ ext{HCO}_3^- + ext{H}_2 ext{O} \rightarrow ext{H}_3 ext{O}^+ + ext{CO}_3^{2-}\]When balancing these equations, it's vital to ensure that both mass and charge are conserved, maintaining the same total charge on both sides of the equation. This practice reflects one of the fundamental laws of chemistry: the law of conservation of mass. Mastering chemical equation balancing is key to predicting the outcomes of reactions and understanding the principles governing them.