Group I (or I A) Elements: (Alkali Metals)

The alkali metals show regular trends in their physical and chemical properties with the increasing atomic number. The atomic, physical and chemical properties of alkali metals are discussed below

Electronic Configuration:-All the alkali metals have one valence electron, ns1

Atomic and Ionic Radii:-

  • Largest sizes in a particular period of the periodic table.
  • With increase in atomic number, the atom becomes larger.
  • The atomic and ionic radii of alkali metals increase on moving down the group

Ionisation Enthalpy

  • The ionisation enthalpies of the alkali metals are considerably low and decrease down the group from Li to Cs.
  • The second ionization enthalpies of alkali metals are very high.

Melting and boiling points: Alkali metals are soft and have low melting and boiling point.

Density: Densities of alkali metals are quite low as compared to other metals

Electropositive or metallic character: All the alkali metals are strongly electropositive or metallic in character

Oxidation states: All the alkali metals predominantly exhibit an oxidation state of +1 in their compounds

Characteristic flame colouration: All the alkali metals and their salts impart characteristic flame colouration

Photoelectric effect: Phenomenon of ejection of electrons when electromagnetic radiation of suitable frequency strikes metal surface is called photo-electric effect. Alkali metals exhibit photo-electric effect

Physical appearance:- All the alkali metals are silvery white, soft and light metals. They have low density which increases down the group from Li to Cs. However, potassium is lighter than sodium. 

Reducing Nature: Alkali metals are strong reducing agents

Chemical Properties

The alkali metals are highly reactive due to their large size and low ionisation enthalpy. The reactivity of these metals increases down the group.

(i) Reactivity towards air: The alkali metals tarnish in dry air due to the formation of their oxides which in turn react with moisture to form hydroxides. They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form superoxides. The superoxide O2 ion is stable only in the presence of large cations such as K, Rb, and Cs.

                       4 Li + O2 →2 Li2O (oxide)   

                      2 Na + O 2 →Na2O2 (peroxide)   

                M + O2 → MO2 (superoxide)(M = K, Rb, Cs).  

(ii) Reactivity towards water: The alkali metals react with water to form hydroxide and hydrogen.          

                    2M + 2H2O → 2 M+ + 2OH- + H2 (M = an alkali metal)

It may be noted that although lithium has most negative E0 value, its reaction with water is less vigorous than that of sodium which has the least negative E0 value among the alkali metals. This behaviour of lithium is attributed to its small size and very high hydration energy. Other metals of the group react explosively with water

(iii) Reactivity towards hydrogen: The alkali metals react with hydrogen at about 673K (lithium at 1073K) to form hydrides. All the alkali metal hydrides are ionic solids with high melting points. 2M + H2 → 2MH.

(iv) Reactivity towards halogens: The alkali metals readily react with halogens to form ionic halides. Lithium iodide is the most covalent in nature.

(v) Reducing nature: The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful.

(vi) Solutions in liquid ammonia: The alkali metals dissolve in liquid ammonia giving deep blue solutions which are conducting in nature. M + (x + y)NH3 → [M(NH3) x]+ + [e(NH3)y]-

The blue colour of the solution is due to the ammoniated electron which absorbs energy in the visible region of light and thus imparts blue colour to the solution

Anomalous Properties of Lithium

The anomalous behaviour of lithium is due to the: 

     (i) Exceptionally small size of its atom and ion, and

     (ii) High polarizing power (i.e.., charge/ radius ratio).

Points of Difference between Lithium and other Alkali Metals

  • Lithium is much harder. Its M.P. and B.P. are higher than the other alkali metals.
  • Lithium is least reactive but the strongest reducing agent among all the alkali metals. On combustion in air it forms mainly monoxide, Li2O and the nitride, Li3N unlike other alkali metals.
  • LiCl is deliquescent and crystallizes as a hydrate, LiCl.2H2O whereas other alkali metal chlorides do not form hydrates.
  • Lithium hydrogen carbonate is not obtained in the solid form while all other elements form solid hydrogen carbonates.
  • Lithium unlike other alkali metals forms no ethynide on reaction with ethyne.
  • Lithium nitrate when heated gives lithium oxide, Li2O, whereas other alkali metal nitrates decompose to give the corresponding nitrite.
  • LiF and Li2O are comparatively much less soluble in water than the corresponding compounds of other alkali metals.

