Question

Write a balanced equation showing how the \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) ion can be either a Bronsted-Lowry acid or a Brønsted-Lowry base.

Short answer

Provide the balanced equations for each case. Answer: The \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) ion can act as a Brønsted-Lowry acid by donating a proton, as shown in the equation: \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-} + \mathrm{H}_{2}\mathrm{O} \rightleftharpoons \mathrm{HPO}_{4}^{2-} + \mathrm{H}_{3}\mathrm{O}^{+}\). It can also act as a Brønsted-Lowry base by accepting a proton, as shown in the equation: \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-} + \mathrm{H}_{2}\mathrm{O} \rightleftharpoons \mathrm{H}_{3} \mathrm{PO}_{4} + \mathrm{OH}^{-}\).

Step-by-step solution

Acting as a Brønsted-Lowry Acid

In this case, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) will donate a proton (H+) to another species (base). Let's choose water as the base: \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-} + \mathrm{H}_{2}\mathrm{O} \rightleftharpoons \mathrm{HPO}_{4}^{2-} + \mathrm{H}_{3}\mathrm{O}^{+}\) Here, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) acts as an acid by donating a proton to water, forming \(\mathrm{HPO}_{4}^{2-}\) and the hydronium ion \(\mathrm{H}_{3}\mathrm{O}^{+}\).

Acting as a Brønsted-Lowry Base

In this case, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) will accept a proton (H+) from another species (acid). Again, let's choose water as the acid: \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-} + \mathrm{H}_{2}\mathrm{O} \rightleftharpoons \mathrm{H}_{3} \mathrm{PO}_{4} + \mathrm{OH}^{-}\) Here, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) acts as a base by accepting a proton from water, forming \(\mathrm{H}_{3} \mathrm{PO}_{4}\) and the hydroxide ion \(\mathrm{OH}^{-}\).

Key concepts

Acid-Base Reaction

In chemistry, an acid-base reaction is a fundamental type of chemical process typified by the exchange of one or more hydrogen ions, H+, between species that are classified as acids or bases. This concept takes on a clear form within the Bronsted-Lowry acid-base theory, which defines an acid as a proton donor and a base as a proton acceptor. For example, the dihydrogen phosphate ion (H_{2}PO_{4}^{-}) can participate in an acid-base reaction either by donating a proton to water, thus acting as a Bronsted-Lowry acid, or by accepting a proton, functioning as a Bronsted-Lowry base.

Understanding these reactions is crucial for students as they reveal how substances interact in solutions to form new products, and can affect pH levels, solubility, and reaction mechanisms. In essence, acid-base reactions are at the heart of many biological and environmental processes.

Proton Transfer

The term proton transfer captures the essence of the Bronsted-Lowry acid-base theory. It's the journey of a proton (H+) from an acid to a base. This transfer is the crux of an acid-base reaction. A terrific illustration of proton transfer is seen when H_{2}PO_{4}^{-} gives up a proton to water, creating HPO_{4}^{2-} and hydronium ion (H_{3}O^{+}), showcasing its ability to act as an acid. Conversely, when it accepts a proton, it forms H_{3}PO_{4} and a hydroxide ion (OH^{-}), showing its base-like behavior.

It's important for students to visualize proton transfer as a central event that changes the identity of the molecules involved, leading to the dynamic nature of chemical processes in solution.

Chemical Equilibrium

The concept of chemical equilibrium is essential when discussing acid-base reactions. It's reached when the rate of the forward reaction, where acids donate protons to bases, equals the rate of the reverse reaction, where the conjugate bases return protons to the conjugate acids. In essence, it's a state of balance where chemical substances do not cease reacting altogether, but do so in a way that their concentrations remain constant over time.

Applying this to our H_{2}PO_{4}^{-} example, whether it's acting as an acid or a base, each reaction reaches a point where no further change in the concentration of reactants and products occurs. The double arrow (rightleftharpoons) signifies this equilibrium in chemical equations. Understanding this concept helps students to comprehend that many reactions in nature are reversible and that equilibrium is a dynamic process of push and pull between products and reactants.

Hydronium Ion

Hydronium ion (H_{3}O^{+}), often represented simply by H+, is a water molecule with an extra proton. It's a key player in the Bronsted-Lowry definition of acids and bases. When an acid donates a proton in an aqueous solution, it's usually a water molecule that picks up this proton to form the hydronium ion. For instance, when H_{2}PO_{4}^{-} acts as an acid, it donates a proton to water, resulting in the formation of a hydronium ion.

This ion is paramount for measuring pH, which is an indication of the acidity or basicity of a solution. The concentration of hydronium ions directly correlates with the pH level; more hydronium ions mean a lower pH and a more acidic solution. By understanding the role of the hydronium ion, students can better grasp the pH scale and the relative strength of acids.

Hydroxide Ion

On the flip side, we have the hydroxide ion (OH^{-}), which is a water molecule that has lost a proton. This ion is what's left when a Bronsted-Lowry base grabs a proton from an acid. Take the scenario where H_{2}PO_{4}^{-} accepts a proton, for instance; the water molecule that donated the proton is transformed into a hydroxide ion. The hydroxide ion is integral to the concept of bases and is equally important in determining pH. An increase in hydroxide ion concentration indicates higher basicity, reflected by a higher pH value.

Comprehending the role of the hydroxide ion helps students to understand not just pH, but also the balancing act in neutralization reactions, where acids and bases react to form water and a salt, essentially the hydroxide ion merging back with a proton to form water again.