Acid-base chemistry involves the transfer of charged particles between compounds. When an acid undergoes a specific transformation, it results in a new substance called the conjugate base. Identifying the conjugate base is a simple mechanical process once the underlying concept of this chemical partnership is understood. Correctly determining the formula and charge of a conjugate base is fundamental for predicting reaction outcomes and understanding chemical equilibrium.
Defining Acids, Bases, and Conjugate Pairs
The relationship between an acid and its conjugate base is defined by the Brønsted-Lowry theory, which focuses on the movement of a proton. An acid is any chemical species capable of donating a proton, represented as a hydrogen ion (\(\text{H}^+\)). Conversely, a base is any species that accepts this proton during a reaction.
When an acid donates its proton, the remnant species is called the conjugate base, which can accept a proton back to reform the original acid. When a base accepts a proton, it becomes the conjugate acid. The acid and its corresponding conjugate base form a pair that differs by only a single proton (\(\text{H}^+\)).
This pairing means the acid and conjugate base are always on opposite sides of a chemical equation. For instance, hydrochloric acid (\(\text{HCl}\)) and the chloride ion (\(\text{Cl}^-\)) constitute a single conjugate pair. Understanding this difference—the presence or absence of one \(\text{H}^+\)—is key to determining the conjugate base.
The Step-by-Step Method for Determining the Conjugate Base
The process for determining the formula of a conjugate base is a straightforward, two-step procedure involving the acid’s chemical formula and electrical charge. The first step involves removing one hydrogen ion (\(\text{H}^+\)) from the acid’s formula. This action represents the acid’s role as a proton donor.
For example, removing one \(\text{H}\) atom from hydrofluoric acid (\(\text{HF}\)) leaves the fluorine atom (\(\text{F}\)). If the starting acid is the ammonium ion (\(\text{NH}_4^+\)), removing one \(\text{H}\) atom yields ammonia (\(\text{NH}_3\)). This establishes the foundational structure of the resulting conjugate base.
The second step is to adjust the overall electrical charge of the resulting molecule. Since the removed particle (\(\text{H}^+\)) carries a single positive charge, the remaining species must have its charge lowered by exactly one unit. If the original acid was neutral, like \(\text{HF}\), the resulting conjugate base (\(\text{F}\)) will now carry a negative one charge (\(\text{F}^-\)).
If the starting acid already carries a charge, such as the bisulfate ion (\(\text{HSO}_4^-\)), the charge adjustment is still applied. Removing \(\text{H}^+\) from \(\text{HSO}_4^-\) results in the sulfate ion (\(\text{SO}_4\)), and the negative one charge decreases by one unit to a negative two charge (\(\text{SO}_4^{2-}\)). These two rules—proton removal and single-unit charge decrease—determine the conjugate base.
Conjugate Bases of Polyprotic Acids
The method is slightly more complex when dealing with polyprotic acids, which can donate more than one proton. These acids do not lose all protons simultaneously; instead, donation occurs sequentially, one proton at a time. Each step produces an intermediate conjugate base that still contains acidic hydrogen and can act as an acid in the next step.
Phosphoric acid (\(\text{H}_3\text{PO}_4\)) is a triprotic acid capable of losing three protons. The first step involves the loss of one proton, yielding the dihydrogen phosphate ion (\(\text{H}_2\text{PO}_4^-\)). This ion is the first conjugate base, but it still retains two acidic hydrogen atoms.
In the second step, the \(\text{H}_2\text{PO}_4^-\) ion loses another proton, resulting in the hydrogen phosphate ion (\(\text{HPO}_4^{2-}\)), the second conjugate base. The third and final step sees \(\text{HPO}_4^{2-}\) donate its last proton to form the phosphate ion (\(\text{PO}_4^{3-}\)), the final conjugate base. Intermediate species, like \(\text{H}_2\text{PO}_4^-\) and \(\text{HPO}_4^{2-}\), are known as amphiprotic because they can both donate and accept a proton.