(i) Because of similatity in electronic configuration, they exhibit similar properties. A regular gradation in their properties with increase in at. no. is observed due to increasing size of atoms/ions and the low binding energy of valency electrons.

(ii) Of all the alkali metals, only sodium and potassium are found in abundance in nature. Francium occurs only in minute quantities as a
radioactive decay product.

See Table for electronic configuration.

(a) 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 `ns^0` configuration). That is why alkali metal salts are colourless and diamagnetic.

(b) 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

`text( Li Na K Rb Cs Fr)`

Atomic radius (pm) `text(152 186 227 248 265 375)`

Ionic radius of `M^(+)` ions (pm) `60` `95` `133` `148` `169` -

(c) 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^(o) C`.


(iii) `K` is lighter than `Na` because of its unusually large atomic size.

(iv) In solid state, they have body centred cubic lattice.


(d) Melting point and Boiling point :

(i) All these elements possess low m.pt and b.pt in comparison to other group members.

`text( Li Na K Rb Cs Fr)`

m.pt (`K`) `453.5` `370.8` `336.2` `312.0` `301.5` `-`

b.pt (`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 m.pt. and b.pt. On moving down the group, the atomic size increases and binding energy of their atoms in crystal lattice decreases which results lowering of m.pts.

(iii) Lattic energy decreases from `Li` to `Cs` and thus m.pt and b.pt also decrease from `Li` to `Cs`.

(e) Ionisation energy & 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`

`IE _1` `quad quad quad quad` `520` `495` `418` `403` `376` `-`

`IE _2` `quad quad quad quad` `7296` `4563` `3069` `2650` `2420` `-`

A jump in `2^(nd)` ionisation energy (huge difference) can be explained as,

`Li : 1s^2 2s^1 undersettext(2s electron) oversettext(removal of) -> Li^+ : 1s^2 undersettext(1s electron) oversettext(removal of)-> Li^(2+) : 1s^1`

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 `ns^1` 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.

(f) Oxidation number and valency :

(i) These elements easily form univalent `+ve` ion by losing solitary `ns^1` electron due to low ionisation energy values.

(ii) Alkali metals are univalent in nature and form ionic compounds. Lithium salts are, however, covalent.

(iii) 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 `M ^(2+)` ion and thus they show `+1` oxidation state.

(iv) Since the electronic configuration of `M^+` ions do not have unpaired electron and thus alkali metal salts are diamagnetic and colouress. Only those alkali metal salts are coloured which have coloured anions e.g. `K_2 Cr_2 O_7` is orange because of orange coloured `Cr_2 O_7 ^(2-)` ion, `KMnO_4` is violet because of violet coloured `MnO_4^( 1-)` ion.


(g) 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 when ions on dissolution water get hydration.

(ii) The hydration is an exothermic process i.e energy is released during hydration.

(iii) 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

`Mg^+ +Aq-> M^+`; `Delta H= -` energy

(iv) Smaller the cation, greater is the degree of hydration.

Hydration energy, `Li ^+ > Na ^+ > K ^+ > Rb ^+ > Cs^+`

(v) `Li ^+` being smallest in size has maximum degree of hydration and that is why lithium salts are mostly hydrated, `LiCl * 2 H_2O`. Also lithuim ion being heavily hydrated, moves very slowly under the influence of electric field and, therefore, is the poorest current conductor 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^+`


(h) Electronegativities :

(i) These metals are highly electropositive and thereby possess low values of electronegativities.

(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 -`

Note : `Fr` being radioactive elements and thus studies on physical properties of this element are limited.

(i) Specific heat : It decreases from `Li` to `Cs`. `Li` `Na` `K` `Rb` `Cs` `Fr`

Specific hem (Cal/g) `0.941` `0.293` `0.17` `0.08` `0.049 `-

(j) Conduction power : All are good conductors of heat & electricity, because of loosely held valence electrons.

(k) Standard oxidation potential and reducting properties

(i) Since alkali metals easily lose `ns^1` electron and thus they have high values of oxidation potential i.e.,


`M+aq->M^+ +e`


(ii) The standard oxidation potentials of a alkali metals (in volts) are listed below,
`quadLi` `Na` `quad quadK` `Rb` `quadCs`

`+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 `H_2` from `H _2O` and `HCl`.

