Alkali and Alkaline Earth Metals
Li, Na, K, Rb, Cs are known as alkalimetals as their oxides on reaction with water to form hydroxides, which are strongly alkaline in nature. They have ns valency shell configuration and exhibit +1 oxidation state. They do not occur in free state due to high reactivity. They are soft, possess low density, melting point, boiling point, ionization potential, electron affinity, electronegativity. They have high electropositive character and metallic nature. Density of these elements increase from Li to Cs. Where as density of K is lessthan that of Na, due to abnormal increase in atomic size, vacant 3d sub-shell and inter atomic distance in the crystal lattice of potassium. They exhibit photo electric effect when they are exposed to light (due to emission of electrons). They give different colours with flame (due to excitation & de excitation of electrons). Li gives crimson red, Na gives golden yellow, K, Rb, Cs give violet colour. Alkali metals are also detected by photometry or atomic absorption spectroscopy.
The extent of hydration of these ions: Li+> Na+ > K+ > Rb+ > Cs+.
Electrical conductance of these ions in water: Li+ < Na+ < K+ < Rb+ < Cs+.
Due to high reactivity, Li is stored in paraffin, Na is stored in kerosene. On reaction with air Li forms monoxide, Na forms peroxide, others form super oxides. They give ionic hydrides with H2.
Order of ionic nature of these hydrides: CsH > RbH > KH > NaH > LiH.
Except Li, other elements give ionic halides with halogens. Sodium decomposes water to liberate H2 gas. Due to small atomic size, high electronegativity, absence of d -orbitals Li exhibit anomalous properties like formation of carbides, Nitrides, hard metal, low solubilities of hydroxide, carbonate, phosphate and flouride. Li is diagonally related with Mg, due to similar atomic size, electronegativity, polarizing power. Li & Mg are slow to react with water, form nitrides, monoxides, their halides are covalent and chlorides are deliquescent, highly hydrated, carbonates, phosphates are sparingly soluble in & H2O.
Sodium is used as reagent in organic reactions (e.g.:Wurtz reaction), as a catalyst in the preparation of rubber, sodium amalgam, in lassaigne's test (to detect S, N, halogens). potassium is used as an electrolyte in storage batteries, in the manufacture of soft soaps, photo electric cells. Na+, K+ help in maintaining osmotic pressure in the cell, essential for the synthesis of proteins, for the metabolism of glucose inside the cell, in producing electrical potential across the cell membrane, in Na+ ion transport from the cell (Sodium pump).
NaCl is used in diet, pickles, in preservation of meat, fish, to prepare freezing mixture with ice, in the manufacture of Na and Cl2.
Other important compound of Na is NaOH. It is known as caustic soda, as it decomposes the muscle proteins and makes pulp. It is used to absorb SO2 near electrical generators, in mercarizing cotton, refining of petroleum, soap, paper, rayon industries. NaOH can be manufactured by
1. Causticizing process (Gossage process): In this method milk of lime is added to 10% warm solution of Na2 CO3 to produce NaOH.
Ca(OH)2 + Na2CO3 
CaCO3 + 2 NaOH
Castner - Kellner (Mercury cathode) process: This cell consists of a large rectangular iron tank. It's divided into three compartments by two slate partitions which are not touching the bottom of the tank, but goes into mercury. Mercury acts as an intermediate electrode due to induction.
Outer compartments are provided with graphite anodes, where Hg acts as cathode. This compartment is filled with brine solution. Reactions taking place in this compartment as follows.
2 NaCl → 2 Na+ + 2 Cl-
At anode: 2 Cl- → Cl2 + 2 e- (Oxidation)
At cathode: 2 Na+ + 2 e- + Hg → Na2Hg (Reduction)
Middle compartment is provided with a series of iron rods, acts as cathode, where Hg acts as anode. It is filled with dilute NaOH. Na2Hg from outer compartment comes into middle compartment due to rocking motion of cell on eccentric wheel. Reactions taking place in the middle compartment are
At anode : Na2Hg → 2 Na+ + 2 e- + Hg (oxidation)
At cathode : 2 H2O + 2 e- → H2 + 2 OH- (Reduction)
2 Na+ + 2 OH- → 2 NaOH
When the concentration of NaOH is reached 20%, it is removed and evaporated in iron tanks at 500°C to get solid NaOH.
Properties of NaOH: It is white, crystalline, deliquescent solid. M.P: 519 K. It forms hydrates NaOH. x H2O (x = 1, 2, 7). It absorbs CO2.
