4.2: Precipitation and Solubility Rules

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Learning Objectives

Predict the solubility of common inorganic compounds by using solubility rules
Write simple net ionic equations for precipitation 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 (see the last module of this chapter).

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. More information on these important concepts is provided in the text chapter on solutions. 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 \(\PageIndex{1}\)).

All compounds that contain the cations Na⁺, K⁺ or NH₄⁺, or the anions NO₃^– or C₂H₃O₂^– are soluble in water. It is best to memorize these. Beyond this, solubilities are normally classified using the anion in the compound. Here are the rules that you will use in Chem 101A:

Table \(\PageIndex{1}\): Solubility Rules for Common Ionic Compounds in Water
An anion from this column...	...combined with a cation from this column produces a soluble compound (a precipitate will NOT form)	... combined with a cation from this column produces an insoluble compound (a precipitate will form)
NO₃^–, C₂H₃O₂^–	All	None
Cl^–, Br^–, I^–	Most	Ag⁺ Pb²⁺ Hg₂²⁺
SO₄^2–	Most	Ag⁺ Pb²⁺ Hg₂²⁺ Ca²⁺ Sr²⁺ Ba²⁺ (the heavier IIA elements)
OH^–	Na⁺ K⁺ (NH₄⁺ reacts with OH^–: see below) Ba²⁺	All others (note: Ag⁺forms an oxide product, rather than hydroxide product)
CO₃^2–, PO₄^3–	Na⁺ K⁺ NH₄⁺	All others
S^2–	Na⁺ K⁺ NH₄⁺ Mg²⁺ Ca²⁺ Sr²⁺ Ba²⁺ (group IIA)	All others (the reactions of sulfide with 3+ ions are not simple precipitations: you do not need to know these)

Note that the reaction of Ag⁺ with OH^– produces Ag₂O (and water), not AgOH. This is a “quirk” of the chemistry of silver ions. The net ionic equation for this reaction is:

\[\ce{2Ag^+}(aq)+\ce{2OH-}(aq)\rightarrow \ce{Ag2O}(s)+\ce{H2O}(l)\]

Note that ammonia (NH₃) dissolves in water to produce a small concentration of hydroxide ions (discussed in a later section.) The resulting hydroxide ions can participate in precipitation reactions. Here is an example, using Mg²⁺:

\[\ce{Mg^2+}(aq)+\ce{2NH3}(aq)+\ce{2H2O}(l)\rightarrow \ce{Mg(OH)2}(s)+\ce{2NH4^+}(aq)\]

Note that the above equation is written in terms of the major species in solution (NH₃and H₂O) as opposed to the minor species (NH₄+ and OH-).

A vivid example of precipitation is observed when solutions of potassium iodide and lead nitrate are mixed, resulting in the formation of solid lead iodide. This observation is consistent with the solubility guidelines given above: The only insoluble combination among all those possible is lead and iodide. The net ionic equation representing this reaction is:

\[\ce{Pb^2+}(aq)+\ce{2I-}(aq)\rightarrow \ce{PbI2}(s)\]

Lead iodide is a bright yellow solid that was formerly used as an artist’s pigment known as iodine yellow (Figure \(\PageIndex{1}\)). The properties of pure PbI₂ crystals make them useful for fabrication of X-ray and gamma ray detectors.

A photograph is shown of a yellow green opaque substance swirled through a clear, colorless liquid in a test tube. — Figure \(\PageIndex{1}\): A precipitate of PbI₂ forms when solutions containing Pb²⁺ and I⁻ are mixed. (credit: Der Kreole/Wikimedia Commons)

The solubility guidelines in Table \(\PageIndex{1}\) may be used to predict whether a precipitation reaction will occur when solutions of soluble ionic compounds are mixed together. One merely needs to identify all the ions present in the solution and then consider if possible cation/anion pairing could result in an insoluble compound. For example, mixing solutions of silver nitrate and sodium chloride will yield a solution containing Ag⁺, \(\ce{NO3-}\), Na⁺, and Cl⁻ ions. Aside from the two ionic compounds originally present in the solutions, AgNO₃ and NaCl, two additional ionic compounds may be derived from this collection of ions: NaNO₃ and AgCl. The solubility guidelines indicate all nitrate salts are soluble but that AgCl is an insoluble combination. A precipitation reaction, therefore, is predicted to occur, as described by the following equation:

\[\ce{Ag+}(aq)+\ce{Cl-}(aq)\rightarrow \ce{AgCl}(s)\hspace{20px}\ce{(net\: ionic)}\]

Example \(\PageIndex{1}\): Predicting Precipitation Products

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 KNO₃ and BaSO₄. The solubility guidelines indicate BaSO₄ 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

\[\ce{Ba^2+}(aq)+\ce{SO4^2-}(aq)\rightarrow \ce{BaSO4}(s)\]

(b) The two possible products for this combination are LiC₂H₃O₂ 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

\[\ce{Ag+}(aq)+\ce{Cl-}(aq)\rightarrow \ce{AgCl}(s)\]

(c) The two possible products for this combination are PbCO₃ and NH₄NO₃. The solubility guidelines indicate PbCO₃ 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

\[\ce{Pb^2+}(aq)+\ce{CO3^2-}(aq)\rightarrow \ce{PbCO3}(s)\]

Exercise \(\PageIndex{1}\)

Which solution could be used to precipitate the barium ion, Ba²⁺, in a water sample: sodium chloride, sodium hydroxide, or sodium sulfate? What is the formula for the expected precipitate?

Answer: sodium sulfate, BaSO₄

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

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

Contributors and Attributions

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).