Occurrence and Preparation of the Elements

Natural occurrence of main group elements

Main group elements (s- and p-block) occur in nature in different forms depending mainly on their reactivity and the conditions on Earth (presence of oxygen, water, carbon dioxide, biological activity, etc.).

Native (elemental) occurrence

Only relatively unreactive main group elements are found as native elements (chemically uncombined):

• Group 18 (noble gases):
• Occur as monoatomic gases in the atmosphere (He, Ne, Ar, Kr, Xe) or in natural gas fields.
• They do not form stable minerals under normal conditions, due to very low reactivity.
• Some group 14 elements:
• Carbon:
• As diamond and graphite (native, crystalline allotropes).
• Also as amorphous carbon (coal, soot).
• Silicon: generally not native; occurs almost exclusively as silicates and silica; native silicon is extremely rare.
• Tin and lead: small amounts can be found in native form in reducing environments, but their main occurrence is as sulfides or oxides.
• Relatively noble metals among the p-block:
• Bismuth (group 15): sometimes native; more commonly in sulfide ores.
• Post-transition elements such as thallium are not usually native; they occur in minerals.

For the highly reactive main group metals (e.g. Na, K, Ca, Mg, Al) native elemental occurrence is essentially absent because they are quickly oxidized or react with water and other substances.

Occurrence as compounds

Most main group elements occur as compounds in minerals, brines, gases, or organic matter.

Typical patterns:

• Oxides and silicates – for strongly lithophilic (rock-forming) elements:
• Silicon: as $ \text{SiO}_2 $ (quartz) and countless silicates (feldspars, micas, olivine, pyroxenes).
• Aluminum: in aluminosilicates and in oxide hydroxides (e.g. bauxite, a mixture of gibbsite $ \text{Al(OH)}_3 $, boehmite $ \gamma\text{-AlO(OH)} $, and diaspore $ \alpha\text{-AlO(OH)} $).
• Calcium, sodium, potassium, magnesium: as silicates (feldspars, pyroxenes, amphiboles), carbonates (calcite $ \text{CaCO}_3 $, dolomite $ \text{CaMg(CO}_3)_2 $), and sulfates (gypsum $ \text{CaSO}_4\cdot 2\text{H}_2\text{O} $).
• Sulfides – in ore deposits formed under reducing conditions:
• Lead: galena $ \text{PbS} $
• Tin: stannite, cassiterite is an oxide ($ \text{SnO}_2 $) but often associated with sulfide ores.
• Antimony: stibnite $ \text{Sb}_2\text{S}_3 $
• Arsenic: arsenopyrite $ \text{FeAsS} $
• Halides – especially in evaporite deposits and brines:
• Sodium: halite $ \text{NaCl} $ (rock salt), in seawater and salt lakes.
• Potassium: sylvite $ \text{KCl} $, carnallite $ \text{KMgCl}_3\cdot 6\text{H}_2\text{O} $ in potash deposits.
• Magnesium: in seawater, salt lakes (e.g. as $ \text{Mg}^{2+} $ plus accompanying ions).
• Carbonates – where $ \text{CO}_2 $-rich water reacts with metal ions:
• Calcium: calcite $ \text{CaCO}_3 $, aragonite (same composition, different crystal structure).
• Magnesium and calcium: dolomite $ \text{CaMg(CO}_3)_2 $
• Sodium and potassium carbonates in alkaline lakes.
• Phosphates and other oxyanions:
• Phosphorus: mainly as apatite, e.g. fluorapatite $ \text{Ca}_5(\text{PO}_4)_3\text{F} $.
• Boron: as borates (borax $ \text{Na}_2\text{B}_4\text{O}_7\cdot 10\text{H}_2\text{O} $, kernite, colemanite) often in evaporite deposits.
• In biological material (bioelements):
• Carbon, hydrogen, oxygen, nitrogen, phosphorus, sulfur (CHNOPS) are essential components of living matter (proteins, nucleic acids, carbohydrates, lipids, etc.).
• They are cycled through the atmosphere, hydrosphere, biosphere, and lithosphere.

