Coupled Equilibria

Learning Outcomes

  • Describe examples of systems involving two (or more) coupled chemical equilibria
  • Calculate reactant and product concentrations for coupled equilibrium systems

As discussed in preceding chapters on equilibrium, coupled equilibria involve two or more separate chemical reactions that share one or more reactants or products. This section of this chapter will address solubility equilibria coupled with acid-base and complex-formation reactions.

An environmentally relevant example illustrating the coupling of solubility and acid-base equilibria is the impact of ocean acidification on the health of the ocean’s coral reefs. These reefs are built upon skeletons of sparingly soluble calcium carbonate excreted by colonies of corals (small marine invertebrates).

 Carbon dioxide in the air dissolves in sea water, forming carbonic acid (H2CO3). The carbonic acid then ionizes to form hydrogen ions and bicarbonate ions [latex]\left({\text{HCO}}_{3}{}^{-}\right)[/latex], which can further ionize into more hydrogen ions and carbonate ions [latex]\left({\text{CO}}_{3}{}^{2-}\right):[/latex]


The excess H+ ions make seawater more acidic. Increased ocean acidification can then have negative impacts on reef-building coral, as they cannot absorb the calcium carbonate they need to grow and maintain their skeletons (Figure 1). This in turn disrupts the local biosystem that depends upon the health of the reefs for its survival. If enough local reefs are similarly affected, the disruptions to sea life can be felt globally. The world’s oceans are presently in the midst of a period of intense acidification, believed to have begun in the mid-nineteenth century, and which is now accelerating at a rate faster than any change to oceanic pH in the last 20 million years.

This figure contains two photographs of coral reefs. In a, a colorful reef that includes hues of purple and pink corals is shown in blue green water with fish swimming in the background. In b, grey-green mossy looking coral is shown in a blue aquatic environment. This photo does not have the colorful appearance or fish that were shown in figure a.

Figure 1. Healthy coral reefs (a) support a dense and diverse array of sea life across the ocean food chain. But when coral are unable to adequately build and maintain their calcium carbonite skeletons because of excess ocean acidification, the unhealthy reef (b) is only capable of hosting a small fraction of the species as before, and the local food chain starts to collapse. (credit a: modification of work by NOAA Photo Library; credit b: modification of work by “prilfish”/Flickr)

Learn more about ocean acidification and how it affects other marine creatures.

You can view the transcript for “Ocean Acidification” here (opens in new window).

This site has detailed information about how ocean acidification specifically affects coral reefs.

The dramatic increase in solubility with increasing acidity described above for calcium carbonate is typical of salts containing basic anions (e.g., carbonate, fluoride, hydroxide, sulfide). Another familiar example is the formation of dental cavities in tooth enamel. The major mineral component of enamel is calcium hydroxyapatite (Figure 2), a sparingly soluble ionic compound whose dissolution equilibrium is

[latex]{\text{Ca}}_{5}{\left({\text{PO}}_{4}\right)}_{3}\text{OH}\left(s\right)\longrightarrow 5{\text{Ca}}^{\text{2+}}\left(aq\right)+3{\text{PO}}_{4}^{3-}\left(aq\right)+{\text{OH}}^{-}\left(aq\right)[/latex]

This figure includes an image of two large light blue apatite crystals in a mineral conglomerate that includes white, grey, and tan crystals. The blue apatite crystals have a dull, dusty, or powdered appearance.

Figure 2. Crystal of the mineral hydroxylapatite, Ca5(PO4)3OH, is shown here. Pure apatite is white, but like many other minerals, this sample is colored because of the presence of impurities.

