Ionic compounds solubility products

Equilibria of Slightly Soluble Ionic Compounds Solubility-Product Constant K )... 

A precipitation reaction occurs when two or more soluble species combine to form an insoluble product that we call a precipitate. The most common precipitation reaction is a metathesis reaction, in which two soluble ionic compounds exchange parts. When a solution of lead nitrate is added to a solution of potassium chloride, for example, a precipitate of lead chloride forms. We usually write the balanced reaction as a net ionic equation, in which only the precipitate and those ions involved in the reaction are included. Thus, the precipitation of PbCl2 is written as... 

Other useful solid-state electrodes are based on silver compounds (particularly silver sulfide). Silver sulfide is an ionic conductor, in which silver ions are the mobile ions. Mixed pellets containing Ag2S-AgX (where X = Cl, Br, I, SCN) have been successfiilly used for the determination of one of these particular anions. The behavior of these electrodes is determined primarily by the solubility products involved. The relative solubility products of various ions with Ag+ thus dictate the selectivity (i.e., kt] = KSp(Agf)/KSP(Aw)). Consequently, the iodide electrode (membrane of Ag2S/AgI) displays high selectivity over Br- and Cl-. In contrast, die chloride electrode suffers from severe interference from Br- and I-. Similarly, mixtures of silver sulfide with CdS, CuS, or PbS provide membranes that are responsive to Cd2+, Cu2+, or Pb2+, respectively. A limitation of these mixed-salt electrodes is tiiat the solubility of die second salt must be much larger than that of silver sulfide. A silver sulfide membrane by itself responds to either S2- or Ag+ ions, down to die 10-8M level. 

Most lanthanide compounds are sparingly soluble. Among those that are analytically important are the hydroxides, oxides, fluorides, oxalates, phosphates, complex cyanides, 8-hydroxyquinolates, and cup-ferrates. The solubility of the lanthanide hydroxides, their solubility products, and the pH at which they precipitate, are given in Table 2. As the atomic number increases (and ionic radius decreases), the lanthanide hydroxides become progressively less soluble and precipitate from more acidic solutions. The most common water-soluble salts are the lanthanide chlorides, nitrates, acetates, and sulfates. The solubilities of some of the chlorides and sulfates are also given in Table 2. 

Most CBD processes start slowly at a specific bath temperature, then accelerate, and eventually slow down again. Nucleation sites appear instantly on the substrates, the moment the solution contacts the substrate. In most cases, it is better to initiate the nucleation sites by inserting the substrate in a metalion solution, instead of a solution mixture containing both cations and anions. When cations and anions are mixed together, the resultant compounds begin to precipitate as soon as the ionic product, also called the solubility product,... 

The equilibrium between an ionic compound like that of formula (1) and an aqueous solution can be described by a solubility product defined by ... 

Sigma (a) bonds Sigma bonds have the orbital overlap on a line drawn between the two nuclei, simple cubic unit cell The simple cubic unit cell has particles located at the corners of a simple cube, single displacement (replacement) reactions Single displacement reactions are reactions in which atoms of an element replace the atoms of another element in a compound, solid A solid is a state of matter that has both a definite shape and a definite volume, solubility product constant (/ p) The solubility product constant is the equilibrium constant associated with sparingly soluble salts and is the product of the ionic concentrations, each one raised to the power of the coefficient in the balanced chemical equation, solute The solute is the component of the solution that is there in smallest amount, solution A solution is defined as a homogeneous mixture composed of solvent and one or more solutes. 

Le Chatelier s principle is a powerful tool for explaining how a reaction at equilibrium shifts when a stress is placed on the system. In this experiment, you can use Le Chatelier s principle to evaluate the relative solubilities of two precipitates. By observing the formation of two precipitates in the same system, you can infer the relationship between the solubilities of the two ionic compounds and the numerical values of their solubility product constants (K ). You will be able to verify your own experimental results by calculating the molar solubilities of the two compounds using the Ksp for each compound. 

As more ions enter the solution, the rate of the reverse change, recrystallisation, increases. Eventually, the rate of recrystallisation becomes equal to the rate of dissolving. As you know, when the forward rate and the backward rate of a process are equal, the system is at equilibrium. Because the reactants and the products are in different phases, the reaction is said to have reached heterogeneous equilibrium. For solubility systems of sparingly soluble ionic compounds, equilibrium exists between the solid ionic compound and its dissociated ions in solution. 

