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4.4: Classifying Chemical Reactions - Acid Base Reactions

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    Learning Objectives
    • Identify common acids and bases
    • Identify when an acid-base reaction will occur
    • Write a balanced chemical equation that describes what happens when a acid-base reaction occurs
    • Solve stoichiometry problems with acid-base reactions

    Humans interact with one another in various and complex ways, and we classify these interactions according to common patterns of behavior. When two humans exchange information, we say they are communicating. When they exchange blows with their fists or feet, we say they are fighting. Faced with a wide range of varied interactions between chemical substances, scientists have likewise found it convenient (or even necessary) to classify chemical interactions by identifying common patterns of reactivity. This module will provide an introduction to acid-base reactions, one of the most prevalent types of chemical reactions.

    Acid Reactions

    An acid-base reaction is one in which a hydrogen ion, H+, is transferred from one chemical species to another. Such reactions are of central importance to numerous natural systems and manufacturing processes, ranging from the chemical transformations inside cells, lakes, and oceans; to the industrial-scale production of fertilizers, pharmaceuticals, and manufacturing materials.

    For purposes of this brief introduction, we will only consider the acid-base reactions that take place in aqueous solutions. In this context, an acid is a substance that dissolves in water to yield hydronium ions, H3O+. As an example, consider the equation shown here for hydrochloric acid:

    \[\ce{HCl}(aq)+\ce{H2O}(aq)\rightarrow \ce{Cl-}(aq)+\ce{H3O+}(aq)\]

    The process represented by this equation shows hydrogen chloride is an acid. When dissolved in water, H+ ions are transferred from HCl molecules to H2O molecules to generate hydronium, H3O+, ions as shown in Figure \(\PageIndex{2}\).

    This figure shows two flasks, labeled a and b. The flasks are both sealed with stoppers and are nearly three-quarters full of a liquid. Flask a is labeled H C l followed by g in parentheses. In the liquid there are approximately twenty space-filling molecular models composed of one red sphere and two smaller attached white spheres. The label H subscript 2 O followed by a q in parentheses is connected with a line to one of these models. In the space above the liquid in the flask, four space filling molecular models composed of one larger green sphere to which a smaller white sphere is bonded are shown. To one of these models, the label H C l followed by g in parentheses is attached with a line segment. An arrow is drawn from the space above the liquid pointing down into the liquid below. Flask b is labeled H subscript 3 O superscript positive sign followed by a q in parentheses. This is followed by a plus sign and C l superscript negative sign which is also followed by a q in parentheses. In this flask, no molecules are shown in the open space above the liquid. A label, C l superscript negative sign followed by a q in parentheses, is connected with a line segment to a green sphere. This sphere is surrounded by four molecules composed each of one red sphere and two white smaller spheres. A few of these same molecules appear separate from the green spheres in the liquid. A line segment connects one of them to the label H subscript 2 O which is followed by l in parentheses. There are a few molecules formed from one central larger red sphere to which three smaller white spheres are bonded. A line segment is drawn from one of these to the label H subscript 3 O superscript positive sign, followed by a q in parentheses.
    Figure \(\PageIndex{2}\): When hydrogen chloride gas dissolves in water, (a) it reacts as an acid, transferring protons to water molecules to yield (b) hydronium ions (and solvated chloride ions)

    The nature of HCl is such that its reaction with water as just described is essentially 100% efficient: Virtually every HCl molecule that dissolves in water will undergo this reaction. Acids that completely react in this fashion are called strong acids, and HCl is one among just a handful of common acid compounds that are classified as strong (Table \(\PageIndex{1}\)).