`2H_2O+2M-> 2MOH+H_2`

`2HCl+2M-> 2 MCl+H_2`

(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 Lithitum

`Li(s)+Li(g); DH_1=` Heat of sublimation, `D H _s`

`ul(Li(g)-> Li_(g)^(+) + e; DH_2=IE_1)`

`Li_(g)^+ ->Li(aq)^(+) DH_3=-`Heat of hydration, `D H_ h`

`Li_(s) + H_2O->Li_(aq)^+ + e; Delta H= Delta H_1 +Delta H_2 + Delta H_3 = Delta H_s + IE_1 - Delta H_h`

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

(l) 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` and `Cs`-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 ^+`

I released : `Li ^+ > Na ^+ > K ^+ > Rb ^+ > Cs ^+`

frequency released : `Li ^+ < Na ^+ < K ^+ < Rb ^+ < Cs ^+`

(a) Occurrence : Alkali metals are very reactive and thus found in combined state Some important ores of alkali metals are given ahead.

(i) Lithuinr : Triphylite, Petalite, lepidolite, Spodumene [`LiAl` `( SiO _3)_3`] Amblygonite [ `Li (Al F )PO _4`]

(ii) Sodium : Chile salt petre ( `NaNO_ 3`), Sodium chloride ( `NaCl `), Sodium sulphate (`Na _2 SO_ 4`),
Borax ( `Na_2 B_4O_ 7 10 H _2 O` ), Glauber salt ( `Na _2 SO_4*10H_ 2 O`)

(iii) Potassium : Sylime (`KCl` ), carnallite ( `KCl *MgCl_2 *6 H _2O` )and felspar (`K _2 O * Al_ 2 O_3 *6Si O_2`)

(iv) Rubidium : Lithuim ores Lepidolite, triphylite contains `0.7` to ` 3%` `Rb _2 O`

(v) Caesium : Lepidolite, Pollucite contains `0.2` to `7% Cs_2 O`

(b) 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 fussion temperature.

Fused NaCl :

`NaCl overset(fusion)-> Na^+ +Cl^(-) `

Electrolysis of fused salt : Anode : ` 2Cl^(-) ->Cl_2 +2e`

cathode: `2Na^(+) +2 e-> 2 Na`

(c) 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.

(d) Formation of oxides and hydroxides :

(i) These are most reactive metals and have strong affinity for `O_2` 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,

`M+O_2-> undersettext(oxide)(M_2O)-> undersettext(peroxide)(M_2O_2)`

(ii) When burnt air (`O_2`), lithium forms lithium oxide (`Li _2 O`) sodium forms sodium peroxide (`Na_2 O_2`) and other alkali metals form super oxide (`MO_2` i.e. `KO _2* RbO _2` or `CsO _2`)

`2Li+1/2 O_2-> Ki_2O; 2Na+O_2-> Na_2O_2; K+O_2-> KO_2`

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 rhus combines with small anion `O^(2-)` to form stable `Li_2O` 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, `O_2^(2-)`to form stable oxides.

The monoxide, peroxides and superoxides have `O^2` and `O_2^(2-)`, `O_2^(1-)` ions respectively. The structures of each are

`[: undersettext(x x)overset(..)O:]^(-2)`, `[: underset(..)overset(..) O - underset(..) overset(..)O:]`, `[:overset(..)O...overset(..)O:]`

The `O_2^(-1)` ion has a three electron covalent bond and has one electron unpaired. It is therefore superoxides are paramagnetic and coloured `KO _2` 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

`M+H_2O-> MOH+1/2 H_2 ; Delta=-ve`

`LiO_2 +H_2O ->2LiOH; Delta H=-ve`

`Na_2O_2+2H_2O->2NaOH+H_2O; Delta H=-ve`

`2KO_2+2H_2O->2KOH +H_2O+O_2; Delta H=-ve`

The peroxides and superoxides act as strong oxidising agents due to formation of `H_2O_2`

(iv) The reactivity of alkali metals towards air and water increases from `Li` to `Cs` that is why Jithjum decomposes `H _2O` very slowly at `25^oC` whereas `Na` does so vigorously, K reacts producing a flame and `Rb`, `Cs` do so explosively.

`M+H_2O->MOH+1/2 H_2`

(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.

