ALKALI METALS AND THEIR COMPOUNDS

The group 1 of the periodic table contains six elements, namely lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). All these elements are
typical metals. Francium is radioactive with longest lived isotope 223 Fr with half life period of only 21 minute. These are usually referred to as alkali metals since their hydroxides form strong bases or alkalies.
 

(1) Electronic configuration

(2) Occurrence : Alkali metals are very reactive and thus found in combined state some important ores of alkali metals are given ahead.
(i) Lithium : Triphylite, Petalite, lepidolite, Spodumene [LiAl(SiO₃)₃], Amblygonite [Li(Al F)PO₄]
(ii) Sodium : Chile salt petre (NaNO₃), Sodium chloride (NaCl), Sodium sulphate (Na₂SO₄), Borax (Na₂B₄O₇10H₂O), Glauber salt (Na₂ SO₄.10H₂O)
(iii) Potassium : Sylime (KCl), carnallite (KCl.MgCl₂.6H₂O) and Felspar (K₂O. Al₂O₃.6SiO₂)
(iv) Rubidium : Lithium ores Lepidolite, triphylite contains 0.7 to 3% Rb₂O
(v) Caesium : Lepidolite, Pollucite contains 0.2 to 7% Cs₂O
(3) Extraction of alkali metals : Alkali metals cannot be extracted by the usual methods for the extraction of metals due to following reasons.
(i) Alkali metals are strong reducing agents, hence cannot be extracted by reduction of their oxides or other compounds.
(ii) Being highly electropositive in nature, it is not possible to apply the method of displacing them from their salt solutions by any other element.

(iii) The aqueous solutions of their salts cannot be used for extraction by electrolytic method because hydrogen ion is discharged at cathode instead of an alkali metal ions as the discharge potentials of alkali metals are high. However, by using Hg as cathode, alkali metal can be deposited. The alkali metal readily combines with Hg to form an amalgam from which its recovery difficult. The only successful method, therefore, is the electrolysis of their fused salts, usually chlorides. Generally, another metal chloride is added to lower their fusion temperature.
Fused NaCl :

                    
(4) Alloys Formation
(i) The alkali metals form alloys among themselves as well as with other metals.
(ii) Alkali metals also get dissolved in mercury to form amalgam with evolution of heat and the amalgamation is highly exothermic.

Physical properties
 

(1) Physical state
(i) All are silvery white, soft and light solids. These can be cut with the help of knife. When freshly cut, they have bright lustre which quickly tarnishes due to surface oxidation.
(ii) These form diamagnetic colourless ions since these ions do not have unpaired electrons, (i.e. M+ has ns0 configuration). That is why alkali metal salts are colourless and diamagnetic.
 

(2) Atomic and ionic radii
(i) The alkali metals have largest atomic and ionic radii than their successive elements of other groups belonging to same period.
(ii) The atomic and ionic radii of alkali metals, however, increases down the group due to progressive addition of new energy shells.
         No doubt the nuclear charge also increases on moving down the group but the influence of addition of energy shell predominates
                                          Li        Na        K        Rb         Cs       Fr
Atomic radius (pm)      152      186       227     248       265    375   
 Ionic radius of M⁺        60        95       133      148       169      – 
  ions (pm)

(3) Density
(i) All are light metals, Li, Na and K have density less than water. Low values of density are because these metals have high atomic volume due to larger atomic size. On moving down the group the atomic size as well as atomic mass both increase but increase in atomic mass predominates over increase in atomic size or atomic volume and therefore the ratio mass/volume i.e. density gradually increases down the groups
(ii) The density increases gradually from Li to Cs, Li is lightest known metal among all.
Li = 0.534, Na = 0.972, K = 0.86, Rb = 1.53 and Cs = 1.87 g/ml at 20^oC.
(iii) K is lighter than Na because of its unusually large atomic size.
(iv) In solid state, they have body centred cubic lattice.