    Table \(\PageIndex{2}\): Common Strong Acids
    Compound Formula Name in Aqueous Solution
    HBr hydrobromic acid
    HCl hydrochloric acid
    HNO3 nitric acid
    H2SO4 sulfuric acid

     

    A far greater number of compounds behave as weak acids and only partially react with water, leaving a large majority of dissolved molecules in their original form and generating a relatively small amount of hydronium ions. Weak acids are commonly encountered in nature.  Examples include citric acid, the tangy taste of citrus fruits, formic acid, the stinging sensation of insect bites, and acetic acid, the main ingredient in vinegar.  The reaction of acetic acid with water to produce acetate ions and hydronium ions is shown below.  :

    \[\ce{CH3CO2H}(aq)+\ce{H2O}(l)\rightleftharpoons \ce{CH3CO2-}(aq)+\ce{H3O+}(aq)\]

    Notice the double arrow used to represent the equilibrium reaction where only some of the acetic acid molecules dissociate.  When dissolved in water under typical conditions, only about 1% of acetic acid molecules are present in the ionized form, \(\ce{CH3CO2-}\). The double-arrow in the equation indicates that the reaction goes in both directions, which is why only some of the molecules are present in the ionized form. Other week acids are shown in Figure \(\PageIndex{3}\).

    This figure contains two images, each with an associated structural formula provided in the lower left corner of the image. The first image is a photograph of a variety of thinly sliced, circular cross sections of citrus fruits ranging in color for green to yellow, to orange and reddish-orange. The slices are closely packed on a white background. The structural formula with this picture shows a central chain of five C atoms. The leftmost C atom has an O atom double bonded above and to the left and a singly bonded O atom below and to the left. This single bonded O atom has an H atom indicated in red on its left side which is highlighted in pink. The second C atom moving to the right has H atoms bonded above and below. The third C atom has a single bonded O atom above which has an H atom on its right. This third C atom has a C atom bonded below it which has an O atom double bonded below and to the left and a singly bonded O atom below and to the right. An H atom appears in red and is highlighted in pink to the right of the singly bonded O atom. The fourth C atom has H atoms bonded above and below. The fifth C atom is at the right end of the structure. It has an O atom double bonded above and to the right and a singly bonded O atom below and to the right. This single bonded O atom has a red H atom on its right side which is highlighted in pink. The second image is a photograph of bottles of vinegar. The bottles are labeled, “Balsamic Vinegar,” and appear to be clear and colorless. The liquid in this bottle appears to be brown. The structural formula that appears with this image shows a chain of two C atoms. The leftmost C atom has H atoms bonded above, below, and to the left. The C atom on the right has a doubly bonded O atom above and to the right and a singly bonded O atom below and to the right. This O atom has an H atom bonded to its right which is highlighted in pink.
    Figure \(\PageIndex{3}\): (a) Fruits such as oranges, lemons, and grapefruit contain the weak acid citric acid. (b) Vinegars contain the weak acid acetic acid. (credit a: modification of work by Scott Bauer; credit b: modification of work by Brücke-Osteuropa/Wikimedia Commons)

    Base Reactions

    A base is a substance that will dissolve in water to yield hydroxide ions, OH−. The most common bases are ionic compounds composed of alkali or alkaline earth metal cations (groups 1 and 2) combined with the hydroxide ion—for example, NaOH and Ca(OH)2. When these compounds dissolve in water, hydroxide ions are released directly into the solution. For example, KOH and Ba(OH)2 dissolve in water and dissociate completely to produce cations (K+ and Ba2+, respectively) and hydroxide ions, OH−. These bases, along with other hydroxides that completely dissociate in water, are considered strong bases.

    Consider as an example the dissolution of lye (sodium hydroxide) in water:

    \[\ce{NaOH}(s)\rightarrow \ce{Na+}(aq)+\ce{OH-}(aq)\]

    This equation confirms that sodium hydroxide is a base. When dissolved in water, NaOH dissociates to yield Na+ and OH− ions. This is also true for any other ionic compound containing hydroxide ions. Since the dissociation process is essentially complete when ionic compounds dissolve in water under typical conditions, NaOH and other ionic hydroxides are all classified as strong bases.