(e) Hydrides :

(i)These metal combines `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 + H_2 -> 2MH`; Reactivity towards `H_2` is `Cs < Rb < K < Na < Li`

(iii) The metal hydrides react with water to give `MOH` & `H_2`; `MH + H_2O -> MOH + H_2`

(iv) The ionic nature of hydrides increases from `Li` to `Cs` because of the fact that hydrogen is present in the these hydrides as `H^-` and the smaller cation will produce more polarisation of anion (according to Fajan rule) and will develop more covalent character.

(v) The electrolysis of fused hydrides give `H_2` at anode. `(NaH_(text(fused))text(Contains)) Na^(+) text(and) H^(-) ` i.e.,

At cathode: `Na^(+) + e -> Na;` At anode: `H^(-) -> 1/2H_2 + e`

(vi) Alkali metals also form hydrides like `NaBH_4,` `Li.A I H_4` which are good reducing agent.

(f) Carbonates and Bicarbonate :

(i) The carbonates `(M_2CO_3)` & bicarbonates `(MHCO_3)` are highly stable to heat, where `M` stands for alkali metals.

(ii) The stability of these salts increases with the increasing electropositive character from Li to Cs. It is therefore Liz `CO_3` decompose on heating, `Li_2CO_3 -> Li_2O + CO_2`

(iii) Bicarbonates are decomposed at relatively low temperature,

`2MHCO_3 overset(300^oC)-> M_2CO_3 +H_2O +CO_2`

(iv) Both carbonates and bicarbonates are soluble in water to give alkaline solution due to hydrolysis of carbonate ions or bicarbonate ions.

(g) 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 (The Fajan's rule) 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 > NaI`

(vi) Halides of potassium, rubidium and caesium have a property of combining with extra halogen atoms forming polyhalides.

`KI + I_2 -> KI_3`; In `KI_(3(aq))` the ions `K^(+)` and `I_3^(-)` are present.

(h) Solubility in liquid `NH_3` :

(i) These metals dissolve in liquid NH 3 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 `NH_3` ceases.

(iii) The metal atom is converted into ammoniated metal in i.e. `M^(+) (NH_3)` and the electron set free combines with `NH_3` molecule to produce ammonia solvated electron .

`Na +(x +y) -> undersettext(Ammoniated cation){NH_3[Na(NH_3)_x]^+} + undersettext(ammoniated electron){[e(NH_3)]^-}`


(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 + 2NH_3 -> 2MNH_2 + H_2` . Sodamide `(NaNH_2)` is a waxy solid, used in preparation of number of sodium compounds.

(i) Nitrates : Nitrates of alkali metals (`MNO_3`) are soluble in water and decompose on heating. `LiNO_3` decomposes to give `NO_2` and `O_2` and rest all give nitrites and oxygen.

`2MNO_3 -> 2MNO_2 + O_2` (except Li) ; `4 LiNO_3 -> 2Li_2O + 4NO_2 + O_2`

(j) Sulphates :

(i) Alkali metals' sulphate have the formula `M_2SO_4`.

(ii) Except `Li_2SO_4,` rest all are soluble water.

(iii) These sulphates on fusing v.rith carbon form sulphides, `M_2SO_4 + 4C -> M_2S + 4CO`

(iv) The sulphates of alkali metals (except Li) form double salts with the sulphate of the trivalent metals like Fe, AI, Cr etc. The double sulphates crystallize with large munber of water molecules as alum. e.g. `K_2SO_4 . Al_2 (SO_4)_3. 24 H_2O`

(k) 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 + H_2 overset(300^oC)-> 2NaH ; 2K + H_2 -> 2KH`

`2Na + Cl_2 -> 2NaC l; 2K + Cl_2 -> 2KCl`

`2Na + S -> Na_2S; 2K + S -> K_2S`

(ii) Li reacts, however directly with carbon and nitrogen to form carbides and nitrides.

`2Li + 2C -> LiCz; 6Li + 2N_2 -> 2 Li_3N`

(iii) The nitrides of these metals on reaction with water give `NH_3 + M_3N + 3H_2O -> 3MOH + NH_3`

(l) Reaction wlth acidic hydrogen : Alkali metals react with acids and other compounds containing aciclic hydrogen (i.e, II atom anached on F,O, N and triply bonded carbon atom, for example, `HF, H_2O, ROH, RNH_2, CHequivCH)` to liberate `H_2`.