Geochemical distribution and reservoirs

Main group elements are unevenly distributed among Earth’s major reservoirs:

• Crust: enriched in lithophilic main group elements:
• $ \text{O}, \text{Si}, \text{Al}, \text{Na}, \text{K}, \text{Ca}, \text{Mg} $ form most common rock-forming minerals.
• Hydrosphere (oceans and waters):
• Major cations: $ \text{Na}^+, \text{Mg}^{2+}, \text{Ca}^{2+}, \text{K}^+ $.
• Major anions: $ \text{Cl}^-, \text{SO}_4^{2-}, \text{HCO}_3^- $.
• Atmosphere:
• Nonmetals predominate: $ \text{N}_2, \text{O}_2, \text{CO}_2, \text{H}_2\text{O} $ vapor, noble gases.
• Biosphere:
• Concentrates light main group elements (H, C, N, O, P, S) and some metals (Na, K, Mg, Ca).

General strategies for preparation of main group elements

“Preparation” here means industrial- or laboratory-scale isolation of an element from its natural compounds. The method chosen depends mainly on:

• Oxidation state and stability of the main natural compounds.
• Position in the electrochemical series (how easily the element is reduced or oxidized).
• Physical properties (boiling point, volatility, solubility, etc.).

Thermal decomposition and distillation

Some elements or their simple compounds can be obtained by heating minerals and then separating products based on volatility.

• Example: sulfur
• Natural occurrence: elemental sulfur deposits, sulfide minerals.
• Preparation:
• Frasch process (for underground sulfur): superheated water melts sulfur, which is pumped to the surface.
• Roasting of sulfide ores can also convert sulfides to $ \text{SO}_2 $ and elemental sulfur in intermediate processes, though the main target is often sulfuric acid.
• Example: phosphorus (partial step)
• Apatite is heated with silica and carbon in an electric furnace (see below); elemental white phosphorus distills off as a vapor and is condensed under water.

Carbothermic reduction of oxides

Less reactive metals that form stable oxides but are below carbon in reducing power can be produced by reduction with carbon (usually as coke):

$$
\text{Metal oxide} + \text{C} \rightarrow \text{Metal} + \text{CO / CO}_2
$$

Typical for some main group metals:

• Lead from galena (via $ \text{PbO} $ intermediate):
1. Roasting:
   $$ 2\text{PbS} + 3\text{O}_2 \rightarrow 2\text{PbO} + 2\text{SO}_2 $$
2. Reduction with carbon:
   $$ \text{PbO} + \text{C} \rightarrow \text{Pb} + \text{CO} $$
• Tin from cassiterite:
• Ore: $ \text{SnO}_2 $
• Reduction:
  $$ \text{SnO}_2 + 2\text{C} \rightarrow \text{Sn} + 2\text{CO} $$
• Silicon (metallurgical grade):
• Raw material: quartz sand $ \text{SiO}_2 $ and coke in an electric arc furnace:
  $$ \text{SiO}_2 + 2\text{C} \rightarrow \text{Si} + 2\text{CO} $$

Carbothermic reduction is not suitable for very reactive metals such as Na, K, Mg, Ca, and Al because they form extremely stable oxides that are not easily reduced by carbon at practical conditions.

Metallothermic reduction

A more reactive metal (often Al, Mg, or Na) reduces a less reactive metal from its compound. This is used when carbon reduction fails or is impractical.

General form:

$$
\text{Oxide or halide of metal A} + \text{metal B} \rightarrow \text{metal A} + \text{compound of B}
$$

Examples involving main group elements:

• Production of silicon with magnesium or aluminum instead of carbon in special grades or laboratory:
  $$ \text{SiO}_2 + 2\text{Mg} \rightarrow \text{Si} + 2\text{MgO} $$
• Production of some alkaline earth metals (e.g. calcium):
• Often by reduction of $ \text{CaI}_2 $ or $ \text{CaCl}_2 $ with sodium:
  $$ \text{CaI}_2 + 2\text{Na} \rightarrow \text{Ca} + 2\text{NaI} $$

Metallothermic processes are important for high-purity or specialty forms of main group elements but are less common as the main bulk route where electrolysis is more economical.

