ACID,BASE
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Solutions of simple salts of metal ions can also be acidic, even though a metal ion cannot donate a proton directly to water to produce H3O+H3O+. Instead, a metal ion can act as a Lewis acid and interact with water, a Lewis base, bycoordinatingtoalonepairofelectronsontheoxygenatomtoformahydratedmetal ion (part (a) in Figure 16.9.116.9.1). A water molecule coordinated to a metal ion is more acidic than a free water molecule for two reasons. First, repulsive electrostatic interactions between the positively charged metal ion and the partially positively charged hydrogen atoms of the coordinated water molecule make it easier for the coordinated water to lose a proton.Loading...
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Solutions of simple salts of metal ions can also be acidic, even though a metal ion cannot donate a proton directly to water to produce H3O+H3O+. Instead, a metal ion can act as a Lewis acid and interact with water, a Lewis base, bycoordinatingtoalonepairofelectronsontheoxygenatomtoformahydratedmetal ion (part (a) in Figure 16.9.116.9.1). A water molecule coordinated to a metal ion is more acidic than a free water molecule for two reasons. First, repulsive electrostatic interactions between the positively charged metal ion and the partially positively charged hydrogen atoms of the coordinated water molecule make it easier for the coordinated water to lose a proton.Loading...
Figure 16.9.116.9.1: Effect of a Metal Ion on
theAcidityofWater(a)Reactionofthemetal ion Al3+
Al3+ with water to form the hydrated metal ion is an example of a Lewis acid–base reaction. (b) The positive charge on the aluminum ion attracts electron density from the oxygen atoms, which shifts electron density away from the O–H bonds. The decrease in electron density weakens the O–H bonds in the water molecules and makes it easier for them to lose a proton.
theAcidityofWater(a)Reactionofthemetal ion Al3+
Al3+ with water to form the hydrated metal ion is an example of a Lewis acid–base reaction. (b) The positive charge on the aluminum ion attracts electron density from the oxygen atoms, which shifts electron density away from the O–H bonds. The decrease in electron density weakens the O–H bonds in the water molecules and makes it easier for them to lose a proton.
Figure 16.9.116.9.1: Effect of a Metal Ion on
theAcidityofWater(a)Reactionofthemetal ion Al3+
Al3+ with water to form the hydrated metal ion is an example of a Lewis acid–base reaction. (b) The positive charge on the aluminum ion attracts electron density from the oxygen atoms, which shifts electron density away from the O–H bonds. The decrease in electron density weakens the O–H bonds in the water molecules and makes it easier for them to lose a proton.
theAcidityofWater(a)Reactionofthemetal ion Al3+
Al3+ with water to form the hydrated metal ion is an example of a Lewis acid–base reaction. (b) The positive charge on the aluminum ion attracts electron density from the oxygen atoms, which shifts electron density away from the O–H bonds. The decrease in electron density weakens the O–H bonds in the water molecules and makes it easier for them to lose a proton.
The magnitude of this effect depends on the following two factors (Figure 16.9.2
16.9.2):
The charge on the metal ion. A divalent ion (M
2+
M2+) has approximately twice as strong an effect on the electron density in a coordinated water molecule as a monovalent ion (M
+
M+) of the same radius.
The radius of the metal ion. For metal ions with the same charge, the smaller the ion, the shorter the internuclear distance to the oxygen atom of the water molecule and the greater the effect of the metal on the electron density distribution in the water molecule.
16.9.2):
The charge on the metal ion. A divalent ion (M
2+
M2+) has approximately twice as strong an effect on the electron density in a coordinated water molecule as a monovalent ion (M
+
M+) of the same radius.
The radius of the metal ion. For metal ions with the same charge, the smaller the ion, the shorter the internuclear distance to the oxygen atom of the water molecule and the greater the effect of the metal on the electron density distribution in the water molecule.
The magnitude of this effect depends on the following two factors (Figure 16.9.2
16.9.2):
The charge on the metal ion. A divalent ion (M
2+
M2+) has approximately twice as strong an effect on the electron density in a coordinated water molecule as a monovalent ion (M
+
M+) of the same radius.
The radius of the metal ion. For metal ions with the same charge, the smaller the ion, the shorter the internuclear distance to the oxygen atom of the water molecule and the greater the effect of the metal on the electron density distribution in the water molecule.
16.9.2):
The charge on the metal ion. A divalent ion (M
2+
M2+) has approximately twice as strong an effect on the electron density in a coordinated water molecule as a monovalent ion (M
+
M+) of the same radius.
The radius of the metal ion. For metal ions with the same charge, the smaller the ion, the shorter the internuclear distance to the oxygen atom of the water molecule and the greater the effect of the metal on the electron density distribution in the water molecule.
Figure 16.9.2
16.9.2: The Effect of the Charge and Radius of a Metal Ion on the Acidity of a Coordinated Water Molecule. The contours show the electron density on the O atoms and the H atoms in both a free water molecule (left) and water molecules coordinated to Na+Na+, Mg2+Mg2+, and Al3+Al3+ ions. These contour maps demonstrate that the smallest, most highly charged metal ion (Al
3+
Al3+) causes the greatest decrease in electron density of the O–H bonds of the water molecule. Due to this effect, the acidity of hydrated metal ions increases as the charge on the metal ion increases and its radius decreases.
Thus aqueous solutions of small, highly charged metal ions, such as Al
3+Al3+ and Fe3+Fe3+, are acidic:
16.9.2: The Effect of the Charge and Radius of a Metal Ion on the Acidity of a Coordinated Water Molecule. The contours show the electron density on the O atoms and the H atoms in both a free water molecule (left) and water molecules coordinated to Na+Na+, Mg2+Mg2+, and Al3+Al3+ ions. These contour maps demonstrate that the smallest, most highly charged metal ion (Al
3+
Al3+) causes the greatest decrease in electron density of the O–H bonds of the water molecule. Due to this effect, the acidity of hydrated metal ions increases as the charge on the metal ion increases and its radius decreases.
Thus aqueous solutions of small, highly charged metal ions, such as Al
3+Al3+ and Fe3+Fe3+, are acidic:
The [Al(H2O)
6]3+[Al(H2O)6]3+ ion has a pKapKa of 5.0, making it almost as strong an acid as acetic acid. Because of the two factors described previously, the most important parameter for predicting the effect of a metal ion on the acidity of coordinated water molecules is the charge-to-radius ratio of the metal ion. A number of pairs of metal ions that lie on a diagonal line in the periodic table, such as Li+Li+ and Mg2+Mg2+ or C2+Ca2+ and Y3+Y3+, have different sizes and charges, but similar charge-to-radius ratios. As a result, these pairs of metal ions have similar effects on the acidity of coordinated water molecules, and they often exhibit other significant similarities in chemistry as well.
6]3+[Al(H2O)6]3+ ion has a pKapKa of 5.0, making it almost as strong an acid as acetic acid. Because of the two factors described previously, the most important parameter for predicting the effect of a metal ion on the acidity of coordinated water molecules is the charge-to-radius ratio of the metal ion. A number of pairs of metal ions that lie on a diagonal line in the periodic table, such as Li+Li+ and Mg2+Mg2+ or C2+Ca2+ and Y3+Y3+, have different sizes and charges, but similar charge-to-radius ratios. As a result, these pairs of metal ions have similar effects on the acidity of coordinated water molecules, and they often exhibit other significant similarities in chemistry as well.