`M +H_2O -> MOH +1/2 H_2 ; M + HX -> MX +1/2 H_2 `

`M + ROH -> ROH +1/2H_2 ; M +RNH_2 -> RNHNa 1/2H_2`

(m) 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.

The group `2` of the periodic table consists of six metallic elements. These are beryllium (`Be`), magnesium (`Mg`), calcium (`Ca`), strontium (`Sr` ), barium (`Ba`) and radium (`Ra`). These (except `Be` ) are known as alkaline earth metals as their oxides are alkaline and occur in earth crust.

Note : Radium was discovered in the ore pitch blende by madam Curie. It is radioactive in nature.

(a) `text(Physical state)` : All are greyish-white, light, malleable and ductile metals with metallic lustre. Their hardness progressively decrease with increase in atomic number. Althought these are fairly soft but relatively harder than alkali metals.

(b) `text(Atomic and ionic radii)` :

(i) The atomic and ionic radii of alkaline earth metals also increase dovm the group due to progressive addition of new energy shells like alkali metals.

`text( Be Mg Ca Sr Ba Ra)`

Atomic radius (pm) `112` `160` `197` `215` `222` `-`

Ionic radius of `M^(2+)` ion (pm) `31` `65` `99` `113` `135` `140`

(ii) The atomic radii of alkaline earth metals are however smaller than their corresponding alkali metal of the same period. This is due to the fact that alkaline earth metals possess a higher nuclear charge than alkali metals which more effectively pulls the orbit electrons towards the nucleus causing a decrease in size.

(c) `text(Density)` :

(i) Density decreases slightly up to `Ca` after which it increases. The decrease in density from `Be` to `Ca` might be due to less packing of atoms in solid lattice of `Mg` and `Ca`.

`Be` `Mg` `Ca` `Sr` `Ba` `Ra`

`1.84` `1.74` `1.55` `2.54` `3.75` `6.00`

(ii) The alkaline earth metals are more denser, heavier and harder than alkali metal The higher density of alkaline earth metals is due to their smaller atomic size and strong intermetallic bonds which provide a more close packing in crystal lattice as compared to alkali metals.

(d) `text(Melting point and Boiling point)` :

(i) Melting points and boiling points of alkaline earth metals do not show any regular trend.

`text(Be Mg Ca Sr Ba Ra)`

m.pt. (`K`) `text(1560 920 1112 1041 1000 973)`

b.pt (`K`) `text(2770 1378 1767 1654 1413 - )`

(ii) The values are, however, more than alkali metals. This might due to close packing of atoms in crystal lattice in alkaline earth metals.

(e) `text(Ionisation energy and electropositive or metallic character)` :

(i) Since the atomic size decreases along the period and the nuclear charge increases and thus the electrons are more tightly held towards nucleus. It is therefore alkaline earth metals have higher ionisation energy in comparison to alkali metals but lower ionisation energies in comparison to p - block elements.

(ii) The ionisation energy of alkaline earth metals decreases from `Be` to `Ba`.

`text(Be Mg Ca Sr Ba Ra)`

First ionisation energy (kJ `text(mol)^(-1)`) `text(899 737 590 549 503 509)`

Second ionisation energy (kJ `text(mol)^(-1)` ) `text(1757 1450 1146 1064 965 979)`

(iii) The higher values of second ionisation energy is due to the fact that removal of one electron from the valence shell, the remaining electrons are more tightly held in which nucleus of cation and thus more energy is required to pull one more electron from monovalent cation.

(iv) No doubt first ionisation energy of alkaline earth metals are higher than alkali metals but a closer look on 2nd ionisation energy of alkali metals and alkaline earth metals reveals that 2nd ionisation energy of alkali metals are more

`Li` `Be`

1st ionisation energy (kJ `text(mol)^(-1)`) `520` `899`

2nd ionisation energy (kj `text(mol)^(-1)`) `7296` `1757`

This may be explained as `Li : 1s^2 , 2s^2 undersettext(electron)oversettext(removal of 1s) -> Li^(+) : 1s^2 undersettext(electron) oversettext(removal of 1s))-> Li^(2+) : 2s^1`

`Be : 1s^2 , 2s^2 undersettext(electron) oversettext(removal of 2s)-> Be^(+) : 1s^2 , 2s^1 undersettext(electron) oversettext(removal of 2s)-> Be^(2+) : 1s^2`

The removal of 2nd electron from alkali metalstakes place from `1s` sub shell which are more closer to nucleus and exert more nuclear charge to hold up `1 s` electron core, whereas removal of 2nd electron from alkaline earth metals takes from `2s` sub shell. More closer are shells to the nucleus, more tightly are held electrons with nucleus and thus more energy is required to remove the electron.