This compound dissolved to yield two different basic ions: triprotic phosphate ions

[latex]{\text{PO}}_{4}{}^{3-}\left(aq\right)+{\text{H}}_{3}{\text{O}}^{+}\longrightarrow {\text{H}}_{2}{\text{PO}}_{4}{}^{2-}+{\text{H}}_{2}\text{O}[/latex]
[latex]{\text{PO}}_{4}{}^{2-}\left(aq\right)+{\text{H}}_{3}{\text{O}}^{+}\longrightarrow {\text{H}}_{2}{\text{PO}}_{4}{}^{-}+{\text{H}}_{2}\text{O}[/latex]
[latex]{\text{H}}_{2}{\text{PO}}_{4}{}^{-}+{\text{H}}_{3}{\text{O}}^{+}\longrightarrow {\text{H}}_{3}{\text{PO}}_{4}+{\text{H}}_{2}\text{O}[/latex]

and monoprotic hydroxide ions:

[latex]{\text{OH}}^{-}\left(aq\right)+{\text{H}}_{3}{\text{O}}^{+}\longrightarrow 2{\text{H}}_{2}\text{O}[/latex]

Of the two basic productions, the hydroxide is, of course, by far the stronger base (it’s the strongest base that can exist in aqueous solution), and so it is the dominant factor providing the compound an acid-dependent solubility. Dental cavities form when the acid waste of bacteria growing on the surface of teeth hastens the dissolution of tooth enamel by reacting completely with the strong base hydroxide, shifting the hydroxyapatite solubility equilibrium to the right. Some toothpastes and mouth rinses contain added NaF or SnF2 that make enamel more acid resistant by replacing the strong base hydroxide with the weak base fluoride:

[latex]\text{NaF}+{\text{Ca}}_{5}{\left({\text{PO}}_{4}\right)}_{3}\text{OH}\rightleftharpoons {\text{Ca}}_{5}{\left({\text{PO}}_{4}\right)}_{3}\text{F}+{\text{Na}}^{+}+{\text{OH}}^{-}[/latex]

The weak base fluoride ion reacts only partially with the bacterial acid waste, resulting in a less extensive shift in the solubility equilibrium and an increased resistance to acid dissolution. See the Chemistry in Everyday Life feature on the role of fluoride in preventing tooth decay for more information.

Role of Fluoride in Preventing Tooth Decay

As we saw previously, fluoride ions help protect our teeth by reacting with hydroxylapatite to form fluorapatite, Ca5(PO4)3F. Since it lacks a hydroxide ion, fluorapatite is more resistant to attacks by acids in our mouths and is thus less soluble, protecting our teeth. Scientists discovered that naturally fluorinated water could be beneficial to your teeth, and so it became common practice to add fluoride to drinking water. Toothpastes and mouthwashes also contain amounts of fluoride (Figure 3).

A tube of toothpaste

Figure 3. Fluoride, found in many toothpastes, helps prevent tooth decay (credit: Kerry Ceszyk).

Unfortunately, excess fluoride can negate its advantages. Natural sources of drinking water in various parts of the world have varying concentrations of fluoride, and places where that concentration is high are prone to certain health risks when there is no other source of drinking water. The most serious side effect of excess fluoride is the bone disease, skeletal fluorosis. When excess fluoride is in the body, it can cause the joints to stiffen and the bones to thicken. It can severely impact mobility and can negatively affect the thyroid gland. Skeletal fluorosis is a condition that over 2.7 million people suffer from across the world. So while fluoride can protect our teeth from decay, the US Environmental Protection Agency sets a maximum level of 4 ppm (4 mg/L) of fluoride in drinking water in the US. Fluoride levels in water are not regulated in all countries, so fluorosis is a problem in areas with high levels of fluoride in the groundwater.

The solubility of ionic compounds may also be increased when dissolution is coupled to the formation of a complex ion. For example, aluminum hydroxide dissolves in a solution of sodium hydroxide or another strong base because of the formation of the complex ion Al(OH)4.Al(OH)4−.

An H atom is bonded to an O atom. The O atom has 2 dots above it and 2 dots below it. The O atom is bonded to an A l atom, which has three additional O atoms bonded to it as well. Each of these additional O atoms has 4 dots arranged around it, and is bonded to an H atom. This entire molecule is contained in brackets, to the right of which is a superscripted negative sign.

The equations for the dissolution of aluminum hydroxide, the formation of the complex ion, and the combined (net) equation are shown below. As indicated by the relatively large value of K for the net reaction, coupling complex formation with dissolution drastically increases the solubility of Al(OH)3.

