Sodium: Occurrence, Extraction from Downs process; properties and uses


Lithium occurs mainly as silicate minerals such as spodumene [LiAl(SiO3)2], lepidolite [(Li,Na,K)2Al2,(SiO3)3 (F,OH)2] etc. It is the 35th most abundant element by weight.

Compounds of sodium and potassium have been known from ancient times. Sodium and potassium are seventh and eighth most abundant elements by weight in the Earth's crust and together make up over 4% of the Earth's crust by weight. NaCl and KCl also occur in large amounts in seawater. Rock salt (NaCl) is the major source of sodium. Potassium occurs mainly as deposits of KCl (sylvite), which is a mixture of KCl and NaCl (sylvinite) and double salt KCl.MgCl2.6H2O (carnallite).

Rubidium and cesium are obtained as a by-product of lithium processing. Francium being radioactive does not occur appreciably in nature.

Ores of Sodium

  • Albite (soda feldspar) NaAlSi3O8
  • Borax (as tincal) Na2B2O7.10H2O
  • Glauber's salt Na2SO4.10H2O
  • Sodium chloride, (common salt) found as rock salt, in sea water and in lakes.
  • Sodium nitrate (NaNO3) as chile saltpetre.
  • Extraction of Sodium

    The metal sodium, was prepared in 1807 by an English chemist,

    Sir Humphery Davy by electrolysis of fused caustic soda. The two methods used for extracting sodium metals are,
    • Castner's process: By the electrolysis of fused caustic soda.
    • Down's process: By the electrolysis of fused sodium chloride.

    Castner's process

    Molten sodium hydroxide (caustic soda, NaOH) is electrolyzed using an iron cathode and nickel anode. A nickel wire gauze cylinder prevents the electrodes from touching each other. On passing electric current through the melt, sodium is liberated at the cathode and oxygen at the anode. The liberated sodium metal floats over the surface (lighter than fused caustic soda), and collects inside the receiving vessel where hydrogen prevents sodium from oxidation. Excess of the gas escapes from the outlet. Sodium is removed from time to time with the help of perforated spoons and kept under kerosene oil. The various reactions taking place during electrolysis are:

    electrolysis of caustic soda

    At cathode:

    At anode:

    Most of the water formed in the reaction gets evaporated, and the rest is electrolyzed into hydrogen and oxygen. Therefore, metallic sodium and hydrogen are liberated at the cathode and oxygen at the anode.

    Castners process

    Fig: 12.2 - Castners process

    Down's process

    In this process, molten sodium chloride (common salt) is electrolyzed using a graphite anode and a ring shaped iron cathode. The two electrodes are separated by a wire gauge partition to avoid the mixing of sodium and chlorine so formed. As Sodium chloride melts at a very high temperature of 1085 K, a mixture containing sodium chloride, potassium chloride and potassium fluoride (NaCl + KCl + KF) is employed. This mixture melts at about 850-875 K.

    The reasons for lowering the temperature are:
    • It is difficult to maintain a high temperature of 1085K.
    • Sodium is volatile at this temperature and so a part of the metal formed may vaporize.
    • At this high temperature chlorine produced as a by product corrodes the vessel.
    • Metal at this temperature will be in a colloidal state and its separation will be difficult.

    On passing electric current, chlorine is liberated at the carbon anode and escapes through the dome shaped steel hood outlet. Sodium rises from the cathode and remains in the wire gauze shell. The sodium produced is in molten state. Being lighter than the electrolyte it rises to the surface. As the level of molten sodium arises, it is forced into the receiver. The process is continuous and fresh salt is introduced to maintain the level of molten electrolyte high enough to allow sodium to rise into the iron pipe. Chemical reactions involved in this process are as follows:

    At cathode:

    At anode:

    apparatus for Down s cell

    Fig:12.3 - Down's cell

    Advantages of Down's process

    • Sodium metal obtained has high degree of purity (99.5%).
    • The starting material, sodium chloride is very cheap.
    • Chlorine is obtained as a useful by-product.
  • Physical properties of lithium and sodium
    • Lithium and sodium are light, soft metals with silvery white lustre. Lithium is harder than sodium.
    • As sodium metal is light and soft, it can be easily cut with a knife. Freshly cut surfaces are shining, but get covered with a layer of oxide or carbonate when kept in contact with air. Lithium though harder than sodium is soft enough to be cut with knife.
    • The densities of lithium and sodium are low (lithium = 0.634 g cm-3 sodium = 0:97 g cm3). Lithium is the lightest metal known.
    • These metals give color to the flame, when heated in the Bunsen burner. Sodium gives golden yellow while lithium gives crimson red. This is due to their low ionization energy, which makes their valence electrons get excited to higher energy levels when heated. On returning to the ground state these elements emit colored radiations in the visible region.
    • Both have high melting and boiling points.
    chemical properties of lithium and sodium
    Both lithium and sodium are extremely reactive metals. Sodium is more reactive than lithium.