The similarity between lithium and magnesium is particularly striking and arises because of their similar sizes: atomic radii, Li = 152 pm, Mg = 160 pm; ionic radii: Li+ = 76 pm, Mg2+= 72 pm. The main points of similarity are:

  • Both lithium and magnesium are harder and lighter than other elements in the respective groups.
  • Lithium and magnesium react slowly with water. Their oxides and hydroxides are much less soluble and their hydroxides decompose on heating. Both form a nitride, Li3N and Mg3N2, by direct combination with nitrogen.
  • The oxides, Li2O and MgO do not combine with excess oxygen to give any superoxide.
  • The carbonates of lithium and magnesium decompose easily on heating to form the oxides and CO2. Solid hydrogen carbonates are not formed by lithium and magnesium.
  • Both LiCl and MgCl2 are soluble in ethanol.
  • Both LiCl and MgCl2 are deliquescent and crystallize from aqueous solution as hydrates, LiCl·2H2O and MgCl2·8H2O.



Sodium Carbonate (Washing Soda),Na2CO3·10H2O

Preparation: Solvay’s process.

The overall process is:

2 NaCl + CaCO3 → Na2CO3 + CaCl2

But it is done in following four steps

  1. NaCl + CO2 + NH3 + H2O → NaHCO3↓ + NH4Cl   (I)
  2. CaCO3 → CO2 + CaO  (II)
  3. 2NH4Cl + CaO → 2NH3 + CaCl2 + H2O  (III)
  4. 2NaHCO3 → Na2CO3 + H2O + CO2  (IV)









Properties: Sodium carbonate is a white crystalline solid which exists as a decahydrate, Na2CO3·10H2O. It is readily soluble in water. On heating, the decahydrate loses its water of crystallisation to form monohydrate. Above 373K, the monohydrate becomes completely anhydrous and changes to a white powder called soda ash.

Na2CO3.10H2ONa2CO3.H2O+9H2O            Na2CO3.H2O>100 oCNa2CO3+H2O  

(i) It is used in water softening, laundering and cleaning.
(ii) It is used in the manufacture of glass, soap, borax and caustic soda.
(iii) It is used in paper, paints and textile industries.
(iv) It is an important laboratory reagent both in qualitative and quantitative analysis

Sodium Hydroxide (Caustic Soda), NaOH manufactured by the electrolysis of brine solution in the Castner –Kellner cell








The cell is divided into three partitions. The graphite anode is fixed on the two outer partitions while the cathode is an iron rod which is fit in the central compartment. A layer of mercury, which serves an intermittent electrode, covers the bottom of the cell. Brine solution is placed in the outer compartments. The equation for the electrolytic process is as given below.

Outer Compartment:-The mercury layer acts as the cathode:-

NaCl® Na+ + Cl

At Cathode:-Na+ + e- ® Na

At Anode:  Cl—  ® Cl +e- and Cl + Cl ® Cl2                 

the sodium, which is formed, dissolves in mercury to form an amalgam. It goes into the central compartment.

Central Compartment: - Here sodium amalgam acts as anode by induction. Sodium reacts with water to form NaOH and H2. Hydrogen gas is liberated at the cathode. The equations are as given below.             2Na/Hg + 2H2O ® 2NaOH + H2 + 2Hg

When the conc. of NaOH reaches 20%, it is removed and the solution is concentrated to get solid NaOH.


Group II ( or II A) Elements: (Alkali Earth Metals)

These (except beryllium) are known as alkaline earth metals. The first element beryllium differs from the rest of the members and shows diagonal relationship to aluminium.

Electronic Configuration: These elements have two electrons in the s orbital of the valence shell [noble gas] ns2.