Zn, Al, C, Si Displace H2 from NaOH
Zn + 2 NaOH → Na2 ZnO2 (sod. Zincate) + H2
2 Al + 6 NaOH → 2 Na3 AlO3 (Sod. Aluminate) + 3 H2
2 C + 6 NaOH → 2 Na + 2 Na2CO3 + 3 H2
Si + 2 NaOH + H2O → Na2SiO3 + 2 H2
NaOH react with ammonium salts to give NH3
NH4Cl + NaOH → NaCl + H2O + NH3
Reaction with halogens:
2 NaOH (cold, dil) + 2 F2 → 2 NaF + OF2 + H2O
4 NaOH (hot, conc.) + 2 F2→ 4 NaF + O2 + 2 H2O
Cl2 + 2 NaOH (cold, dil) → NaCl + NaOCl (Sod. hypochlorite) + H2O
3 Cl2 + 6 NaOH (hot, conc.) → 5 NaCl + NaClO3 (Sod. Chlorate) + 3H2O
Other reactions:
4 S + 6 NaOH → Na2S2O3 + 2 Na2S + 3 H2O
4 P + 3 NaOH + 3 H2O → 3 Na H2PO2 (Sod. hypophosphite) + PH3
2 NaOH + CO2→ Na2CO3 + H2O
3 NaOH + FeCl3→ Fe(OH)3 [reddish brown ppt.] + 3 NaCl
2 NaOH + FeSO4 → Fe(OH)2 [Palegreen ppt.] + Na2SO4
ZnSO4 + 2 NaOH → Na2SO4 + Zn(OH)2
Zn(OH)2 + 2 NaOH (excess) → Na2ZnO2 (soluble) + 2 H2O
AlCl3 + 3 NaOH → Al(OH)3 + 3 NaCl
Al(OH)3 + NaOH (excess) → NaAlO2 (Sod. meta aluminate) + 2 H2O
(Soluble)
NaOH + HCl → NaCl + H2O
2 AgNO3 + 2 NaOH → 2 AgOH + 2 NaNO3
2 AgOH → Ag2O + H2O
Sodium Carbonate: Na2CO3. 10 H2O known as washing soda, Na2CO3 is called as soda ash. It can be prepared by Leblanc or Solvay process.
Leblanc Process: Raw material NaCl, conc. H2SO4, lime stone react together to give Na2CO3. In this method HCl, Na2S are by products.
NaCl + H2SO4 → NaHSO4 + HCl
NaHSO4 + NaCl → HCl + Na2SO4
Na2SO4 + 4 C → Na2S + 4 CO
Solvay (Ammonia - Soda) Process: Solvay process involves five stages.
Ist Stage: Saturation of brine with NH3: 30% brine solution is saturated with NH3 and small amount of CO2. Mg, Fe, Ca impurities are precipitated and removed.
NH3 + H2O → NH4OH
2 NH3 + H2O + CO2→ (NH4) 2CO3
CaCl2 + 2 NH4OH → Ca(OH)2 ↓ 2 NH4Cl
CaCl2 + (NH4)2 CO3→ CaCO3 ↓ + 2 NH4Cl
2nd Stage: Carbonation: Sodium bicarbonate is formed due to the reaction between
ammonical brine of Ist stage and CO2.
CaCO3
CaO + CO2
NH3 + CO2 + H2O → NH4HCO3
NH4HCO3 + NaCl → NaHCO3 + NH4Cl
3rd Stage: Filtration: NaHCO3 is filtered by rotary vacuum filter and filtrate is pumped to ammonia recovery tower.
4th Stage: NH3 regeneration: Filtrate is mixed with Ca(OH)2 and steam to get NH3
NH4 HCO3 NH3 + H2O + CO2
2 NH4Cl + Ca(OH)2 2 NH3 + CaCl2 + 2H2O
5th Stage: Calcination: NaHCO3 on calcination gives Na2CO3
2 NaHCO3 Na2CO3 + CO2+ H2O
Properties of Na2CO3:
It is white, crystalline, efflorescent solid. Its M.P.: 1125 K. It give alkaline solution due to anionic hydrolysis.
CO3-2 + 2 H2O → H2CO3 + 2 OH-
Acids liberate CO2 from Na2CO3, gives Na2S2O3 with S and SO2. Na2S2O3 reacts with SiO2 to give Na2SiO3.
Uses: Used to remove hardness of water, as a reagent in qualitative & quantitative analysis, refining petroleum, in glass, dye, paper industry, laundries, to prepare "ultra marines" [ultramarine is aluminosilicate and is used as pigment.
e.g.: Na3(AlO2)6 (SiO2)6 Cl2 - sodalite], to prepare fusion mixture (Na2CO3 + K2CO3 )
Sodium bi Carbonate: It is a white, crystalline solid. It does not give any colouration with phenolphthalein but gives pink colour with methyl orange indicator. It is used as antacid, in fire extinguishers, in baking cakes, effervescent drinks, reagent in laboratory.
Alkaline Earth Metals
Mg, Ca, Ba, Sr, Ra (except Be) are known as alkaline earth metals as their oxides and hydroxides are alkaline in nature & their oxides are found in thin rockey outer layer of the earth's crust. General electronic configuration is ns, exhibit +2 oxidation state. Ca gives brick red, Ba gives apple green, Sr. gives Crimson red colour to the flame.
