Question

Hydrazine, \(\mathrm{N}_{2} \mathrm{H}_{4}\), is a weaker Brønsted-Lowry base than ammonia. In the following reaction, would the position of equilibrium lie to the left or to the right? Justify your answer $$ \mathrm{N}_{2} \mathrm{H}_{5}^{+}+\mathrm{NH}_{3} \rightleftharpoons \mathrm{N}_{2} \mathrm{H}_{4}+\mathrm{NH}_{4}^{+} $$

Short answer

The position of equilibrium would lie to the right.

Step-by-step solution

Understanding the Concept of Brønsted-Lowry Acids and Bases

Recall the definition of Brønsted-Lowry acids and bases. A Brønsted-Lowry acid is a substance that can donate a proton (H+), while a Brønsted-Lowry base is a substance that can accept a proton. A stronger base has a higher affinity for protons and will more readily accept them.

Compare the Basicity of the Species Involved

Since hydrazine (N_2H_4) is a weaker base than ammonia (NH_3), hydrazine has less affinity for protons. Therefore, hydrazine is less likely to accept a proton compared to ammonia.

Analyze the Reaction and Equilibrium Position

The equation shows hydrazinium ion (N_2H_5^{+}) reacting with ammonia (NH_3) to form hydrazine (N_2H_4) and ammonium ion (NH_4^{+}). Since hydrazine is the weaker base, it is less likely to recombine with a proton (from NH_4^{+}) to form N_2H_5^{+}. Instead, the equilibrium will favor the side with the weaker base, which means it shifts to the right, producing more N_2H_4 and NH_4^{+}.

Key concepts

Chemical Equilibrium

Chemical equilibrium is a state in which the rate of the forward reaction is equal to the rate of the reverse reaction, resulting in no net change in the concentration of the reactants and products over time. It is represented by the equilibrium constant, K, which quantifies the concentrations of reactants and products at equilibrium. In the context of acid-base reactions, such as the one between hydrazinium ion and ammonia, knowing which side the equilibrium favors is critical. It will favor the formation of the weaker Brønsted-Lowry base, hydrazine, over the stronger base, ammonia, indicating that the reaction lies to the right. This direction is predicted by Le Chatelier's principle, which states that an equilibrium will shift to counteract any changes imposed on the system. In this case, the presence of the stronger base, ammonia, pushes the equilibrium towards the weaker base, hydrazine, to minimize the effect of the added ammonia.

Understanding chemical equilibrium helps in predicting the outcome of reactions and the concentrations of species present at equilibrium. This concept is crucial in many fields, including industrial chemistry, where maximizing the yield of a desired product is essential.

Acid-Base Reactions

Acid-base reactions are processes in which a proton (H+) is transferred from one substance to another. According to the Brønsted-Lowry theory, an acid is a proton donor, whereas a base is a proton acceptor. The reaction between hydrazinium ion (N_2H_5^+) and ammonia (NH_3) is an example of an acid-base interaction, where N_2H_5^+donates a proton to NH_3, yielding hydrazine (N_2H_4) and ammonium ion (NH_4^+).

In analyzing acid-base reactions, it is also important to consider the relative strength of the acids and bases involved. The strength of an acid or base is determined by its ability to donate or accept protons, which reflects the stability of the resulting species after the transfer of the H+. A stronger base will more readily accept a proton, leading to a higher position on the basicity scale, whereas a stronger acid easily donates its proton. The reaction's equilibrium will often favor the formation of the weaker acid and base, as is the case in the reaction being discussed.

Basicity of Compounds

The basicity of compounds refers to their ability to accept protons. Factors influencing basicity include the availability of electron pairs for bonding, the stability of the base once it has accepted a proton, and the molecules' structure and charge. In the presented reaction, comparing the basicity of hydrazine and ammonia is key to understanding the equilibrium position. Hydrazine, being a weaker base than ammonia, is less inclined to accept a proton, which is directly correlated to its lower basicity.

Basicity can be quantified using a pKb value, which is derived from the equilibrium constant of the base's reaction with water. A lower pKb indicates a stronger base, and vice versa. Alterations in the molecular structure, such as electron-donating or withdrawing groups, can impact the basicity of a compound by affecting electron density. The exercise showcases that knowing the basicity of the reactants can be crucial in predicting the direction of the equilibrium in acid-base reactions.