Comprehensive Study Guide to Aqueous Acids, Bases, and Their Chemical Theories

Properties of Aqueous Solutions of Acids and Bases

Aqueous acidic solutions exhibit several distinct physical and chemical properties. From a sensory perspective, they are characterized by a sour taste. Chemically, they have the ability to alter the colors of various indicators; for instance, acids will turn blue litmus paper red and will change the color of bromothymol blue from blue to yellow. Furthermore, acidic solutions react with metals to generate hydrogen gas, represented as $H_2(g)$. They also react with metal oxides and hydroxides to produce salts and water. In the presence of salts derived from weaker acids, a stronger acid will react to form the weaker acid and the salt of the stronger acid. Finally, aqueous acidic solutions are electrolytic, meaning they conduct electricity.

Aqueous basic solutions, on the other hand, possess a bitter taste and a characteristically slippery feeling to the touch. Like acids, bases interact with indicators, turning red litmus paper blue and changing bromothymol blue from yellow to blue. Bases participate in neutralization reactions by reacting with acids to form salts and water. Similar to acidic solutions, aqueous basic solutions are capable of conducting electricity.

The Arrhenius Theory

Presented in 1884 by Svante Augustus Arrhenius, this theory provides a fundamental definition of acids and bases. According to Arrhenius, acids are substances that contain hydrogen and produce hydrogen ions, represented as $H^+$, when dissolved in aqueous solutions. Bases are defined as substances that contain the hydroxyl group, $OH$, and produce hydroxide ions, $OH^-$, in aqueous solutions. Within this framework, neutralization reactions are understood as the combination of $H^+$ (often represented as the hydronium ion, $H_3O^+$) with $OH^-$ to produce water, $H_2O$.

Acids and bases are further classified by their strength based on their level of ionization in water. Strong acids are those that ionize $100\%$ in aqueous solutions. The recognized list of aqueous strong acids includes hydrochloric acid ($HCl$), hydrobromic acid ($HBr$), hydroiodic acid ($HI$), sulfuric acid ($H_2SO_4$), nitric acid ($HNO_3$), perchloric acid ($HClO_4$), and chloric acid ($HClO_3$). Similarly, strong bases ionize $100\%$ in water. The list of aqueous strong bases includes lithium hydroxide ($LiOH$), sodium hydroxide ($NaOH$), potassium hydroxide ($KOH$), rubidium hydroxide ($RbOH$), cesium hydroxide ($CsOH$), calcium hydroxide ($Ca(OH)_2$), strontium hydroxide ($Sr(OH)_2$), and barium hydroxide ($Ba(OH)_2$).

The BrØnsted-Lowry Theory

Developed in 1923 by J.N. BrØnsted and T.M. Lowry, this theory offers a more generalized perspective than the Arrhenius model. It defines an acid as a proton ($H^+$) donor and a base as a proton acceptor. Consequently, acid-base reactions are characterized as the transfer of a proton from an acid to a base. It is important to note that coordinate covalent bonds are frequently formed during these reactions. One of the primary advantages of this theory over the Arrhenius theory is that reactions are not required to occur in aqueous solutions, and bases are not required to contain hydroxide groups.

A central component of this theory is the concept of conjugate acid-base pairs, which are two species that differ only by a single proton. In the reaction $HNO_3 + H_2O \rightarrow H_3O^+ + NO_3^-$, the reactant acid is $HNO_3$ and its conjugate base is $NO_3^-$. Conversely, the reactant base is $H_2O$ and its conjugate acid is $H_3O^+$. In the standard format for writing these pairs, such as in the equilibrium $HF + H_2O \rightleftharpoons H_3O^+ + F^-$, the pairs are labeled with subscripts: pair 1 consists of the acid $HF$ and its conjugate base $F^-$, while pair 2 consists of the base $H_2O$ and its conjugate acid $H_3O^+$.

The theory also dictates a relationship between the relative strengths of these pairs. For example, since ammonia ($NH_3$) is a weak base, its conjugate acid, the ammonium ion ($NH_4^+$), must be a strong acid because it readily gives up a proton to reform $NH_3$. In contrast, because sodium hydroxide ($NaOH$) is a strong base that ionizes $100\%$ in solution, its conjugate acid ($Na^+$) must be a very weak acid, as it does not recombine to form $NaOH$.