(4) Melting point and Boiling point
(i) All these elements possess low melting point and boiling point in comparison to other group members.
                                   Li          Na         K         Rb               Cs         Fr
melting point (K)    453.5    370.8     336.2      312.0         301.5       -
boiling point (K)     1620    1154.4   1038.5     961.0         978.0        -

(ii) The lattice energy of these atoms in metallic crystal lattice relatively low due to larger atomic size and thus possess low melting point and boiling point on moving down the group, the atomic size increases and binding energy of their atoms in crystal lattice decreases which results lowering of melting point.
(iii) Lattice energy decreases from Li to Cs and thus melting point and boiling also decreases from Li to Cs.
 

(5) Ionisation energy and electropositive or metallic character
(i) Due to unpaired lone electron in ns sub-shell as well as due to their larger size, the outermost electron is far from the nucleus, the removal of electron is easier and these low values of ionisation energy. (I.E.)
(ii) Ionisation energy of these metal decreases from Li to Cs.
Ionisation energy                 Li        Na         K          Rb         Cs        Fr
                                  IE₁       520    495       418       403       376       –
                                 IE₂       7296   4563    3069    2650     2420     –
A jump in 2nd ionisation energy (huge difference) can be explained as,

Removal of 1s electrons from Li⁺ and that too from completely filled configuration requires much more energy and a jump in 2^nd ionisation is noticed.
(iii) Lower are ionisation energy values, greater is the tendency to lose ns1 electron to change in M⁺ ion (i.e. M M⁺+e^–) and therefore stronger is electropositive character.
(iv) Electropositive character increases from Li to Cs.
 Due to their strong electropositive character, they emit electrons even when exposed to light showing photoelectric effect. This property is responsible for the use of Cs and K in photoelectric cell.
 

(6) Oxidation number and valency
(i) Alkali metals are univalent in nature due to low ionisation energy values and form ionic compounds. Lithium salts are, however, covalent.
(ii) Further, the M⁺ ion acquires the stable noble gas configuration. It requires very high values of energy to pull out another electron from next to outer shell of M⁺ ion and that is why their second ionisation energy is very high. Consequently, under ordinary conditions, it is not possible for these metals to form M²⁺ ion and thus they show +1 oxidation state.

(iii) Since the electronic configuration of M⁺ ions do not have unpaired electron and thus alkali metal salts are diamagnetic and colourless. Only those alkali metal salts are coloured which have coloured anions
e.g. K₂Cr₂O₇ is orange because of orange coloured Cr₂O₇₂^- ion, KMnO₄ is violet because of violet coloured MnO₄1^- ion.
 

(7) Hydration of Ions
(i) Hydration represents for the dissolution of a substance in water to get adsorb water molecule by weak valency force. Hydration of ions is the exothermic process (i.e energy is released during hydration) when ions on dissolution water get hydration.
(ii) The energy released when 1 mole of an ion in the gaseous state is dissolved in water to get it hydrated is called hydration energy 

                                         
(iii) Smaller the cation, greater is the degree of hydration. Hydration energy is in the order 
        of, Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

(iv) Li⁺ being smallest in size has maximum degree of hydration and that is why lithium salts are mostly hydrated, LiCl. 2H₂O also lithium ion being heavily hydrated, moves very slowly under the influence of electric field and, therefore, is the poorest conductor current among alkali metals ions It may, therefore, be concluded that it is the degree of hydration as well as the size of ion is responsible for the current carried by an ion.
Relative ionic radii                      Cs⁺ > Rb⁺ > K⁺ > Na⁺ > Li⁺
Relative hydrated ionic radii    Li⁺ >  Na⁺ > K⁺ > Rb⁺ > Cs⁺
Relative conducting power       Cs⁺ > Rb⁺ > K⁺ > Na⁺ > Li⁺
 

(8) Electronegativity, Electro positivity and metallic character.
(i) These metals are highly electropositive and thereby possess low values of electronegativities. Metallic character and electro positivity increase from Li to Cs (Li < Na < K < Rb < Cs)
(ii) Electronegativity of alkali metals decreases down the group as the trend of numerical values of electronegativity given below on Pauling scale suggests.
                                  Li     Na      K       Rb     Cs     Fr
Electronegativity     0.98  0.93   0.82   0.82   0.79    –
Fr being radioactive elements and thus studies on physical properties of this element are limited.

