Step 1: Understanding the Concept:
The Brønsted-Lowry theory of acids and bases provides a versatile framework for understanding chemical reactions in terms of proton (\(H^{+}\)) transfer. In this model, an acid is defined as a chemical species capable of donating a proton, while a base is a species capable of accepting a proton. This transfer process creates "conjugate pairs." When a base accepts a proton, it transforms into its "conjugate acid." Conversely, when an acid loses a proton, it becomes its "conjugate base." These two species differ from each other by exactly one proton and a unit of charge. Understanding these pairs is essential for determining the strength of acids and bases and calculating the pH of various buffer systems.
Step 2: Key Formula or Approach:
To find the Conjugate Acid of any given chemical species, the standard procedure is to add one proton (\(H^{+}\)) to the original formula. The general relationship is:
\[ \text{Base} + H^{+} \rightarrow \text{Conjugate Acid} \]
This process involves two simultaneous changes:
1. Increasing the number of Hydrogen atoms by one.
2. Increasing the total net electrical charge of the species by \(+1\).
Step 3: Detailed Explanation:
The species provided in the question is the amide ion, \(NH_{2}^{-}\). To determine its conjugate acid, we treat \(NH_{2}^{-}\) as a Brønsted-Lowry base that will accept a proton.
Let's apply the rules:
1. Hydrogen atom count:
The amide ion has 2 hydrogen atoms. Adding one proton brings the count to \(2 + 1 = 3\) hydrogen atoms. The chemical skeleton remains centered around Nitrogen (\(N\)), so the resulting molecular formula starts with \(NH_{3}\).
2. Electrical charge calculation:
The initial charge on the amide ion is \(-1\). A proton carries a charge of \(+1\). When the proton binds to the ion, the new charge is:
\[ \text{Final Charge} = (-1) + (+1) = 0 \]
Since the final charge is zero, the resulting species is a neutral molecule.
Combining these findings, the resulting species is neutral ammonia, \(NH_{3}\). The chemical equilibrium for this process can be represented as:
\[ NH_{2}^{-} + H^{+} \rightleftharpoons NH_{3} \]
In this reaction, \(NH_{2}^{-}\) is the base and \(NH_{3}\) is its corresponding conjugate acid. Note that if we were asked for the conjugate acid of ammonia (\(NH_{3}\)), we would add another proton to get the ammonium ion (\(NH_{4}^{+}\)). However, for the amide ion, the immediate conjugate acid is ammonia.
Step 4: Final Answer:
The conjugate acid of the amide ion (\(NH_{2}^{-}\)) is ammonia, represented by the formula \(NH_{3}\).