(v) All these possess strong electropositive character which increases from `Be` to `Ba`.

(vi) These have less electropositive character than alkali metals as the later have low values of ionisation energy.

(f) `text(Oxidation number and valency)`

(i) The `IE_1` of the these metals are much lower than `IE_1` and thus it appears that these metals should form univalent ion rather than divalent ions but in actual practice, all these give bivalent ions. This is due to the fact that `M^(2+)` ion possesses a higher degree of hydration or `M^(2+)` ions are extensively hydrated to form `[M(H_2O)_x]^(2+)` , a hydrated ion. This involves a large amount of energy evolution which counter balances the higher value of second ionisation energy.

`M -> M^(2+) , Delta H = IE_1 + E_2`

`M^(2+) + xH_2O -> [M(H_2O)_x]^(2+); Delta H = -` hydration energy.

(ii) The tendency of these metals to exist as divalent cation can thus be accounted as,

(A) Divalent cation of these metals possess noble gas or stable configuration.

(B) The formation of divalent cation lattice leads to evolution of energy due to strong lattice structure of divalent cation which easily compensates for the higher values of second ionisation energy of these metals.

(C) The higher heats of hydration of divalent cation which accounts for the existence of the divalent ions of these metals in solution state.

`(g)` `text(Hydration of ions)`

(i) The hydration energies of alkaline earth metals divalent cation are much more than the hydration energy of monovalent cation.

`Mg^(+)` `Mg^(2+)`

Hydration energy or Heat of hydration (kJ `text(mol)^(-1)`) `353` `1906`

The abnormally higher values of heat of hydration for divalent cations of alkaline earth metals are responsible for their divalent nature. `MgCl_2` formation occurs with more amotmt of heat evolution and thus `MgCl_2` is more stable.

(ii) The hydration energies of `M^(2+)` ion decreases with increase in ionic radii.

`text( Be^(+2) Mg^(+2) Ca^(+2) Sr^(+2) Ba^(+2) Ra^(+2))`

Heat of hydration kJ `text(mol)^(-1)` `2382` `1906` `1651` `1484` `1275`

(iii) Heat of hydration are larger than alkali metalsions and thus alkaline earth metals compounds aremore extensively hydrated than those of alkali metals e.g `MgCl_2` and `CaCl_2` exists as `Mg Cl_2 .6H_2O` and `CaCl_2. 6H_2O` which NaCl and KCI do not form such hydrates.

(iv) The ionic mobility, therefore, increases from b`Be^(2+)` to `Ba^(2+)` , as the size of hydrated ion decreases.

(h) `text(Electronegativities)`

(i) The electronegativities of alkaline earth metals are also small but are higher than alkali metals.

(ii) Electronegalivity decreases from `Be` to `Ba` as shown below,

`text( Be Mg Ca Sr Ba)`

Electronegativity `1.57` `1.31` `1.00` `0.95` `0.89`

(i) `text(Conduction power)` : Good conductor of heat and electricity.

(j) `text(Standard oxidation potential and reducing properties)` :

(i) The standard oxidation potential (in volts) are,

`text(Be Mg Ca Sr Ba)`

`text(1.69 2.35 2.87 2.89 2.90)`

(ii) All these metals possess tendency to lose two electrons to give `M^(2+)` ion and are used as reducing agent.

(iii) The reducing character increases from Be to Ba, however, these are less powerful reducing agent than alkali metals.

(iv) Beryllium having relatively lower oxidation potential and thus does not liberate `H_2` from acids.

(k) `text(Characteristic flame colours)` :

(i) Tile characteristic flame colour shown are : `Ca` brick red; `Sr` -crimson; `Ba`-apple green and `Ra` crimson.

(ii) Alkaline earth metals except `Be` and `Mg` produce charactetistic colour to flame due to easy excitation of electrons to higher energy levels.

(iii) `Be` and `Mg` atoms due to their small size, bind their electrons more strongly (because of higher effective nuclear charge) Hence these requires high excitation energy and are not excited by the energy of flame with the result that no flame colour is shown by them.

