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16.1: Acids and Bases: A Brief Review

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    25225
  • Acids and bases have been known for a long time. When Robert Boyle characterized them in 1680, he noted that acids dissolve many substances, change the color of certain natural dyes (for example, they change litmus from blue to red), and lose these characteristic properties after coming into contact with alkalis (bases). In the eighteenth century, it was recognized that acids have a sour taste, react with limestone to liberate a gaseous substance (now known to be CO2), and interact with alkalis to form neutral substances. In 1815, Humphry Davy contributed greatly to the development of the modern acid-base concept by demonstrating that hydrogen is the essential constituent of acids. Around that same time, Joseph Louis Gay-Lussac concluded that acids are substances that can neutralize bases and that these two classes of substances can be defined only in terms of each other. The significance of hydrogen was reemphasized in 1884 when Carl Axel Arrhenius defined an acid as a compound that dissolves in water to yield hydrogen cations (now recognized to be hydronium ions) and a base as a compound that dissolves in water to yield hydroxide anions.

    Acids and bases are common solutions that exist everywhere. Almost every liquid that we encounter in our daily lives consists of acidic and basic properties, with the exception of water. They have completely different properties and are able to neutralize to form H2O, which will be discussed later in a subsection. Acids and bases can be defined by their physical and chemical observations (Table \(\PageIndex{1}\)).

    ACIDS BASES
    Table \(\PageIndex{1}\): General Properties of Acids and Bases
    produce a piercing pain in a wound. give a slippery feel.
    taste sour. taste bitter.
    are colorless when placed in phenolphthalein (an indicator). are pink when placed in phenolphthalein (an indicator).
    are red on blue litmus paper (a pH indicator). are blue on red litmus paper (a pH indicator).
    have a pH<7. have a pH>7.
    produce hydrogen gas when reacted with metals.  
    produce carbon dioxide when reacted with carbonates.  
    Common examples: Lemons, oranges, vinegar, urine, sulfuric acid, hydrochloric acid

    Common Examples: Soap, toothpaste, bleach, cleaning agents, limewater, ammonia water, sodium hydroxide.

    Acids and bases in aqueous solutions will conduct electricity because they contain dissolved ions. Therefore, acids and bases are electrolytes. Strong acids and bases will be strong electrolytes. Weak acids and bases will be weak electrolytes. This affects the amount of conductivity.

    The Arrhenius Definition of Acids and Bases

    In 1884, the Swedish chemist Svante Arrhenius proposed two specific classifications of compounds, termed acids and bases. When dissolved in an aqueous solution, certain ions were released into the solution. The Arrhenius definition of acid-base reactions is a development of the "hydrogen theory of acids". It was used to provide a modern definition of acids and bases, and followed from Arrhenius's work with Friedrich Wilhelm Ostwald in establishing the presence of ions in aqueous solution in 1884. This led to Arrhenius receiving the Nobel Prize in Chemistry in 1903.

    An Arrhenius acid is a compound that increases the concentration of \(H^+\) ions that are present when added to water. These \(H^+\) ions form the hydronium ion (\(H_3O^+\)) when they combine with water molecules. This process is represented in a chemical equation by adding H2O to the reactants side.

    \[ HCl_{(aq)} \rightarrow H^+_{(aq)} + Cl^-_{(aq)} \]

    In this reaction, hydrochloric acid (\(HCl\)) dissociates into hydrogen (\(H^+\)) and chlorine (\(Cl^-\)) ions when dissolved in water, thereby releasing H+ ions into solution. Formation of the hydronium ion equation:

    \[ HCl_{(aq)} + H_2O_{(l)} \rightarrow H_3O^+_{(aq)} + Cl^-_{(aq)} \]

    The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions and refer to the concentration of the solvated ions. Under this definition, pure \(H_2SO_4\) or \(HCl\) dissolved in toluene are not acidic, despite the fact that both of these acids will donate a proton to toluene. In addition, under the Arrhenius definition, a solution of sodium amide (\(NaNH_2\)) in liquid ammonia is not alkaline, despite the fact that the amide ion (\(NH^−_2\)) will readily deprotonate ammonia. Thus, the Arrhenius definition can only describe acids and bases in an aqueous environment.

    Limitation of the Arrhenius Definition of Acids and Bases

    The Arrhenius definition can only describe acids and bases in an aqueous environment.

    In chemistry, acids and bases have been defined differently by three sets of theories: One is the Arrhenius definition defined above, which revolves around the idea that acids are substances that ionize (break off) in an aqueous solution to produce hydrogen (\(H^+\)) ions while bases produce hydroxide (\(OH^-\)) ions in solution. The other two definitions are discussed in detail alter in the chapter and include the Brønsted-Lowry definition the defines acids as substances that donate protons (\(H^+\)) whereas bases are substances that accept protons and the Lewis theory of acids and bases states that acids are electron pair acceptors while bases are electron pair donors.

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