Ionic compounds dissociation

Molecular compounds ionize, ionic compounds dissociate. 

Solid ionic compound Dissociated ions in solution... 

Q Why do ionic compounds dissociate into atoms in the gas phase ... 

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. 

Some ionic compounds dissolve readily in water and some barely dissolve at all. Sodium chloride, or table salt, is typical of the soluble ionic compounds. On dissolving, all ionic compounds dissociate into ions. 

When water is added, the polar water molecules surround the sodium and chloride ions, and the ionic compound dissociates. Because water molecules are polar and have a negative end and a positive end, they are attracted to both the positively charged sodium ions and the negatively charged chloride ions. Water molecules surround both types of ions. The opposite charges attract. 

Ionic compounds dissociate when they dissolve in water. 

Any substance whose aqueous solution contains ions is called an electrolyte. Any substance that forms a solution containing no ions is a nonelectrolyte. Electrolytes that are present in solution entirely as ions are strong electrolytes, whereas those that are present partly as ions and partly as molecules are weak electrolytes. Ionic compounds dissociate into ions when they dissolve, and they are strong electrolytes. The solubility of ionic substances is made possible by solvation, the interaction of ions with polar solvent molecules. Most molecular compounds are nonelectrolytes, although some are weak electrolytes, and a few are strong electrolytes. When representing the ionization of a weak electrolyte in solution, half-arrows in both directions are used, indicating that the forward and reverse reactions can achieve a chemical balance called a chemical equilibrium. 

Finally, we generally assume that ionic compounds dissociate completely when they dissolve, but this assumption is not always valid. When MgF2 dissolves, for example, it yields not only Mg " and F ions but also MgF ions. 

Frequently, the compounds involved in aqueous chemistry are ionic, as is true for both HNO3 and NH4NO3 in this example. These ionic compounds dissociate when they dissolve in water, so the actual solution does not contain intact molecules of these compounds. This dissociation is emphasized if we write the total ionic equation rather than the molecular equation. This form emphasizes what is actually present in the reacting mixture by writing dissociated compounds as separated ions in the solution. The total ionic equation for our example reaction from above is thus... 

Ionic compounds such as NaOH are composed of positive and negative ions. In solution, soluble ionic compounds dissociate into their component ions. Molecular compounds containing an OH group, such as methanol CH3OH, do not dissociate and therefore do not act as bases. 

Soluble ionic compounds dissociate into their component ions when dissolved in an excess of water. The concentrations of the individual ions will depend on the amounts of these ions when the substance (salt, base or alkali) dissolves. In a 4-0 mol dm aqueous solution of aluminium nitrate, for example, the concentration of the aluminium ions is 4.0 mol dm but the concentration of the nitrate ions is 12.0 mol dm. ... 

Many chemical reactions that occur in water solutions are reactions involving ions. Soluble ionic compounds dissociate into ions when they dissolve, and some molecular compounds,... 

Because each AgCl unit contains only one Ag and one CF ion, its solubility product expression is particularly simple to write. Many ionic compounds dissociate into more than two ions. Table 17.4 lists a number of slightly soluble ionic compounds along with equations representing their dissolution equilibria and their solubility product constants. (Compounds deemed soluble by the solubility rules in Chapter 4 are not listed for the same reason we did not list values for the strong acids in Table 16.6.) In general, the magnitude of indicates the solubility of an ionic compound—the smaller the value, the less soluble the compound. To make a direct... 

The solubility product constant (Kgp) is the equilibrium constant that indicates to what extent a slightly soluble ionic compound dissociates in water. 

In a solution of sodium chloride, all the dissolved compound exists in the form of ions. Thus, NaCl, which is an ionic compound [ Section 2.7], is said to have dissociated completely. An electrolyte that dissociates completely is known as a strong electrolyte. All water-soluble ionic compounds dissociate completely upon dissolving, so all water-soluble ionic compounds are strong electrolytes. 

In the absence of die polyether, potassium fluoride is insoluble in benzene and unreactive toward alkyl halides. Similar enhancement of solubility and reactivity of other salts is observed in the presence of crown ethers The solubility and reactivity enhancement result because the ionic compound is dissociated to a tightly complexed cation and a naked anion. Figure 4.13 shows the tight coordination that can be achieved with a typical crown ether. The complexed cation, because it is surrounded by the nonpolar crown ether, has high solubility in the nonpolar media. To maintain electroneutrality, the anion is also transported into the solvent. The cation is shielded from interaction with the anion as a... 

H2SO4.Z2H2O, are known with = 1, 2, 3, 4 (mps 8.5", -39.5". -36.4" and -28.3% respectively). Other compounds in the H2O/SO3 system are H2S2O7 (mp 36") and H2S4O13 (mp 4"). Anhydrous H2SO4 is a remarkable compound with an unusually high dielectric constant, and a very high electrical conductivity which results from the ionic self-dissociation (autoprotolysis) of the compound coupled with a proton-switch mechanism for the rapid... 

A final point to bear in mind is that, when a reaction involves fully dissociated ionic compounds in solution, then the equilibrium constant should be written for the net ionic equation, by using the activity for each type of ion. 

For a given molecule and a given intemuclear separation a would have a definite value, such as to make the energy level for P+ lie as low as possible. If a happens to be nearly 1 for the equilibrium state of the molecule, it would be convenient to say that the bond is an electron-pair bond if a is nearly zero, it could be called an ionic bond. This definition is somewhat unsatisfactory in that it does not depend on easily observable quantities. For example, a compound which is ionic by the above definition might dissociate adiabatically into neutral atoms, the value of a changing from nearly zero to unity as the nuclei separate, and it would do this in case the electron affinity of X were less than the ionization potential of M. HF is an example of such a compound. There is evidence, given bdow, that the normal molecule approximates an ionic compound yet it would dissociate adiabatically into neutral F and H.13... 

Ionization refers to the process in which a molecular compound, such as HC1, separates or reacts with water to form ions in solution. Dissociation refers to the process in which a solid ionic compound, such as NaCl, separates into its ions in aqueous solution. 

The electrochemical behaviour of the compounds containing bonds between silicon and other group-14-metals is also interesting. Mochida et al. reported the electrochemical oxidation potentials of group-14-dimetals [66], As shown in Table 8, there is a good correlation between the oxidation potentials and the ionization potentials which decrease in the order Si-Si > Si-Ge > Ge-Ge > Si-Sn > Ge-Sn > Sn-Sn in accord with the metal-metal ionic bond dissociation energy. 

Disregarding such chemical-specific properties as dissociation constants (in the case of ionic compounds), particle size, and polymorphism, as well as side effects of viscosity, binding to vehicle components, complex formulation, and the like, the following formulation principles arise ... 

Many of the reactions that you will study occur in aqueous solution. Water readily dissolves many ionic compounds as well as some covalent compounds. Ionic compounds that dissolve in water (dissociate) form electrolyte solutions— solutions that conduct electrical current due to the presence of ions. We may classify electrolytes as either strong or weak. Strong electrolytes dissociate (break apart or ionize) completely in solution, while weak electrolytes only partially dissociate. Even though many ionic compounds dissolve in water, many do not. If the attraction of the oppositely charged ions in the solid is greater than the attraction of the water molecules to the ions, then the salt will not dissolve to an appreciable amount. 

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. 

For example, suppose that you slowly add an ionic compound, such as magnesium sulfate, MgS04, to water in a beaker. The following equation represents the dissociation. 

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. 

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