Textbook content produced by OpenStax is licensed under a Creative Commons Attribution License. We recommend using aĪuthors: Paul Flowers, Klaus Theopold, Richard Langley, William R. Use the information below to generate a citation. Then you must include on every digital page view the following attribution: If you are redistributing all or part of this book in a digital format, Then you must include on every physical page the following attribution: If you are redistributing all or part of this book in a print format, Want to cite, share, or modify this book? This book uses the This predictive strategy and related calculations are demonstrated in the next few example exercises. Q sp K sp: the reaction proceeds in the reverse direction (solution is supersaturated precipitation will occur) For the specific case of solubility equilibria: The comparison of Q sp to K sp to predict precipitation is an example of the general approach to predicting the direction of a reaction first introduced in the chapter on equilibrium. If the ion concentrations yield a reaction quotient greater than the solubility product, then precipitation will occur, lowering those concentrations until equilibrium is established ( Q sp = K sp). If the concentrations of calcium and carbonate ions in the mixture do not yield a reaction quotient, Q sp, that exceeds the solubility product, K sp, then no precipitation will occur. Consider, for example, mixing aqueous solutions of the soluble compounds sodium carbonate and calcium nitrate. It is important to realize that this equilibrium is established in any aqueous solution containing Ca 2+ and CO 3 2– ions, not just in a solution formed by saturating water with calcium carbonate. Solution ( a ) AgI ( s ) ⇌ Ag + ( a q ) + I − ( a q ) K sp = ( b ) CaCO 3 ( s ) ⇌ Ca 2+ ( a q ) + CO 3 2− ( a q ) K sp = ( c ) Mg ( OH ) 2 ( s ) ⇌ Mg 2+ ( a q ) + 2OH − ( a q ) K sp = 2 ( d ) Mg ( NH 4 ) PO 4 ( s ) ⇌ Mg 2+ ( a q ) + NH 4 + ( a q ) + PO 4 3− ( a q ) K sp = ( e ) Ca 5 ( PO 4 ) 3 OH ( s ) ⇌ 5Ca 2+ ( a q ) + 3PO 4 3− ( a q ) + OH − ( a q ) K sp = 5 3 ( a ) AgI ( s ) ⇌ Ag + ( a q ) + I − ( a q ) K sp = ( b ) CaCO 3 ( s ) ⇌ Ca 2+ ( a q ) + CO 3 2− ( a q ) K sp = ( c ) Mg ( OH ) 2 ( s ) ⇌ Mg 2+ ( a q ) + 2OH − ( a q ) K sp = 2 ( d ) Mg ( NH 4 ) PO 4 ( s ) ⇌ Mg 2+ ( a q ) + NH 4 + ( a q ) + PO 4 3− ( a q ) K sp = ( e ) Ca 5 ( PO 4 ) 3 OH ( s ) ⇌ 5Ca 2+ ( a q ) + 3PO 4 3− ( a q ) + OH − ( a q ) K sp = 5 3 Ĭheck Your LearningWrite the dissolution equation and the solubility product for each of the following slightly soluble compounds:ĬaCO 3 ( s ) ⇌ Ca 2+ ( a q ) + CO 3 2− ( a q ) K s p = 8.7 × 10 − 9 CaCO 3 ( s ) ⇌ Ca 2+ ( a q ) + CO 3 2− ( a q ) K s p = 8.7 × 10 − 9 (e) Ca 5(PO 4) 3OH, the mineral apatite, a source of phosphate for fertilizers (d) Mg(NH 4)PO 4, magnesium ammonium phosphate, an essentially insoluble substance used in tests for magnesium (c) Mg(OH) 2, magnesium hydroxide, the active ingredient in Milk of Magnesia (b) CaCO 3, calcium carbonate, the active ingredient in many over-the-counter chewable antacids (a) AgI, silver iodide, a solid with antiseptic properties Write the dissolution equation and the solubility product expression for each of the following slightly soluble ionic compounds: Writing Equations and Solubility Products For example, a saturated solution of silver chloride is one in which the equilibrium shown below has been established. A solute with finite solubility can yield a saturated solution when it is added to a solvent in an amount exceeding its solubility, resulting in a heterogeneous mixture of the saturated solution and the excess, undissolved solute. Recall from the chapter on solutions that the solubility of a substance can vary from essentially zero ( insoluble or sparingly soluble) to infinity ( miscible). This section applies previously introduced equilibrium concepts and tools to systems involving dissolution and precipitation. An understanding of the factors affecting compound solubility is, therefore, essential to the effective management of these processes. These equilibria underlie many natural and technological processes, ranging from tooth decay to water purification. Solubility equilibria are established when the dissolution and precipitation of a solute species occur at equal rates. Carry out equilibrium computations involving solubility, equilibrium expressions, and solute concentrations.Write chemical equations and equilibrium expressions representing solubility equilibria.By the end of this section, you will be able to:
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