Ionic solutes in water

If the attractive strength of the water molecules is not sufficient to break the ionic bonds, then the compound does not dissolve. While most ionic compounds dissolve in water, a number of them do not. Precipitation formation in a chemical reaction is related to these concepts. If, after mixing two separate solutions together, there are two kinds of ions present that form a stronger ionic bond than the strength of the solvation bonds to water molecules, the ionic bonds will form, and an undissolved compound (or precipitate) will result. 

FIGURE 10.2 An illustration of an ionic compound, such is sodium chloride, dissolving in water. The water molecules, by virtue of their polar nature, pull the positive and negative ions from the ionic crystal array, and the compound dissolves. (From Kenkel, J., Kelter, P., and Hage, D., Chemistry An Industry-Based Introduction with CD-ROM, CRC Press, Boca Raton, FL, 2001. With Permission.) 

In general, we can make the following statements about common ionic compounds dissolving in water  

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Since the properties of an ionic solution (that is, a solution containing ions) differ in important ways from those of nonconducting solutions, it is important to be able to predict which substances are likely to form ionic solutions in water. The periodic table guides us. 

Using the periodic table as a guide, predict which of the following compounds form ionic solutions in water silicon carbide, SiC magnesium bromide, MgBr2 carbon tetrabromide, CBr chromic chloride, CrCl3. 

Hereafter in this chapter we shall be concerned exclusively with substances that form ionic solutions in water. Since each substance is electrically neutral before it dissolves, it must form ions of positive charge and, as well, ions of negative charge. Ions with positive charges are called cations. Ions with negative charges are called anions. A conducting solution is electrically neutral it contains both anions and cations. 

Nevertheless, many ionic solutes in water display standard hydrophobic behavior such as closed loop coexistence curves (Xu etal., 1991 Weingartner and Steinle, 1992). So the possibilities for ion hydration range from hydrophobic to primitive chemical interactions, and leave lots of room for molecular-scale complexity in between. 

Figure 13.5 shows enthalpy diagrams for dissolving three ionic solutes in water. The first, NaCl, has a small positive heat of solution (AZ/join = 3.9 kJ/mol). 

Explain how the dissolving of an ionic solute in water represents an equilihrium process. 

The interface distribution coeflBcient is a measure of the distortion the solute imposes on the lattice of the solid (148). The distortion is zero if k = 1. For ionic solutes in water, it is always very much smaller than one. This means that the solute weakens the binding between atoms, increasing their interatomic distances. Consequently, the melting point of the solvent is lowered. 

The smallness of the apparent distribution coeflBcients of ionic solutes in water suggests that all of them impose a considerable distortion on the ice lattice. Ionic distribution curves such as those discussed above show that there are differences in degree. Furthermore, the distortion imposed by a given ion depends also on its counterion. Thus, the ammonium ion in combination with the chloride ion is largely rejected, but in combination with the fluoride ion it is absorbed more readily than any other ion combination investigated. The freezing potential indicates that, even in NH4F, F" is somewhat more readily taken into the ice structure than NH4 ... 

Consider the following representations of an ionic solute in water. Which flask contains MgS04, and which flask contains NaCl How can you tell ... 

When yon come to a homework or exam question that involves dissolving an ionic solute in water or the saturated solution equihbrium process, caU your mental model into your working memory, and then nse your model to help answer the question. In this way, you are learning to think as a chemist. 

Some General Solubility Rules for Common Ionic Solutes in Water... 

Dissolution of ionic and ionizable solutes in water is favored by ion—dipole bonds between ions and water. Figure 6 illustrates a hydrated sodium ion,... 

Technology Description Hydrolysis is the process of breaking a bond in a molecule (which is ordinarily not water-soluble) so that it will go into ionic solution with water. Hydrolysis can be achieved by the addition of chemicals (e.g., acid hydrolysis), by irradiation (e.g., photolysis) or by biological action (e.g., enzymatic bond cleavage). The cloven molecule can then be further treated by other means to reduce toxicity. 

As we noted in Chapter 4, the solubility of ionic compounds in water varies tremendously from one solid to another. The extent to which solution occurs depends on a balance between two forces, both electrical in nature ... 

In Chapter 6 we saw that the chemistry of sodium can be understood in terms of the special stability of the inert gas electron population of neon. An electron can be pulled away from a sodium atom relatively easily to form a sodium ion, Na+. Chlorine, on the other hand, readily accepts an electron to form chloride ion, Cl-, achieving the inert gas population of argon. When sodium and chlorine react, the product, sodium chloride, is an ionic solid, made up of Na+ ions and Cl- ions packed in a regular lattice. Sodium chloride dissolves in water to give Na+(aq) and C (aq) ions. Sodium chloride is an electrolyte it forms a conducting solution in water. 

Table 1.1 summarizes the solubility patterns of common ionic compounds in water. Notice that all nitrates and all common compounds of the Group 1 metals are soluble so they make useful starting solutions for precipitation reactions. Any spectator ions can be used, provided that they remain in solution and do not otherwise react. For example, Table 1.1 shows that mercury(I) iodide, Hg2I2, is insoluble. It is formed as a precipitate when solutions containing Hg22+ ions and I ions are mixed ... 

Valence and oxidation state are directly related to the valence-shell electron configuration of a group. Binary hydrides are classified as saline, metallic, or molecular. Oxides of metals tend to be ionic and to form basic solutions in water. Oxides of nonmetals are molecular and many are the anhydrides of acids. 