    Unlike ionic hydroxides, some compounds produce hydroxide ions when dissolved by chemically reacting with water molecules. In all cases, these compounds react only partially and so are classified as weak bases. These types of compounds are also abundant in nature and important commodities in various technologies. For example, global production of the weak base ammonia is typically well over 100 million metric tons annually, being widely used as an agricultural fertilizer, a raw material for chemical synthesis of other compounds, and an active ingredient in household cleaners (Figure \(\PageIndex{4}\)). When dissolved in water, ammonia reacts partially to yield hydroxide ions, as shown here:

    \[\ce{NH3}(aq)+\ce{H2O}(l)\rightleftharpoons \ce{NH4+}(aq)+\ce{OH-}(aq)\]

    This is, by definition, an acid-base reaction, in this case involving the transfer of H+ ions from water molecules to ammonia molecules. Under typical conditions, only about 1% of the dissolved ammonia is present as \(\ce{NH4+}\) ions.

    This photograph shows a large agricultural tractor in a field pulling a field sprayer and a large, white cylindrical tank which is labeled “Caution Ammonia.”
    Figure \(\PageIndex{4}\): Ammonia is a weak base used in a variety of applications. (a) Pure ammonia is commonly applied as an agricultural fertilizer. (b) Dilute solutions of ammonia are effective household cleansers. (credit a: modification of work by National Resources Conservation Service; credit b: modification of work by pat00139)

    Acid-Base Reactions

    The chemical reactions described in which acids and bases dissolved in water produce hydronium and hydroxide ions, respectively, are, by definition, acid-base reactions. In these reactions, water serves as both a solvent and a reactant. A neutralization reaction is a specific type of acid-base reaction in which the reactants are an acid and a base, the products are often a salt and water, and neither reactant is the water itself:

    \[\mathrm{acid+base\rightarrow salt+water}\]

    To illustrate a neutralization reaction, consider what happens when a typical antacid such as milk of magnesia (an aqueous suspension of solid Mg(OH)2) is ingested to ease symptoms associated with excess stomach acid (HCl):

    \[\ce{Mg(OH)2}(s)+\ce{2HCl}(aq)\rightarrow \ce{MgCl2}(aq)+\ce{2H2O}(l).\]

    Note that in addition to water, this reaction produces a salt, magnesium chloride.

    Example \(\PageIndex{2}\): Writing Equations for Acid-Base Reactions

    Write balanced chemical equations for the acid-base reactions described here:

    1. the weak acid hydrogen hypochlorite reacts with water
    2. a solution of barium hydroxide is neutralized with a solution of nitric acid

    Solution

    (a) The two reactants are provided, HOCl and H2O. Since the substance is reported to be an acid, its reaction with water will involve the transfer of H+ from HOCl to H2O to generate hydronium ions, H3O+ and hypochlorite ions, OCl−.

    \[\ce{HOCl}(aq)+\ce{H2O}(l)\rightleftharpoons \ce{OCl-}(aq)+\ce{H3O+}(aq) \nonumber \]

    A double-arrow is appropriate in this equation because it indicates the HOCl is a weak acid that has not reacted completely.

    (b) The two reactants are provided, Ba(OH)2 and HNO3. Since this is a neutralization reaction, the two products will be water and a salt composed of the cation of the ionic hydroxide (Ba2+) and the anion generated when the acid transfers its hydrogen ion \(\ce{(NO3- )}\).

    \[\ce{Ba(OH)2}(aq)+\ce{2HNO3}(aq)\rightarrow \ce{Ba(NO3)2}(aq)+\ce{2H2O}(l) \nonumber \]

    Exercise \(\PageIndex{21}\)

    Write the net ionic equation representing the neutralization of any strong acid with an ionic hydroxide. (Hint: Consider the ions produced when a strong acid is dissolved in water.)

    Answer

    \[\ce{H3O+}(aq)+\ce{OH-}(aq)\rightarrow \ce{2H2O}(l) \nonumber\]

    In the 18th century, the strength (actually the concentration) of vinegar samples was determined by noting the amount of potassium carbonate, K2CO3, which had to be added, a little at a time, before bubbling ceased. The greater the weight of potassium carbonate added to reach the point where the bubbling ended, the more concentrated the vinegar.

    We now know that the effervescence that occurred during this process was due to reaction with acetic acid, CH3CO2H, the compound primarily responsible for the odor and taste of vinegar. Acetic acid reacts with potassium carbonate according to the following equation:

    \[\ce{2CH3CO2H}(aq)+\ce{K2CO3}(s)\rightarrow 2 \ce{KCH3CO3}(aq)+\ce{CO2}(g)+\ce{H2O}(l) \nonumber \]

    The bubbling was due to the production of CO2.

