1: Precipitation Reactions

The independent behavior of each type of ion in solution was illustrated in Chemical Bonding by means of precipitation reactions. Precipitation is a process in which a solute separates from a supersaturated solution. In a chemical laboratory it usually refers to a solid crystallizing from a liquid solution, but in weather reports it applies to liquid or solid water separating from supersaturated air.

A typical precipitation reaction occurs when an aqueous solution of barium chloride is mixed with one containing sodium sulfate. The equation

\[\ce{BaCl2(aq) + Na2SO4(aq) -> BaSO4(s) + 2NaCl(aq)}\label{1}\]

can be written to describe what happens, and such an equation is useful in making chemical calculations. However, Equation \(\ref{1}\) does not really represent the microscopic particles (that is, the ions) present in the solution. Thus we might write

\[\ce{Ba^{2+}(aq) + 2Cl^{-}(aq) + 2Na^{+}(aq) + SO4^{2-}(aq) -> BaSO4(s) + 2Na^{+} + 2Cl^{-}(aq)}\label{2}\]

Figure \(\PageIndex{1}\): The figure on the left shows all ions present in the solution before the reaction occurs. After the reaction occurs, Na⁺ and Cl^- ions remain present in the solution, un-reacted (spectator ions). Ba²⁺ and SO₄^2- ions however, react to form the insoluble, white BaSO₄ (s), which precipitates out of the solution.

Equation \(\ref{2}\) is rather cumbersome and includes so many different ions that it may be confusing. In any case, we are often interested in the independent behavior of ions, not the specific compound from which they came. A precipitate of BaSO₄(s) will form when any solution containing Ba²⁺(aq) is mixed with any solution containing SO₄^2–(aq) (provided concentrations are not extremely small). This happens independently of the Cl^–(aq) and Na⁺(aq) ions in Eq. \(\ref{2}\). These ions are called spectator ions because they do not participate in the reaction (see the figure above). When we want to emphasize the independent behavior of ions, a net ionic equation is written, omitting the spectator ions. For precipitation of BaSO₄ the net ionic equation is

\[\ce{Ba^{2+}(aq) + SO4^{2-}(aq) -> BaSO{4}(s)}\label{3} \]

Example \(\PageIndex{1}\): Precipitation Reaction

When a solution of AgNO₃ is added to a solution of CaCl₂, insoluble AgCl precipitates. Write three equations to describe this process.

Solution

Both AgNO₃ and CaCl₂ are soluble ionic compounds, and so they are strong electrolytes. The three equations are

\(\ce{2AgNO3(aq) + CaCl2(aq) -> 2AgCl(s) + Ca(NO3)2(aq)}\)

\(\ce{2Ag^+(aq) + 2NO3^{-}(aq) + Ca^{2+}(aq) + Cl^{-}(aq) -> 2AgCl(s) + Ca^{2+}(aq) + 2NO3^{-}(aq)}\)

\(\ce{Ag^+(aq) + Cl^{-} (aq) -> AgCl(s)}\)

The occurrence or nonoccurrence of precipitates can be used to detect the presence or absence of various species in solution. BaCl₂ solution, for instance, is often used as a test for SO₄^2–(aq) ion. There are several insoluble salts of Ba, but they all dissolve in dilute acid except for BaSO₄. Thus, if BaCl₂ solution is added to an unknown solution which has previously been acidified, the occurrence of a white precipitate is proof of the presence of the SO₄^2– ion. AgNO₃ solution is often used in a similar way to test for halide ion.

If AgNO₃ solution is added to an acidified unknown solution, a white precipitate indicates the presence of Cl^– ions, a cream-colored precipitate indicates the presence of Br^– ions, and a yellow precipitate indicates the presence of I^– ions. Further tests can then be made to see whether perhaps a mixture of these ions is present. When AgNO₃ is added to tap water, a white precipitate is almost always formed. The Cl^– ions in tap water usually come from the Cl₂ which is added to municipal water supplies to kill microorganisms.

Precipitates are also used for quantitative analysis of solutions, that is, to determine the amount of solute or the mass of solute in a given solution. For this purpose it is often convenient to use the first of the three types of equations described above. Then the rules of stoichiometry may be applied.