[latex]\text{Al(OH)}_3 (s) \rightleftharpoons \text{Al}^{3+} (aq) + 3\text{OH}^- (aq)[/latex]        [latex]K_{sp} = 2 \times 10^{-32}[/latex]

[latex]\text{Al}^{3+} (aq) + 4\text{OH}^- (aq) \rightleftharpoons \text{Al(OH)}_4^{-} (aq)[/latex]        [latex]K_{f} = 1.1 \times 10^{33}[/latex]

[latex]\text{Net: Al(OH)}_3 (s) + \text{OH}^- (aq) \rightleftharpoons \text{Al(OH)}_4^{-} (aq)[/latex]        [latex]K = K_{sp} K_f = 22[/latex]

Example 1: Increased solubility in acidic solutions

Compute and compare the molar solublities for aluminum hydroxide, Al(OH)3, dissolved in (a) pure water and (b) a buffer containing 0.100 M acetic acid and 0.100 M sodium acetate.


Check Your Learning

What is the solubility of aluminum hydroxide in a buffer comprised of 0.100 M formic acid and 0.100 M sodium formate?


Example 2: Multiple Equilibria

Unexposed silver halides are removed from photographic film when they react with sodium thiosulfate (Na2S2O3, called hypo) to form the complex ion [latex]\text{Ag}{\left({\text{S}}_{2}{\text{O}}_{3}\right)}_{2}{}^{3-}[/latex] (Kf = 4.7 [latex]\times [/latex] 1013).

A chemical reaction is shown using structural formulas. On the left, A g superscript plus is followed by a plus sign, the number 2, and a structure in brackets. The structure is composed of a central S atom which has O atoms single bonded above, right, and below. A second S atom is single bonded to the left. Each of these bonded atoms has 6 dots around it. Outside the brackets is a superscript 2 negative. Following a bidirectional arrow is a structure in brackets with a central A g atom. To the left and right, S atoms are single bonded to the A g atom. Each of these S atoms has four dots around it, and an S atom connected with a single bond moving out from the central A g atom, forming the ends of the structure. Each of these atoms has three O atoms attached with single bonds above, below, and at the end of the structure. Each O atom has six dots around it. Outside the brackets is a superscript 3 negative.


What mass of Na2S2O3 is required to prepare 1.00 L of a solution that will dissolve 1.00 g of AgBr by the formation of [latex]\text{Ag}{\left({\text{S}}_{2}{\text{O}}_{3}\right)}_{2}{}^{3-}?[/latex]

Check Your Learning

AgCl(s), silver chloride, is well known to have a very low solubility:

[latex]\text{AgCl}(\text{s})\rightleftharpoons\text{Ag}^{+}(\text{aq}) + \text{Cl}^{-}(\text{aq})\qquad{K}_{\text{sp}}=1.77\times{10}^{-10}[/latex]

Adding ammonia significantly increases the solubility of AgCl because a complex ion is formed:


What mass of NH3 is required to prepare 1.00 L of solution that will dissolve 2.00 g of AgCl by formation of Ag(NH3)2+?

Key Concepts and Summary

Systems involving two or more chemical equilibria that share one or more reactant or product are called coupled equilibria. Common examples of coupled equilibria include the increased solubility of some compounds in acidic solutions (coupled dissolution and neutralization equilibria) and in solutions containing ligands (coupled dissolution and complex formation). The equilibrium tools from other chapters may be applied to describe and perform calculations on these systems.