    The test of vinegar with potassium carbonate is one type of quantitative analysis—the determination of the amount or concentration of a substance in a sample. In the analysis of vinegar, the concentration of the solute (acetic acid) was determined from the amount of reactant that combined with the solute present in a known volume of the solution. In other types of chemical analyses, the amount of a substance present in a sample is determined by measuring the amount of product that results.

    Titration

    The described approach to measuring vinegar strength was an early version of the analytical technique known as titration analysis. A typical titration analysis involves the use of a buret (Figure \(\PageIndex{1}\)) to make incremental additions of a solution containing a known concentration of some substance (the titrant) to a sample solution containing the substance whose concentration is to be measured (the analyte). The titrant and analyte undergo a chemical reaction of known stoichiometry, and so measuring the volume of titrant solution required for complete reaction with the analyte (the equivalence point of the titration) allows calculation of the analyte concentration. The equivalence point of a titration may be detected visually if a distinct change in the appearance of the sample solution accompanies the completion of the reaction. The halt of bubble formation in the classic vinegar analysis is one such example, though, more commonly, special dyes called indicators are added to the sample solutions to impart a change in color at or very near the equivalence point of the titration. Equivalence points may also be detected by measuring some solution property that changes in a predictable way during the course of the titration. Regardless of the approach taken to detect a titration’s equivalence point, the volume of titrant actually measured is called the end point. Properly designed titration methods typically ensure that the difference between the equivalence and end points is negligible. Though any type of chemical reaction may serve as the basis for a titration analysis, precipitation and acid-base reactions are most common.

    Two pictures are shown. In a, a person is shown pouring a liquid from a small beaker into a buret. The person is wearing goggles and gloves as she transfers the solution into the buret. In b, a close up view of the markings on the side of the buret is shown. The markings for 10, 15, and 20 are clearly shown with horizontal rings printed on the buret. Between each of these whole number markings, half markings are also clearly shown with horizontal line segment markings.
    Figure \(\PageIndex{1}\): (a) A student fills a buret in preparation for a titration analysis. (b) A typical buret permits volume measurements to the nearest 0.1 mL. (credit a: modification of work by Mark Blaser and Matt Evans; credit b: modification of work by Mark Blaser and Matt Evans)
    Example \(\PageIndex{1}\): Titration Analysis

    The end point in a titration of a 50.00-mL sample of aqueous HCl was reached by addition of 35.23 mL of 0.250 M NaOH titrant. The titration reaction is:

    \[\ce{HCl}(aq)+\ce{NaOH}(aq)\rightarrow \ce{NaCl}(aq)+\ce{H2O}(l) \nonumber \]

    What is the molarity of the HCl?

    Solution

    As for all reaction stoichiometry calculations, the key issue is the relation between the molar amounts of the chemical species of interest as depicted in the balanced chemical equation. The approach outlined in previous modules of this chapter is followed, with additional considerations required, since the amounts of reactants provided and requested are expressed as solution concentrations.

    For this exercise, the calculation will follow the following outlined steps:

    This figure shows four rectangles. The first is shaded lavender and is labeled, “Volume of N a O H.” This rectangle is followed by an arrow pointing right which is labeled, “Molar concentration,” to a second rectangle. This second rectangle is shaded pink and is labeled, “Moles of N a O H.” This rectangle is followed by an arrow pointing right which is labeled, “Stoichiometric factor,” to a third rectangle which is shaded pink and is labeled, “Moles of H C l.” This rectangle is followed by an arrow labeled, “Solution volume,” which points right to a fourth rectangle. This fourth rectangle is shaded lavender and is labeled, “Concentration of H C l.”