(9) Specific heat : It decreases from Li to Cs.
                                       Li       Na        K       Rb        Cs        Fr
Specific heat (Cal/g)  0.941   0.293    0.17   0.08     0.049      –
(10) Conduction power : All are good conductors of heat and electricity, because of loosely held valence electrons.
(11) Standard oxidation potential and reduction properties
(i) Since alkali metals easily lose ns¹ electron and thus they have high values of oxidation potential i.e.,

                                        
(ii) The standard oxidation potentials of a alkali metals (in volts) are listed below,
                             Li            Na           K               Rb              C
                          +3.05      +2.71      +2.93       +2.99         +2.99
(iii) More is oxidation potential, more is the tendency to get oxidized and thus more powerful is reducing nature in aqueous medium that is why alkali metals liberate H₂ from H₂O and HCl.
                        2H₂O + 2M → 2MOH + H₂ ; 2HCl + 2M  → 2MCl + H₂

(iv) However, an examination of ionisation energy for alkali metals reveals that Li should have the minimum tendency to lose electron and thus its reducing nature should be minimum. The greatest reducing nature of Li in aq. medium is accounted due to the maximum hydration energy of Li⁺ ion. For Lithium

ΔHh for Li > ΔHh for Na. Therefore, large negative ΔH values are observed in case of Li and this explains for more possibility of Li to get itself oxidized or have reducing nature.

(12) Characteristic flame colours : The alkali metals and their salts give characteristic colour to Bunsen flame. The flame energy causes and excitation of the outermost electron which on reverting back to its initial position gives out the absorbed energy as visible light. These colour differ from each other Li – crimson, Na – Golden yellow, K – Pale violet, Rb - Red violet and Cs - Blue violet. These different colours are due to different ionisation energy of alkali metals. The energy released is minimum in the case of Li⁺ and increases in the order.
                 Energy released : Li⁺< Na⁺< K⁺< Rb⁺< Cs⁺
                           λ released : Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺
          Frequency released : Li⁺< Na⁺< K⁺< Rb⁺< Cs⁺

^{Chemical properties}

(1) Formation of oxides and hydroxides
(i) These are most reactive metals and have strong affinity for O₂ quickly tranish in air
due to the formation of a film of their oxides on the surface. These are, therefore, kept under kerosene or paraffin oil to protect them from air,                               

(ii) When burnt air (O₂), lithium forms lithium oxide (Li₂O) sodium forms sodium peroxide (Na₂O₂) and other alkali metals form super oxide (Mo₂ i.e. KO₂, RbO₂ or CsO₂)

                                
The reactivity of alkali metals towards oxygen to form different oxides is due to strong positive field around each alkali metal cation. Li⁺ being smallest, possesses strong positive field and thus combines with small anion O₂^– to form stable Li₂O compound. The Na⁺ and K⁺ being relatively larger thus exert less strong positive field around them and thus reacts with larger oxygen anion i.e,

 to form stable oxides. The monoxide, peroxides and superoxides have O2 and   ions respectively. The structures of each are, 

The O₂^–1 ion has a three electron covalent bond and has one electron unpaired. It is therefore superoxides are paramagnetic and coloured KO₂ is light yellow and paramagnetic substance.

