Orgánica

Most elements are found in nature not as free elements but in the form of ions. Hence, we begin our study of the chemistry of the elements with some chemical properties of their ions. In this chapter, we will investigate the interaction of some common cations and monoatomic anions of the elements with water and see how the periodic trends in these reaction tendencies can be related to the atomicproperties reviewed in chapter 1. Your instructor may choose to have you begin you study of the chemistry of these ions with a laboratory investigation (or classroom demonstration or video and discussion) of the process. You will find, upon analyzing the results, that even a simple reaction involves some important chemistry.
You instructor may now assign Experiment A.1. Then, we will construct aphysical model of what is happening during these reactions. You will find that the principles you derive, and the classification scheme that is developed, apply to far more than just the reaction of a cation or an anion with the humble water molecule; ions react similarly with many other chemical species. The ways in which you begin looking at positively and negatively charged species in thischapter will be useful in subsequent chapters.
Hydration of Cations.
When writing chemical equations for reactions of ions in solution we often write ions as if they were simple particles in solution; for example, we may write the sodium ion as Na+ or perhaps as Na+(aq). But there are definite reactions between ions and molecules of the polar solvent, water. Since the oxygen atom of the watermolecule is much more electronegative than the hydrogen atoms, each H-O bond is a polar covalent bond in which the bond electrons are (on the average) closer to the oxygen atom than the hydrogen atom, giving rise to a partial negative charge on oxygen and a partial positive charge on hydrogen. Since H2O is not a linear molecule, it has a partial negatively charged end (its oxygen end) and a partialpositively charged end (the hydrogen end). Since opposite charges attract, a positive ion (cation) placed in water surrounds itself with water molecules, with the oxygen ends inward toward the ion (Fig. 2.1). Conversely, a negative ion surrounds itself with water molecules, with the hydrogen ends inward. The products of these reactions are called hydrated ions.
The attraction of opposite charges isreally quite a strong force. If we were to plunge 1 mol of gaseous cations into water, they mould form hydrated ions and release a large amount of energy, which we call the hydration energy of the cation. Hydration energies of a number of cations are listed in Table 2.1; by any normal chemical standard these are large energies. The data in Table 2.1 show that the hydration energy of a cationdepends on the charge and the radius of the cation, as expected qualitatively from Coulomb´s law, and also depends on the electronegativity of the element. Latimer observed that if the electronegativity of the metal is not too great, the hydration energies of metal ions are given approximately by the equation.
ΔHhyd= -60,900Z2/(r+50)KJmol-1
Where Z is the charge on the cation and r is the cationicradius (in picometers). (We can loosely equate the constant added to the radius of the cation with the radius of the oxygen in the water.)
No attempt is made in Latimer´s equation to include the effects of electronegativity, but examination of the data for metals of Pauling electronegativities (Xp) the greater than 1.5 (on the right side of Table 2.1) shows that their hydration energies aresubstantially higher that those of ions of comparable radius and charge on the left side of the table. Such metals have electronegativities within about two units of that of oxygen, which suggests that for these metals there is not just an electrostatic attraction between the metal ion and the negative end of the water molecules, but that there also may be some degree of covalent bond formation, in...