Electrolytic production of active metals

Very electropositive main group metals (alkali and alkaline earth metals, and aluminum) cannot be extracted from their compounds by chemical reducing agents at reasonable cost. They are prepared almost exclusively by electrolysis of molten salts or occasionally aqueous solutions (where possible).

Molten salt electrolysis of halides

Used for metals that are more easily obtained from halides than from oxides and cannot be deposited from water because they would react with it.

General anode and cathode reactions for a metal $ \text{M} $:

• Cathode (metal formation):
  $$ \text{M}^+ + e^- \rightarrow \text{M} $$
• Anode (halogen formation):
  $$ 2\text{X}^- \rightarrow \text{X}_2 + 2e^- $$

Examples:

• Sodium:
• Electrolysis of molten sodium chloride (Downs process):
  $$ \text{NaCl(l)} \rightarrow \text{Na(l)} + \tfrac{1}{2}\text{Cl}_2(\text{g}) $$
• Magnesium:
• Electrolysis of molten $ \text{MgCl}_2 $ (from seawater or brines), producing Mg and $ \text{Cl}_2 $.
• Calcium and other alkaline earth metals:
• Electrolysis of molten $ \text{CaCl}_2 $ or mixtures with other salts to reduce the melting point.

Electrolysis of molten oxides in ionic melts

For aluminum, the oxide is dissolved in a molten salt that conducts electricity:

• Alumina ($ \text{Al}_2\text{O}_3 $) is dissolved in molten cryolite ($ \text{Na}_3\text{AlF}_6 $).
• In the Hall–Héroult process:
• Cathode:
  $$ \text{Al}^{3+} + 3e^- \rightarrow \text{Al(l)} $$
• Anode (consumable carbon anodes):
  $$ \text{C} + \text{O}_2^{2-} \rightarrow \text{CO}_2 + 4e^- $$

Although details belong in electrochemistry and industrial chemistry, the key idea here is: very reactive main group metals are obtained by providing electrical energy to drive highly endergonic reductions that cannot be achieved by chemical reagents.

Electrolysis of aqueous solutions

Some less reactive main group metals (and nonmetals such as hydrogen) can be deposited from aqueous solutions, but there are important limitations:

• Metals more reactive than hydrogen cannot usually be obtained from water because they reduce water instead of being deposited.
• Metals like aluminum, sodium, potassium, and calcium therefore require molten salt electrolysis.

Aqueous electrolysis is more important for nonmetal elements of the main groups (e.g. chlorine, hydrogen) or for ions of relatively noble p-block metals that do not undergo vigorous reaction with water.

Displacement reactions (cementation)

A less noble metal can displace a more noble metal from its salt solution. This simple redox method is often used for purification or small-scale recovery of main group metals but is less central for their primary production.

Generic reaction:

$$
\text{M}_\text{less noble} + \text{M}_\text{more noble}^{n+} \rightarrow \text{M}_\text{less noble}^{m+} + \text{M}_\text{more noble}
$$

Example:

• Iron displacing tin or lead from solution:
  $$ \text{Fe} + \text{Sn}^{2+} \rightarrow \text{Fe}^{2+} + \text{Sn} $$

Such processes are more frequently applied to transition metals but can be relevant when recovering e.g. tin or lead from waste streams.

