Which ion precipitates first

CaC 2 O 4 does not appear in this expression because it is a solid. Water does not appear because it is the solvent. Check Your Learning If a solution contains 0.

Neglect any increase in volume upon adding the solid silver nitrate. It is sometimes useful to know the concentration of an ion that remains in solution after precipitation. We can use the solubility product for this calculation too: If we know the value of K sp and the concentration of one ion in solution, we can calculate the concentration of the second ion remaining in solution. The calculation is of the same type as that in Example 9 —calculation of the concentration of a species in an equilibrium mixture from the concentrations of the other species and the equilibrium constant.

However, the concentrations are different; we are calculating concentrations after precipitation is complete, rather than at the start of precipitation. Solution The dissolution of Mn OH 2 is described by the equation:.

From that, we calculate the pH. At equilibrium:. If the person doing laundry adds a base, such as the sodium silicate Na 4 SiO 4 in some detergents, to the wash water until the pH is raised to Due to their light sensitivity, mixtures of silver halides are used in fiber optics for medical lasers, in photochromic eyeglass lenses glass lenses that automatically darken when exposed to sunlight , and—before the advent of digital photography—in photographic film.

When two anions form slightly soluble compounds with the same cation, or when two cations form slightly soluble compounds with the same anion, the less soluble compound usually, the compound with the smaller K sp generally precipitates first when we add a precipitating agent to a solution containing both anions or both cations.

When the K sp values of the two compounds differ by two orders of magnitude or more e. This is an example of selective precipitation , where a reagent is added to a solution of dissolved ions causing one of the ions to precipitate out before the rest. 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 5.

Specifically, selective precipitation is used to remove contaminants from wastewater before it is released back into natural bodies of water. 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. One common way to remove phosphates from water is by the addition of calcium hydroxide, known as lime, Ca OH 2.

The lime is converted into calcium carbonate, a strong base, in the water. As the water is made more basic, the calcium ions react with phosphate ions to produce hydroxylapatite, Ca 5 PO4 3 OH , which then precipitates out of the solution:. The precipitate is then removed by filtration and the water is brought back to a neutral pH by the addition of CO 2 in a recarbonation process. Other chemicals can also be used for the removal of phosphates by precipitation, including iron III chloride and aluminum sulfate.

View this site for more information on how phosphorus is removed from wastewater. Selective precipitation can also be used in qualitative analysis. In this method, reagents are added to an unknown chemical mixture in order to induce precipitation. Certain reagents cause specific ions to precipitate out; therefore, the addition of the reagent can be used to determine whether the ion is present in the solution.

View this simulation to study the process of salts dissolving and forming saturated solutions and precipitates for specific compounds, or compounds for which you select the charges on the ions and the K sp. Precipitation of Silver Halides A solution contains 0. AgNO 3 is gradually added to this solution. Which forms first, solid AgI or solid AgCl? If the solution contained about equal concentrations of Cl — and I — , then the silver salt with the smallest K sp AgI would precipitate first.

Check Your Learning If silver nitrate solution is added to a solution which is 0. As we saw when we discussed buffer solutions, the hydronium ion concentration of an aqueous solution of acetic acid decreases when the strong electrolyte sodium acetate, NaCH 3 CO 2 , is added.

Because sodium acetate and acetic acid have the acetate ion in common, the influence on the equilibrium is called the common ion effect. The common ion effect can also have a direct effect on solubility equilibria.

Suppose we are looking at the reaction where silver iodide is dissolved:. If we were to add potassium iodide KI to this solution, we would be adding a substance that shares a common ion with silver iodide. In this example, there would be an excess of iodide ions, so the reaction would shift toward the left, causing more silver iodide to precipitate out of solution. View this simulation to see how the common ion effect work with different concentrations of salts.

The K sp of CdS is 1. Solution The first thing you should notice is that the cadmium sulfide is dissolved in a solution that contains cadmium ions. We need to use an ICE table to set up this problem and include the CdBr 2 concentration as a contributor of cadmium ions:.

