18.2 Occurrence and Preparation of the Representative Metals

Occurrence and Extraction of Representative Metals

Most representative metals don't exist as pure elements in nature. Instead, they're locked inside minerals and ores, and getting them out requires specific metallurgical techniques. The method you choose depends on how reactive the metal is: highly reactive metals need electrolysis, while less reactive ones can be extracted through chemical reduction.

Representative Metals in Minerals

Representative metals naturally occur in minerals and ores. Understanding the difference between these two terms matters:

• Minerals are naturally occurring inorganic solids with a defined chemical composition and crystalline structure. Examples include halite and magnesite.
• Ores are rocks or sediments containing minerals in high enough concentration to make extraction economically worthwhile. Bauxite and galena are common examples.

Here are the key minerals and ores you should know for representative metals:

• Sodium: halite ($\text{NaCl}$) and trona ($\text{Na}_3(\text{CO}_3)(\text{HCO}_3) \cdot 2\text{H}_2\text{O}$)
• Magnesium: dolomite ($\text{CaMg(CO}_3)_2$) and magnesite ($\text{MgCO}_3$)
• Aluminum: bauxite ($\text{Al}_2\text{O}_3 \cdot n\text{H}_2\text{O}$)
• Potassium: sylvite ($\text{KCl}$) and carnallite ($\text{KMgCl}_3 \cdot 6\text{H}_2\text{O}$)
• Calcium: limestone ($\text{CaCO}_3$) and gypsum ($\text{CaSO}_4 \cdot 2\text{H}_2\text{O}$)
• Tin: cassiterite ($\text{SnO}_2$)
• Lead: galena ($\text{PbS}$)

Metallurgy: Extraction and Purification of Metals

Metallurgy is the broad field covering how we extract and purify metals from ores. There are three main branches, each suited to different metals and situations:

• Pyrometallurgy uses high-temperature processes like smelting. Heat drives the chemical reactions that separate the metal from its ore.
• Hydrometallurgy uses aqueous (water-based) solutions to dissolve and recover metals from ores or concentrates.
• Electrometallurgy uses electrical energy to drive redox reactions that extract or refine metals.

The choice of method depends largely on the metal's reactivity. Highly reactive metals hold onto their electrons tightly in compounds, so they need more energy-intensive methods like electrolysis. Less reactive metals can be freed with simpler chemical reduction.

Electrolytic Isolation of Reactive Metals

Electrolysis forces nonspontaneous redox reactions to occur by supplying electrical energy. During the process, metal cations migrate to the cathode, where they gain electrons (are reduced) and deposit as pure metal. This approach is necessary for very reactive metals like sodium and aluminum because no common chemical reducing agent is strong enough to do the job.

Extraction of Sodium: The Downs Process

1. A molten mixture of $\text{NaCl}$ and $\text{CaCl}_2$ is prepared. The $\text{CaCl}_2$ is added to lower the melting point of the mixture, making the process more energy-efficient.
2. Electrical current passes through the molten salt. Sodium ions are reduced to liquid sodium metal at the cathode, while chloride ions are oxidized to chlorine gas at the anode:

$\text{NaCl} \xrightarrow{\text{electrolysis}} \text{Na} + \frac{1}{2}\text{Cl}_2$

Extraction of Aluminum: The Hall-Héroult Process

1. Pure alumina ($\text{Al}_2\text{O}_3$) is dissolved in molten cryolite ($\text{Na}_3\text{AlF}_6$). Cryolite acts as a solvent and lowers the melting point from about 2072°C for pure alumina to roughly 1000°C.
2. Electrolysis reduces aluminum ions to liquid aluminum metal at the cathode.
3. Oxygen produced at the anode reacts with the carbon anode to form $\text{CO}_2$, which means the carbon anodes gradually get consumed and need regular replacement.

The overall reaction:

$\text{Al}_2\text{O}_3 \xrightarrow{\text{electrolysis}} 2\text{Al} + \frac{3}{2}\text{O}_2$

Chemical Reduction for Metal Production

For metals that are less reactive than sodium or aluminum, chemical reduction works well. A more reactive element or compound donates electrons to reduce the metal compound to its elemental form.

Magnesium Production: The Pidgeon Process

1. Magnesium oxide ($\text{MgO}$) is heated with ferrosilicon ($\text{FeSi}$) at high temperatures:

$2\text{MgO} + \text{FeSi} \xrightarrow{\Delta} 2\text{Mg} + \text{Fe} + \text{SiO}_2$

1. Magnesium vaporizes at these temperatures, so it's collected by condensation, which also helps drive the reaction forward by continuously removing product.

Zinc Production: Roasting and Reduction

This is a two-step process because zinc's most common ore is a sulfide, not an oxide:

1. Roasting: Zinc sulfide ($\text{ZnS}$) is heated in air to convert it to zinc oxide ($\text{ZnO}$):

$2\text{ZnS} + 3\text{O}_2 \xrightarrow{\Delta} 2\text{ZnO} + 2\text{SO}_2$

1. Reduction: The zinc oxide is then reduced with carbon monoxide in a blast furnace:

$\text{ZnO} + \text{CO} \xrightarrow{\Delta} \text{Zn} + \text{CO}_2$

The roasting step is necessary because oxide ores are much easier to reduce than sulfide ores.

Tin Production: Carbothermic Reduction

1. Cassiterite ($\text{SnO}_2$) is reduced with carbon at high temperature in a blast furnace:

$\text{SnO}_2 + 2\text{C} \xrightarrow{\Delta} \text{Sn} + 2\text{CO}$

1. Because tin has a relatively low melting point (232°C), it collects as molten metal at the bottom of the furnace and is easily tapped off.

Tin's low reactivity is what makes this straightforward single-step reduction possible, unlike zinc, which requires roasting first.