4.4: Classifying Chemical Reactions - Acid Base Reactions
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- Scott Van Bramer
- Widener University
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, H 3 O + . As an example, consider the equation shown here for hydrochloric acid:
The process represented by this equation shows hydrogen chloride is an acid. When dissolved in water, H + ions are transferred from HCl molecules to H 2 O molecules to generate hydronium, H 3 O + , ions as shown in Figure \(\PageIndex{2}\).
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 \(\PageIndex{1}\)).
| Compound Formula | Name in Aqueous Solution |
|---|---|
| HBr | hydrobromic acid |
| HCl | hydrochloric acid |
| HNO 3 | nitric acid |
| H 2 SO 4 | 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. :
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 \(\PageIndex{3}\).
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 Ba 2 + , 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:
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:
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.
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:
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):
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:
- the weak acid hydrogen hypochlorite reacts with water
- a solution of barium hydroxide is neutralized with a solution of nitric acid
Solution
(a) The two reactants are provided, HOCl and H 2 O. Since the substance is reported to be an acid, its reaction with water will involve the transfer of H + from HOCl to H 2 O to generate hydronium ions, H 3 O + 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 HNO 3 . Since this is a neutralization reaction, the two products will be water and a salt composed of the cation of the ionic hydroxide (Ba 2 + ) 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
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\[\ce{H3O+}(aq)+\ce{OH-}(aq)\rightarrow \ce{2H2O}(l) \nonumber\]
In the 18th century, the strength (actually the concentration) of vinegar samples was determined by noting the amount of potassium carbonate, K 2 CO 3 , 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, CH 3 CO 2 H, 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 CO 2 .
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.
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:
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 milli moles of solute per milli liter 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, H 2 C 2 O 4 , was titrated with a 0.09113- M solution of potassium permanganate, KMnO 4 .
\[\ce{2MnO4-}(aq)+\ce{5H2C2O4}(aq)+\ce{6H+}(aq)\rightarrow \ce{10CO2}(g)+\ce{2Mn^2+}(aq)+\ce{8H2O}(l) \nonumber \]
A volume of 23.24 mL was required to reach the end point. What is the oxalic acid molarity?
- Answer
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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 H 3 O + 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
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Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors. Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/85abf193-2bd...a7ac8df6@9.110 ).