We can solve this equation using the quadratic formula, but we can also make an assumption to make this calculation much simpler. Since the K sp value is so small compared with the cadmium concentration, we can assume that the change between the initial concentration and the equilibrium concentration is negligible, so that 0. Going back to our K sp expression, we would now get:. Therefore, the molar solubility of CdS in this solution is 1. 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 solubility product of a slightly soluble electrolyte can be calculated from its solubility; conversely, its solubility can be calculated from its K sp , provided the only significant reaction that occurs when the solid dissolves is the formation of its ions. A slightly soluble electrolyte begins to precipitate when the magnitude of the reaction quotient for the dissolution reaction exceeds the magnitude of the solubility product.

Precipitation continues until the reaction quotient equals the solubility product. A reagent can be added to a solution of ions to allow one ion to selectively precipitate out of solution. The common ion effect can also play a role in precipitation reactions. Sea water has a density of 1. What mass, in kilograms, of Ca OH 2 is required to precipitate Find the K sp.

There is no change. A solid has an activity of 1 whether there is a little or a lot. The solubility of silver bromide at the new temperature must be known. Normally the solubility increases and some of the solid silver bromide will dissolve. MnCO 3 will form first, since it has the smallest K sp value it is the least soluble. MnCO 3 will be the last to precipitate, it has the largest K sp value. Skip to content Chapter Equilibria of Other Reaction Classes.

Learning Objectives By the end of this section, you will be able to:. Write chemical equations and equilibrium expressions representing solubility equilibria Carry out equilibrium computations involving solubility, equilibrium expressions, and solute concentrations. Example 1 Writing Equations and Solubility Products Write the ionic equation for the dissolution and the solubility product expression for each of the following slightly soluble ionic compounds: a AgI, silver iodide, a solid with antiseptic properties b CaCO 3 , calcium carbonate, the active ingredient in many over-the-counter chewable antacids c Mg OH 2 , magnesium hydroxide, the active ingredient in Milk of Magnesia d Mg NH 4 PO 4 , magnesium ammonium phosphate, an essentially insoluble substance used in tests for magnesium e Ca 5 PO 4 3 OH, the mineral apatite, a source of phosphate for fertilizers Hint: When determining how to break d and e up into ions, refer to the list of polyatomic ions in the section on chemical nomenclature.

Before any Hg 2 Cl 2 dissolves, Q is zero, and the reaction will shift to the right to reach equilibrium. Determine x and equilibrium concentrations. Concentrations and changes are given in the following ICE table:. Hg 2 Cl 2 is a pure solid, so it does not appear in the calculation. Solve for x and the equilibrium concentrations. We substitute the equilibrium concentrations into the expression for K sp and calculate the value of x :.

Tabulated K sp values can also be compared to reaction quotients calculated from experimental data to tell whether a solid will precipitate in a reaction under specific conditions: Q equals K sp at equilibrium; if Q is less than K sp , the solid will dissolve until Q equals K sp ; if Q is greater than K sp , precipitation will occur at a given temperature until Q equals K sp.

The equation that describes the equilibrium between solid calcium carbonate and its solvated ions is:. We can establish this equilibrium either by adding solid calcium carbonate to water or by mixing a solution that contains calcium ions with a solution that contains carbonate ions. The reaction shifts to the left and the concentrations of the ions are reduced by formation of the solid until the value of Q equals K sp.

A saturated solution in equilibrium with the undissolved solid will result. If the concentrations are such that Q is less than K sp , then the solution is not saturated and no precipitate will form. Note: Since all forms of equilibrium constants are temperature dependent, we will assume a room temperature environment going forward in this chapter unless a different temperature value is explicitly specified.

The first step in the preparation of magnesium metal is the precipitation of Mg OH 2 from sea water by the addition of lime, Ca OH 2 , a readily available inexpensive source of OH — ion:.

The reaction shifts to the left if Q is greater than K sp. Calculation of the reaction quotient under these conditions is shown here:. Mg OH 2 s forms until the concentrations of magnesium ion and hydroxide ion are reduced sufficiently so that the value of Q is equal to K sp.

Does silver chloride precipitate when equal volumes of a 2. The equation for the equilibrium between solid silver chloride, silver ion, and chloride ion is:. The solubility product is 1. AgCl will precipitate if the reaction quotient calculated from the concentrations in the mixture of AgNO 3 and NaCl is greater than K sp.

The volume doubles when we mix equal volumes of AgNO 3 and NaCl solutions, so each concentration is reduced to half its initial value. The reaction quotient, Q , is momentarily greater than K sp for AgCl, so a supersaturated solution is formed:.