(iii) The oxides of alkali metals and metal itself give strongly alkaline solution in water with evolution of heat

The peroxides and superoxides act as strong oxidising agents due to formation of H₂O₂
(iv) The reactivity of alkali metals towards air and water increases from Li to Cs that is why lithium decomposes H₂O very slowly at 25^oC whereas Na does so vigorously, K reacts producing a flame and Rb, Cs do so explosively.

                                

(v) The basic character of oxides and hydroxides of alkali metals increases from Li to Cs. This is due to the increase in ionic character of alkali metal hydroxides down the group which leads to complete dissociation and leads to increase in concentration of OH^– ions.
(2) Hydrides
(i) These metals combine with H to give white crystalline ionic hydrides of the general of the formula MH.
(ii) The tendency to form their hydrides, basic character and stability decreases from Li to Cs since the electropositive character decreases from Cs to Li.
       2M + H₂  2MH ; Reactivity towards H₂ is Cs < Rb < K < Na < Li.
(iii) The metal hydrides react with water to give MOH and H₂ ; MH + H₂ O  MOH + H₂
(iv) The ionic nature of hydrides increases from Li to Cs because of the fact that hydrogen is present in these hydrides as H^– and the smaller cation will produce more polarisation of anion (according to Fajans rule) and will develop more covalent character.
(v) The electrolysis of fused hydrides give H₂ at anode. NaH fusedContains Na⁺ and H^- i.e.,
At cathode: Na⁺ + e^– Na; At anode: H⁻  H₂ + e⁻
(vi) Alkali metals also form hydrides like NaBH₄, LiAlH₄ which are good reducing agent.

(3) Carbonates and Bicarbonates
(i) The carbonates (M₂CO₃) & bicarbonates (MHCO₃) are highly stable to heat, where
stands for alkali metals.
(ii) The stability of these salts increases with the increasing electropositive character from Li to Cs. It is therefore Li₂CO₃ decompose on heating,
     Li₂CO₃ Li₂O + CO₂
(iii) Bicarbonates are decomposed at relatively low temperature, 

     
(iv) Both carbonates and bicarbonates are soluble in water to give alkaline solution due to hydrolysis of carbonate ions or bicarbonate ions.
(4) Halides
(i) Alkali metals combine directly with halogens to form ionic halide M +X^{− .
(ii) The ease with which the alkali metals form halides increases from Li to Cs due to increasing electropositive character from Li to Cs.
(iii) Lithium halides however have more covalent nature. Smaller is the cation, more is deformation of anion and thus more is covalent nature in compound. Also among lithium halides, lithium iodide has maximum covalent nature because of larger anion which is easily deformed by a cation. Thus covalent character in lithium halides is, LiI  > LiBr  > LiCl  > LiF}

(iv) These are readily soluble in water. However, lithium fluoride is sparingly soluble. The low solubility of LiF is due to higher forces of attractions among smaller Li⁺ and smaller F^– ions (high lattice energy).
(v) Halides having ionic nature have high m.pt. and good conductor of current. The melting points of halides shows the order, NaF > NaCl > NaBr > Nal
(vi) Halides of potassium, rubidium and caesium have a property of combining with extra halogen atoms forming polyhalides.
KI + I₂ KI₃ ; In KI₃(aq) the ions K + and I –3 are present
(5) Solubility in liquid NH₃
(i) These metals dissolve in liquid NH₃ to produce blue coloured solution, which conducts electricity to an appreciable degree.
(ii) With increasing concentration of ammonia, blue colour starts changing to that of metallic copper after which dissolution of alkali metals in NH₃ ceases.
(iii) The metal atom is converted into ammoniated metal in i.e. M⁺ (NH₃) and the electron set free combines with NH₃ molecule to produce ammonia solvated electron.

(iv) It is the ammoniated electron which is responsible for blue colour, paramagnetic nature and reducing power of alkali metals in ammonia solution. However, the increased conductance nature of these metals in ammonia is due to presence of ammoniated cation and ammonia solvated electron.

