Aqueous equilibria complex ions

The bond dissociation energy of the hydrogen-fluorine bond in HF is so great that the above equilibrium lies to the left and hydrogen fluoride is a weak acid in dilute aqueous solution. In more concentrated solution, however, a second equilibrium reaction becomes important with the fluoride ion forming the complex ion HFJ. The relevant equilibria are ... 

Even though the equilibrium constant for the formation of Au3- from gold is very unfavorable, the reaction proceeds because any Au3+ ions formed are immediately complexed by Cl- ions and removed from the equilibrium. In a process widely used in the refining of the metal, gold also reacts with sodium cyanide in an aerated aqueous solution to form the complex ion [Au(CN)2] ... 

In Section 16-6, we describe how metal cations in aqueous solution can form bonds to anions or neutral molecules that have lone pairs of electrons. This leads to formation of complex ions and to chemical equilibria involving complexation. The complexation equilibrium between Ag and NH3 is an example ... 

Data at 25°C. Free energy of formation, equilibrium constants and re-lated data at 25°C exist for a wide range of minerals, other solids, gases and aqueous species, including ions and complexes (see later discussion on data sources). Availability of data at this reference temperature is usually not a major stumbling block. 

Bromine is moderately soluble in water, 33.6 g/L at 25°C. It gives a crystalline hydrate having a formula of <7.9H20 (6). The solubilities of bromine in water at several temperatures are given in Table 2. Aqueous bromine solubility increases in the presence of bromides or chlorides because of complex ion formation. This increase in the presence of bromides is illustrated in Figure 1. Equilibrium constants for the formation of the tribromide and pentabromide ions at 25°C have been reported (11). 

The solubility product is the equilibrium constant for the dissolution of a solid salt into its constituent ions in aqueous solution. The common ion effect is the observation that, if one of the ions of that salt is already present in the solution, the solubility of a salt is decreased. Sometimes, we can selectively precipitate one ion from a solution containing other ions by adding a suitable counterion. At high concentration of ligand, a precipitated metal ion may redissolve by forming soluble complex ions. In a metal-ion complex, the metal is a Lewis acid (electron pair acceptor) and the ligand is a Lewis base (electron pair donor). 

As another example of how a change in concentration affects an equilibrium, let s consider the reaction in aqueous solution of iron(III) and thiocyanate (SCN -) ions to give an equilibrium mixture that contains the Fe-N bonded red complex ion FeNCS2+ ... 

Addition of aqueous HgCl2 also eliminates the red color because Hg2+ reacts with SCN ions to form the stable Hg-S bonded complex ion Hg( SCN)42-. Removal of free SCN (aq) shifts the equilibrium Ve3+(aq) + SCN (aq) FeNCS2+(ag) from right to left to replenish the SCN ions. 

The solubility of an ionic compound increases dramatically if the solution contains a Lewis base that can form a coordinate covalent bond (Section 7.5) to the metal cation. Silver chloride, for example, is insoluble in water and in acid, but it dissolves in an excess of aqueous ammonia, forming the complex ion Ag(NH3)2 + (Figure 16.13). A complex ion is an ion that contains a metal cation bonded to one or more small molecules or ions, such as NH3, CN-, or OH-. In accord with Le Chatelier s principle, ammonia shifts the solubility equilibrium to the right by tying up the Ag+ ion in the form of the complex ion ... 

The stability of a complex ion is measured by its formation constant Kf (or stability constant), the equilibrium constant for formation of the complex ion from the hydrated metal cation. The large value of Kf for Ag(NH3)2+ means that this complex ion is quite stable, and nearly all the Ag+ ion in an aqueous ammonia solution is therefore present in the form of Ag(NH3)2+ (see Worked Example 16.12). 

Mass action equations. The first step in any calculation is to collate the mass action expressions that define the formation of the species. The way in which the formation constants can be used can be illustrated by considering a metal such as aluminium in an aqueous medium. Aluminium ions can undergo a number of hyrolysis reactions in water to form several hydroxy-metal complexes. The reactions can be written as the overall hydrolysis reactions and their associated equilibrium formation constants are shown below. 

First law of thermodynamics the energy of the universe is constant same as the law of conservation of energy. (9.1) Fission the process of using a neutron to split a heavy nucleus into two nuclei with smaller mass numbers. (21.6) Formal charge the charge assigned to an atom in a molecule or polyatomic ion derived from a specific set of rules. (13.12) Formation constant (stability constant) the equilibrium constant for each step of the formation of a complex ion by the addition of an individual ligand to a metal ion or complex ion in aqueous solution. (8.9)... 