Try It

  1. A saturated solution of a slightly soluble electrolyte in contact with some of the solid electrolyte is said to be a system in equilibrium. Explain. Why is such a system called a heterogeneous equilibrium?
  2. Calculate the equilibrium concentration of Ni2+ in a 1.0-M solution [Ni(NH3)6](NO3)2.
  3. Calculate the equilibrium concentration of Zn2+ in a 0.30-M solution of [latex]\text{Zn}{\left(\text{CN}\right)}_{4}{}^{2-}[/latex].
  4. Calculate the equilibrium concentration of Cu2+ in a solution initially with 0.050 M Cu2+ and 1.00 M NH3.
  5. Calculate the equilibrium concentration of Zn2+ in a solution initially with 0.150 M Zn2+ and 2.50 M CN.
  6. Calculate the Fe3+ equilibrium concentration when 0.0888 mole of K3[Fe(CN)6] is added to a solution with 0.0.00010 M CN.
  7. Calculate the Co2+ equilibrium concentration when 0.100 mole of [Co(NH3)6](NO3)2 is added to a solution with 0.025 M NH3. Assume the volume is 1.00 L.
  8. The equilibrium constant for the reaction [latex]{\text{Hg}}^{\text{2+}}\left(aq\right)+2{\text{Cl}}^{-}\left(aq\right)\rightleftharpoons {\text{HgCl}}_{2}\left(aq\right)[/latex] is 1.6 [latex]\times [/latex] 1013. Is HgCl2 a strong electrolyte or a weak electrolyte? What are the concentrations of Hg2+ and Cl in a 0.015-M solution of HgCl2?
  9. Calculate the molar solubility of Sn(OH)2 in a buffer solution containing equal concentrations of NH3 and [latex]{\text{NH}}_{4}{}^{+}[/latex].
  10. Calculate the molar solubility of Al(OH)3 in a buffer solution with 0.100 M NH3 and 0.400 M [latex]{\text{NH}}_{4}{}^{+}[/latex].
  11. What is the molar solubility of CaF2 in a 0.100-M solution of HF? Ka for HF = 7.2 [latex]\times [/latex] 10–4.
  12. What is the molar solubility of BaSO4 in a 0.250-M solution of NaHSO4? Ka for [latex]{\text{HSO}}_{4}{}^{-}[/latex] = 1.2 [latex]\times [/latex] 10–2.
  13. What is the molar solubility of Tl(OH)3 in a 0.10-M solution of NH3?
  14. What is the molar solubility of Pb(OH)2 in a 0.138-M solution of CH3NH2?
  15. A solution of 0.075 M CoBr2 is saturated with H2S ([H2S] = 0.10 M). What is the minimum pH at which CoS begins to precipitate?
    [latex]\text{CoS}\left(s\right)\rightleftharpoons {\text{Co}}^{\text{2+}}\left(aq\right)+{\text{S}}^{2-}\left(aq\right){K}_{\text{sp}}=4.5\times {10}^{-27}[/latex][latex]{\text{H}}_{2}\text{S}\left(aq\right)+2{\text{H}}_{2}\text{O}\left(l\right)\rightleftharpoons 2{\text{H}}_{3}{\text{O}}^{\text{+}}\left(aq\right)+{\text{S}}^{2-}\left(aq\right)K=1.0\times {10}^{\text{-26}}[/latex]
  16. A 0.125-M solution of Mn(NO3)2 is saturated with H2S ([H2S] = 0.10 M). At what pH does MnS begin to precipitate?[latex]\text{MnS}\left(s\right)\rightleftharpoons {\text{Mn}}^{\text{2+}}\left(aq\right)+{\text{S}}^{2-}\left(aq\right){K}_{\text{sp}}=4.3\times {10}^{-22}[/latex][latex]{\text{H}}_{2}\text{S}\left(aq\right)+2{\text{H}}_{2}\text{O}\left(l\right)\rightleftharpoons 2{\text{H}}_{3}{\text{O}}^{\text{+}}\left(aq\right)+{\text{S}}^{2-}\left(aq\right)K=1.0\times {10}^{-26}[/latex]
  17. Calculate the molar solubility of BaF2 in a buffer solution containing 0.20 M HF and 0.20 M NaF.
  18. Calculate the molar solubility of CdCO3 in a buffer solution containing 0.115 M Na2CO3 and 0.120 M NaHCO3
  19. To a 0.10-M solution of Pb(NO3)2 is added enough HF(g) to make [HF] = 0.10 M.
    1. Does PbF2 precipitate from this solution? Show the calculations that support your conclusion.
    2. What is the minimum pH at which PbF2 precipitates?
  20. Calculate the concentration of Cd2+ resulting from the dissolution of CdCO3 in a solution that is 0.250 M in CH3CO2H, 0.375 M in NaCH3CO2, and 0.010 M in H2CO3.
