Problem 15 - Introductory Chemistry Atoms First

Question

In the presence of water, the bicarbonate ion, \(\mathrm{HCO}_{3}^{-}\), can be either a BronstedLowry acid or a Bronsted-I.owry base. Write both equilibrium equations.

Step-by-Step Solution

Answer

The bicarbonate ion (HCO₃⁻) can act as a Bronsted-Lowry acid or base in the presence of water. As an acid, the equilibrium equation is \( HCO_{3}^{-} + H_{2}O \rightleftharpoons H_{3}O^{+} + CO_{3}^{2-} \). As a base, the equilibrium equation is \( HCO_{3}^{-} + H_{2}O \rightleftharpoons OH^{-} + H_{2}CO_{3} \).

Step 1: Understand Bronsted-Lowry Acid-Base Definition

A Bronsted-Lowry acid is a substance that can donate a proton (H⁺), while a Bronsted-Lowry base is a substance that can accept a proton (H⁺). In this case, we are given that the bicarbonate ion (HCO₃⁻) can act as both an acid and a base when dissolved in water.

Step 2: Write the Acid Equilibrium Equation

When the bicarbonate ion (HCO₃⁻) acts as an acid, it donates a proton (H⁺) to a water molecule and forms hydronium ions (H₃O⁺) and carbonate ions (CO₃²⁻). The acid equilibrium equation can be written as: \[ HCO_{3}^{-} + H_{2}O \rightleftharpoons H_{3}O^{+} + CO_{3}^{2-} \]

Step 3: Write the Base Equilibrium Equation

When the bicarbonate ion (HCO₃⁻) acts as a base, it accepts a proton (H⁺) from a water molecule and forms hydroxide ions (OH⁻) and carbonic acid (H₂CO₃). The base equilibrium equation can be written as: \[ HCO_{3}^{-} + H_{2}O \rightleftharpoons OH^{-} + H_{2}CO_{3} \] So, the two equilibrium equations representing bicarbonate ion (HCO₃⁻) as a Bronsted-Lowry acid and as a Bronsted-Lowry base in the presence of water are: \[ HCO_{3}^{-} + H_{2}O \rightleftharpoons H_{3}O^{+} + CO_{3}^{2-} \] \[ HCO_{3}^{-} + H_{2}O \rightleftharpoons OH^{-} + H_{2}CO_{3} \]

Key Concepts

Bicarbonate Ion

The bicarbonate ion, denoted as \( \mathrm{HCO}_{3}^{-} \), is an interesting and versatile species in chemistry. It plays a crucial role in buffering systems, which help maintain pH stability in various environments such as blood plasma and natural waters.
In the Bronsted-Lowry acid-base theory, the bicarbonate ion is notable for its dual personality. It can switch roles, acting either as an acid or a base, depending on the situation. This ability to interchange between donating and accepting protons makes bicarbonate ions a central component in many chemical processes.
When acting as an acid, the bicarbonate ion can donate a proton generating carbonate ions \( \mathrm{CO}_{3}^{2-} \), while as a base, it can accept a proton to form carbonic acid \( \mathrm{H}_{2}\mathrm{CO}_{3} \). This dual functionality is a practical examples of chemical equilibrium in action.

Equilibrium Equations

Equilibrium equations help us understand the reversible reactions where substances can both convert and revert into different products. In the context of bicarbonate ions, these form the basis for depicting how \( \mathrm{HCO}_{3}^{-} \) interacts with water in dual roles.
When bicarbonate acts as an acid, the equilibrium equation is:

• \[ \mathrm{HCO}_{3}^{-} + \mathrm{H}_{2}\mathrm{O} \rightleftharpoons \mathrm{H}_{3}\mathrm{O}^{+} + \mathrm{CO}_{3}^{2-} \]

This shows the bicarbonate ion donating a proton to form hydronium \( \mathrm{H}_{3}\mathrm{O}^{+} \) and carbonate ions.
When it acts as a base, the equation becomes:

• \[ \mathrm{HCO}_{3}^{-} + \mathrm{H}_{2}\mathrm{O} \rightleftharpoons \mathrm{OH}^{-} + \mathrm{H}_{2}\mathrm{CO}_{3} \]

Here, bicarbonate accepts a proton to produce hydroxide ions \( \mathrm{OH}^{-} \) and carbonic acid. These equations neatly illustrate the transformation process depending on its function.

Proton Transfer

Proton transfer is an essential concept in Bronsted-Lowry acid-base reactions. It involves the movement of protons, \( \mathrm{H}^{+} \), from one molecule to another.
In our example, the bicarbonate ion \( \mathrm{HCO}_{3}^{-} \) transfers protons in two primary ways:

• **As a Proton Donor (Acid):** It gives up a proton to water, forming hydronium ions \( \mathrm{H}_{3}\mathrm{O}^{+} \).
• **As a Proton Acceptor (Base):** It accepts a proton from water, resulting in hydroxide ions \( \mathrm{OH}^{-} \).

The efficiency and speed of these proton transfers are influenced by surrounding conditions such as pH, temperature, and the presence of catalysts. Understanding proton transfer reactions is key to predicting the direction and position of chemical equilibria.

Hydronium and Hydroxide Ions

In aqueous chemistry, hydronium \( \mathrm{H}_{3}\mathrm{O}^{+} \) and hydroxide \( \mathrm{OH}^{-} \) ions are fundamental in defining acidity and basicity. These ions arise prominently in reactions where water serves as either a base or an acid.

• **Hydronium Ion:** Formed when water accepts a proton. It represents the characteristic cation in acidic solutions. For instance, when the bicarbonate ion acts as an acid, it donates a proton to produce \( \mathrm{H}_{3}\mathrm{O}^{+} \).
• **Hydroxide Ion:** Formed when water donates a proton. This ion is significant in basic solutions. In the bicarbonate base reaction, \( \mathrm{OH}^{-} \) formation signifies an increase in basicity as \( \mathrm{HCO}_{3}^{-} \) accepts a proton.

In the context of Bronsted-Lowry acid-base theory, these ions help to understand the splitting and recombination of molecules making them essential in analyzing pH change and acid-base balance.