Chemistry » Acid-Base and Redox Reactions » Acids And Bases Continued

Models for Acids and Bases

Models for acids and bases

For the acids you will encounter this year, an acid is a molecule that donates a \(\text{H}^{+}\) ion.

Substances that will act as a base include hydroxides, oxides, carbonates or hydrogen carbonate, among others. Bases often release free hydroxide ions (\(\text{OH}^{-}\)) when dissociating in water.

Arrhenius model for acids and bases

A number of models for acids and bases have been developed over the years. One of the earliest was the Arrhenius definition. In 1884, Arrhenius discovered that water dissociates (splits up) into hydronium \((\text{H}_{3}\text{O}^{+})\) and hydroxide \((\text{OH}^{-})\) ions according to the following equation:

\(2\text{H}_{2}\text{O}(\text{l})\) \(\leftrightharpoons {\color{red}{\text{H}_{3}{\text{O}}^{+}{\text{(aq)}}}} + {\color{blue}{\text{OH}^{-}{\text{(aq)}}}}\)

Another way of writing this is:

\(\text{H}_{2}\text{O}(\text{l})\) \(\leftrightharpoons {\color{red}{\text{H}^{+}{\text{(aq)}}}} + {\color{blue}{\text{OH}^{-}{\text{(aq)}}}}\)

Arrhenius described an acid as a compound that forms \(\text{H}_{3}\text{O}^{+}\) when added to water. An Arrhenius acid therefore increases the concentration of \(\text{H}_{3}\text{O}^{+}\) ions ( [\(\text{H}_{3}\text{O}^{+}\)] ) in water. Arrhenius described a base as a compound that dissociates in water to form \(\text{OH}^{-}\) ions. An Arrhenius base therefore increases the concentration of \(\text{OH}^{-}\) ions ([\(\text{OH}^{-}\)] ) in water.

Definition: Arrhenius acids and bases

An Arrhenius acid forms \(\text{H}_{3}\text{O}^{+}\) in water (increases [\(\text{H}_{3}\text{O}^{+}\)]). An Arrhenius base forms \(\text{OH}^{-}\) in water (increases [\(\text{OH}^{-}\)]).

\(\color{red}{\textbf{Arrhenius acid}}\)

\(\text{increases }[\color{red}{\text{H}_{3}\text{O}^{+}}]\)

\(\color{blue}{\textbf{Arrhenius base}}\)

\(\text{increases }[\color{blue}{\text{OH}^{-}}]\)

Table: The Arrhenius definition of acids and bases.

Look at the following examples showing the dissociation of hydrochloric acid and sodium hydroxide:

  1. \({\color{red}{\text{HCl(aq)}}}\) + \(\text{H}_{2}\text{O}(\text{l})\) \(\to {\color{red}{\text{H}_{3}{\text{O}}^{+}{\text{(aq)}}}}\) + \(\text{Cl}^{-}(\text{aq})\)

    \({\color{red}{\text{Hydrochloric acid}}}\) in water increases the concentration of \({\color{red}{\text{H}_{3}{\text{O}}^{+}}}\) ions and is therefore an \({\color{red}{\textit{acid}}}\).

  2. \({\color{blue}{\text{NaOH(s)}}}\) + \(\text{H}_{2}\text{O}(\text{l})\) \(\to\) \(\text{Na}^{+}(\text{aq})\) + \({\color{blue}{\text{OH}^{-}{\text{(aq)}}}}\) + \(\text{H}_{2}\text{O}(\text{l})\)

    \({\color{blue}{\text{Sodium hydroxide}}}\) in water increases the concentration of \({\color{blue}{\text{OH}^{-}}}\) ions and is therefore a \({\color{blue}{\textit{base}}}\).

However, this definition can only be used for acids and bases in water. Since there are many reactions which do not occur in water, it was important to come up with a much broader definition for acids and bases.

Brønsted-Lowry model for acids and bases

In 1923, Lowry and Brønsted took the work of Arrhenius further to develop a broader definition for acids and bases. The Brønsted-Lowry model defines acids and bases in terms of their ability to donate or accept protons (se table below).

Fact:

Hydrogen atoms contain only one proton. \(\text{H}^{+}\) is a hydrogen atom that has lost its electron and is often called a proton.

Definition: Brønsted-Lowry acids and bases

An acid is a substance that donates (gives away) protons (\(\text{H}^{+}\)). A base is a substance that accepts (takes) protons.

Under the Brønsted-Lowry definition: an \(\color{red}{\text{acid}}\) is a \(\color{red}{\textbf{proton donor}}\); a \(\color{blue}{\text{base}}\) is a \(\color{blue}{\textbf{proton}}\) \(\color{blue}{\textbf{acceptor}}\).

\(\color{red}{\textbf{Brønsted-Lowry acid}}\)

\(\text{donates H}^{+}\)

\(\text{proton }\color{red}{\textbf{donor}}\)

\(\color{blue}{\textbf{Brønsted-Lowry base}}\)

\(\text{accepts H}^{+}\)

\(\text{proton }\color{blue}{\textbf{acceptor}}\)

Table: The Brønsted-Lowry definition of acids and bases.