How do you predict whether a given concentration of ions will result in the precipitation of an ionic compound How can you tell if a solution is saturated You substitute the concentrations of the ions into an expression that is identical to the solubility product expression. Because these concentrations may not be the same as the concentrations that the equilibrium system would have, however, the expression has a different name the ion product. 

You calculate Qp by substituting the concentration of each ion into the expression. If Qp is larger than K p, the product of the concentrations of the ions is greater than it would be at equilibrium. For the system to attain equilibrium, some of the ions must leave the solution by precipitation. Conversely, if Qp is less than IQp, the product of the concentration of the ions is smaller than it is at equilibrium. Therefore, the solution is not yet saturated and more ions can be added to the solution without any precipitation. The relationship between Qsp and K p for the dissociation of a slightly soluble ionic compound is summarized on the next page. Use the following general equation as a reference. 

You can use the relationship between the ion product expression and the solubility product expression to predict whether a precipitate will form in a given system. One common system involves mixing solutions of two soluble ionic compounds, which react to form an ionic compound with a very low solubility. If Qsp > Kp. based on the initial concentrations of the ions in solution, the sparingly soluble compound will form a precipitate. 

How do you know which ionic compounds are soluble and which are not In your previous chemistry course, you learned a set of solubility guidelines. Table 9.3 summarizes these guidelines. Remember the higher guideline takes precedence. (For instance, guideline 3 says that carbonates have very low solubility. Sodium carbonate is soluble, however, because guideline 1 says that ionic compounds containing sodium are soluble.) Chemists do not usually work with solubility products for soluble compounds. Thus, you will not find soluble ionic compounds listed in tables. If you see a compound in a A p table, you know that it has a low solubility relative to compounds such as sodium chloride. 

In Chapter 9, as in most of Unit 4, you learned about equilibrium reactions. In this section, you analyzed precipitation reactions. You mainly examined double-displacement reactions—reactions in which two soluble ionic compounds react to form a precipitate. You used the solubility product constant, Ksp, to predict whether or not a precipitate would form for given concentrations of ions. In Unit 5, you will learn about a class of reactions that will probably be new to you. You will see how these reactions interconvert chemical and electrical energy. 

When equilibrium is reached, solubility product constants are used to describe saturated solutions of ionic compounds of relatively low solubility. When the ion concentration in solution reaches saturation, equilibrium between the solid and dissolved ions is established. 

Ionic liquids display good solubility for some organic compounds, typically aromatics, but poor solubility for many saturated hydrocarbons, and solubilities of gases also depend on their properties. It has therefore been possible to run chemical reactions with reactants that are more soluble in the ionic liquids than products. 

Solubility data are presented for practically all entries. Quantitative data are also given for some compounds at different temperatures. In general, ionic substances are soluble in water and other polar solvents while the non-polar, covalent compounds are more soluble in the non-polar solvents. In sparingly soluble, slightly soluble or practically insoluble salts, degree of solubility in water and occurrence of any precipitation process may be determined from the solubility product, Ksp, of the salt. The smaller the Ksp value, the less its solubility in water. 

The first condition necessary to form a solid in a solution is to exceed its solubility. In ionic solutions, the products of the concentrations (activities) of the actual reactants must be higher than required by the solubility product of the resulting compound at a given temperature. 

Chemistry is often conducted in aqueous solutions. Soluble ionic compounds dissolve into their component ions, and these ions can react to form new products. In these kinds of reactions, sometimes only the cation or anion of a dissolved compound reacts. The other ion merely watches the whole affair, twiddling its charged thumbs in electrostatic boredom. These uninvolved ions cire called spectator ions. 

Precipitation reactions are processes in which soluble reactants yield an insoluble solid product that drops out of the solution. Formation of this stable product removes material from the aqueous solution and provides the driving force for the reaction. Most precipitations take place when the anions and cations of two ionic compounds change partners. For example, an aqueous solution of lead(II) nitrate reacts with an aqueous solution of potassium iodide to yield an aqueous solution of potassium nitrate plus an insoluble yellow precipitate of lead iodide ... 

The solubility of an ionic compound can be described quantitatively by a value called the solubility product constant, Ksp. For the general solubility process AaBi, a An+ +b Bm, Ksp = [A, +HBm"]6. The brackets refer to concentrations in moles per liter. 

