Question

The charge/size ratio of a cation determines its polarizing power. Which one of the following sequences represents the increasing order of polarizing power?

• A. ${Ca}^{2+}<{Mg}^{2+}<{Be}^{2+}<K^{+}$
• B. ${Mg}^{2+}<{Be}^{2+}<K^{+}<{Ca}^{2+}$
• C. ${Be}^{2+}<K^{+}<{Ca}^{2+}<{Mg}^{2+}$
• D. $K^{+}<{Ca}^{2+}<{Mg}^{2+}<{Be}^{2+}$

Answer 1

The correct option is D

$K^{+}<{Ca}^{2+}<{Mg}^{2+}<{Be}^{2+}$

Polarizing power ($\varphi$) = $\frac{Cationcharge}{Cationsize}$

Among the given options,

Easily, $K^{+}$ has the least charge and the largest cationic radius. Thus $K^{+}$ is the least polarizing. With twice as much charge and slightly smaller cationic radius, $C{a}^{2+}$ has higher polarizing character than the $K^{+}$ ion. As we go down the group of Alkaline Earth Metals, the cationic radius increases doesn't it? With all the ions - $B{e}^{2+}$, $M{g}^{2+}$ & $C{a}^{2+}$ all having the same charge, the increasing order is:

$\therefore Theorderis{K}^{+}<C{a}^{2+}<M{g}^{2+}<B{e}^{2+}$