Water as a Lewis base

Although in this chapter we are mainly concerned with Bronsted acids and bases, it is important not to lose sight of the definition of Lewis acids and bases. Relevant to this chapter is the fact that water functions as a Lewis base when it acts as a solvent. 

A Lewis acid is an electron acceptor, and a Lewis base is an electron donor. 

Hydration is the specific case of solvation when the solvent is water. 

In the aqueous chemistry of cations, hydrolysis refers to the reversible loss of from an aqua species. The term hydrolysis is, however, also used in a wider context, e.g. the reaction  

PCI3 -f 3H2O H3PO3 + 3HC1 is a hydrolysis process. 

VIA Fig. 6.6 If the plane of each water molecule in 9 [M(H20)g]+ makes an angle of 50° with the M O axis, it suggests that the metal-oxygen interaction involves the use of an oxygen lone pair. 

Spectroscopic studies suggest that the hydration of other halide ions is similar to that of Cr, but for more complex anions, very few data are available. For a limited number of hydrated cations, tracer methods and electronic and NMR spectroscopies provide reliable information about coordination number and stoichiometry. 

The configurations 7.6 and 7.7 have been established in the first hydration shell for dilute solutions of LiCl and NaCl by neutron diffraction studies. In concentrated solutions, the plane of the water molecule in 7.6 makes an angle of up to 50° with the O axis (Fig. 7.6) implying interaction of the cation with a lone pair of electrons rather than an ion-dipole interaction. 

For both the cation and anion in NaCl, there are six H2O molecules in the primary hydration shell (Fig. 7.5). Spectroscopic studies suggest that the hydration of other halide irais is similar to that of CP. Experimental techniques that are used to investigate the hydration shells around metal icMis include X-ray and neutron diffraction extended X-ray absorption fine structure (EXAFS) spectroscopy and NMR (particularly O NMR) spectroscopy. Modem computatiOTial methods (e.g. molecular dynamics) are also invaluable.  

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Lewis acid and water as a Lewis base, and a structure of defined coordination number with several Mn+ — OI I2 bonds results an experimentally-determined M—O—H angle of 130° is consistent with involvement of a lone pair on the approximately tetrahedral oxygen. The coordinate bond is at the core of all natural and synthetic complexes. 

Water is also a Lewis base, because an HzO molecule can donate one of its lone pairs to a Lewis acid and form complexes such as Fe(H20)63+. Water s ability to act as a Lewis base is also the origin of its ability to hydrolyze substances. The reaction between water and phosphorus pentachloride mentioned in Section 15.2 is an example. 

B Boric acid acts as a Lewis acid. The boron atom in B(OH)3 has an incomplete octet and forms a bond by accepting a lone pair of electrons from a water molecule, which is acting as a Lewis base. The complex formed is a weak Bronsted acid in which an acidic proton can be lost from the H20 molecule in the complex. 

P Pr3)2. The protonation of OsHMe(CO)2(P Pr3)2 with a diethyl ether solution of HBF4 in the presence of acetone leads to the quantitative formation of the cationic hydride complex [OsH(CO)2 ii1-OCMe2 (P,Pr3)2]BF4. If water instead of acetone is used as a Lewis base, the aquo hydride compound [0sH(C0)2(H20) (P Pr3)2]BF4 is obtained also in excellent yield. 

A water molecule acting as a Lewis base donates an electron pair to the carbocation (a Lewis acid). This gives the cationic carbon eight electrons. 

According to Lewis theory, water can act as a Lewis base (electron pair donor). Water contains an unshared pair of electrons that is utilized in accepting a proton to form the hydronium ion. 

In reactions involving coordination compounds, the metal acts as the Lewis acid (electron-pair acceptor), while the ligand acts as a Lewis base (electron-pair donor). In the reaction above, the ammonia ligand displaced the water ligand from the chromium complex because nitrogen is a better electron-pair donor (less electronegative) than oxygen. 

In the adsorption of water molecules on metal electrodes in aqueous solutions, unpaired electrons in the frontier orbital of oi en atoms in water molecules form covalent bonds with surface metal atoms. Then, the adsorbate water molecules act as a Lewis base (covalent-electron providers) and the adsorbent surface metal atoms act as a Lewis acid (covalent-electron receivers). Since the bond energy (0.4 to 0.7 eV) of water molecules with the surface metal atoms is close to the energy of hydrogen bond (0.2 to 0.4 eV) between water molecules, the adsorbed water molecule is combined not only with the metallic surface atoms but also with the acijacent water molecules to form a bi-molecular layer rather than a monomer layer as shown in Fig. 5-31. 

