Question

Write a chemical equation showing how each of the following species behaves as a base when dissolved in water. a. \(\mathrm{NH}_{3}\) b. \(\mathrm{NH}_{2}^{-}\) c. \(\mathrm{HSO}_{3}^{-}\) d. \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\)

Short answer

a. \(\mathrm{NH}_{3(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)} \rightleftharpoons \mathrm{NH}_{4}^{+}_{(aq)} + \mathrm{OH}^{-}_{(aq)}\) b. \(\mathrm{NH}_{2}^{-}_{(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)} \rightleftharpoons \mathrm{NH}_{3(aq)} + \mathrm{OH}^{-}_{(aq)}\) c. \(\mathrm{HSO}_{3}^{-}_{(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)} \rightleftharpoons \mathrm{H}_{2}\mathrm{SO}_{3(aq)} + \mathrm{OH}^{-}_{(aq)}\) d. \(\mathrm{HC}_{2}\mathrm{H}_{3}\mathrm{O}_{2(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)} \rightleftharpoons \mathrm{H}_{2}\mathrm{C}_{2}\mathrm{H}_{3}\mathrm{O}_{2}^{+}_{(aq)} + \mathrm{OH}^{-}_{(aq)}\)

Step-by-step solution

a. Ammonia (\(\mathrm{NH}_{3}\))

When ammonia (\(\mathrm{NH}_{3}\)) is dissolved in water, it acts as a base by accepting a proton from a water molecule to form ammonium ion (\(\mathrm{NH}_{4}^{+}\)) and a hydroxide ion (\(\mathrm{OH}^{-}\)). The chemical equation for this process is: \[\mathrm{NH}_{3(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)} \rightleftharpoons \mathrm{NH}_{4}^{+}_{(aq)} + \mathrm{OH}^{-}_{(aq)}\]

b. Amide ion (\(\mathrm{NH}_{2}^{-}\))

When the amide ion (\(\mathrm{NH}_{2}^{-}\)) is dissolved in water, it acts as a base by accepting a proton from a water molecule to form ammonia (\(\mathrm{NH}_{3}\)) and a hydroxide ion (\(\mathrm{OH}^{-}\)). The chemical equation for this process is: \[\mathrm{NH}_{2}^{-}_{(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)} \rightleftharpoons \mathrm{NH}_{3(aq)} + \mathrm{OH}^{-}_{(aq)}\]

c. Bisulfite ion (\(\mathrm{HSO}_{3}^{-}\))

When the bisulfite ion (\(\mathrm{HSO}_{3}^{-}\)) is dissolved in water, it acts as a base by accepting a proton from a water molecule to form sulfurous acid (\(\mathrm{H}_{2}\mathrm{SO}_{3}\)) and a hydroxide ion (\(\mathrm{OH}^{-}\)). The chemical equation for this process is: \[\mathrm{HSO}_{3}^{-}_{(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)} \rightleftharpoons \mathrm{H}_{2}\mathrm{SO}_{3(aq)} + \mathrm{OH}^{-}_{(aq)}\]

d. Acetic acid (\(\mathrm{HC}_{2}\mathrm{H}_{3}\mathrm{O}_{2}\))

When acetic acid (\(\mathrm{HC}_{2}\mathrm{H}_{3}\mathrm{O}_{2}\)) is dissolved in water, it can act as a base by accepting a proton from a water molecule to form a dihydrogen acetate ion (\(\mathrm{H}_{2}\mathrm{C}_{2}\mathrm{H}_{3}\mathrm{O}_{2}^{+}\)) and a hydroxide ion (\(\mathrm{OH}^{-}\)). The chemical equation for this process is: \[\mathrm{HC}_{2}\mathrm{H}_{3}\mathrm{O}_{2(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)} \rightleftharpoons \mathrm{H}_{2}\mathrm{C}_{2}\mathrm{H}_{3}\mathrm{O}_{2}^{+}_{(aq)} + \mathrm{OH}^{-}_{(aq)}\]

Key concepts

Chemical equations

Chemical equations are essential for describing chemical reactions in a clear and concise way. They show how reactants transform into products over a reaction.
It's like a language that conveys all the actions taking place during chemical reactions. In these equations, each chemical formula represents a substance involved, while symbols indicate whether the process is happening in aqueous, liquid, solid, or gaseous states.

Chemical equations help visualize important concepts, such as:

• Conservation of mass: The number of each type of atom must be equal on both sides of the equation.
• Balance: In balancing equations, coefficients are adjusted to ensure the same amount of each element on both the reactant and product sides.
• Direction of reaction: Reversible reactions are indicated by a double arrow (\[\rightleftharpoons\]), reflecting the dynamic equilibrium where reactants and products are continuously formed.

This open system allows for an understanding of how reactions are influenced by various conditions like concentration, temperature, and pressure.
As shown above, chemical equations make it easier to observe how specific substances act as bases in water, each by accepting a proton.

Bronsted-Lowry theory

The Bronsted-Lowry theory is a fundamental concept in acid-base chemistry. It broadens the definition of acids and bases beyond hydrogen ions and hydroxide ions.
According to this theory:

• Acids are proton donors (substances that can release a hydrogen ion or \( \mathrm{H}^{+} \)).
• Bases are proton acceptors (substances that can take in a hydrogen ion).

In the reactions presented, each species acts as a Bronsted-Lowry base by accepting a proton from water. This process results in the formation of a hydroxide ion (\( \mathrm{OH}^{-} \)).
This concept allows for an explanation of acid-base reactions in a wider variety of chemical contexts, making it invaluable wherever hydrogen transfer occurs.

You can think of it as a hand-off—the base accepts a proton from an acid, much like receiving a ball in a relay race. The Bronsted-Lowry theory helps explain the behavior of these species in aqueous solutions and supports predictions about how chemical species will interact in varied environments.

Amphoteric substances

Amphoteric substances are fascinating because they have the dual ability to behave both as acids and bases. Such substances bring flexibility in reaction mechanisms, as they can either donate or accept protons based on the circumstances.
Water (\( \mathrm{H}_{2}\mathrm{O} \)) is a classic example—it can donate a proton to form hydroxide (\( \mathrm{OH}^{-} \)), acting as an acid, or accept a proton to become hydronium (\( \mathrm{H}_{3}\mathrm{O}^{+} \)), behaving as a base.

Amphoterism is particularly important in acidic and basic reactions:

• These substances can help neutralize excess acids or bases, acting as buffers to maintain the pH in solutions.
• In complex reactions, amphoteric substances provide pathways for both forwards and reverse reactions, supporting equilibrium processes.

The dual nature of amphoteric substances makes them crucial in biological systems and industrial applications, such as controlling pH in environmental engineering or the manufacturing of various materials.