Both sodium hydrogen carbonate (sodium bicarbonate) and sodium hydroxide can be used to neutralize acid spills. What is the pH of \(1.00 \mathrm{M} \mathrm{NaHCO}_{3}(\mathrm{aq})\) and of \(1.00 \mathrm{M} \mathrm{NaOH}(\mathrm{aq}) ?\) On a per-liter basis, do these two solutions have an equal capacity to neutralize acids? Explain. On a per-gram basis, do the two solids, \(\mathrm{NaHCO}_{3}(\mathrm{s})\) and \(\mathrm{NaOH}(\mathrm{s}),\) have an equal capacity to neutralize acids? Explain. Why do you suppose that \(\mathrm{NaHCO}_{3}\) is often preferred to \(\mathrm{NaOH}\) in neutralizing acid spills?

Using the formula: pH = -log[H+], and knowing that a 1.00 M \(\mathrm{NaOH}\) solution is a strong base and it will completely dissociate into \(\mathrm{Na}^{+}\) and \(\mathrm{OH}^{-}\) ions resulting in [OH-] = 1.00 M. The relationship between [H+] and [OH-] is represented as Kw=[H+][OH-], where Kw is the ion product of water, equal to \(1 * 10^{-14}\). Solving for [H+], we get [H+] = Kw/[OH-] = \(1 * 10^{-14}/1.00\). Using this value to solve for pH, we get pH = -log[H+] = 14.

Considering \(\mathrm{NaHCO}_{3}\) as a weak base, and knowing that it partially dissociates in water to form bicarbonate ions \(HCO_{3}^{-}\) and hydronium ions \(H_{3}O^{+}\). The reaction can be written as: \(\mathrm{NaHCO}_{3}\) + H2O <-> \(H_{3}O^{+}\) + \(\mathrm{HCO}_{3}^{-}\). Now the equation for the bicarbonate ion acting as an acid can be written as: \(\mathrm{HCO}_{3}^{-}\) <-> \(H_{3}O^{+}\) + \(\mathrm{CO}_{3}^{2-}\). By setting up an equilibrium constant expression for this reaction with the known \(Ka\) and \(Kb\) values for the bicarbonate ion, the [H+] concentration can be solved. Using this \(H^{+}\) concentration, the pH of the solution can be calculated, which results to be approximately 8.3.

To neutralize an acid, a base must provide OH- ions. A 1.00 M solution of NaOH will provide a greater number of OH- ions (and hence a greater acid-neutralizing capacity) than a 1.00 M solution of NaHCO3, because the NaOH dissociates completely while NaHCO3 only partially dissociates to provide OH- ions.

To neutralize an acid, a base must provide OH- ions. Comparing the molar masses (NaHCO3: 84 g/mol, NaOH: 40 g/mol), we find that per gram, NaOH provides a greater number of OH- ions (and hence a greater acid-neutralizing capacity) than NaHCO3.

\(\mathrm{NaHCO}_{3}\) is often preferred over \(\mathrm{NaOH}\) in neutralizing acid spills because it is a weaker base as compared to \(\mathrm{NaOH}\), so it is inherently safer to handle and less likely to cause injury due to excessive neutralization (over-shooting the pH to highly basic). It also releases carbon dioxide gas when it reacts with acids, providing an indicator of reaction progress.