Use the Bronsted-Lowry definition to explain why \(\mathrm{NH}_{3}\) is a base in water.

The theory of acids and bases is fundamental in understanding how various substances interact in an aqueous solution. According to the Bronsted-Lowry theory, acids are proton donors, while bases are proton acceptors. This definition shifts the focus from the substance itself to the reaction it undergoes. For instance, a substance can act as an acid in one reaction and a base in another, emphasizing the importance of context.

Specifically, in aqueous solutions, water can facilitate these proton transfer reactions. An acid when dissolved in water, will donate a proton to a water molecule, whereas a base will receive a proton from a water molecule, altering the composition of the solution. This interaction is essential for processes in both chemistry and biology, as the formation and breakage of these bonds have vast implications, from industrial synthesis to cellular metabolism.

Chemical equilibrium occurs when the rates of the forward and reverse reactions equalize, and the concentrations of reactants and products remain constant over time. It's a dynamic balance rather than a static one - reactions continue to occur, but with no net change in concentration of substances. For the reaction involving ammonia in water, \(\mathrm{NH_3 (aq) + H_2O (l) \rightleftharpoons NH_4^+ (aq) + OH^- (aq)}\), equilibrium implies that the speed at which ammonia accepts protons from water to form ammonium and hydroxide ions is equal to the speed at which ammonium ions donate protons back to hydroxide ions to form water and ammonia.

Grasping the concept of chemical equilibrium is crucial because it helps predict the behavior of substances in various chemical reactions and understand the conditions required to shift the equilibrium to favor the production of desired products.

Proton transfer reactions are at the heart of the Bronsted-Lowry acid-base theory. They are characterized by the movement of a single proton (\(H^+\)) from one molecule to another. These reactions are ubiquitous and fundamental in chemistry - from simple acid-base reactions you observe in a high school lab to complex enzymatic reactions that sustain life.

When ammonia (\(\mathrm{NH_3}\)) reacts with water, it undergoes a proton transfer reaction. In this process, water (\(\mathrm{H_2O}\)), which acts as a Bronsted-Lowry acid, donates a proton to the ammonia, demonstrating ammonia's act as a Bronsted-Lowry base. The outcome of this reaction is the formation of ammonium (\(\mathrm{NH_4^+}\)) and hydroxide ions (\(\mathrm{OH^-}\)), both of which are significant ions in chemistry and biochemistry.

In the context of the Bronsted-Lowry theory, ammonia (\(\mathrm{NH_3}\)) is a classic example of a weak base when dissolved in water. It readily accepts a proton from water, forming an ammonium ion (\(\mathrm{NH_4^+}\)). As ammonia takes on a proton, it contributes to the increase in hydroxide ion concentration, creating a more basic environment.

This behavior of ammonia can be used to explain its role in various biological and industrial processes, such as in the formation of fertilizers, where its reactivity and ability to form bonds with acids are capitalized on. Understanding the base behavior of ammonia is also necessary for comprehending pH regulation in biological systems, as it plays a role in buffering changes in pH by reacting with acids and bases present in the body.