(v) The stability of metal _ ammonia solution decreases from Li to Cs.
(vi) The blue solution on standing or on heating slowly liberates hydrogen, 2M + 2NH₃ 2MNH₂ + H₂. Sodamide (NaNH₂) is a waxy solid, used in preparation of number of sodium compounds.
(6) Nitrates : Nitrates of alkali metals (MNO₃) are soluble in water and decompose on heating. LiNO₃ decomposes to give NO₂ and O₂ and rest all give nitrites and oxygen.
        2MNO₃ 2MNO₂ + O₂ (except Li)
       4 LiNO₃  2Li₂O + 4NO₂ + O₂
(7) Sulphates
(i) Alkali metals’ sulphate have the formula M₂SO₄ .
(ii) Except Li₂SO₄, rest all are soluble in water.
(iii) These sulphates on fusing with carbon form sulphides, M₂SO₄ + 4C  M₂S + 4CO

(iv) The sulphates of alkali metals (except Li) form double salts with the sulphate of the trivalent metals like Fe, Al, Cr etc. The double sulphates crystallize with large number of water molecules as alum. e.g. K₂SO₄ . Al₂(SO₄)₃. 24 H₂O.
(8) Reaction with non-metals
(i) These have high affinity for non-metals. Except carbon and nitrogen, they directly react with hydrogen, halogens, sulphur, phosphorus etc. to form corresponding compounds on heating.
2Na + H₂ 2NaH  ;  2K + H₂  2KH
2Na + Cl₂ 2NaCl               ;  2K + Cl₂ 2KCl
2Na + S  Na₂S                   ;  2K + S  K₂S
3Na + P  Na₃P                   ;  3K + P  K₃P
(ii) Li reacts, however directly with carbon and nitrogen to form carbides and nitrides.
2Li + 2C   LiC₂ ; 6Li + 2N₂

 2 Li₃N
(iii) The nitrides of these metals on reaction with water give NH₃.
M₃N + 3H₂O  3MOH + NH₃
(9) Reaction with acidic hydrogen : Alkali metals react with acids and other compounds containing acidic hydrogen (i.e, H atom attached on F,O, N and triply bonded carbon atom, for example, HF, H₂O, ROH, RNH₂, CH º CH) to liberate H₂.

(10) Complex ion formation : A metal shows complex formation only when it possesses the following characteristics, (i) Small size (ii) High nuclear charge (iii) Presence of empty orbitals in order to accept electron pair ligand. Only Lithium in alkali metals due to small size forms a few complex ions Rest all alkali metals do not possess the tendency to form complex ion.
Anomalous behaviour of Lithium
Anomalous behaviour of lithium is due to extremely small size of lithium its cation on account of small size and high nuclear charge, lithium exerts the greatest polarizing effect out of all alkali metals on negative ion. Consequently lithium ion possesses remarkable tendency towards solvation and develops covalent character in its compounds. Li differs from other alkali metals in the following respects,
(1) It is comparatively harder than other alkali metals. Li can’t be stored in kerosene as it floats to the surface, due to its very low density. Li is generally kept wrapped in parrafin wax.
(2) It can be melted in dry air without losing its brilliance.

(3) Unlike other alkali metals, lithium is least reactive among all. It can be noticed by the following properties,
(i) It is not affected by air.
(ii) It decomposes water very slowly to liberate H₂.
(iii) It hardly reacts with bromine while other alkali metals react violently.
(4) Lithium is the only alkali metal which directly reacts with N₂ ^{to form Lithium nitride (Li}3N)
(5) Lithium when heated in NH₃ forms amide, Li₂^{NH while other metals form amides, MNH}₂^{.
(6) When burnt in air, lithium form Li}2^{O sodium form Na}₂^{O and Na}₂^O₂ other alkali metals form monoxide, peroxide and superoxide.
(7) Li₂O is less basic and less soluble in water than other alkali metals.
(8) LiOH is weaker base than NaOH or KOH and decomposes on heating.
2LiOH  Li₂O + H₂O
9) LiHCO₃ is liquid while other metal bicarbonates are solid.
(10) Only Li₂CO₃ decomposes on heating
Li₂CO₃  Li₂O + CO₂.
Na₂CO₃, K₂CO₃ etc. do not decompose on heating.
(11) LiNO₃ and other alkali metal nitrates give different products on heating
4LiNO₃ = 2Li₂O + 4NO₂ + O₂ ; 2NaNO₃ = 2NaNO₂ + O₂