  21. Both AgCl and AgI dissolve in NH3.
    1. What mass of AgI dissolves in 1.0 L of 1.0 M NH3?
    2. What mass of AgCl dissolves in 1.0 L of 1.0 M NH3?
  22. Calculate the volume of 1.50 M CH3CO2H required to dissolve a precipitate composed of 350 mg each of CaCO3, SrCO3, and BaCO3.
  23. Even though Ca(OH)2 is an inexpensive base, its limited solubility restricts its use. What is the pH of a saturated solution of Ca(OH)2?
  24. What mass of NaCN must be added to 1 L of 0.010 M Mg(NO3)2 in order to produce the first trace of Mg(OH)2?
  25. Magnesium hydroxide and magnesium citrate function as mild laxatives when they reach the small intestine. Why do magnesium hydroxide and magnesium citrate, two very different substances, have the same effect in your small intestine. (Hint: The contents of the small intestine are basic.)
  26. The following question is taken from a Chemistry Advanced Placement Examination and is used with the permission of the Educational Testing Service. Solve the following problem: [latex]{\text{MgF}}_{2}\left(s\right)\rightleftharpoons {\text{Mg}}^{\text{2+}}\left(aq\right)+2{\text{F}}^{-}\left(aq\right)[/latex]In a saturated solution of MgF2 at 18 °C, the concentration of Mg2+ is 1.21 [latex]\times [/latex] 10–3M. The equilibrium is represented by the equation above.
    1. Write the expression for the solubility-product constant, Ksp, and calculate its value at 18 °C.
    2. Calculate the equilibrium concentration of Mg2+ in 1.000 L of saturated MgF2 solution at 18 °C to which 0.100 mol of solid KF has been added. The KF dissolves completely. Assume the volume change is negligible.
    3. Predict whether a precipitate of MgF2 will form when 100.0 mL of a 3.00 [latex]\times [/latex] 10–3M solution of Mg(NO3)2 is mixed with 200.0 mL of a 2.00 [latex]\times [/latex] 10–3M solution of NaF at 18 °C. Show the calculations to support your prediction.
    4. At 27 °C the concentration of Mg2+ in a saturated solution of MgF2 is 1.17 [latex]\times [/latex] 10–3M. Is the dissolving of MgF2 in water an endothermic or an exothermic process? Give an explanation to support your conclusion.
  27. Which of the following compounds, when dissolved in a 0.01-M solution of HClO4, has a solubility greater than in pure water: AgBr, BaF2, Ca3(PO4)3, ZnS, PbI2? Explain your answer.
  28. What is the effect on the amount of solid Mg(OH)2 that dissolves and the concentrations of Mg2+ and OH when each of the following are added to a mixture of solid Mg(OH)2 and water at equilibrium?
    1. MgCl2
    2. KOH
    3. HClO4
    4. NaNO3
    5. Mg(OH)2
  29. What is the effect on the amount of CaHPO4 that dissolves and the concentrations of Ca2+ and [latex]{\text{HPO}}_{4}{}^{-}[/latex] when each of the following are added to a mixture of solid CaHPO4 and water at equilibrium?
    1. CaCl2
    2. HCl
    3. KClO4
    4. NaOH
    5. CaHPO4
  30. Identify all chemical species present in an aqueous solution of Ca3(PO4)2 and list these species in decreasing order of their concentrations. (Hint: Remember that the [latex]{\text{PO}}_{4}{}^{\text{3-}}[/latex] ion is a weak base.)
  31. A volume of 50 mL of 1.8 M NH3 is mixed with an equal volume of a solution containing 0.95 g of MgCl2. What mass of NH4Cl must be added to the resulting solution to prevent the precipitation of Mg(OH)2?


coupled equilibrium: system characterized by more than one state of balance between a slightly soluble ionic solid and an aqueous solution of ions working simultaneously