Problem 34 - Introductory Chemistry Atoms First

Question

(a) Show how \(\mathrm{HCO}_{3}^{-}\) can act as a weak acid. (b) Show how \(\mathrm{HCO}_{3}^{-}\) can act as a weak base. (c) Since \(\mathrm{HCO}_{3}^{-}\) can act as either an acid or a base, can it be its own conjugate?

Step-by-Step Solution

Answer

(a) As a weak acid, HCO3- donates a proton to a water molecule, forming H2CO3 and OH-. The reaction is: \[ \mathrm{HCO}_{3}^{-} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{2} \mathrm{CO}_{3} + \mathrm{OH}^{-} \]. (b) As a weak base, HCO3- accepts a proton from a water molecule, forming CO3^2- and H3O+. The reaction is: \[ \mathrm{HCO}_{3}^{-} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{CO}_{3}^{2-} + \mathrm{H}_{3}\mathrm{O}^{+} \]. (c) HCO3- cannot be its own conjugate acid or base, as it forms different species (CO3^2- and H2CO3) when it acts as a base and as an acid, respectively. However, it is an amphiprotic species, since it can both donate and accept protons in different reactions.

Step 1: Write down the Bronsted-Lowry definition of an acid

An acid is a substance that donates a proton (H+) to another molecule or ion.

Step 2: Show HCO3- donating a proton

When HCO3- acts as a weak acid, it donates a proton (H+) to a water molecule, forming H2CO3 (carbonic acid) and OH- (hydroxide ion). The equation for this reaction can be written as: \[ \mathrm{HCO}_{3}^{-} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{2} \mathrm{CO}_{3} + \mathrm{OH}^{-} \] #b) HCO3- acting as a weak base#

Step 1: Write down the Bronsted-Lowry definition of a base

A base is a substance that accepts a proton (H+) from another molecule or ion.

Step 2: Show HCO3- accepting a proton

When HCO3- acts as a weak base, it accepts a proton from a water molecule, forming CO3^2- (carbonate ion) and H3O+ (hydronium ion). The equation for this reaction can be written as: \[ \mathrm{HCO}_{3}^{-} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{CO}_{3}^{2-} + \mathrm{H}_{3}\mathrm{O}^{+} \] #c) Can HCO3- be its own conjugate? #

Step 1: Define conjugate acid and conjugate base

A conjugate acid is the species that is formed after a base accepts a proton, and a conjugate base is the species that is formed after an acid donates a proton.

Step 2: Identify the conjugate acid and base of HCO3-

In the reaction where HCO3- acts as a weak base, it forms CO3^2-, so CO3^2- is the conjugate base of HCO3-. In the reaction where HCO3- acts as a weak acid, it forms H2CO3, so H2CO3 is the conjugate acid of HCO3-.

Step 3: Determine if HCO3- can be its own conjugate

Since HCO3- forms different species (CO3^2- and H2CO3) when it acts as a base and as an acid, respectively, it cannot be considered its own conjugate acid or base. However, it can be considered as an amphiprotic species, as it can both donate and accept protons in different reactions.

Key Concepts

Bronsted-Lowry acid

In the language of chemistry, an acid is any substance capable of donating a proton (H⁺) to another molecule or ion. This concept is central to the Bronsted-Lowry theory. It helps chemists understand and predict the behavior of acids in various reactions.

Take the bicarbonate ion, \( ext{HCO}_3^- \). It can act as a Bronsted-Lowry acid in certain conditions. This means it donates a proton to another substance.

For instance, when \( ext{HCO}_3^- \) encounters water, it releases a proton to form carbonic acid \( ext{H}_2 ext{CO}_3 \) and leaves behind a hydroxide ion, \( ext{OH}^- \). The reaction is:

• \[ ext{HCO}_3^- + ext{H}_2 ext{O} \rightleftharpoons ext{H}_2 ext{CO}_3 + ext{OH}^- \]

This ability to donate a proton shows its role as a weak acid, highlighting the amphiprotic nature of bicarbonate in various chemical processes.

Bronsted-Lowry base

On the flip side, a base in the Bronsted-Lowry sense is a molecule or ion that accepts a proton. This definition broadens the understanding of bases beyond the traditional hydroxide ions in water.

Consider again \( ext{HCO}_3^- \). It has the capacity to act as a Bronsted-Lowry base by accepting a proton. When interacting with water, \( ext{HCO}_3^- \) takes a proton to form carbonate ion \( ext{CO}_3^{2-} \) and the hydronium ion \( ext{H}_3 ext{O}^+ \). The reaction looks like this:

• \[ ext{HCO}_3^- + ext{H}_2 ext{O} \rightleftharpoons ext{CO}_3^{2-} + ext{H}_3 ext{O}^+ \]

In this instance, bicarbonate is acting as a weak base, absorbing protons and illustrating its dual nature.

Conjugate acid-base pair

The concept of conjugate acid-base pairs is an extension of the Bronsted-Lowry acid and base theory. A conjugate acid is formed when a base gains a proton. Conversely, a conjugate base arises when an acid loses a proton.

Let's focus on \( ext{HCO}_3^- \):

• As a base (when accepting a proton), it transforms into \( ext{CO}_3^{2-} \), which is its conjugate base.
• As an acid (when donating a proton), it turns into \( ext{H}_2 ext{CO}_3 \), which is its conjugate acid.

This dynamic shows how substances can shift roles, balancing between acid and base states in chemical reactions. In the case of bicarbonate, it forms different species depending on its function in a chemical interaction.

Therefore, a single substance can't be both its own conjugate acid and base since they result from different proton transfer events.

Carbonate chemistry

Understanding carbonate chemistry is crucial, especially given the role carbon dioxide and carbonates play in natural systems. Carbonate chemistry often involves amphiprotic substances like bicarbonate \( ext{HCO}_3^- \), which can act as both acids and bases.

In carbonate systems, interactions usually involve:

• The conversion between carbonic acid \( ext{H}_2 ext{CO}_3 \), bicarbonate \( ext{HCO}_3^- \), and carbonate ions \( ext{CO}_3^{2-} \).
• The equilibrium dynamics that maintain pH balance in environments like oceans and biological systems.

For example, in aquatic systems, bicarbonate acts as a buffer, helping maintain stable pH levels despite changes in environmental conditions. This adaptability is due to its role as both an acid, releasing protons, and a base, accepting protons, ready to neutralize pH fluctuations in its surroundings.

This buffering capability showcases how pivotal carbonate chemistry is across various fields, including climate science and biology.