(a) `text(Occurrence)` : These are found mainly in combined state such as oxides, carbonates and sulphates `Mg` and `Ca` are found in abundance in nature. `Be` is not very abundant, `Sr` and `Ba` are less abundant. `Ra` is rare element. Some important ores of alkaline earth metals are given below,

(i) `text(Baryllium :)` Beryl `(3BeO.Ai_2O_3.6SiO_2);` Phenacite `(Be_2SiO_4)`

(ii) `text(Magnesium :)` Magnesite `(MgCO_3);` Dolomite `(CaCO_3. MgCO_3);` Epsomite `(MgSO_4. 7H_2O);` Carnallite `(MgCi_2.KC L 6H_2O)` ; Asbestos `[CaMg_3(SiO_3)_4)`

(iii) `text(Calcium :)` Limestone `(CaCO_3);` Gypsum `(CaSO_4.2H_2O),` Anhydrite `(CaSO_4);` Fluorapatite `[(3Ca_3(PO < O_2.Caf_2))` Phosphorite rock `[Ca_3(PO_4)_2]`

(iv) `text(Barium :)` Batytes `(BaSO_4)` ; witherite `(BaCO_3)`

(v) `text(Radium :)` Pitch blende `(U_3O_8);` (Ra in traces); other radium rich minerals are carnotite `(K_2UO_2))` `(V0_4)_2 8H_2O` and antamite `[Ca(UO_2)_2)`

(b) `text(Extraction of alkaline earth metals)`

(i) `Be` and `Mg` are obtained by reducing their oxides carbon,

`BeO + C -> Be + CO ; MgO + C -> Mg + CO`

(ii) The extraction of a lkaline earth metals can also be made by the reduction of their oxides by alkali metals or by electrolysing their fused salts.

(c) `text(Alloy formation)` : These dissolve in mercury and form amalgams.

(d) `text(Formation of oxides and hydroxides)`

(i) The elements (except `Ba` and `Ra`) when burnt in air give oxides of ionic nature `M^(2+)O^(2-)` which are crystalline in nature. `Ba` and `Ra` however give peroxide. The tendency to form higher oxides increases from `Be` to `Ra`.

`2M + O_2 -> 2MO` (`M` is `Be`, `Mg` or `Ca`)

`2M + O_2 -> MO_2` (`M` is `Ba` or `Sr`)

(ii) Their less reactivity than the alkali metals is evident by the fact that they are slowly oxidized on exposure to air, However the reactivity of these metals towards oxygen increases on moving down the group.

(iii) The oxides of these metals are very stable due to high lattice energy.

(iv) The oxides of the metal (except `BeO` and `MgO`) dissolve in water to form basic hydroxides and evolve a large amount of heat. `BeO` and `MgO` possess high lattice energy and thus insoluble in water.

(v) `BeO` dissolves both in acid and alkalies to give salts i.e. `BeO` possesses amphoteric nature.

`BeO + 2NaOH -> undersettext(Sod. beryllate) (Na_2BeO_2) + H_2O ; BeO + 2HCl -> undersettext(Beryllium chloride)(BeCl_2) + H_2O`

(vi)The basic nature of oxides of alkaline earth metals increases from `Be` to `Ra` as the electropositive Character increases from `Be` to `Ra`.

(vii)The tendency of these metal to react with water increases with increase in electropositive character i.e. `Be` to `Ra`.

(viii) Reaction of `Be` with water is not certain, magnesium reacts only with hot water, while other metals react with cold water but slowly and less energetically than alkali metals.

(ix) The inertness of `Be` and `Mg` towards water is due to the formation of protective, thin layer of hydroxide on the surface of the metals.

(x) The basic nature of hydroxides increase from `Be` to `Ra`. It is because of increase in ionic radius down the group which results in a decrease in strength of `M -O` bond in `M -(OH)_2` to show more dissociation of hydroxides and greater basic character.

(xi) The solubility of hydroxides of alkaline earth metals is relatively less than their corresponding alkali metal hydroxides Furthermore, the solubility of hydroxides of alkaline earth metals increases from `Be` to `Ba`. `Be (OH)_2` and `Mg (OH)_2` are almost insoluble, `Ca (OH)_2` (often called lime water) is sparingly soluble whereas `Sr(OH)_2` and `Ba (OH)_2` (often called baryta water) are more soluble.