In order to verify the differences observed in 4f-5f complex stability, we used an additional solvent extraction method based on the competition between a soluble organic chelatant (TTA] and the aqueous soluble azide ions for binding the metal ions. The extraction equilibrium is shown by equation (2] and the distribution coefficients CD] of the metal at constant pH are correlated with the formation constants by equation (3]. 

The strength of association between the ions in solution is expressed by various equilibrium constants. Stability (formation) constants refer to complex ions and ion pairs hydrolysis (deprotonation) constants refer to the loss of H+ from the water ligands surrounding central cations. Solubility products refer to the aqueous ion activities in equilibrium with solid phases. Some constants are reported in the literature in terms of concentrations rather than activities. Such constants are misnamed, since they depend both on the concentration and on the nature of other ions in solution. Converting concentrations to activities gives a much more useful value. 

The waters of the hydration sphere are usually ignored in the equilibrium constant, because excess water is present in aqueous solutions, and the energy of the Al-F bond must be greater than that of the ion-water bonds for the complex to form. The concentration of the AIF2+ complex ion increases with increasing concentrations of Al3+ andF-. 

Consider just a few cases of aqueous equilibria. The magnificent formations i n limestone caves and the vast expanses of oceanic coral reefs result from subtle shifts in carbonate solubility equilibria. Carbonates also influence soil pH and prevent acidification of lakes by acid rain. Equilibria involving carbon dioxide and phosphates help organisms maintain cellular pH within narrow limits. Equilibria involving clays in soils control the availability of ionic nutrients for plants. The principles of ionic equilibrium also govern how water is softened, how substances are purified by precipitation of unwanted ions, and even how the weak acids in wine and vinegar influence the delicate taste of a fine French sauce. In this chapter, we explore three aqueous ionic equilibrium systems acid-base buffers, slightly soluble salts, and complex ions. 

The final type of aqueous ionic equilibrium we consider involves a different kind of ion than we ve examined up to now. A simple ion, such as Na" " or S04 , consists of one or a few bound atoms, with an excess or deficit of electrons. A complex ion consists of a central metal ion covalently bonded to two or more anions or molecules, called ligands. Hydroxide, chloride, and cyanide ions are some ionic ligands water, carbon monoxide, and ammonia are some molecular ligands. In the complex ion Cr(NH3)6, for example, Cr is the central metal ion and six NH3 molecules are the ligands, giving an overall 3-1- charge (Figure 19.13). 

Chapter 15 Applications of Aqueous Equilibria" was a long chapter, which dealt with difficult material. In order to make this chapter more manageable for students, we have split this chapter in two "Chapter 15 Acid-Base Equilibrium" and "Chapter 16 Solubility and Complex Ion Equilibrium."... 

In several of the studies of aqueous chemistry of aluminum that have been made since about 1950, polynuclear complexing mechanisms have been proposed to identify and describe the dissolved aluminum hydroxide complex species (3, JO, 11). The formulae proposed have generally been based on stoichiometric considerations and pH measurements assuming the polynuclear species were ionic, and that equilibrium was attained. The complex ions reported by Hsu and Bates (8) were single six-mem-bered rings Ale (OH) 12 or multiples of this unit. Johansson (JO) identified a structural unit containing 13 aluminum and 40 oxygen atoms with various numbers of protons in crystalline basic aluminum sulfate. Because this solid formed readily, the same structural unit of aluminum was proposed as a solute species. Most of the proposed formulae for polynuclear complexes, however, have not been derived from structural considerations. 

Formation constant (stability constant) the equilibrium constant for each step of the formation of a complex ion by the addition of an individual ligand to a metal ion or complex ion in aqueous solution. (8.9)... 

An assembly of a metal ion and the Lewis bases bonded to it, such as Ag(NH3)2, is called a complex ion. The stability of a complex ion in aqueous solution can be judged by the size of the equilibrium constant for its formation from the hydrated metal ion. For example, the equilibrium constant for Equation 17.24 is... 

The use of formation constants to determine aqueous equilibrium in solutions that exhibit complex ion formation is illustrated in Example 12.16. 

A metal ion combined with a Lewis base, such as Ag(NHj>2 (aq), is called a complex ion. The equilibrium constant for the formation of a complex ion from a metal ion in aqueous solution is called the formation constant (A ) of the complex ion. The higher the value of Af the more stable is the complex ion. Values of the formation constants for some metal complex ions are given in Appendix IV. 

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