(12) LiCl and LiNO₃ are soluble in alcohol and other organic solvents. These salts of other alkali metals are, however, insoluble in organic solvents.
(13) LiCl is deliquescent while NaCl, KBr etc. are not. Lithium chloride crystals contain two molecules of water of crystallisation (LiCl. 2H₂O). Crystals of NaCl KBr, KI etc do not contain water of crystallisation.
(14) Li₂SO₄ does not form alums like other alkali metals.
(15) Li reacts with water slowly at room temperature Na reacts vigorously Reaction with K. Rb and Cs is violent.
(16) Li reacts with Br₂ slowly. Reaction of other alkali metals with Br₂ is fast.
(17) Li₂ CO₃ Li₂C₂O₄, LiF , Li₃PO₄ are the only alkali metal salts which are insoluble or sparingly soluble in water.
Diagonal Relationship of Li with Mg
        Due to its small size lithium differs from other alkali metals but resembles with Mg as its size is closer to Mg Its resemblance with Mg is known as diagonal relationship. Generally the periodic properties show either increasing or decreasing trend along the group and vice versa along the period which brought the diagonally situated elements to closer values.

Following are the characteristic to be noted.
  Period         Group I            Group II
      2

     3           
(1) Both Li and Mg are harder and higher m.pt than the other metals of their groups.
(2) Due to covalent nature, chlorides of both Li and Mg are deliquescent and soluble in alcohol and pyridine while chlorides of other alkali metals are not so.
(3) Fluorides, phosphates of Li and Mg are sparingly soluble in water whereas those of other alkali metals are soluble in water.
(4) Carbonates of Li and Mg decompose on heating and liberate CO₂ Carbonates of other alkali metals are stable towards heat and decomposed only on fusion.
Li₂CO₃ Li₂O + CO₂ ; Mg CO₃  MgO + CO₂
(5) Hydroxides and nitrates of both Li and Mg decompose on heating to give oxide. Hydroxides of both Li and Mg are weak alkali.
   4 LiNO₃  2Li₂O + 4NO₂ + O₂
   2Mg(NO₃)₂ 2MgO + 4NO₂ + O₂
   2LiOH  Li₂O + H₂O ; Mg(OH)₂  MgO + H₂O

Hydroxides of other alkali metals are stable towards heat while their nitrates give O₂ and nitrite.
2KNO₃ 2KNO₂ + O₂
(6) Both Li and Mg combine directly with N₂ to give nitrides Li₃N and Mg₃N₂. Other alkali metals combine at high temperature, 6Li + N₂ 2Li₃N; 3Mg + N₂ Mg₃N₂. Both the nitrides are decomposed by water to give NH₃
Li₃N + 3H₂O

 3LiOH + NH₃ ;
Mg₃N₂ + 6H₂O   3Mg(OH)₂ + 2NH₃
(7) Bicarbonates of Li and Mg are more soluble in water than carbonates whereas carbonates of alkali metals are more soluble.
(8) Both Li and Mg combine with carbon on heating.
      2Li + 2C  Li₂C₂ ; Mg + 2C  Mg C₂

(10) Both have high polarizing power.