The trend of the solubility of these hydroxides depends on the values of lattice energy and hydration energy of these hydroxides. The magnitude of hydration energy remains almost same whereas lattice energy decreases appreciably down the group leading to more `-ve` values for `Delta H` solution down the group.

`Delta H_(text(solution)) = Delta H_(text(lattice energy)) + Delta H_(text(hydration energy))` More negative is `Delta H_(text(solution))` more is solubility of compounds.

(xii) The basic character of oxides and hydroxides of alkaline earth metals is lesser than their corresponding alkali metal oxides and hydroxides.

(xiii) Aqueous solution of lime water `[Ca(OH)_2]` or baryta water `[Ba(OH)_2]` are used to qualitative identification and quantative estimation of carbon dioxide, as both of them gives white precipitate with `CO_2` due to formation of insoluble `CaCO_3` or `BaCO_3`

`Ca(OH)_2 + CO_2 -> undersettext(white ppt)(CaCO_3) + H_2O ; Ba(OH)_2 + CO_2 -> undersettext(white ppt)(BaCO_3) + H_2O`

`Note` : q `SO_2` also give white ppt of `CaSO_3` and `BaSO_3` on passing through lime water or baryta water However on passing `CO_2` in excess, the white turbidity of insoluble carbonates dissolve to give a clear solution again due to the formation of soluble
bicarbonates,

(e) `text(Hydrides)` :

(i) Except `Be`, all alkaline earth metals form hydrides `(MH_2)` on heating directly with `H_2. M^(+) H_2 -> MH_2`.

(ii) `BeH_2` is prepared by the action of `Li A I H_4` On `BeCl_2 ; 2BeCl_2 + LiAlH_4 -> 2BeH_2 + LiCl + AlCl_3.`

(iii) `BeH_2` and `MgH_2` are covalent while other hydrides are ionic.

(iv) The ionic hydrides of `Ca`, `Sr`, `Ba` liberate `H_2` at anode and metal at cathode.

`CaH_2 oversettext(fusion)⇋ Ca^(2+) + 2H^(-)`

Anode : `2H^(-) -> H_2+ 2e` Cathode : `Ca^(2+) + 2e -> Ca`

(v) The stability of hydtides decreases from `Be` to `Ba`.

(vi) The hydrides having higher reactivity for water, dissolves readily and produce hydrogen gas.

`CaH_(2(s)) + 2H_2O -> Ca(OH)_2 + 2H_2`

`(f)` `text(Carbonates and Bicarbonates)`

(i) All these metal carbonates `(MCO_3)` are insoluble in neutral medium but soluble in acid medium. These are precipitated by the addition of alkali metal or ammonium carbonate solution to the solution of these metals.

`(NH_4)_2 CO_3 + CaCl_2 -> 2NH_4 Cl + CaCO_3 ; Na_2CO_3 + BaClz -> 2NaCl + BaCO_3`

(ii) Alkaline earth metal carbonates are obtained as white precipitates when calculated amount of carbon dioxide is passed through the solution of the alkaline metal hydroxides.

`M(OH)_2 (aq) + CO_2(g) -> MCO_3(s) + H_2O(l)`

and sodium or ammonium carbonate is added to the solution of the alkaline earth metal salt such as `CaCl_2`.

`CaCl_2 (aq) + Na_2CO_3 (aq) -> CaCO_3(s) +2 NaCl(aq)`

(iii) Solubility of carbonates of these metals also decreases downward in the group due to the decrease of hydration energy as the lattice energy remains almost unchanged as in case of sulphates.

(vi) The carbonates of these metals decompose on heating to give the oxides, the temperature of decomposition increasing from Re to Ra. Reryllium carbonate is unstable.

(g) `text(Halides)` :

(i) The alkaline earth metals combine directly with halogens at appropriate temperatures forming halides, `MX_2` . These halides can also be prepared by the action of halogen acids `(HX)` on metals, metal oxides, hydroxides and carbonates.

`M + 2HX -> MX__2 + Hz ; MO + 2HX -> MX_2 + H_2O`

`M(OH)_2 + 2HX -> MX_2 +2H_20 ; MCO_3 + 2HX -> MX_2 + CO_2 + H_2O`

Beryllium chloride is however, conveniently obtained from oxide

(ii) `BeCl_2` is essentially covalent, the chlorides `MgCl_2, CaCl_2 , SrCl_2` and `BaCl_2` are ionic; the ionic character increases as the sizeof the metal ion increases. The evidence is provided by the following facts,

(A) Beryllium chloride is relatively low melting and volatile whereas `BaCl_2` has high melting and stable.

