Chemistry » Metals, Metalloids, and Nonmetals » Occurrence and Preparation of the Representative Metals

Chemical Reduction

Chemical Reduction

It is possible to isolate many of the representative metals by chemical reduction using other elements as reducing agents. In general, chemical reduction is much less expensive than electrolysis, and for this reason, chemical reduction is the method of choice for the isolation of these elements.

For example, it is possible to produce potassium, rubidium, and cesium by chemical reduction, as it is possible to reduce the molten chlorides of these metals with sodium metal. This may be surprising given that these metals are more reactive than sodium; however, the metals formed are more volatile than sodium and can be distilled for collection. The removal of the metal vapor leads to a shift in the equilibrium to produce more metal (see how reactions can be driven in the discussions of Le Châtelier’s principle in the tutorial on fundamental equilibrium concepts).

The production of magnesium, zinc, and tin provide additional examples of chemical reduction.

The Preparation of Magnesium

The Pidgeon process involves the reaction of magnesium oxide with elemental silicon at high temperatures to form pure magnesium:

\(\text{Si}(s)+\text{2MgO}(s)\;\stackrel{\phantom{\rule{0.4em}{0ex}}\text{Δ}\phantom{\rule{0.4em}{0ex}}}{\to }\;{\text{SiO}}_{2}(s)+\text{2Mg}(g)\)

Although this reaction is unfavorable in terms of thermodynamics, the removal of the magnesium vapor produced takes advantage of Le Châtelier’s principle to continue the forward progress of the reaction. Over 75% of the world’s production of magnesium, primarily in China, comes from this process.

The Preparation of Zinc

Zinc ores usually contain zinc sulfide, zinc oxide, or zinc carbonate. After separation of these compounds from the ores, heating in air converts the ore to zinc oxide by one of the following reactions:

\(\text{2ZnS}(s)+{\text{3O}}_{2}(g)\;\stackrel{\phantom{\rule{0.4em}{0ex}}\text{Δ}\phantom{\rule{0.4em}{0ex}}}{\to }\;\text{2ZnO}(s)+{\text{2SO}}_{2}(g)\)

\({\text{ZnCO}}_{3}(s)\;\stackrel{\phantom{\rule{0.4em}{0ex}}\text{Δ}\phantom{\rule{0.4em}{0ex}}}{\to }\;\text{ZnO}(s)+{\text{CO}}_{2}(g)\)

Carbon, in the form of coal, reduces the zinc oxide to form zinc vapor:

\(\text{ZnO}(s)+\text{C}(s)\;⟶\;\text{Zn}(g)+\text{CO}(g)\)

The zinc can be distilled (boiling point 907 °C) and condensed. This zinc contains impurities of cadmium (767 °C), iron (2862 °C), lead (1750 °C), and arsenic (613 °C). Careful redistillation produces pure zinc. Arsenic and cadmium are distilled from the zinc because they have lower boiling points. At higher temperatures, the zinc is distilled from the other impurities, mainly lead and iron.

The Preparation of Tin

The ready reduction of tin(IV) oxide by the hot coals of a campfire accounts for the knowledge of tin in the ancient world. In the modern process, the roasting of tin ores containing SnO2 removes contaminants such as arsenic and sulfur as volatile oxides. Treatment of the remaining material with hydrochloric acid removes the oxides of other metals. Heating the purified ore with carbon at temperature above 1000 °C produces tin:

\({\text{SnO}}_{2}(s)+\text{2C}(s)\;\stackrel{\phantom{\rule{0.4em}{0ex}}\text{Δ}\phantom{\rule{0.4em}{0ex}}}{\to }\;\text{Sn}(s)+\text{2CO}(g)\)

The molten tin collects at the bottom of the furnace and is drawn off and cast into blocks.

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