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.
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.
Example 2 Calculation of K sp from Equilibrium Concentrations We began the chapter with an informal discussion of how the mineral fluorite Introduction to Chapter 15 is formed. Answer: 8. Answer: 1. Example 5 Determination of K sp from Gram Solubility Many of the pigments used by artists in oil-based paints Figure 2 are sparingly soluble in water.
Figure 2. Oil paints contain pigments that are very slightly soluble in water. Using Barium Sulfate for Medical Imaging Various types of medical imaging techniques are used to aid diagnoses of illnesses in a noninvasive manner. Figure 3. The suspension of barium sulfate coats the intestinal tract, which allows for greater visual detail than a traditional X-ray. Example 9 Precipitation of Calcium Oxalate Blood will not clot if calcium ions are removed from its plasma. Figure 4.
Answer: 4. Answer: The Role of Precipitation in Wastewater Treatment 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. Figure 5. Wastewater treatment facilities, such as this one, remove contaminants from wastewater before the water is released back into the natural environment. Example 11 Precipitation of Silver Halides A solution contains 0.
What additional information do we need to answer the following question: How is the equilibrium of solid silver bromide with a saturated solution of its ions affected when the temperature is raised? Because these compounds are only slightly soluble, assume that the volume does not change on dissolution and calculate the solubility product for each. Show that changes in the initial concentrations of the common ions can be neglected. Show that it is not appropriate to neglect the changes in the initial concentrations of the common ions.
Explain why the changes in concentrations of the common ions in Chapter Calculate the solubility of aluminum hydroxide, Al OH 3 , in a solution buffered at pH Refer to Appendix J for solubility products for calcium salts.
Determine which of the calcium salts listed is most soluble in moles per liter and which is most soluble in grams per liter. Most barium compounds are very poisonous; however, barium sulfate is often administered internally as an aid in the X-ray examination of the lower intestinal tract Figure 3. This use of BaSO 4 is possible because of its low solubility. Calculate the molar solubility of BaSO 4 and the mass of barium present in 1.
What mass of this salt will dissolve in 1. Assuming that no equilibria other than dissolution are involved, calculate the concentrations of ions in a saturated solution of each of the following see Appendix J for solubility products.
See Appendix J for K sp values. Calculate the concentration of sulfate ion when BaSO 4 just begins to precipitate from a solution that is 0. Calculate the concentration of F — required to begin precipitation of CaF 2 in a solution that is 0. A volume of 0. Does BaSO 4 precipitate? Explain your answer. Perform these calculations for nickel II carbonate.
Iron concentrations greater than 5. A solution is 0. A listing of solubility product constants for several sparingly soluble compounds is provided in Appendix J. The K sp of a slightly soluble ionic compound may be simply related to its measured solubility provided the dissolution process involves only dissociation and solvation, for example:. For cases such as these, one may derive K sp values from provided solubilities, or vice-versa.
What is the solubility product of fluorite? Substituting the ion concentrations into the K sp expression gives. Substituting the equilibrium concentration terms into the solubility product expression and solving for x yields.
Since the dissolution stoichiometry shows one mole of copper I ion and one mole of bromide ion are produced for each moles of Br dissolved, the molar solubility of CuBr is 7. Substituting terms for the equilibrium concentrations into the solubility product expression and solving for x gives.
As defined in the ICE table, x is the molarity of calcium ion in the saturated solution. The dissolution stoichiometry shows a relation between moles of calcium ion in solution and moles of compound dissolved, and so, the molar solubility of Ca OH 2 is 6. Before calculating the solubility product, the provided solubility must be converted to molarity:. Substituting the equilibrium concentration terms into the solubility product expression and solving for x gives.
Various types of medical imaging techniques are used to aid diagnoses of illnesses in a noninvasive manner. One such technique utilizes the ingestion of a barium compound before taking an X-ray image.
A suspension of barium sulfate, a chalky powder, is ingested by the patient. Since the K sp of barium sulfate is 2. Barium-coated areas of the digestive tract then appear on an X-ray as white, allowing for greater visual detail than a traditional X-ray Figure Visit this website for more information on how barium is used in medical diagnoses and which conditions it is used to diagnose. The equation that describes the equilibrium between solid calcium carbonate and its solvated ions is:.
Consider, for example, mixing aqueous solutions of the soluble compounds sodium carbonate and calcium nitrate. 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.
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 potassium carbonate is used to selectively precipitate one of the cations while leaving the other cation in solution, which cation will precipitate first? But I am not sure. Borek Mr. Correct thinking.
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