(B) Beryllium chloride is soluble in organic solvents.

(iii) The halides of the members of this group are soluble in water and produce neutral solutions from which the hydrates such : `MgCl_2 6H_2O`, `CaCl_2*6H_2O, BaCl_2*2H_2O` can be crystallised. The tendency to form hydrated halides decreases with increasing size of the metal ions.

(iv) `BeCl_2` is readily hydrolysed with water to form acid solution, `BeCl_2 + 2H_2O -> Be (OH)_2 + 2HCl`.

(v) The fluorides are relatively less soluble than the chlorides due to high lattice energies. Except `BeCl_2` and `MgCl_2` the chlorides of alkaline earth metals impart characteristic colours to flame.

`CaCl_2` `SrCl_2` `BaCl_2`

Brick red colour Crimson colour Grassy green colour

Structure of `BeCl_2` In the solid phase polymeric chain structure with three centre 2 electron bonding with `Be-Cl-Be` bridged structure.

In the vapour phase it tends to form a chlorobridged dimer which dissociates into the linear triatomic monomer at high temperature at nearly `1200` `K`.

(h) `text(Solubility in liquid ammonia :)` Like alkali metals, alkaline earth metals also dissolve in liquid ammonia to form coloured solutions When such a solution is evaporated, hexammoniate, `M(NH3)_6` is formed.

(i) `text(Nitrides)`

(i) All the alkaline earth metals directs combine with `N_2` give nitrides, `M_3N_2`.

(ii) The ease of formation of nitrides however decreases from `Be` to `Ba`.

(iii) These nitrides are hydrolysed water to liberate `NH_3 + M_3N_2 + 6H_2O -> 3M(OH)_2 + 2NH_3`

(j) `text(Sulphates)` :

(i) All these form sulphate of the type `M SO_4` by the action of `H_2SO_4` on metals, their oxides, carbonates or hydroxides.

`M + H_2SO_4 -> MSO_4 + H_2 ; MO + H_2SO_4 -> MSO_4 + H_2O`

`MCO_3+ H_2SO_4 -> MSO_4 + H_2O+CO_2 ; M(OH)_2 + H_2SO_4 -> MSO_4 + 2H_2O`

(ii) The solubility of sulphates in water decreases on moving down the group `BeSO_4` and `MgSO_4` are fairly soluble in water while `BaSO_4` is completely insoluble. This is due to increases in lattice energy of sulphates down the group which predominates over hydration energy.

(ii) Sulphate are quite stable to heat however reduced to sulphide on heating with carbon.

`MSO_4 + 2C -> MS+2CO_2`

(k) `text(Action with carbon :)` Alkaline metals (except Be, Mg) when heated with carbon form carbides of the type `MC_2` These carbides are also called acetylides as on hydrolysis they evolve acetylene.

`MC_2 + 2H_2O -> M(OH)_2 + C_2H_2`

(l) `text(Action with sulphur and phosphorus :)`

Alkaline earth metals directly combine with sulphur and phosphorus when heated to form sulphides of the type `MS` and phosphides of the type `M_3P_2` respectively.

`M + S -> MS ; 3M + 2P -> M_3P_2`

Sulphides on hydrolysis liberate `H_2S` while phosphides on hydrolysis evolve phosphine.

`MS + text(dil acid) -> H_2S; M_3P_2 + text(dil acid) -> PH_3`

Sulphides are phosphorescent and are decomposed by water

`2MS + 2H_20 -> M(OH)_2 + M(HS)_2`

(m) `text(Nitrates :)` Nitrates of these metals are soluble in water. On heating they decompose into their corresponding oxides with evolution of a mixture of nitrogen dioxide and oxygen.

(n) `text(Formation of complexes)` :

(i) Tendency to show complex ion formation depends upon smaller size, high nuclear charge and vacant orbitals to accept electron. Since alkaline metals too do not possess these characteristics and thus are unable to form complex ion.

(ii) However, `Be^(2+)` on account of smaller size forms many complex such as `(BeF_3)^(-1); (BeF_4)^(2-)`