For a heterogeneous equilibrium involving the slightly soluble solid M pX q and its ions M m+ and X n–: The equilibrium constant for an equilibrium involving the precipitation or dissolution of a slightly soluble ionic solid is called the solubility product, K sp, of the solid. The solution is already saturated, though, so the concentrations of dissolved magnesium and hydroxide ions will remain the same: (d) Adding more solid magnesium hydroxide will increase the amount of undissolved compound in the mixture. (c) The added compound does not contain a common ion, and no effect on the magnesium hydroxide solubility equilibrium is expected. (b) Adding a common ion, OH –, will increase the concentration of this ion and shift the solubility equilibrium to the left, decreasing the concentration of magnesium ion and increasing the amount of undissolved magnesium hydroxide. (a) Adding a common ion, Mg 2+, will increase the concentration of this ion and shift the solubility equilibrium to the left, decreasing the concentration of hydroxide ion and increasing the amount of undissolved magnesium hydroxide. The reaction shifts to the right direction.
(c) The concentration of OH – is reduced as the OH – reacts with the acid. At the new equilibrium, is greater and is less than in the solution of Mg(OH) 2 in pure water. Mg(OH) 2 forms until the reaction quotient again equals K sp. (b) The reaction shifts to the left to relieve the stress of the additional OH – ion. At the new equilibrium, is less and is greater than in the solution of Mg(OH) 2 in pure water. In quantitative terms, the added Mg 2+ causes the reaction quotient to be larger than the solubility product ( Q > K sp), and Mg(OH) 2 forms until the reaction quotient again equals K sp. (a) The reaction shifts to the left to relieve the stress produced by the additional Mg 2+ ion, in accordance with Le Châtelier’s principle. As the water is made more basic, the calcium ions react with phosphate ions to produce hydroxylapatite, Ca 5(PO4) 3OH, which then precipitates out of the solution: One common way to remove phosphates from water is by the addition of calcium hydroxide, or lime, Ca(OH) 2. (credit: “eutrophication&hypoxia”/Wikimedia Commons) Figure 7.2.5 – Wastewater treatment facilities, such as this one, remove contaminants from wastewater before the water is released back into the natural environment. An abundance of phosphate causes excess algae to grow, which impacts the amount of oxygen available for marine life as well as making water unsuitable for human consumption. For example, phosphate ions are often present in the water discharged from manufacturing facilities. Specifically, selective precipitation is used to remove contaminants from wastewater before it is released back into natural bodies of water. Solubility equilibria are useful tools in the treatment of wastewater carried out in facilities that may treat the municipal water in your city or town (Figure 7.2.5). The Role of Precipitation in Wastewater Treatment 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.
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