Important minerals of Be are Beryl (3 BeO. Al2O3. 6 SiO2), Phenacite (2 BeO. SiO2), Ca are Dolomite (CaCO3.MgCO3), Gypsum (CaSO4 . 2 H2O), Sr are Celestite (SrSO4), Strontianite (SrCO3), Ba are Berytes (BaSO4), Witherite (BaCO3), Mg are Magnesite (MgCO3), Carnallite (KCl. MgCl2 .6 H2O)
General Characteristics of Compounds of Alkaline Earth Metals: BeO, Be(OH)2 are amphoteric in nature (reacts with acids and bases). Thermal stability, solubility increases from Mg(OH)2 to Ba(OH)2. Be halides are covalent, other halides are ionic. Carbonates of this group elements are insoluble in water. They decompose on heating. BeSO4, MgSO4 are soluble in water. The solubility decreases from CaSO4 to BaSO4. Nitrates decompose on heating Alkaline earth metal carbonates, sulphates solubility decreases down the group as the hydration enthalpies decrease down the group.
Anomalous behaviour of Be: Due to small atomic size, high electronegativity Be shows anomalous behaviour.
Be compounds are covalent, not effected by dry air, amphoteric, do not respond to flame test, forms complexes, maximum covalency is 4 (for other elements is 6).
Diagonal relationship with Al: Both Be & Al forms covalent compounds, undergo hydrolysis, forms complexes and passive by conc. HNO3 loss of initial reactivity of Be, Al, Fe, Cr with conc. HNO3 is known as "passivity" Be & Al are amphoteric metals. Carbides on hydrolysis give methane, and they are called "Methanides".
Be2C + 4 H2O → 2 Be (OH)2 + CH4
Al4C3 + 12 H2O → 4 Al(OH)3 + 3 CH4
Where as CaC2 is acetylide. [CaC2 + 2 H2O → C2H2 + Ca(OH)2]
II A elements react with H2 to form ionic hydrides.
Order of ionic nature: BaH2 > SrH2 > CaH2 > MgH2 > BeH2
Be forms covalent, hygroscopic halides, which can fume in air on hydrolysis. Other elements form ionic halides. Be is used in making electrodes in neon signs, Cu alloys; Mg is reducing agent and deoxidiser in metallurgy, to prepare alloys like electron magnalium.
Compounds of Calcium: CaO is called as quick lime, Ca (OH)2 is slaked lime, CaSO4 . 2 H2O is gypsum, CaSO4 . 
H2O is plaster of paris. A mixture of 1 part of slaked lime, Water and 3 parts of sand is known as "mortar" sand makes the mass porous and prevents the formation of cracks.
setting of mortar is due to the formation of CaSiO3
Ca(OH)2 + SiO2→ CaSiO3 + H2O
When mortar is mixed with Cement, it is called "Cement mortar". Lime Stone and Clay on heating gives "Hydraulic Mortar". Which is an antiseptic and bleaching agent. Gypsum on heating upto 393 K gives plaster of paris, above 393 K gives "dead burnt plaster".
CaSO4 . 2 H2O 
CaSO4 . H2O 
CaSO4
(Gypsum) (P.O.P.) (Dead burnt)
Plaster of Paris set in to hard mass in 2 stages.
CaSO4 . H2O + 1 H2O 
CaSO4 . 2 H2O
Orthorhombic dihydrate
CaSO4 . 2 H2O 
CaSO4 . 2 H2O
Orthohombic dihydrate Monoclinic dihydrate
P.O.P. is used in building industry, to form plaster over the bone fracture or sprain, in dentistry, in ornamental works, in casts of busts and statues.
Gypsum (2-3%) is mixed with Cement to delay setting time.
Biological importance of Ca & Mg: An adult has 1200 g of Ca and 25 g of Mg.
Role of Ca+2: It is necessary for blood clotting, muscle contraction to maintain regular heart beat, in bones and teeth. Ca in plasma are regulated by Calcitonin & parathyroid hormones.
Role of Mg+2: Chlorophyll in plants contain Mg+2. It is cofactor for the enzymes that utilise ATP in Phosphate transfer. Enzymes "Phosphohydrolase" & "Phosphotransferases" contain
Mg+2. It is present in animal cells.
CEMENT
Cement is an important building material, introduced by Joseph Aspdin. The average composition of cement is
CaO: 50 - 60%
SiO2: 20 - 25%
Al2O3: 5 - 10%
MgO: 2 - 3%
Fe2O3: 1 - 2%
For good quality of cement, the ratio of SiO2 to Al2O3 must be between 2.5 - 4.0. Cement is basically a mixture of clinker (formed by clay and lime on strong heating) and 2-3% Gypsum (to regulate the setting time). Cement is used in concrete and plastering. It is used in the construction of dams, bridges, roads, tunnels etc.
