Solubility Rules
- Page ID
- 238
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)When a substance is added to a solvent, several outcomes are possible depending on its solubility, which is the maximum amount of solute that can dissolve in a given quantity of solvent under specified conditions. Solubility rules provide general guidelines for predicting which substances will dissolve readily in a solvent and which will remain largely undissolved.
Solubility Effects on Reactions
Depending on the solubility of a solute, there are three possible results:
- if the solution has less solute than the maximum amount that it is able to dissolve (i.e., its solubility) then it is a dilute solution
- if the amount of solute is exactly the same amount as its solubility then it is saturated
- if there is more solute than is able to be dissolved, the excess solute separates from the solution.
If this separation process includes crystallization, it forms a precipitate (a solid at the bottom of the solution). Precipitation lowers the concentration of the solute to the saturation to increase the stability of the solution.
Solubility Rules in Water
The following are the solubility rules for common ionic solids. If there two rules appear to contradict each other, the preceding rule takes precedence.
- Salts containing Group I elements (\(\ce{Li+}\), \(\ce{Na+}\), \(\ce{K+}\), \(\ce{Cs+}\), \(\ce{Rb+}\)) are soluble. There are few exceptions to this rule. Salts containing the ammonium ion (\(\ce{NH4+}\)) are also soluble.
- Salts containing nitrate ion (\(\ce{NO3-}\)) are generally soluble.
- Salts containing \(\ce{Cl-}\), \(\ce{Br-}\), or \(\ce{I-}\) are generally soluble. Important exceptions to this rule are halide salts of \(\ce{Ag+}\), \(\ce{Pb^{2+}}\), and \(\ce{(Hg2)^{2+}}\). Thus, \(\ce{AgCl}\), \(\ce{PbBr2}\), and \(\ce{Hg2Cl2}\) are insoluble.
- Most silver salts are insoluble. \(\ce{AgNO3}\) and \(\ce{Ag(C2H3O2}\)) are common soluble salts of silver -- virtually all others are insoluble.
- Most sulfate salts are soluble. Important exceptions to this rule include \(\ce{CaSO4}\), \(\ce{BaSO4}\), \(\ce{PbSO4}\), \(\ce{Pb(SO4)2}\), \(\ce{Ag2SO4}\), \(\ce{SrSO4 }\), \(\ce{Hg2SO4 }\), and \(\ce{HgSO4 }\).
- Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and \(\ce{Al^{3+}}\) are insoluble. Thus, \(\ce{Fe(OH)3}\), \(\ce{Al(OH)3}\), \(\ce{Co(OH)2}\) are not soluble.
- Most sulfides of transition metals are highly insoluble, including \(\ce{CdS}\), \(\ce{FeS}\), \(\ce{ZnS}\), and \(\ce{Ag2S}\). Arsenic, antimony, bismuth, and lead sulfides are also insoluble.
- Carbonates are frequently insoluble. Group II carbonates (\(\ce{CaCO3}\), \(\ce{SrCO3}\), and \(\ce{BaCO3}\) are insoluble, as are \(\ce{FeCO3}\) and \(\ce{PbCO3}\).
- Chromates are frequently insoluble. Examples include \(\ce{PbCrO4}\) and \(\ce{BaCrO4}\).
- Phosphates such as \(\ce{Ca3(PO4)2}\) and \(\ce{Ag3PO4}\) are frequently insoluble.
- Fluorides such as \(\ce{BaF2}\), \(\ce{MgF2}\), and \(\ce{PbF2}\) are frequently insoluble.
Is \(\ce{FeCO3}\) soluble in water?
- Answer
-
According to Rule #5, carbonates tend to be insoluble. Therefore, \(\ce{FeCO3}\) is likely to form a precipitate.
Does \(\ce{ClO4^{-}}\) tend to form a precipitate?
- Answer
-
This is perchlorate, which according to Rule #2 is likely to be soluble. Therefore, it will not form a precipitate.
Which of these substances is likely to form a precipitate?
- \(\ce{CaSO4}\)
- table salt
- \(\ce{AgBr}\)
- Answer
-
Letters a and c are both likely to form precipitates.
For option a: \(\ce{CaSO4}\), although sulfates tend to be soluble, Rule #5 indicates that calcium sulfate is an important exception to this rule.
For option b: Rule #1 indicates that table salt (\(\ce{NaCl}\)) is soluble because it is a salt of an alkali metal.
For option c: is an example of two rules contradicting each other. Rule #4 states that bromides are usually soluble, but Rule #3 states that salts of silver are insoluble. Because Rule #3 precedes Rule #4, the compound is insoluble and will form a precipitate.
Predict whether a precipitate will form as a result of this reaction:
\[\ce{2AgNO_3 + Na_2S \rightarrow Ag_2S + 2NaNO_3} \nonumber \]
The products of the reaction must be examined; if either of the substances formed in the reaction is insoluble, a precipitate will form.
- Answer
-
Considering \(\ce{NaNO3}\), Rule #3 states that nitrates tend to be soluble. A precipitate of this compound will not form.
Next, consider \(\ce{Ag2S}\). According to Rule #5, that sulfides tend to be insoluble. Therefore, because of this compound, a precipitate will form in the course of this reaction.
Predict if a precipitate will form as a result of this reaction:
\[\ce{2NaOH + K_2CrO_4 \rightarrow KOH + Na_2CrO_4 } \nonumber\]
- Answer
-
Consider again the products of the reaction: if either is insoluble, a precipitate will form.
The first product, \(\ce{KOH}\), is an example of two rules contradicting each other. Although Rule #5 says that hydroxides tend to be insoluble, Rule #1 states that salts of alkali metal cations tend to be soluble, and Rule #1 precedes Rule #5. Therefore, this compound will not contribute to any precipitation being formed.
The second product, \(\ce{Na2CrO4}\), also adheres to Rule #1, which states that salts of alkali metals tend to be soluble.
Because both products are soluble, no precipitate form as a result of this reaction.
References
- Petrucci, Ralph H., F. Geoffrey Herring, Jeffrey D. Madura and Carey Bissonnette. General Chemistry: Principles and Modern Applications. 10th ed. Upper Saddle River, New Jersey: Pearson Education, 2011. Print.
- Nathan, Harold D., and Charles Henrickson. Chemistry. New York: Wiley, 2001. Print.