The solubility product, Ksp, for an ionic compound is the equilibrium constant for dissolution of the compound in water. The solubility of the compound and Ksp are related by the equilibrium equation for the dissolution reaction. The solubility of an ionic compound is (1) suppressed by the presence of a common ion in the solution (2) increased by decreasing the pH if the compound contains a basic anion, such as OH-, S2-, or CO32- and (3) increased by the presence of a Lewis base, such as NH3, CN-, or OH-, that can bond to the metal cation to form a complex ion. The stability of a complex ion is measured by its formation constant, Kf. 

When solutions of soluble ionic compounds are mixed, an insoluble compound will precipitate if the ion product (IP) for the insoluble compound exceeds its fCsp. The IP is defined in the same way as /equilibrium concentrations. Certain metal cations can be separated by selective precipitation of metal sulfides. Selective precipitation is important in qualitative analysis, a procedure for identifying the ions present in an unknown solution. 

The product of the reaction is a salt and has the structure shown. The properties given in the problem (soluble in polar solvents, high melting point) are typical of those of an ionic compound. 

What happens if both products are soluble ionic compounds Both ionic compounds will be ions dissolved in the water. If neither product precipitates out, no reaction occurs. Try the following problem to practise writing the products of double displacement reactions and predicting their states. 

Suppose that you did not have any information about the solubility of various compounds, but you did have access to a large variety of ionic compounds. What would you need to do before predicting the products of the displacement reactions above Outline a brief procedure. 

This equation does not show the change that occurs, however. It shows the reactants and products as intact compounds. In reality, soluble ionic compounds dissociate into their respective ions in solution. So chemists often use a total ionic equation to show the dissociated ions of the soluble ionic compounds. 

In aqueous solutions of ionic compounds, the ions act independently of each other. Soluble ionic compounds are written as their separate ions. We must be familiar with the solubility rules presented in Chapter 8 and recognize that the following types of compounds are strong electrolytes strong acids in solution, soluble metallic hydroxides, and salts. (Salts, which can be formed as the products of reactions of acids with bases, include all ionic compounds except strong acids and bases and metalhc oxides and hydroxides.) Compounds must be both ionic and soluble to be written in the form of their separate ions. (Section 9.1)... 

Net ionic equations (Chapter 9), like all other balanced chemical equations, give the mole ratios of reactants and products. Therefore, any calculations that require mole ratios may be done with net ionic equations as well as with total equations. However, a net ionic equation does not yield mass data directly because part of each soluble ionic compound is not given. For example, we can tell how many moles of silver ion are required to produce a certain number of moles of a product. 

Iron of inorganic dissolved compounds (bicarbonates, sulfates, chlorides, fluosilicates, etc.) may enter into the dissolved form of iron of inorganic origin (Fcj"" ), but their existence is governed by an acid environment with a pH not higher than 3. As a rule the pH in sea water is close to 8 ( 0.5). Under these conditions iron compounds are easily hydrolyzed and converted into hydroxides, which form colloidal solutions in sea water. In appropriate conditions colloidal hydroxide condenses to clots of gel and converts to the suspended state. Therefore there are practically no ionic forms of iron (Fe "" proper). As early as 1937 Cooper (1937) concluded, on the basis of the solubility product and activity of ferrous and ferric iron ions and FeOH ions, that until equilibrium is reached sea water may contain about 10 jiig/1 of iron ions in true solution at pH = 8.5 the amount of ionic Fe in ferric form is still less—10 which corresponds to the extremely... 

In dilute aqueous solutions, it has been demonstrated experimentally for poorly soluble ionic salts (solubilities less than 0.01 molL ) that the mathematical product of the total molar concentrations of the component ions is a constant at constant temperature. This product, is called the solubility product. Thus for a saturated solution of a simple ionic compound AB in water, we have the dynamic equilibrium ... 

The product of and the concentration of the undissolved solid creates a new constant called the solubility product constant, The solubility product constant is an equilihrium constant for the dissolving of a sparingly soluble ionic compound in water. The solubility product constant expression is... 

The solubility product constants for some ionic compounds are listed in Table 18-3. Note that they are all small numbers. Solubility product constants are measured and recorded only for sparingly soluble compounds. 

You have learned that the solubility product constant can be used to determine the molar solubility of an ionic compound. You can apply this information as you do the CHEMLAB at the end of this chapter. also can be used to find the concentrations of the ions in a saturated solution. 

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