    The molar amount of HCl is calculated to be:

    \[\mathrm{35.23\:\cancel{mL\: NaOH}\times \dfrac{1\:\cancel{L}}{1000\:\cancel{mL}}\times \dfrac{0.250\:\cancel{mol\: NaOH}}{1\:\cancel{L}}\times \dfrac{1\: mol\: HCl}{1\:\cancel{mol\: NaOH}}=8.81\times 10^{-3}\:mol\: HCl} \nonumber \]

    Using the provided volume of HCl solution and the definition of molarity, the HCl concentration is:

    \[\begin{align*}
    M&=\mathrm{\dfrac{mol\: HCl}{L\: solution}}\\
    M&=\mathrm{\dfrac{8.81\times 10^{-3}\:mol\: HCl}{50.00\: mL\times \dfrac{1\: L}{1000\: mL}}}\\
    M&=0.176\:M
    \end{align*} \nonumber \]

    Note: For these types of titration calculations, it is convenient to recognize that solution molarity is also equal to the number of millimoles of solute per milliliter of solution:

    \[M=\mathrm{\dfrac{mol\: solute}{L\: solution}\times \dfrac{\dfrac{10^3\:mmol}{mol}}{\dfrac{10^3\:mL}{L}}=\dfrac{mmol\: solute}{mL\: solution}} \nonumber \]

    Using this version of the molarity unit will shorten the calculation by eliminating two conversion factors:

    \[\mathrm{\dfrac{35.23\:mL\: NaOH\times \dfrac{0.250\:mmol\: NaOH}{mL\: NaOH}\times \dfrac{1\:mmol\: HCl}{1\:mmol\: NaOH}}{50.00\:mL\: solution}=0.176\: \mathit M\: HCl} \nonumber \]

    Exercise \(\PageIndex{1}\)

    A 20.00-mL sample of aqueous oxalic acid, H2C2O4, was titrated with a 0.09113-M solution of potassium permanganate, KMnO4.

    \[\ce{2MnO4-}(aq)+\ce{5H2C2O4}(aq)+\ce{6H+}(aq)\rightarrow \ce{10CO2}(g)+\ce{2Mn^2+}(aq)+\ce{8H2O}(l) \nonumber \]

    A volume of 23.24 mL was required to reach the end point. What is the oxalic acid molarity?

    Answer

    0.2648 M

    Summary

    Chemical reactions are classified according to similar patterns of behavior. A large number of important reactions are included in three categories: precipitation, acid-base, and oxidation-reduction (redox). Precipitation reactions involve the formation of one or more insoluble products. Acid-base reactions involve the transfer of hydrogen ions between reactants. Redox reactions involve a change in oxidation number for one or more reactant elements. Writing balanced equations for some redox reactions that occur in aqueous solutions is simplified by using a systematic approach called the half-reaction method.

    Glossary

    acid
    substance that produces H3O+ when dissolved in water
    acid-base reaction
    reaction involving the transfer of a hydrogen ion between reactant species
    base
    substance that produces OH− when dissolved in water
    neutralization reaction
    reaction between an acid and a base to produce salt and water
    strong acid
    acid that reacts completely when dissolved in water to yield hydronium ions
    strong base
    base that reacts completely when dissolved in water to yield hydroxide ions
    weak acid
    acid that reacts only to a slight extent when dissolved in water to yield hydronium ions
    weak base
    base that reacts only to a slight extent when dissolved in water to yield hydroxide ions
    analyte
    chemical species of interest
    buret
    device used for the precise delivery of variable liquid volumes, such as in a titration analysis
    end point
    measured volume of titrant solution that yields the change in sample solution appearance or other property expected for stoichiometric equivalence (see equivalence point)
    equivalence point
    volume of titrant solution required to react completely with the analyte in a titration analysis; provides a stoichiometric amount of titrant for the sample’s analyte according to the titration reaction
    indicator
    substance added to the sample in a titration analysis to permit visual detection of the end point
    quantitative analysis
    the determination of the amount or concentration of a substance in a sample
    titrant
    solution containing a known concentration of substance that will react with the analyte in a titration analysis
    titration analysis
    quantitative chemical analysis method that involves measuring the volume of a reactant solution required to completely react with the analyte in a sample

    Contributors and Attributions

    This page titled 4.4: Classifying Chemical Reactions - Acid Base Reactions is shared under a CC BY license and was authored, remixed, and/or curated by Scott Van Bramer.