Laboratory preparation routes

On the laboratory scale, methods differ from industrial processes because safety, cost, and scale are different. Some general strategies for main group elements:

• Reduction with hydrogen
• Certain metal oxides or halides of p-block elements (e.g. some oxides of tin or germanium) can be reduced with hydrogen:
  $$ \text{SnO}_2 + 2\text{H}_2 \rightarrow \text{Sn} + 2\text{H}_2\text{O} $$
• Thermal decomposition of unstable compounds
• E.g. decomposition of ammonium salts or hydrides to give elemental nonmetals or lower oxidation state compounds.
• Disproportionation / comproportionation
• Some nonmetals (e.g. phosphorus, sulfur) and metalloids can be prepared or purified through redox rearrangements where the same element shifts between oxidation states.
• Chemical vapor transport and sublimation
• Useful to obtain pure samples of elements with appreciable vapor pressures (e.g. iodine, sulfur) or to purify metalloids like silicon and germanium using carrier gases (e.g. $ \text{HCl} $, $ \text{Cl}_2 $).

These methods are usually geared toward purity and convenience rather than bulk production.

Examples by group

The following overview shows characteristic occurrence and preparation strategies for different main group element families, emphasizing patterns rather than exhaustive detail.

Group 1: Alkali metals

• Occurrence:
• Never native; extremely reactive.
• As soluble salts (mainly chlorides, sulfates, carbonates) in seawater, salt lakes, and evaporite deposits.
• Common minerals: halite $ (\text{NaCl}) $, sylvite $ (\text{KCl}) $, complex silicates (feldspars: $ \text{NaAlSi}_3\text{O}_8, \text{KAlSi}_3\text{O}_8 $).
• Preparation:
• Industrially by electrolysis of molten halides:
• Sodium: molten NaCl (Downs process).
• Potassium: molten KCl or KCl/NaCl mixtures.
• Lithium: molten $ \text{LiCl/KCl} $ mixture.
• Some laboratory preparations use reduction of alkali halides with even more reducing metals, but these are not common industrially.

Group 2: Alkaline earth metals

• Occurrence:
• As carbonates (calcite, dolomite) and sulfates (gypsum, anhydrite), in silicate minerals, and in seawater.
• No native occurrence; too reactive.
• Preparation:
• Electrolysis of molten halides:
• Magnesium: molten $ \text{MgCl}_2 $ from seawater or brines.
• Calcium, strontium, barium: molten chlorides (often mixed salts to lower melting point).
• Some metallothermic reductions:
• $ \text{CaCl}_2 $ reduced by sodium on smaller scales.

Group 13: Boron group

• Boron:
• Occurrence: as borates in evaporites (borax, kernite, colemanite).
• Preparation:
• Reduction of boron oxides or halides with Mg or Na.
• High-purity boron via gas-phase reactions (e.g. from $ \text{BCl}_3 $ and $ \text{H}_2 $) and chemical vapor deposition.
• Aluminum:
• Occurrence: as aluminosilicates and oxide hydroxides (bauxite).
• Preparation:
• Purification of alumina from bauxite (Bayer process).
• Electrolysis of alumina dissolved in molten cryolite (Hall–Héroult process).
• Gallium, indium, thallium:
• Occurrence: in trace amounts in aluminum, zinc, and lead ores.
• Preparation: recovered as by-products via hydrometallurgical processes, cementation, and electrolysis.

Group 14: Carbon group

• Carbon:
• Occurrence:
• Elemental as diamond, graphite, amorphous carbon.
• In carbonates (limestones), organic matter, fossil fuels.
• Preparation:
• Graphite by heating carbon-rich precursors.
• Synthetic diamonds by high-pressure high-temperature or chemical vapor deposition processes.
• Silicon:
• Occurrence: in $ \text{SiO}_2 $ (quartz) and silicates; major component of Earth’s crust.
• Preparation:
• Metallurgical silicon via carbothermic reduction of $ \text{SiO}_2 $.
• High-purity silicon for electronics by further refining (e.g. via volatile chlorosilanes and zone melting).
• Germanium, tin, lead:
• Occurrence: in sulfide and oxide ores (e.g. cassiterite $ \text{SnO}_2 $, galena $ \text{PbS} $).
• Preparation:
• Roasting of sulfides to oxides.
• Carbothermic reduction or other metallothermic methods.
• Electrolytic refining for high purity.