    Action of air

    Both lithium and sodium remain unaffected by dry air but get readily tarnished in moist air forming a film of oxide. These oxides react with moisture of the air give the corresponding hydroxide and finally carbonate.

    action of lithium on air

    action of lithium hydroxide with CO2

    action of sodium on air

    action of sodium hydroxide with CO2

    This is the reason that sodium and potassium are stored under kerosene.

    With oxygen

    Lithium tarnishes slowly in moist air while dry air has no effect. When heated in air or oxygen at about 200°C, it burns with a brilliant white light forming lithium monoxide.

    action of lithium with oxygen

    When heated in air or oxygen, sodium burns with a golden yellow flame forming a mixture of oxide and peroxide.

    action of sodium with oxygen

    formation of sodium peroxide

    Action of water

    Both, lithium and sodium decompose cold water vigorously liberating hydrogen. Sodium reacts with water more vigorously than lithium.

    action of lithium with water

    action of sodium with water

    The hydroxides of lithium and sodium are strong alkalies. They are highly soluble in water and their aqueous solutions contain hydroxyl ions:

    llithium hydroxide as strong alkali

    sodium hydroxide as strong alkali

    Sodium hydroxide is a stronger base than lithium hydroxide.

    Action with Non-metals

    Both lithium and sodium combine directly with hydrogen, sulphur, halogens and other non-metals on heating.

    With hydrogen

    action of lithium with hydrogen

    action of sodium with nitrogen

    These hydrides are ionic hydrides.

    With chlorine

    action of lithium with chlorine

    action of sodium wih chlorine

    These halides are ionic crystalline halides. Lithium halides are less ionic than sodium halides.

    With sulphur

    action of phosphorus with sulhur

    action of sodium with sulphur

    With phosphorous

    formation of sodium phosphide

    With nitrogen

    It may be noted that lithium reacts with N2 to form lithium nitride but sodium does not form sodium nitride.

    action of lithium with nitrogen

    With ammonia

    Sodium and potassium give the corresponding amide when heated in ammonia gas.

    action of lithium with ammonia

    action of sodium with ammonia

    These amides act as reducing agents and reduce many oxides.

    Reducing action

    Lithium and Sodium act as strong reducing agents. So, they reduce some metallic chlorides and oxides into metals. This property is applied in the preparation of some metals. For example, beryllium, uranium etc., can be obtained by the reduction of the corresponding halides using sodium.

    formation of berillium

    formation of uranium

    formation of aluminium

    formation of silicon

    formation of sodium bicarbonate formation

    Action with acids

    Sodium reacts vigorously with acids evolving hydrogen.

    action of sodium with HCl

    action of potassium with HCl

    Action with mercury

    With mercury, sodium form amalgams of varying composition e.g., NaHg, Na2Hg, Na3Hg etc.

    Solubility in liquid ammonia

    Sodium dissolves in liquid ammonia to give conducting blue-colored solution due to the presence of ammoniated electrons in solutions.

    Solubility in liquid ammonia

    The blue color of the solution is due to ammoniated electrons, which absorb energy corresponding to red region of visible light, for their excitation to higher energy levels. The transmitted light thus imparts blue color to the solution.

    The solution is made conducting in nature by both ammoniated cation and ammoniated electron.

    Uses of sodium


    • For refining copper and nickel.
    • For producing thermo nuclear energy for propelling rockets and guided missiles.
    • For the manufacture of various alloys such as lithium lead alloy, lithium aluminium alloy etc.
    • Lithium and its compounds are used as reagents in a number of organic syntheses.
    • Lithium alimunium hydride is a powerful reducing agent commonly used in organic synthesis.
    • Lithium is used as a important systemic anti-depressant drug.


    • In the manufacture of sodium peroxide, sodamide, sodium cyanide etc. About 50 per cent of sodium extracted is used in the manufacture of tetraethyl lead (C2H5)4Pb), which is used as anti-knock agent in petrol.
    • For producing amalgams and alloys. Amalgams are used as reducing agents.
    • Sodium also finds application in illumination engineering and in sodium vapor discharge lamps.
    • Acts as a catalyst in the preparation of artificial rubber.
    • As a deoxidizer in the preparation of light alloys.
    • Liquid sodium or its alloy with potassium is used as a coolant in nuclear reactors.
    • Is used as a laboratory reagent (Lassaigne's extract).


    5. Although lithium and sodium are difficult to handle yet these metals are widely used in laboratories, why?


    Lithium and Sodium act as strong reducing agents. So, they reduce some metallic chlorides and oxides into metals easily. This property is applied in the preparation of some metals. For example, beryllium, uranium etc., can be obtained by the reduction of the corresponding halides using sodium.

    Lithium and its compounds find important use as reagents in a number of organic syntheses. Lithium aluminium hydride (LiAlH4) is a powerful reducing agent commonly used in organic synthesis.

    Sodium is used as a reagent to detect the presence of nitrogen, sulphur and halogens by Lassaigne's extract method. It is also used for the preparation of sodamide (NaNH2), sodium peroxide (Na2O2) and sodium cyanide (NaCN), which are useful reagents for a number of synthetic processes.


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