Charged metals (cations) in water behave as Lewis acids (willing to accept electrons). Water on the other hand, because it is willing to share its two unshared oxygen-associated pair of electrons, behaves as a Lewis base. Strong H2Q-metal (Lewis base-Lewis acid) interactions allow H+ on the water molecule to dissociate, hence, low pH water is produced. The degree of dissociation of water interacting with a cation (Mn+) is described by the metal hydrolysis constant (Table 2A)... 

Water is an excellent binding molecule because it can act as a Lewis base by coordinating to a meted cation, and as a Lewis acid by coordinating to a given anion through its hydrogen atoms. 

Our studies have shown that almost all metal atoms will form an adduct with water and in many instances undergo further reaction. Theory and experiment (1 -7 ) support the concept of water acting as a Lewis base which donates electron density to the metal. Theory has also shown for Be and Mg that donation occurs extensively from the 3ai (a lone pair) orbital of water to the metal. This is of interest since this orbital is largely responsible for the nonlinearity of the water molecule. It is known, for instance, that ionization from this orbital causes water to become linear and results in a decrease in its bending frequency by approximately 700 cm . This is consistent with our finding that only the V2 bending mode of water decreases measurably upon adduct formation. The observed Av2 changes for metal-water adducts are compared and discussed below. 

The structure of the metal-water adduct has been determined experimentally for Li(, 7) and theoretically for Li, Na( ), Be, Mg(3,4), Ca(5), and Al(2). In all cases water is acting as a Lewis base and donating electron density to the metal through the oxygen. For all metals but Be the minimum energy geometry is planar. For Be, the Be-0 bond is calculated to be symmetrical... 

The final acid-base theory that we shall consider was proposed by chemist Gilbert Lewis in the early 1920s. The Lewis Theory is the most general, including more substances under its definitions than the other theories of acids and bases. A Lewis acid is a substance that accepts a pair of electrons to form a covalent bond. A Lewis base is a substance that provides a pair of electrons to form a covalent bond. In order for a substance to act as a Lewis base, it must have a pair of unshared electrons in its valence shell. An example of this is seen when a hydrogen ion attaches to the unpaired electrons of oxygen in a water molecule, as shown here ... 

The American chemist G. N. Lewis formulated such a definition. He defined what we now call a Lewis base as a substance that can donate a pair of electrons. A. Lewis acid is a substance that can accept a pair of electrons. For example, in the protonation of ammonia, NH3 acts as a Lewis base because it donates a pair of electrons to the proton H, which acts as a Lewis acid by accepting the pair of electrons. A Lewis acid-base reaction, therefore, is one that involves the donation of a pair of electrons from one species to another. Such a reaction does not produce a salt and water. 

Phosphonium halides, PH4X, are formed by treating PH3 with HX but only the iodide is stable under ambient conditions. The chloride is unstable above 243 K and the bromide decomposes at 273 K. The [PH4] ion is decomposed by water (equation 14.29). Phosphine acts as a Lewis base and a range of adducts (including those with low oxidation state rf-block metal centres) are known. Examples include H3B-PH3, C13B-PH3, Ni(PH3)4 (decomposes above 243 K) and Ni(CO)2(PH3)2. Combustion of PH3 5delds H3PO4. 

The Lewis acid-base definition focuses on the donation or acceptance of an electron pair to form a new covalent bond in an adduct, the product of an acid-base reaction. Lewis bases donate the electron pair, and Lewis acids accept it. Thus, many species that do not contain El are Lewis acids. Molecules with polar double bonds act as Lewis acids, as do those with electron-deficient atoms. Metal ions act as Lewis acids when they dissolve in water, which acts as a Lewis base, to form an adduct, a hydrated cation. Many metal ions function as Lewis acids in biomolecules. 

As we discussed in Section 18.8, all complex ions are Lewis adducts. The metal ion acts as a Lewis acid (accepts an electron pair) and the ligand acts as a Lewis base (donates an electron pair). The acidic hydrated metal ions that we discussed in Section 18.6 are complex ions with water molecules as ligands. In Chapter 22, we discuss the transition metals and the structures and properties of the numerous complex ions they form. Our focus here is on equilibria of hydrated ions with ligands other than water. 

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