Group 15: Nitrogen group

• Nitrogen:
• Occurrence: $ \text{N}_2 $ (about 78% of the atmosphere).
• Preparation:
• Fractional distillation of liquefied air (industrial).
• For laboratory use, often by removing $ \text{O}_2 $ from air chemically or by gas generators.
• Phosphorus:
• Occurrence: in phosphate minerals (apatite) and biological materials.
• Preparation:
• Heating phosphate rock with silica and coke in an electric furnace, producing $ \text{P}_4 $ vapor that is condensed under water.
• Heavier group 15 elements (As, Sb, Bi):
• Occurrence: mainly as sulfides and other minerals.
• Preparation: roasting and reduction similar to other p-block metals; refinement by electrolysis or distillation.

Group 16: Chalcogens

• Oxygen:
• Occurrence: as $ \text{O}_2 $ in air (about 21%) and as oxide in nearly all minerals.
• Preparation:
• Fractional distillation of liquefied air.
• Electrolysis of water is used when high purity is needed or in smaller scale.
• Sulfur:
• Occurrence: elemental sulfur deposits, sulfide minerals (pyrite $ \text{FeS}_2 $, galena, sphalerite).
• Preparation:
• Frasch process for underground deposits.
• Recovery from sour natural gas and petroleum via processes that convert $ \text{H}_2\text{S} $ into elemental sulfur.
• Selenium, tellurium:
• Occurrence: as minor components in sulfide ores.
• Preparation: recovered as by-products of refining (especially of copper).

Group 17: Halogens

• Occurrence:
• As halide ions in seawater, salt lakes, and various minerals:
• Chloride in halite $ (\text{NaCl}) $ and other salts.
• Bromide and iodide in seawater and brines.
• Fluoride in fluorite $ (\text{CaF}_2) $ and apatite.
• Preparation:
• Primarily via oxidation of halide ions:
• Electrolysis of brine for chlorine:
• Anode:
  $$ 2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^- $$
• Bromine and iodine: oxidation of $ \text{Br}^- $ or $ \text{I}^- $ with chlorine or other oxidizing agents.
• Fluorine: electrolysis of molten $ \text{KF}\cdot \text{HF} $ mixtures (since aqueous electrolysis would produce $ \text{O}_2 $ and $ \text{F}_2 $ is too reactive with water).

Group 18: Noble gases

• Occurrence:
• In the atmosphere (He, Ne, Ar, Kr, Xe) and in some natural gas reservoirs (He especially).
• Radon as a radioactive decay product in rocks.
• Preparation:
• Fractional distillation of liquefied air:
• Different boiling points of $ \text{N}_2, \text{O}_2, \text{Ar}, \text{Ne}, \text{Kr}, \text{Xe} $ allow separation.
• Helium:
• Separation from natural gas through low-temperature distillation or selective membrane processes.

From occurrence to industrial process: general steps

For many main group elements, the route from natural occurrence to elemental product follows a sequence:

1. Mining or extraction of raw material
• E.g. mining of bauxite, phosphate rock, salt, or sulfur; pumping of brine; separation of air.
2. Concentration and beneficiation
• Physical methods: crushing, grinding, flotation, magnetic separation.
• Chemical methods: leaching, precipitation to remove impurities.
3. Chemical conversion
• Transforming the element into a compound more suitable for reduction or electrolysis (e.g. converting ores to oxides or chlorides).
4. Reduction or electrolysis
• Application of carbothermic, metallothermic, or electrolytic methods to yield the elemental form.
5. Refining and purification
• Distillation, electrorefining, zone refining, or recrystallization to achieve the desired purity for specific applications.

Understanding where an element occurs and in which chemical form is essential for choosing an efficient preparation route and for evaluating the environmental and energy costs linked to its production.