Halide Ions (Cl⁻, Br⁻, I⁻)

Last updated
Save as PDF

Page ID: 97282

\( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \) \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)\(\newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\) \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\) \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\) \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\) \( \newcommand{\Span}{\mathrm{span}}\) \(\newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\) \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\) \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\) \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\) \( \newcommand{\Span}{\mathrm{span}}\)\(\newcommand{\AA}{\unicode[.8,0]{x212B}}\)

Acid Equilibria

These ions are all very weak bases since they are the conjugate bases of very strong acids. Hence, they undergo negligible hydrolysis.

Solubility

Most halide salts are soluble. Exceptions are the halide salts of silver, lead(II), and mercury(I). For example, the solubility of the silver salts is indeed very low, as shown by their solubility product constants:

\[\ce{AgCl(s) <=> Ag^{+}(aq) + Cl^{-}(aq)} \nonumber \]

with \(K_{sp} = 1.2 \times 10^{-10}\)

\[\ce{AgBr(s) <=> Ag^{+}(aq) + Br^{-}(aq)} \nonumber \]

with \(K_{sp} = 4.8 \times 10^{-13}\)

\[\ce{AgI(s) <=> Ag^{+}(aq) + I^{-}(aq)} \nonumber \]

with \(K_{sp} = 1.4 \times 10^{-16}\)

The silver halide solubility can be increased by addition of ammonia in appropriate concentrations, due to complex ion formation:

\[\ce{Ag^{+}(aq) + 2NH3(aq) <=> [Ag(NH3)2]^{+}(aq)} \nonumber \]

with \(K_f = 1.5 \times 10^7\).

The less soluble the silver halide, the greater the concentration of ammonia needed to dissolve the silver halide. \(\ce{AgCl}\) dissolves in 6 M \(\ce{NH3(aq)}\), while \(\ce{AgBr}\) dissolves in 15 M \(\ce{NH3(aq)}\) (the concentrated reagent). \(\ce{AgI}\) will not dissolve even in 15 M \(\ce{NH3(aq)}\). Thus, adding an appropriate concentration of aqueous ammonia can be used to separate the silver halides from one another.

Oxidation Reduction

As reducing agents, the halide ions follow the trend in reducing strength:

\[\ce{I^{-} > Br^{-} > Cl^{-}}. \nonumber \]

Conversely, the halogens follow the opposite order of oxidizing strength:

\[\ce{Cl2 > Br2 > I2} \nonumber \]

Thus pale green \(\ce{Cl2}\) oxidizes \(\ce{Br^{-}}\) to red \(\ce{Br2}\) and \(\ce{I^{-}}\) to violet \(\ce{I2}\). These colors are better observed if the halogens are extracted into a small amount of hexane. CAUTION: hexane is flammable.

Acid Equilibria

Solubility

Oxidation ­Reduction

Oxidation Reduction