Atomic and ionic radii of alkaline earth metals increases down the group and are smaller than the corresponding members of the alkali metals



The alkaline earth metals have low ionisation enthalpies

Since the atomic size increases down the group, their ionisation enthalpy decreases. Although IE1 values of alkaline earth metals are higher than those of alkali metals, the IE2 values of alkaline earth metals are much smaller than those of alkali metals

Melting and boiling points: Alkaline earth metals have higher melting and boiling points than the corresponding alkali metals

Electropositive or metallic character: The electropositive character increases down the group i.e., from Be to Ba but alkaline earth metals are not as strongly electropositive as the alkali metals

Oxidation states: All the alkaline earth metal exhibits an oxidation state of +2 in their compounds



Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group.

             Be2+> Mg2+ > Ca2+ > Sr2+ > Ba2+



The alkaline earth metals, in general, are silvery white, lustrous and relatively soft but harder than the alkali metals

Chemical Properties: The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group.

(i) Reactivity towards air and water: Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface. However, powdered beryllium burns brilliantly on ignition in air to give BeO and Be3N2

(ii) Reactivity towards the halogens: All the alkaline earth metals combine with halogen at elevated temperatures forming their halides.

                               M + X 2 ® MX2 ( X = F, Cl, Br, l)

(iii) Reactivity towards hydrogen: All the elements except beryllium combine with hydrogen upon heating to form their hydrides,MH2. .BeH2, however, can be prepared by the reaction of BeCl2 with LiAlH4 as follows:                                                                             

                              2BeCl 2 + LiAlH4 ® 2BeH2 + LiCl + AlCl3

(iv) Reactivity towards acids: The alkaline earth metals readily react with acids liberating hydrogen.

                                         M + 2HCl ® MCl2 + H2

(v) Reducing nature: alkaline earth metals are strong reducing agents

(vi) Solutions in liquid ammonia: alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.                 

M + (x+ y) NH3 ® [M(NH3) x]2+ + 2[e(NH3)x]-




(i) Oxides and Hydroxides:

     The alkaline earth metals burn in oxygen to form the monoxide, MO

(ii) Halides: Except for beryllium halides, all other halides of alkaline earth metals are ionic in nature. Beryllium chloride has a chain structure in the solid state as shown

(iii) Carbonates: Carbonates of alkaline earth metals are insoluble in water and can be precipitated by addition of a sodium or ammonium carbonate solution to a solution of a soluble salt of these metals. The solubility of carbonates in water decreases as the atomic number of the metal ion increases. All the carbonates decompose on heating to give carbon dioxide and the oxide. Beryllium carbonate is unstable and can be kept only in the atmosphere of CO2.

(iv) Sulphates: The sulphates of the alkaline earth metals are all white solids and stable to heat. BeSO4, and MgSO4 are readily soluble in water the solubility decreases from CaSO4 to BaSO4



Beryllium shoes anomalous behaviour as compared to magnesium and rest of the members. Further, it shows diagonal relationship to aluminium

(i) Beryllium has exceptionally small atomic and ionic sizes and thus does not compare well with other members of the group. Because of high ionisation enthalpy and small size it forms compounds which are largely covalent and get easily hydrolysed.

(ii) Beryllium does not exhibit coordination number more than four as in its valence   shells there are only four orbitals. The remaining members of the group can have a coordination number of six by making use of d-orbitals

(iii) The oxide and hydroxide of beryllium, unlike the hydroxides of other elements    in the group, are amphoteric in nature.

Diagonal Relationship between Beryllium and Aluminium

The ionic radius of Be2+ is estimated to be 31 pm; the charge/radius ratio is nearly the same as that of the Al3+ ion.:

(i) Like aluminium, beryllium is not readily attacked by acids because of the presence of an oxide film on the surface of the metal.

(ii) Beryllium hydroxide dissolves in excess of alkali to give a beryllate ion, Be(OH)4]2– just as aluminium hydroxide gives aluminate ion, [Al(OH)4].

(iii) The chlorides of both beryllium and Aluminium has Cl bridged chloride structure in vapour phase. Both the chlorides are soluble in organic solvents and are strong Lewis acids. They are used as Friedel Craft catalysts.

(iv) Beryllium and aluminium ions have strong tendency to form complexes, BeF42–, AlF63–.


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