4.3: Classifying Chemical Reactions - Precipitation Reactions
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- Scott Van Bramer
- Widener University
Learning Objectives
- Predict the solubility of common inorganic compounds by using solubility rules
- Identify when a precipitation reaction will occur
- Write a balanced chemical equation that describes what happens when a precipitation reaction occurs
- Solve stoichiometry problems with precipitation 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 precipitation reactions, one of the most prevalent types of chemical reactions.
Precipitation Reactions and Solubility Rules
A precipitation reaction is one in which dissolved substances react to form one (or more) solid products. Many reactions of this type involve the exchange of ions between ionic compounds in aqueous solution and are sometimes referred to as double displacement , double replacement , or metathesis reactions. These reactions are common in nature and are responsible for the formation of coral reefs in ocean waters and kidney stones in animals. They are used widely in industry for production of a number of commodity and specialty chemicals. Precipitation reactions also play a central role in many chemical analysis techniques, including spot tests used to identify metal ions and gravimetric methods for determining the composition of matter.
The extent to which a substance may be dissolved in water, or any solvent, is quantitatively expressed as its solubility , defined as the maximum concentration of a substance that can be achieved under specified conditions. Substances with relatively large solubilities are said to be soluble . A substance will precipitate when solution conditions are such that its concentration exceeds its solubility. Substances with relatively low solubilities are said to be insoluble , and these are the substances that readily precipitate from solution. For purposes of predicting the identities of solids formed by precipitation reactions, one may simply refer to patterns of solubility that have been observed for many ionic compounds (Table \(\PageIndex{1}\)).
| Always Soluble compounds contain | |
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| Usually Soluble compounds contain | Except if they also contain |
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| Usually Insoluble compounds contain | Exceptions include |
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A vivid example of precipitation is shown in Figure \(\PageIndex{1}\) when solutions of potassium iodide and lead nitrate are mixed, resulting in the formation of solid lead iodide. Lead iodide is a bright yellow solid that was formerly used as an artist’s pigment known as iodine yellow. The properties of pure PbI 2 crystals make them useful for fabrication of X-ray and gamma ray detectors. The balanced chemical reaction is:
\[\ce{2KI}(aq)+\ce{Pb(NO3)2}(aq)\rightarrow \ce{PbI2}(s)+\ce{2KNO3}(aq)\]
The formation of the precipitate observed in Figure \(\PageIndex{1}\) is consistent with the solubility guidelines: The only insoluble compound among all those involved is lead iodide, one of the exceptions to the general solubility of iodide salts.
The net ionic equation for this reaction is:
\[\ce{Pb^2+}(aq)+\ce{2I-}(aq)\rightarrow \ce{PbI2}(s)\]
The solubility guidelines in Table \(\PageIndex{1}\) will predict if a precipitation reaction occurs when solutions of soluble ionic compounds are mixed together. To do this, first identify all the ions present in the solution and then consider if any possible cation/anion combination forms an insoluble compound. For example, mixing solutions of silver nitrate and sodium fluoride will yield a solution containing Ag + , \(\ce{NO3-}\), Na + , and F − ions. The possible combinations include the two ionic compounds originally present in the solutions, AgNO 3 and NaF, which are both soluble in water. And two additional ionic compounds NaNO 3 and AgF. The solubility guidelines state all nitrate salts are soluble, so NaNO 3 will not form a precipitate. However, Ag + is an exception for halide ions so AgF forms a precipitate. The reaction is described by the following equations:
\[\ce{NaF}(aq)+\ce{AgNO3}(aq)\rightarrow \ce{AgF}(s)+\ce{NaNO3}(aq)\hspace{20px}\ce{(molecular)}\]
\[\ce{Na+}(aq)+\ce{F-}(aq)+\ce{Ag+}(aq)+\ce{NO3-}(aq)\rightarrow \ce{AgF}(s)+\ce{Na+}(aq)+\ce{NO3-}(aq)\hspace{20px}\ce{(complete\: ionic)}\]
Example \(\PageIndex{1}\): P redicting Precipitation Reactions
Predict the result of mixing reasonably concentrated solutions of the following ionic compounds. If precipitation is expected, write a balanced net ionic equation for the reaction.
- potassium sulfate and barium nitrate
- lithium chloride and silver acetate
- lead nitrate and ammonium carbonate
Solution
(a) The two possible products for this combination are KNO 3 and BaSO 4 . The solubility guidelines indicate BaSO 4 is insoluble, and so a precipitation reaction is expected. The net ionic equation for this reaction, derived in the manner detailed in the previous module, is
(b) The two possible products for this combination are LiC 2 H 3 O 2 and AgCl. The solubility guidelines indicate AgCl is insoluble, and so a precipitation reaction is expected. The net ionic equation for this reaction, derived in the manner detailed in the previous module, is
(c) The two possible products for this combination are PbCO 3 and NH 4 NO 3 . The solubility guidelines indicate PbCO 3 is insoluble, and so a precipitation reaction is expected. The net ionic equation for this reaction, derived in the manner detailed in the previous module, is
\[\ce{Pb^2+}(aq)+\ce{CO3^2-}(aq)\rightarrow \ce{PbCO3}(s) \nonumber\]
Exercise \(\PageIndex{1}\)
Which solution could be used to precipitate the barium ion, Ba 2 + , in a water sample: sodium chloride, sodium hydroxide, or sodium sulfate? What is the formula for the expected precipitate?