Below are some examples:

  1. \(\text{HCl}(\text{aq}) + \text{NH}_{3}(\text{g})\) \(\to\) \(\text{NH}_{4}^{+}(\text{aq}) + \text{Cl}^{-}(\text{aq})\)

    In order to decide which substance is a \({\color{red}{\text{proton donor}}}\) and which is a \({\color{blue}{\text{proton acceptor}}}\), we need to look at what happens to each reactant. The reaction can be broken down as follows:

    \(\text{HCl}(\text{aq})\) \(\to\) \({\color{darkgreen}{\text{H}^{+}}}\)(aq) + \(\text{Cl}^{-}(\text{aq})\)

    • \(\text{HCl}\) donates a \(\color{darkgreen}{\text{proton}}\).

      It is a \({\color{red}{\textit{proton donor}}}\) and is therefore the \({\color{red}{\textbf{acid}}}\).

    \(\text{NH}_{3}(\text{g})\) + \({\color{darkgreen}{\text{H}^{+}}}{\text{(aq)}} \to\) \(\text{NH}_{4}^{+}(\text{aq})\)

    • \(\text{NH}_{3}\) accepts a \(\color{darkgreen}{\text{proton}}\).

      It is a \(\color{blue}{\textit{proton acceptor}}\) and is therefore the \(\color{blue}{\textbf{base}}\).

  2. \(\text{CH}_{3}\text{COOH}(\text{aq}) + \text{H}_{2}\text{O}(\text{l})\) \(\to\) \(\text{H}_{3}\text{O}^{+}(\text{aq}) + \text{CH}_{3}\text{COO}^{-}(\text{aq})\)

    The reaction can be broken down as follows:

    \(\text{CH}_{3}\text{COOH}(\text{aq})\) \(\to\) \(\text{CH}_{3}\text{COO}^{-}(\text{aq})\) + \(\color{darkgreen}{\text{H}^{+}}\)(aq)

    • \(\text{CH}_{3}\text{COOH}\) donates a \(\color{darkgreen}{\text{proton}}\).

      It is a \(\color{red}{\textit{proton donor}}\) and is therefore the \(\color{red}{\textbf{acid}}\).

    \(\text{H}_{2}\text{O}(\text{l})\) + \(\color{darkgreen}{\text{H}^{+}}\text{(aq)} \to\) \(\text{H}_{3}\text{O}^{+}(\text{aq})\)

    • Water accepts a \(\color{darkgreen}{\text{proton}}\).

      It is a \(\color{blue}{\textit{proton acceptor}}\) and is therefore the \(\color{blue}{\textbf{base}}\).

  3. \(\text{NH}_{3}(\text{g}) + \text{H}_{2}\text{O}(\text{l})\) \(\to\) \(\text{NH}_{4}^{+}(\text{aq}) + \text{OH}^{-}(\text{aq})\)

    The reaction can be broken down as follows:

    \(\text{H}_{2}\text{O}(\text{l})\) \(\to\) \(\text{OH}^{-}(\text{aq})\) + \(\color{darkgreen}{\text{H}^{+}}\)(aq)

    • Water donates a \(\color{darkgreen}{\text{proton}}\).

      It is a \(\color{red}{\textit{proton donor}}\) and is therefore the \(\color{red}{\textbf{acid}}\).

    \(\text{NH}_{3}(\text{g})\) + \(\color{darkgreen}{\text{H}^{+}}\text{(aq)} \to\) \(\text{NH}_{4}^{+}(\text{aq})\)

    • Ammonia accepts a \(\color{darkgreen}{\text{proton}}\).

      It is a \({\color{blue}{\textit{proton acceptor}}}\) and is therefore the \({\color{blue}{\textbf{base}}}\).

    Notice that in example \(\text{2}\) \(\color{darkgreen}{\textbf{water}}\) acted as a \({\color{blue}{\textbf{base}}}\), while in example \(\text{3}\) \(\color{darkgreen}{\textbf{water}}\) acted as an \({\color{red}{\textbf{acid}}}\). Water can act as both an acid and a base depending on the reaction. A substance that can act as either an acid or a base is called amphoteric.

Definition: Amphoteric

An amphoteric substance is one that can act as an acid in one reaction, or a base in another reaction.

Fact:

An amphoteric substance that contains both acidic and basic functional groups is called an ampholyte.

An amphiprotic substance is an amphoteric substance that can donate a proton in one reaction (a Brønsted-Lowry acid), or accept a proton in another reaction (a Brønsted-Lowry base).

Definition: Amphiprotic

An amphiprotic substance can donate a proton in one reaction, or accept a proton in another reaction.

Substances such as ammonia (\(\text{NH}_{3}\)), zinc oxide (\(\text{ZnO}\)), and beryllium hydroxide (\(\text{Be}(\text{OH})_{2}\)) are amphoteric. Water and ammonia are also amphiprotic.

An acid that releases only one \({\color{darkgreen}{\text{H}^{+}}}\) ion per molecule of acid (e.g. \(\color{darkgreen}{\text{H}}\)Cl) is referred to as monoprotic. An acid that can release two \({\color{darkgreen}{\text{H}^{+}}}\) ions per molecule of acid (e.g. \(\color{darkgreen}{\text{H}_{2}}\)\(\text{SO}_{4}\)) is referred to as diprotic. Any acid than can donate more than one \({\color{darkgreen}{\text{H}^{+}}}\) ion per molecule of acid can be referred to as polyprotic (this means that diprotic acids are also polyprotic).

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