Since supersaturated solutions are unstable, AgCl will precipitate from the mixture until the solution returns to equilibrium, with Q equal to K sp. Will KClO 4 precipitate when 20 mL of a 0. Remember to calculate the new concentration of each ion after mixing the solutions before plugging into the reaction quotient expression.

Thus, if we know the concentration of one of the ions of a slightly soluble ionic solid and the value for the solubility product of the solid, then we can calculate the concentration that the other ion must exceed for precipitation to begin. To simplify the calculation, we will assume that precipitation begins when the reaction quotient becomes equal to the solubility product constant.

Blood will not clot if calcium ions are removed from its plasma. For this reaction Table E3 :. CaC 2 O 4 does not appear in this expression because it is a solid. Water does not appear because it is the solvent. If a solution contains 0. Neglect any increase in volume upon adding the solid silver nitrate.

It is sometimes useful to know the concentration of an ion that remains in solution after precipitation. We can use the solubility product for this calculation too: If we know the value of K sp and the concentration of one ion in solution, we can calculate the concentration of the second ion remaining in solution.

However, the concentrations are different; we are calculating concentrations after precipitation is complete, rather than at the start of precipitation. From that, we calculate the pH. At equilibrium:. If the person doing laundry adds a base, such as the sodium silicate Na 4 SiO 4 in some detergents, to the wash water until the pH is raised to The first step in the preparation of magnesium metal is the precipitation of Mg OH 2 from sea water by the addition of Ca OH 2.

Due to their light sensitivity, mixtures of silver halides are used in fiber optics for medical lasers, in photochromic eyeglass lenses glass lenses that automatically darken when exposed to sunlight , and—before the advent of digital photography—in photographic film. When two anions form slightly soluble compounds with the same cation, or when two cations form slightly soluble compounds with the same anion, the less soluble compound usually, the compound with the smaller K sp generally precipitates first when we add a precipitating agent to a solution containing both anions or both cations.

When the K sp values of the two compounds differ by two orders of magnitude or more e. 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. One common way to remove phosphates from water is by the addition of calcium hydroxide, or lime, Ca OH 2.

As the water is made more basic, the calcium ions react with phosphate ions to produce hydroxylapatite, Ca 5 PO4 3 OH, which then precipitates out of the solution:. Because the amount of calcium ion added does not result in exceeding the solubility products for other calcium salts, the anions of those salts remain behind in the wastewater. The precipitate is then removed by filtration and the water is brought back to a neutral pH by the addition of CO 2 in a recarbonation process.

Other chemicals can also be used for the removal of phosphates by precipitation, including iron III chloride and aluminum sulfate. View this site for more information on how phosphorus is removed from wastewater. If the solution contained about equal concentrations of Cl — and Br — , then the silver salt with the smaller K sp AgBr would precipitate first.

Note the chloride ion concentration of the initial mixture was significantly greater than the bromide ion concentration, and so silver chloride precipitated first despite having a K sp greater than that of silver bromide. Compared with pure water, the solubility of an ionic compound is less in aqueous solutions containing a common ion one also produced by dissolution of the ionic compound.

Consider the dissolution of silver iodide:. In solutions that already contain either of these ions, less AgI may be dissolved than in solutions without these ions. This effect may also be explained in terms of mass action as represented in the solubility product expression:.

View this simulation to explore various aspects of the common ion effect. The solution is already saturated, though, so the concentrations of dissolved magnesium and hydroxide ions will remain the same.

Thus, changing the amount of solid magnesium hydroxide in the mixture has no effect on the value of Q , and no shift is required to restore Q to the value of the equilibrium constant.

The molar solubility of CdS in this solution is 1. As an Amazon Associate we earn from qualifying purchases. Want to cite, share, or modify this book?

This book is Creative Commons Attribution License 4. Skip to Content Go to accessibility page. Chemistry 2e My highlights. Table of contents. Answer Key. By the end of this section, you will be able to: Write chemical equations and equilibrium expressions representing solubility equilibria Carry out equilibrium computations involving solubility, equilibrium expressions, and solute concentrations.

Figure Calculate the molar solubility of copper bromide. Calculate the molar solubility of calcium hydroxide.