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Chemistry LibreTexts

17: Acids and Bases

  • Page ID
    219264
    • 17.1: Heartburn
      Heartburn is caused by a buildup of excessive amounts of stomach acid, particularly HCl. This acid is used to digest the food we eat, but it can often back up into the esophagus causing that burning sensation many of us are familiar with.
    • 17.2: The Nature of Acids and Bases
      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 and the Lewis theory.
    • 17.3: Definitions of Acids and Bases
      A compound that can donate a proton (a hydrogen ion) to another compound is called a Brønsted-Lowry acid. The compound that accepts the proton is called a Brønsted-Lowry base. The species remaining after a Brønsted-Lowry acid has lost a proton is the conjugate base of the acid. The species formed when a Brønsted-Lowry base gains a proton is the conjugate acid of the base. Thus, an acid-base reaction occurs when a proton is transferred from an acid to a base.
    • 17.4: Acid Strength and the Acid Dissociation Constant (Ka)
      Acid–base reactions always contain two conjugate acid–base pairs. Each acid and each base has an associated ionization constant that corresponds to its acid or base strength. Two species that differ by only a proton constitute a conjugate acid–base pair. The magnitude of the equilibrium constant for an ionization reaction can be used to determine the relative strengths of acids and bases.
    • 17.5: Autoionization of Water and pH
      Water is amphiprotic: it can act as an acid by donating a proton to a base to form the hydroxide ion, or as a base by accepting a proton from an acid to form the hydronium ion ( H3O+ ). The autoionization of liquid water produces OH− and H3O+ ions. The equilibrium constant for this reaction is called the ion-product constant of liquid water (Kw) and is defined as Kw=[H3O+][OH−] . At 25°C, Kw is 1.01×10−14 ; hence pH+pOH=pKw=14.00 .
    • 17.6: Finding the [H3O+] and pH of Strong and Weak Acid Solutions
      Acid–base reactions always contain two conjugate acid–base pairs. Each acid and each base has an associated ionization constant that corresponds to its acid or base strength. Two species that differ by only a proton constitute a conjugate acid–base pair. The magnitude of the equilibrium constant for an ionization reaction can be used to determine the relative strengths of acids and bases.
    • 17.7: Base Solutions
    • 17.8: The Acid-Base Properties of Ions and Salts
      A salt can dissolve in water to produce a neutral, a basic, or an acidic solution, depending on whether it contains the conjugate base of a weak acid as the anion ( A−A− ), the conjugate acid of a weak base as the cation ( BH+ ), or both. Salts that contain small, highly charged metal ions produce acidic solutions in water. The reaction of a salt with water to produce an acidic or a basic solution is called a hydrolysis reaction.
    • 17.9: Polyprotic Acids
      An acid that contains more than one ionizable proton is a polyprotic acid. The protons of these acids ionize in steps. The differences in the acid ionization constants for the successive ionizations of the protons in a polyprotic acid usually vary by roughly five orders of magnitude. As long as the difference between the successive values of Ka of the acid is greater than about a factor of 20, it is appropriate to break down the calculations of the concentrations sequentially.
    • 17.10: Acid Strength and Molecular Structure
      Inductive effects and charge delocalization significantly influence the acidity or basicity of a compound. The acid–base strength of a molecule depends strongly on its structure. The weaker the A–H or B–H+ bond, the more likely it is to dissociate to form an H+H+ ion. In addition, any factor that stabilizes the lone pair on the conjugate base favors the dissociation of H+H+ , making the conjugate acid a stronger acid.
    • 17.11: Lewis Acids and Bases
      Lewis proposed that the electron pair is the dominant actor in acid-base chemistry. An Lewis acid is a substance that accepts a pair of electrons, and in doing so, forms a covalent bond with the entity that supplies the electrons. A Lewis base is a substance that donates an unshared pair of electrons to a recipient species with which the electrons can be shared. Lewis acis/base theory is a powerful tool for describing many chemical reactions used in organic and inorganic chemistry.
    • 17.12: Acid rain
      The damaging effects of acid rain have led to strong pressure on industry to minimize the release of harmful reactants. Acid rain is rainfall whose pH is less than 5.6, the value typically observed, due to the presence of dissolved carbon dioxide. Acid rain is caused by nitrogen oxides and sulfur dioxide produced by both natural processes and the combustion of fossil fuels. Eventually, these oxides react with oxygen and water to give nitric acid and sulfuric acid.

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