Complete and write the overall equation, the complete ionic equation, and the net ionic equation for the following acid-base reactions. If the substance is a weak acid or base, leave it in its molecular form in writing the equations. (a) \(\mathrm{HCl}\) (aq) \(+\mathrm{NaOH}\) (aq) \(\rightarrow\) (b) \(\mathrm{NH}_{3}\) (aq) \(+\mathrm{HNO}_{3}\) (aq) \(\rightarrow\) (c) \(\mathrm{CH}_{3} \mathrm{NH}_{2}\) (aq) \(+\mathrm{Hl}\) (aq) \(\rightarrow\)

Acid-base reactions are fundamental chemical interactions where an acid and a base react to form water and a salt.

An acidic substance is one that is capable of donating a proton (\text{H}^+), according to the Bronsted-Lowry definition. An example is hydrochloric acid (\text{HCl}). Bases, on the other hand, are substances that can accept protons, such as sodium hydroxide (\text{NaOH}).

When these substances are in aqueous solution, they react according to the equation: \[ \text{Acid} + \text{Base} \rightarrow \text{Water} + \text{Salt} \]
For example, the reaction of hydrochloric acid with sodium hydroxide can be written as:\[ \text{HCl} + \text{NaOH} \rightarrow \text{H}_2\text{O} + \text{NaCl} \]
It's critical to remember that the strength of an acid or base varies. Strong acids and bases dissociate completely in solution, whereas weak acids and bases do not. When writing reaction equations, strong acids and bases are written as ions, while weak ones are left in their molecular form.

Diving into the realm of the complete ionic equation, it's important to differentiate it from a simple balanced chemical equation.

A complete ionic equation shows all of the ions present in the solution before and after the chemical reaction. To write a complete ionic equation, one must consider the dissociation of soluble ionic compounds into their respective ions.

For the reaction between \text{HCl} and \text{NaOH}, the complete ionic equation would look like this:\[ \text{H}^+_{(aq)} + \text{Cl}^-_{(aq)} + \text{Na}^+_{(aq)} + \text{OH}^-_{(aq)} \rightarrow \text{H}_2\text{O}_{(l)} + \text{Na}^+_{(aq)} + \text{Cl}^-_{(aq)} \]
Each soluble strong electrolyte is written separately as ions. However, insoluble salts, weak acids, or bases as well as liquids, solids, and gases are shown in their complete form. Understanding this distinction is crucial for accuracy in representing chemical reactions.

Now, let's turn the spotlight to spectator ions, an interesting component in the chemical drama of reactions.

Spectator ions are like the bystanders of a reaction—present but not directly involved in the chemical change. These are ions that remain unchanged on both the reactant and the product side of a complete ionic equation.

In the context of our example with \text{HCl} and \text{NaOH}, the sodium (\text{Na}^+) and chloride (\text{Cl}^-) ions do not participate in the reaction that forms water. They start and finish in the same form, merely observing the reaction take place. Thus, when we write the net ionic equation, we leave these spectator ions out and focus on the truly reactive species:\[ \text{H}^+_{(aq)} + \text{OH}^-_{(aq)} \rightarrow \text{H}_2\text{O}_{(l)} \]
Identifying and omitting spectator ions simplifies the equation and allows us to see the actual chemical change occurring, known as the net ionic reaction. This is an essential step for understanding reactions at the ionic level.