Example \(\PageIndex{2}\): Concentration

When a solution of 0.1 M AgNO₃is added to 50.0 cm³ of a CaCl₂ solution of unknown concentration, 2.073 g AgCl precipitates. Calculate the concentration of the unknown solution.

Solution

We know the volume of the unknown solution, and so only the amount of solute is needed to calculate the concentration. This can be found using Eq. (2a) in Example 1. From the equation the stoichiometric ratio S(CaCl₂/AgCl) may be obtained. A road map to the solution of the problem is

\(m_{AgCl} \underset{\text{M_{AgCl}}}{\mathop{\rightarrow}}\, n_{AgCl} \underset{\text{M_{CaCL_2}}}{\mathop{\rightarrow}}\, n_{CaCl_2}\)

\(n_{CaCl-2} = 2.073 g AgCl * \frac{1mol AgCl}{143.32 g AgCl} * \frac{1 mol CaCl_2}{2 mol AgCl} = 7.23 * 10^{-3} \text{mol CaCl}_2\)

\(c_{CaCl_2} = \frac{n_{CaCl_2}}{V_{soln}} = \frac{7.23 * 10^{-3} mol CaCl_2}{50.0 cm^3} * \frac{10^3 cm^3}{1 dm^3} = 0.145 \frac{mol}{dm^3}\)

Thus the concentration of the unknown solution is 0.145 M.

Because of the general utility of precipitates in chemistry, it is worth having at least a rough idea of which common classes of compounds can be precipitated from solution and which cannot. Table \(\PageIndex{1}\) gives a list of rules which enable us to predict the solubility of the most commonly encountered substances. Use of this table is illustrated in the following example.

Table \(\PageIndex{1}\): Solubility. Rules

Soluble in Water	Important Exceptions (insoluble)
All Group IA and NH4^+ salts	CaSO4, BaSO4, SrSO4, PbSO4
All nitrates, chlorates, perchlorates and acetates	AgX, Hg2X2, PbX2 (X= Cl, Br, or I)
All sulfates
All chlorides, bromides, and iodides
Slightly Soluble in Water	Important Exceptions (soluble)
All carbonates and phosphates	Group IA and NH4^+ salts
All hydroxides	Group IA and NH4^+ salts; Ba^2^+, Sr^2^+, Ca^2^+ sparingly soluble
All sulfides	Group IA, IIA and NH4^+ salts; MgS, CaS, BaS sparingly soluble
All oxalates	Group IA and NH4^+ salts
The following electrolytes are of only moderate solubility in water: CH3COOAg, Ag2SO4, KClO4 They will precipitate only if rather concentrated solutions are used.

Example \(\PageIndex{3}\): Net Ionic Equation

Write balanced net ionic equations to describe any reactions which occur when the following solutions are mixed:

1. 0.1 M Na₂SO₄ + 0.1 M NH₄I
2. 0.1 M K₂CO₃ + 0.1 M SrCl₂
3. 0.1 M FeSO₄ + 0.1 M Ba(OH)₂

Solution

a) If any precipitate forms, it will be either a combination of Na⁺ ions and I^– ions, namely, NaI, or a combination of ammonium ions, NH₄⁺, and sulfate ions, SO₄^2–, namely, (NH₄)₂SO₄. From Table 11.2 we find that NaI and (NH₄)₂SO₄ are both soluble. Thus no precipitation reaction will occur, and there is no equation to write.

b) Possible precipitates are KCl and SrCO₃. From Table 11.2 we find that SrCO₃ is insoluble. Accordingly we write the net ionic equation as

\(\ce{Sr^{2+}(aq) + CO3^{2-}(aq) -> SrCO3(s)}\)

omitting the spectator ions K⁺ and Cl^–.

c) Possible precipitates are Fe(OH)₂ and BaSO₄. Both are insoluble. The net ionic equation is thus

\(\ce{Fe^{2+}(aq) + SO4^{2-}(aq) + Ba^{2+}(aq) + 2OH^{-}(aq) -> Fe(OH)2(s) + BaSO4(s)}\)