Question

In theory, most metals should easily corrode in air. Why?

A group of metals called noble metals is relatively difficult to corrode in the air. Some noble metals include gold, platinum, and silver. Reference Table 11.1 to come up with a possible reason why the noble metals are relatively difficult to corrode.

Short answer

The highly corrosive nature of certain metals is due to their lower reduction potential whereas the ones with higher reduction potential do not corrode easily.

Step-by-step solution

 Step 1: Explaining why most metals corrode easily

Corrosion is basically oxidation of the metal. It involves two reactions,

$$\begin{gathered} {\mathrm{O}}_2+4{\mathrm{H}}^{+}+4{\mathrm{e}}^{-}\to 2{\mathrm{H}}_2\mathrm{O}{\mathrm{E}}^{\mathrm{o}}=1.23\mathrm{V} \\ \mathrm{M}\to {\mathrm{M}}^{+}+{\mathrm{e}}^{-}{\mathrm{E}}^{\mathrm{o}}={\mathrm{E}}_1 \end{gathered}$$

If the E^ofor sum of these reactions is positive reaction will be spontaneous and corrosion will take place. For example

$$\begin{gathered} {\mathrm{O}}_2+4{\mathrm{H}}^{+}+4{\mathrm{e}}^{-}\to 2{\mathrm{H}}_2\mathrm{O}{\mathrm{E}}^{\mathrm{o}}=1.23\mathrm{V} \\ \mathrm{Fe}\to {\mathrm{Fe}}^{2+}+2{\mathrm{e}}^{-}{\mathrm{E}}^{\mathrm{o}}=-0.44\mathrm{V} \\ \mathrm{Overall} \mathrm{reaction}: \\ {\mathrm{O}}_2+4{\mathrm{H}}^{+}+2\mathrm{Fe}\to 2{\mathrm{H}}_2\mathrm{O}+{\mathrm{Fe}}^{2+} \end{gathered}$$

$\begin{array}{rcl}& & {{\text{E}}^{\text{o}}}_{\text{cell}}\text{=}{{\text{E}}^{\text{o}}}_{\text{cathode}}\text{-}{{\text{E}}^{\text{o}}}_{\text{anode}}\\ & & \text{= 1.23V - ( - 0.44V)}\\ & & \text{= 1.67V}\end{array}$

As E^o_cell is positive, corrosion will take place.

Explanation

Metals like gold (Au) have a higher reduction potential (around 1.52 V). They do not result in a positive standard value of E And do not corrode easily.