Polarizing Power = Ionic charge / (ionic radius)₂.
(11) Li and Mg Form only monooxide on heating in oxygen.
4Li + O₂ 2 Li₂O ; 2Mg + O2  2 MgO
(12) Li₂SO₄ like MgSO₄ does not form alums.
(13) The bicarbonates of Li and Mg do not exist in solid state, they exist in solution only.
(14) Alkyls of Li and Mg (R. Li and R.MgX) are soluble in organic solvent.
(15) Lithium chloride and MgCl₂ both are deliquescent and separate out from their aqueous solutions as hydrated crystals, LiCl. 2H₂O and MgCl₂ . 2H₂O.
Uses of Lithium
(1) It is used as a deoxidiser in metallurgy of Cu and Ni.
(2) It is used as an alloying metal with
(i) Pb to give toughened bearings.
(ii) Al to give high strength Al _alloy for aircraft industry.
(iii) Mg (14% Li) to give extremely tough and corrosion resistant alloy which is used for armour plate in aerospace components.

Sodium and its compounds
(1) Ores of sodium : NaCl (common salt), NaNO₃ (chile salt petre), Na₂SO₄ .10 H₂O (Glauber's salt), borax (sodium tetraborate or sodium borate) (Na₂B₄O₇ .10H₂O) .
(2) Extraction of sodium : It is manufactured by the electrolysis of fused sodium chloride in the presence of CaCl₂ and KF using graphite anode and iron cathode. This process is called  Down process.
NaCl ⇌ Na⁺ + Cl⁻ .
At cathode : Na⁺ + e⁻ Na ;
At anode : Cl⁻ Cl⁺ e⁻ ; Cl⁺ Cl  Cl₂ ↑
Sodium cannot be extracted from aqueous NaCl because E₀ H₂O/H₂ (–0.83V) is more than E₀ Na⁺ / Na (– 2.71V).
Anode and cathode are separated by means of a wire gauze to prevent the reaction between Na and Cl₂ .

(3) Compound of sodium
(i) Sodium chloride : It is generally obtained by evaporation of sea water by sun light. However NaCl so obtained contains impurities like CaSO₄, CaCl₂ and MgCl₂ which makes the salt deliquescent.It is then purified by allowing HCl gas to pass through the impure saturated solution of NaCl .The concentration of Cl −ions increases and as a result pure NaCl gets precipitated due to common ion effect.
(ii) Sodium hydroxide NaOH (Caustic soda)
Preparation
(a) Gossage process :

(b) Electrolytic method : Caustic soda is manufactured by the electrolysis of a concentrated solution of NaCl.
           At anode: Cl⁻ discharged; At cathode: Na⁺ discharged
(c) Castner ‐ Kellener cell (Mercury cathode process) :
NaOH obtained by electrolysis of aq. solution of brine. The cell comprises of rectangular iron tank divided into three compartments.
Outer compartment – Brine solution is electrolysed ; Central compartment – 2% NaOH solution and H₂

Properties : White crystalline solid, highly soluble in water, It is only sparingly soluble in alcohol.
(a) Reaction with salt :

Zn, Al, Sb, Pb, Sn and As forms insoluble hydroxide which dissolve in excess of NaOH (amphoteric hydroxide).
NH₄ Cl + NaOH  NaCl + NH₃ H₂O
(b) Reaction with halogens : 

(c) Reaction with metals : Weakly electropositive metals like Zn, Al and Sn etc.
Zn + 2NaOH  Na₂ZnO₂ + H₂ 
(d) Reaction with sand, SiO₂:
2NaOH + SiO₂   Na₂SiO₃ +  H₂O
                               Sod. silicate (glass)
(e) Reaction with CO: 

^{NaOH breaks down the proteins of the skin flesh to a pasty mass, therefore it is commonly known as caustic soda.
Caustic property : sodium hydroxide breaks down the proteins of the skin flesh to a pasty mass, therefore, it is commonly known as caustic soda.
Uses : Sodium hydroxide is used :
(a) in the manufacture of soidum metal, soap (from oils and fats), rayon, paper, dyes and drugs,
(b) for mercurinzing cotton to make cloth unshrinkable and
(c) as a reagent in the laboratory.}