Autoionization and Amphoterism

Water undergoes a process known as autoionization, though it does not occur extensively. In pure water, the concentrations of hydronium and hydroxide ions are equal, such that $[H_3O^+] = [OH^-]$, and both are approximately equal to $1.0 \times 10^{-7}\,M$. This autoionization process serves as the foundational basis for the $pH$ scale.

Certain species possess the ability to behave as both an acid and a base, a property known as amphoterism. Species with this dual capability are described as amphoteric. When specifically referring to proton transfer reactions where a species can either donate or accept a proton, the term amphiprotic is used.

Strengths of Acids

For binary acids, the strength of the acid is determined by the strength of the $H-X$ bond. As the bond strength decreases, the acid strength increases. This is observed in the hydrohalic binary acids. The periodic trend for bond strength is $HF \gg HCl > HBr > HI$. Consequently, the trend for acid strength is the reverse: $HF \ll HCl < HBr < HI$. This allows for a relative ranking of acid and base strengths based on BrØnsted-Lowry theory.

According to Table 10-2, acids like $HClO_4$, $HI$, $HBr$, $HCl$, and $HNO_3$ are $100\%$ ionized in dilute aqueous solutions and have negligible conjugate base strength. Acids such as $H_3O^+$, $HF$, $CH_3COOH$, $HCN$, and $NH_4^+$ exist in an equilibrium mixture of nonionized molecules and conjugate bases. The strongest bases, such as the amide ion ($NH_2^-$), react completely with water to form $OH^-$ and cannot exist in aqueous solution.

The strength of ternary acids (oxyacids) typically increases with an increasing number of Oxygen ($O$) atoms attached to the central atom, which correlates with an increasing oxidation state of the central atom. Each additional oxygen atom increases the oxidation state of the central atom by $2$. This phenomenon explains why ternary acid strength follows these specific trends.

The Lewis Theory

The Lewis Theory provides another alternative for defining acid-base behavior based on electron pairs rather than protons. An example of a Lewis acid-base interaction is the ionization of ammonia: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$, where ammonia acts as the base. Another classic example involves Boron trifluoride ($BF_3$), which is a strong Lewis acid and can react with ammonia ($NH_3$).

The Preparation of Acids

Binary acids can be prepared by reacting a nonmetallic element directly with hydrogen gas ($H_2$). For example, the reaction $H_2(g) + Cl_2(g) \rightarrow 2HCl(g)$ is performed in the presence of ultraviolet ($UV$) light. Volatile acids, which are those with low boiling points, are prepared by reacting salts with a nonvolatile acid such as concentrated sulfuric acid ($H_2SO_4$) or phosphoric acid ($H_3PO_4$). Examples include: $NaCl(s) + H_2SO_4(conc.) \rightarrow NaHSO_4(s) + HCl(g)$$NaF(s) + H_2SO_4(conc.) \rightarrow NaHSO_4(s) + HF(g)$

To prepare hydrobromic ($HBr$) and hydroiodic ($HI$) acids, phosphoric acid must be used instead of sulfuric acid to avoid unwanted oxidation reactions: $NaBr(s) + H_3PO_4(conc.) \rightarrow NaH_2PO_4(s) + HBr(g)$$NaI(s) + H_3PO_4(conc.) \rightarrow NaH_2PO_4(s) + HI(g)$

Ternary acids are synthesized by reacting nonmetal oxides (also known as acid anhydrides) with water. Representative reactions include: $SO_2(g) + H_2O(l) \rightarrow H_2SO_3(aq)$$N_2O_5(g) + H_2O(l) \rightarrow 2HNO_3(aq)$

Additionally, some nonmetal halides and oxyhalides react with water to yield both a binary and a ternary acid simultaneously: $PCl_5(s) + 4H_2O(l) \rightarrow H_3PO_4(aq) + 5HCl(aq)$$POCl_3(l) + 3H_2O(l) \rightarrow H_3PO_4(aq) + 3HCl(aq)$

Questions & Discussion

Question: Identify the reactant acid and base in the reaction $HNO_3 + H_2O \rightarrow H_3O^+ + NO_3^-$. Response: In this reaction, $HNO_3$ is the reactant acid and $H_2O$ is the reactant base.

Question: Identify the species that differs from the acid by a proton in the reaction $HNO_3 + H_2O \rightarrow H_3O^+ + NO_3^-$. Response: The species that differs from the acid $HNO_3$ by a proton is the conjugate base, $NO_3^-$.