The presence of ions in solution is what gives a sodium chloride solution the ability to conduct electricity. If positively and negatively charged wires are dipped into the solution, the ions in the solution respond to the charges on the wires. Chloride anions move toward the positive wire, and sodium cations move toward the negative wire. This directed movement of ions in solution is a flow of electrical current. Pure water, which has virtually no dissolved ions, does not conduct electricity. Any solution formed by dissolving an ionic solid in water conducts electricity. Ordinary tap water, for example, contains Ionic Impurities that make It an electrical conductor. 

In the first reaction, two ionic compounds in water are mixed. The AgCl formed by the swapping of anions is insoluble, causing the reaction to proceed. The solid AgCl formed from solution is an example of a precipitate. In the second reaction, a covalent compound, HzO, is formed from its ions in solution, H+ and OH, causing the reaction to proceed. In the third reaction, a solid reacts with the acid in solution to produce two covalent compounds. 

The titration curve of penicillamine hydrochloride at 25 °C revealed the presence of three ionizable groups with pKa values of 1.8 (carboxyl group), 7.9 (oc-amino group), and 10.5 (/J-thiol group). Recently, the ionization constants for the acidic functions of (D)-penicillamine were verified by pH titration at 37 °C and 0.15 M ionic strength [2], A 1% solution in water has a pH of 4.5-5.5 [3],... 

Carell and Olin (58) were the first to derive thermodynamic functions relating to beryllium hydrolysis. They determined the enthalpy and entropy of formation of the species Be2(OH)3+ and Be3(OH)3+. Subsequently, Mesmer and Baes determined the enthalpies for these two species from the temperature variation of the respective equilibrium constants. They also determined a value for the species Be5(OH) + (66). Ishiguro and Ohtaki measured the enthalpies of formation of Be2(OH)3+ and Be3(OH)3+ calorimetrically in solution in water and water/dioxan mixtures (99). The agreement between the values is satisfactory considering the fact that they were obtained with different chemical models and ionic media. 

By batch description trials Organo- and inorganically- modified zealot was subjected up to 24 hr in distilled water, tap water and 2% Nalco aqueous solutions in laboratory shaken machine to demonstrate how strongly the examined oxyanions are bound on the modified zeolite. While only slightly chromate desorption in the maximum extent about 20 mg/L was observed, approximately one order higher arsenate desorption was found, corresponding to increased ionic strength in waters. However, in both cases ODA-clinoptilolite exhibited the lowest desorption characteristics. Here, the... 

Another solution to the problem of catalyst/product separation is the biphasic catalysis. The liquid biphasic catalysis became an attractive technology for potential commercial application of enantioselective homogeneous catalysis. The most important features of such systems are related to the fact that both reaction rate and e.s. may be influenced by the number of ionic groups in water-soluble ligand or by addition of surfactants. Descriptions of water-soluble ligands and the recent results in the rapidly progressing area of biphasic enantioselective catalysis are available in recent reviews [255,256],... 

The non-ionic surfactants do not produce ions in aqueous solution. The solubility of non-ionic surfactants in water is due to the presence of functional groups in the molecules that have a strong affinity for water. Similarly to the anionic surfactants, and any other group of surfactants, they also show the same general property of these products, which is the reduction of the surface tension of water. 

To understand the dissolution of ionic solids in water, lattice energies must be considered. The lattice enthalpy, A Hh of a crystalline ionic solid is defined as the energy released when one mole of solid is formed from its constituent ions in the gas phase. The hydration enthalpy, A Hh, of an ion is the energy released when one mole of the gas phase ion is dissolved in water. Comparison of the two values allows one to determine the enthalpy of solution, AHs, and whether an ionic solid will dissolve endothermically or exothermically. Figure 1.4 shows a comparison of AH and A//h, demonstrating that AgF dissolves exothermically. 

Breitschwerdt, K. G. Structural relaxation in water and ionic solutions. In Structure of water and aqueous solutions, pp. 473—490 (Luck, W. A. P., ed.). Weinheim Verlag Chemie 1974. 

In this chapter, you will continue your study of acid-base reactions. You will find out how ions in aqueous solution can act as acids or bases. Then, by applying equilibrium concepts to ions in solution, you will be able to predict the solubility of ionic compounds in water and the formation of a precipitate. 

The solubility of ionic solids in water covers a wide range of values. Knowing the concentration of ions in aqueous solution is important in medicine and in chemical analysis. In this section, you will continue to study equilibrium. You will examine the solubility equilibria of ionic compounds in water. 

In this section, you determined the solubility product constant, Kgp, based on solubility data. You obtained your own solubility data and used these data to calculate a value for Kgp. You determined the molar solubility of ionic solutions in pure water and in solutions of common ions, based on their Ksp values. In section 9.3, you will further explore the implications of Le Chatelier s principle. You will use a reaction quotient, Qsp, to predict whether a precipitate forms. As well, you will learn how selective precipitation can be used to identify ions in solution. 

In this chapter we discuss water and ionic solutions, in Chapter 3, structure of metals and metal surfaces and in Chapter 4, the formation and structure of the metal-solution interface. Discussion is limited to those topics that are directly relevant to the electrodeposition processes and the properties of electrodeposits. 

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