(iii) Sodium carbonate or washing soda, Na₂CO₃
It exists in various forms, namely anhydrous sodium carbonate Na₂CO₂ (soda‐ash); monohydrate Na₂CO₃.H₂O (crystal carbonate); hyptahydrate Na₂CO₃.7H₂O and decahydrate Na₂CO₃.10 H₂O (washing soda or sal soda).
Preparation :
(a) Solvay process : In this process, brine (NaCl) , NH₃ and CO₂ are the raw materials.
                 NH₃ + CO₂ + H₂O   NH₄HCO₃  

                      
CaCl₂ so formed in the above reaction is a by product of solvay process.
Properties
(a) 
      Na₂ CO₃ + H₂O    Na₂CO₃
It does not decompose on further heating even to redness (m.pt. 853°C)

^{(b) It is soluble in water with considerable evolution of heat.}

(c) It is readily decomposed by acids with the evolution of CO₂ gas.
(d) Na₂CO₃ + H₂O + CO₂ 2NaHCO₃
Uses : In textile and petroleum refining, Manufacturing of glass, NaOH soap powders etc..
  (iv) Sodium peroxide (Na₂O₂)
Preparation : It is manufactured by heating sodium metal on aluminium trays in air (free fromCO₂)
2Na + O₂ (air)  Na₂O₂
Properties : (a) When pure it is colourless. The faint yellow colour of commercial product is due to presence of small amount of superoxide (NaO₂).
(b) On coming with moist air it become white due to formation of NaOH and Na₂CO₃ .
2Na₂O₂ + 2H₂O   4 NaOH + O₂ ;
2NaOH + CO₂   Na₂CO₃ + H₂O
(c) It is powerful oxidising agent. It oxidises Cr (III) hydroxide to sodium chromate, Mn (II) to sodium manganate and sulphides to sulphates.

Uses :  As a bleaching agent and it is used for the purification of air in confined spaces such as submarines because it can combine with CO₂ to give Na 2CO₃ and oxygen, 2CO₂ + 2Na₂O₂   2Na₂CO₃ + O₂ .
(v) Micro cosmic salt [Na (NH₄)HPO₄.4H₂O]
Prepared by dissolving equimolar amounts of Na₂HPO₄ and NH₄Cl in water in 1 : 1 ratio followed by crystallization
Crystallization

Chemical properties :
On heating M.C.S, NaPO₃ is formed. NaPO₃ forms coloured beads with oxides of transition metal cloudySiO₂

NaPO₃ + CoO  CoNaPO₃ (blue bead)
NaPO₃ + MnO  NaMnO₃ (blue bead)
Uses :
(a) For the formation of sodium meta phosphate and copper sodium phosphate
(b) It is used for the detection of colured ion
(c) It is espacially used for testing silica with which a cloudy bead containing floating properties of silica is obtained.
(vi) Sodium bi Carbonate (NaHCO₃, Baking soda)
Preparation : It is an inter mediate compound in manufacture of sodium carbonate by the solvay’s process

Uses :
(a) Baking powder contains NaHCO₃,Ca(H₂ PO₄)₂ and starch.
Improved Baking powder contains 40% starch 30% NaHCO₃, 20% NaAl (SO₄)₂ and 10% CaH₂ (P0₄)
(b) In pharmaceutical industry (Antacids etc.)
(c) Fire extingerishers.
(vii) Sodium Sulphate Na₂SO₄ or salt cake
Preparation : It is the by‐product of HCl industry

Properties : When aqueous solution of Na₂SO₄ is cooled below 32^oC Glauber’s salt
(Na₂SO₄10H₂O)gets crystallised and if cooled to 12^oC, Na₂SO₄7H₂O crystals are formed.

^{Uses : Na}2^SO4 ^{finds use in paper industry detergent and glass manufacturing.}