Problem 36 - Chemistry The Molecular Nature of Matter

Question

Write ionic equations that illustrate how each pair of compounds can serve as a buffer pair. (a) \(\mathrm{H}_{2} \mathrm{CO}_{3}\) and \(\mathrm{NaHCO}\) (the "carbonate" buffer in blood) (b) \(\mathrm{NaH}_{2} \mathrm{PO}_{4}\) and \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) (the "phosphate" buffer inside body cells) (c) \(\mathrm{NH}_{4} \mathrm{Cl}\) and \(\mathrm{NH}_{3}\) (d) Phenol and sodium phenolate

Step-by-Step Solution

Answer

a) \(\mathrm{H}_{2}\mathrm{CO}_{3} \leftrightarrow \mathrm{H}^+ + \mathrm{HCO}_{3}^-\), b) \(\mathrm{NaH}_{2}\mathrm{PO}_{4} \leftrightarrow \mathrm{H}^+ + \mathrm{Na}_{2}\mathrm{HPO}_{4}\), c) \(\mathrm{NH}_{4}^+ \leftrightarrow \mathrm{H}^+ + \mathrm{NH}_{3}\), d) \(\text{Phenol} \leftrightarrow \mathrm{H}^+ + \text{Sodium phenolate}\)

Step 1: Identify the Acid and its Conjugate Base

In each buffer pair, identify the weak acid (proton donor) and its conjugate base (proton acceptor). For (a) the weak acid is \(\mathrm{H}_{2}\mathrm{CO}_{3}\) and its conjugate base is \(\mathrm{HCO}_{3}^-\). For (b), the weak acid is \(\mathrm{NaH}_{2}\mathrm{PO}_{4}\) and the conjugate base is \(\mathrm{HPO}_{4}^{2-}\). For (c), the weak acid is \(\mathrm{NH}_{4}^+\) and its conjugate base is \(\mathrm{NH}_{3}\). For (d), the weak acid is phenol and its conjugate base is sodium phenolate.

Step 2: Write Weak Acid Dissociation

Write the dissociation reaction for the weak acid. For example, \(\mathrm{H}_{2}\mathrm{CO}_{3}\) dissociates as \(\mathrm{H}_{2}\mathrm{CO}_{3} \leftrightarrow \mathrm{H}^+ + \mathrm{HCO}_{3}^-\). Do this for each weak acid.

Step 3: Write Conjugate Base Reaction

Write the reaction for the conjugate base with \(\mathrm{H}^+\) to form the weak acid. For example, \(\mathrm{HCO}_{3}^- + \mathrm{H}^+ \leftrightarrow \mathrm{H}_{2}\mathrm{CO}_{3}\). This represents how the buffer can absorb excess protons. Do this for each conjugate base.

Step 4: Combine the Reactions

Combine the dissociation reaction of the weak acid with the reaction of the conjugate base to fully represent the buffer action. This gives the overall buffer equation.

Key Concepts

Ionic Equations

In chemistry, ionic equations are used to illustrate the reactions where ions participate. Unlike molecular equations, which show all reactants and products as compounds, ionic equations break down aqueous compounds into their respective ions.
For example, when discussing buffer solutions, we consider the dissociation of a weak acid into ions and its interaction with its conjugate base. This is essential to understand how buffers work. A weak acid like carbonic acid, goes through partial dissociation in water to form and bicarbonate ion . By representing such reactions as ionic equations, we make it clear which species are present in a solution and which are responsible for buffering action.

Acid-Base Buffer Systems

Acid-base buffer systems are remarkable because they maintain pH levels within a narrow range, crucial for many biological and chemical processes. A buffer system is generally composed of a weak acid and its conjugate base. The power of a buffer lies in its ability to neutralize added acids (protons) or bases (hydroxide ions) with little change in the overall pH of the solution.
Using the blood's carbonate buffer as an example, we see that carbonic acid can neutralize added hydroxide ions, while bicarbonate can neutralize additional protons. This dual action keeps the blood's pH tightly regulated. It's this equilibrium between the weak acid and its conjugate base that makes the buffer system efficient and fundamental to life.

Chemical Equilibrium

Chemical equilibrium is a state in a chemical reaction where the rate of the forward reaction equals the rate of the reverse reaction. At equilibrium, the concentrations of the reactants and products remain constant over time but are not necessarily equal to each other. This concept is critical when analyzing buffer solutions.
Take, for instance, the weak acid dissociation in a buffer system. Initially, the acid dissociates into ions, but as soon as the concentration of ions increases, the reverse reaction becomes significant, and a point of balance—or equilibrium—is established. This dynamism allows buffers to adjust to changes in pH by shifting the equilibrium slightly to the left or right to either produce more ions or reassociate them into the weak acid.

Weak Acid Conjugate Base Pair

A weak acid conjugate base pair is central to the functioning of buffer systems. This pair consists of a weak acid, which donates protons, and its conjugate base, which accepts protons. The acid and base are related by the loss or gain of a single proton.
For every weak acid, like acetic acid, there is a conjugate base, acetate in this case, that forms when the acid loses a proton. This pair works together to resist changes in pH. When an acid is added to the buffer, the conjugate base neutralizes it. Conversely, if a base is added, the weak acid donates a proton to neutralize the base. This interplay is crucial for maintaining the stability of the buffer's pH, thus allowing the buffer to effectively 'buffer' against pH changes in the solution.