Acids and Bases

An acid is any compound that increases the amount of hydrogen ion ([latex]\ce{H^{+}}[/latex]) in an aqueous solution. The chemical opposite of an acid is a base. The equivalent definition of a base is that a base is a compound that increases the amount of hydroxide ion ([latex]\ce{OH^{-}}[/latex]) in an aqueous solution. These original definitions were proposed by Arrhenius (the same person who proposed ion dissociation) in 1884, so they are referred to as the Arrhenius definitions of an acid and a base, respectively. These definitions simply state that a compound that generates hydrogen ions when dissolved in water is an acid, and a compound that generates hydroxyl ions when dissolved in water is a base.

You may recognise that based on the description of a hydrogen atom, an [latex]\ce{H^{+}}[/latex] ion is a hydrogen atom that has lost its lone electron; that is, [latex]\ce{H^{+}}[/latex] is simply a proton. Do we really have bare protons moving about in aqueous solution? No. What is more likely is that the [latex]\ce{H^{+}}[/latex] ion has attached itself to one (or more) water molecule(s). To represent this chemically, we define the hydronium ion [latex]\ce{H_{3}O^{+}}{(aq)}[/latex], a water molecule with an extra hydrogen ion attached to it, as [latex]\ce{H_{3}O^{+}}[/latex], which represents an additional proton attached to a water molecule. We use the hydronium ion as the more logical way a hydrogen ion appears in an aqueous solution, although in many chemical reactions, [latex]\ce{H^{+}}[/latex] and [latex]\ce{H_{3}O}^{+}[/latex] are treated equivalently.

You will learn about acids and bases in detail in chapter 6.

Neutralisation Reactions

The reaction of an acid and a base is called a neutralisation reaction. Although acids and bases have their own unique chemistries, the acid and base cancel each other’s chemistry to produce a rather innocuous substance—water. In fact, the general reaction between an acid and a base is

[latex]\ce{acid} + \ce{base} → \ce{water} + \ce{salt}[/latex]

where the term salt is generally used to define any ionic compound (soluble or insoluble) that is formed from a reaction between an acid and a base. It should be noted that, in chemistry, the word salt refers to more than just table salt. For example, the balanced chemical equation for the reaction between HCl(aq) and KOH(aq) is

[latex]\ce{HCl}{(aq)} + \ce{KOH}{(aq)} → \ce{H_{2}O}{(ℓ)} + \ce{KCl}{(aq)}[/latex]

where the salt is [latex]\ce{KCl}[/latex]. By counting the number of atoms of each element, we find that only one water molecule is formed as a product. However, in the reaction between [latex]\ce{HCl}{(aq)}[/latex] and [latex]\ce{Mg(OH)_{2}}{(aq)}[/latex], additional molecules of [latex]\ce{HCl}[/latex] and [latex]\ce{H_{2}O}[/latex] are required to balance the chemical equation:

[latex]\ce{2HCl}{(aq)} + \ce{Mg(OH)_{2}}{(aq)} → \ce{2H_{2}O}{(ℓ)} + \ce{MgCl_{2}}{(aq)}[/latex]

Here, the salt is [latex]\ce{MgCl_{2}}[/latex]. This is one of several reactions that take place when a type of antacid—a base—is used to treat stomach acid.

Neutralisation reactions are one type of chemical reaction that proceeds even if one reactant is not in the aqueous phase. For example, the chemical reaction between [latex]\ce{HCl}{(aq)}[/latex] and [latex]\ce{Fe(OH)}_{3}{(s)}[/latex] still proceeds according to the equation:

[latex]\ce{3HCl}{(aq)} + \ce{Fe(OH)}_{3}{(s)} → \ce{3H_{2}O}{(ℓ)} + \ce{FeCl}_{3}{(aq)}[/latex]

even though [latex]\ce{Fe(OH)_{3}}[/latex] is not soluble. When one realises that [latex]\ce{Fe(OH)}_{3}{(s)}[/latex] is a component of rust, this explains why some cleaning solutions for rust stains contain acids—the neutralisation reaction produces products that are soluble and wash away.

Complete and net ionic reactions for neutralisation reactions will depend on whether the reactants and products are soluble, even if the acid and base react. For example, in the reaction of [latex]\ce{HCl}{(aq)}[/latex] and [latex]\ce{NaOH}{(aq)}[/latex]:

[latex]\ce{HCl}{(aq)} + \ce{NaOH}{(aq)} → \ce{H_{2}O}{(ℓ)} + \ce{NaCl}{(aq)}[/latex]

The complete ionic reaction is:

[latex]\ce{H^{+}}{(aq)} + \ce{Cl}^{-}{(aq)} + \ce{Na}^{+}{(aq)} + \ce{OH}^{-}{(aq)} → \ce{H_{2}O}{(ℓ)} + \ce{Na}^{+}{(aq)} + \ce{Cl}^{-}{(aq)}[/latex]

The [latex]\ce{Na}^{+}{(aq)}[/latex] and [latex]\ce{Cl}^{-}{(aq)}[/latex] ions are spectator ions, so we can remove them to have:

[latex]\ce{H}^{+}{(aq)} + \ce{OH}^{-}{(aq)} → \ce{H_{2}O}{(ℓ)}[/latex]

as the net ionic equation. If we wanted to write this in terms of the hydronium ion, [latex]\ce{H_{3}O}^{+}{(aq)}[/latex], we would write it as:

[latex]\ce{H_{3}O}^{+}{(aq)} + \ce{OH}^{-}{(aq)} → \ce{2H_{2}O}{(ℓ)}[/latex]

With the exception of the introduction of an extra water molecule, these two net ionic equations are equivalent.

However, for the reaction between [latex]\ce{HCl}{(aq)}[/latex] and [latex]\ce{Cr(OH)_{2}}{(s)}[/latex], because chromium(II) hydroxide is insoluble, we cannot separate it into ions for the complete ionic equation:

[latex]\ce{2H}^{+}{(aq)} + \ce{2Cl}^{-}{(aq)} + \ce{Cr(OH)_{2}}{(s)} → \ce{2H_{2}O}{(ℓ)} + \ce{Cr}^{2+}{(aq)} + \ce{2Cl}^{-}{(aq)}[/latex]

The chloride ions are the only spectator ions here, so the net ionic equation is:

[latex]\ce{2H}^{+}{(aq)} + \ce{Cr(OH)_{2}}{(s)} → \ce{2H_{2}O}{(ℓ)} + \ce{Cr}^{2+}{(aq)}[/latex]