Models and conjugate pairs

Understanding of the behavior of acids and bases developed decisively in the late 19th and early 20th centuries. Consider the substances $\mathrm{HCl}$ and $\mathrm{NH}_3$ at $25\,^{\circ}\mathrm{C}$.

1. a) Justify why hydrogen chloride behaves as an acid according to both Arrhenius's theory and the Brønsted-Lowry theory, writing the corresponding chemical equation in aqueous solution.
2. b) Explain, with reasoning, why a strict historical interpretation based only on the dissociation of hydroxides does not directly describe ammonia. Write its reaction with water and explain how proton transfer produces hydroxide ions, consistent with the extended aqueous definition of a base.

The concept of neutralization has complementary interpretations depending on the theoretical framework used.

1. a) Describe the neutralization reaction in aqueous solution between solutions of sodium hydroxide, $\mathrm{NaOH}$, and nitric acid, $\mathrm{HNO}_3$, writing the net ionic equation and justifying the process according to Arrhenius's definition.
2. b) Now consider the reaction between the gases ammonia, $\mathrm{NH}_3(\mathrm{g})$, and hydrogen chloride, $\mathrm{HCl}(\mathrm{g})$, to form the white solid ammonium chloride, $\mathrm{NH}_4\mathrm{Cl}(\mathrm{s})$, in the complete absence of solvent. Explain whether this process is a neutralization according to Arrhenius and according to Brønsted-Lowry, indicating the acid, the base and the proton transferred.

In the Brønsted-Lowry framework, two chemical species form a conjugate acid-base pair if they are related by the transfer of exactly one proton ($\mathrm{H}^+$), which involves a change of one unit in their net electric charge.

1. a) Write the molecular or ionic formula of the conjugate base of each of the following acidic chemical species: hydrogen sulfate ion, $\mathrm{HSO}_4^-$, and hydronium ion, $\mathrm{H}_3\mathrm{O}^+$.
2. b) Write the molecular or ionic formula of the conjugate acid corresponding to each of the following bases: hydrogen phosphate ion, $\mathrm{HPO}_4^{2-}$, and methylamine, $\mathrm{CH}_3\mathrm{NH}_2$.

Amphiprotic substances are those that can donate or accept protons depending on the chemical species they react with. The bicarbonate ion, $\mathrm{HCO}_3^-$, is an amphiprotic species of great importance in aqueous solution at $25\,^{\circ}\mathrm{C}$.

1. a) Write the balanced reaction in which the bicarbonate ion acts as a Brønsted-Lowry acid toward the hydroxide anion, $\mathrm{OH}^-$, identifying the two conjugate acid-base pairs.
2. b) Write the balanced reaction in which the bicarbonate ion acts as a Brønsted-Lowry base toward the hydronium ion, $\mathrm{H}_3\mathrm{O}^+$, likewise identifying the two conjugate pairs.

Two independent aqueous solutions of two monoprotic acids, $\mathrm{HA}$ and $\mathrm{HB}$, are prepared at the same initial analytical concentration $c_0 = 0.10\,\mathrm{mol\times L^{-1}}$ and at $25\,^{\circ}\mathrm{C}$. The diagram shows a microscopic sample of identical volume for both containers, in which a total equivalent of $6$ acid molecules was initially introduced into each solution. Observe the distribution of particles at equilibrium illustrated below:

1. a) From the count of species present at equilibrium (undissociated molecules and ions produced), determine the percentage or degree of dissociation of each acid.
2. b) Deduce, with reasoning, which of the two acids is stronger, justifying whether two solutions with exactly the same initial molar concentration of solute necessarily have the same acidity. Use the ideal model and neglect the autoprotolysis of water where that approximation is valid.