- Answer
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sodium sulfate, BaSO 4
Solubility of Ionic Compounds - Video
Video Topics
This video discusses how to predict if a given ionic compound is soluble or insoluble in water. Solubility rules make these prediction based on the species making up the ionic compound. If the ionic compound is soluble it will disassociate in water to form strong electrolyte aqueous solution.
Link to Video
Predicting the Solubility of Ionic Compounds: https://youtu.be/U3QNwnfmvGU
Determining the Products of a Precipitation Reaction - Video
Video Topics
During a precipitation reaction, also called a double substitution reaction, the ions of two soluble ionic compounds recombine to form an insoluble ionic compound and a precipitate (solid) forms. This video discusses how to determining the molecular formula of the precipitate given the ionic species present in the reactants. In particular, combining different charged species to create a neutral compounds will be discussed.
Link to Video
Determining the Products for Precipitation Reactions: https://youtu.be/r0kYeZVuTAM
Determining the Net Ionic Equation - Video
Determining Net Ionic Equation for a Precipitation Reaction: https://youtu.be/AMJz1Sdz8IA
Gravimetric Analysis
A gravimetric analysis is one in which a sample is subjected to some treatment that causes a change in the physical state of the analyte that permits its separation from the other components of the sample. Mass measurements of the sample, the isolated analyte, or some other component of the analysis system, used along with the known stoichiometry of the compounds involved, permit calculation of the analyte concentration. Gravimetric methods were the first techniques used for quantitative chemical analysis, and they remain important tools in the modern chemistry laboratory.
The required change of state in a gravimetric analysis may be achieved by various physical and chemical processes. For example, the moisture (water) content of a sample is routinely determined by measuring the mass of a sample before and after it is subjected to a controlled heating process that evaporates the water. Also common are gravimetric techniques in which the analyte is subjected to a precipitation reaction of the sort described earlier in this chapter. The precipitate is typically isolated from the reaction mixture by filtration, carefully dried, and then weighed (Figure \(\PageIndex{2}\)). The mass of the precipitate may then be used, along with relevant stoichiometric relationships, to calculate analyte concentration.
Example \(\PageIndex{2}\): Gravimetric Analysis
A 0.4550-g solid mixture containing MgSO 4 is dissolved in water and treated with an excess of Ba(NO 3 ) 2 , resulting in the precipitation of 0.6168 g of BaSO 4 .
\[\ce{MgSO4}(aq)+\ce{Ba(NO3)2}(aq)\rightarrow \ce{BaSO4}(s)+\ce{Mg(NO3)2}(aq) \nonumber \]
What is the concentration (percent) of MgSO 4 in the original 0.4550 g solid mixture?
Solution
The plan for this calculation is similar to others used in stoichiometric calculations, the central step being the connection between the moles of BaSO 4 and MgSO 4 through their stoichiometric factor. Once the mass of MgSO 4 is computed, it may be used along with the mass of the sample mixture to calculate the requested percentage concentration.
The mass of MgSO 4 that would yield the provided precipitate mass is
\[\mathrm{0.6168\:\cancel{g\: BaSO_4}\times \dfrac{1\:\cancel{mol\: BaSO_4}}{233.43\:\cancel{g\: BaSO_4}}\times \dfrac{1\:\cancel{mol\: MgSO_4}}{1\:\cancel{mol\: BaSO_4}}\times \dfrac{120.37\:g\: MgSO_4}{1\:\cancel{mol\: MgSO_4}}=0.3181\:g\: MgSO_4} \nonumber \]
The concentration of MgSO 4 in the sample mixture is then calculated to be
\[\begin{align*}
\ce{percent\: MgSO4}&=\ce{\dfrac{mass\: MgSO4}{mass\: sample}}\times100\%\\
\mathrm{\dfrac{0.3181\: g}{0.4550\: g}}\times100\%&=69.91\%
\end{align*} \nonumber \]
Exercise \(\PageIndex{2}\)
What is the percent of chloride ion in a sample if 1.1324 g of the sample produces 1.0881 g of AgCl when treated with excess Ag + ?
\[\ce{Ag+}(aq)+\ce{Cl-}(aq)\rightarrow \ce{AgCl}(s) \nonumber \]
- Answer
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23.76%
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
- insoluble
- of relatively low solubility; dissolving only to a slight extent
precipitate
- insoluble product that forms from reaction of soluble reactants
- precipitation reaction
- reaction that produces one or more insoluble products; when reactants are ionic compounds, sometimes called double-displacement or metathesis
- salt
- ionic compound that can be formed by the reaction of an acid with a base that contains a cation and an anion other than hydroxide or oxide
- single-displacement reaction
- (also, replacement) redox reaction involving the oxidation of an elemental substance by an ionic species
- soluble
- of relatively high solubility; dissolving to a relatively large extent
- solubility
- the extent to which a substance may be dissolved in water, or any solvent
analyte
- chemical species of interest
- gravimetric analysis
- quantitative chemical analysis method involving the separation of an analyte from a sample by a physical or chemical process and subsequent mass measurements of the analyte, reaction product, and/or sample
- quantitative analysis
- the determination of the amount or concentration of a substance 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 ).
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Prof. Steven Farmer ( Sonoma State University )