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.
(1) Electronic configuration
Radium was discovered in the ore pitch blende by madam Curie. It is radioactive in nature.
(2) 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) Baryllium : Beryl (3BeO.Al2O3.6SiO2); Phenacite (Be2SiO4)
(ii) Magnesium : Magnesite (MgCO3); Dolomite (CaCO3. MgCO3);
Epsomite(MgSO4. 7H2O); Carnallite (MgCl2.KCl. 6H2O); Asbestos [CaMg3(SiO3)4]
(iii) Calcium : Limestone (CaCO3); Gypsum : (CaSO4.2H2O), Anhydrite (CaSO4); Fluorapatite [(3Ca3(PO4)2.CaF2)] Phosphorite rock [Ca3(PO4)2]
(iv) Barium : Barytes (BaSO4) ; witherite (BaCO3)
(v) Radium : Pitch blende (U3O8); (Ra in traces); other radium rich minerals are carnotite [K2UO2)] (VO4)2 8H2O and antamite [Ca(UO2)2]
(3) 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 alkaline earth metals can also be made by the reduction of their oxides by alkali metals or by electrolysing their fused salts.
(4) Alloy formation : These dissolve in mercury and form amalgams.
(1) Physical state : All are greyish‐white, light, malleable and ductile metals with metallic lustre. Their hardness progressively decrease with increase in atomic number. Although these are fairly soft but relatively harder than alkali metals.
(2) Atomic and ionic radii
(i) The atomic and ionic radii of alkaline earth metals also increase down the group due to progressive addition of new energy shells like alkali metals.
Be Mg Ca Sr Ba Ra
Atomic radius (pm) 112 160 197 215 222 –
Ionic radius of M2+ 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.
(i) Density decreases slightly upto 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.
(4) Melting point and Boiling point
(i) Melting points and boiling points of alkaline earth metals do not show any regular trend.
Be Mg Ca Sr Ba Ra
melting points (K) 1560 920 1112 1041 1000 973
boiling point (K) 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.
(5) 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.
Be Mg Ca Sr Ba Ra
First ionisation energy (kJ mol‐1) 899 737 590 549 503 509
Second ionisation energy (kJ mol‐1) 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
1st ionisation energy (kJ mol–1) 520 899
2nd ionisation energy (kJ mol–1) 7296 1757
This may be explained as,
The removal of 2nd electron from alkali metals takes place from 1s sub shell which are more closer to nucleus and exert more nuclear charge to hold up 1s 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.
(6) Oxidation number and valency
(i) The IE1 of the these metals are much lower than IE1 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 M2+ ion possesses a higher degree of hydration or M2+ ions are extensively hydrated to form [M(H2O)x]2+, a hydrated ion. This involves a large amount of energy evolution which counter balances the higher value of second ionisation energy.
M → M2+ , ΔH = IE1 + E2
M2++ + xH2O → [M(H2O)x]2+; ΔH = – 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.
(7) Hydration of ions
(i) The hydration energies of alkaline earth metals divalent cation are much more than the hydration energy of monovalent cation.
Hydration energy or Heat of hydration (kJ 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. MgCl2 formation occurs with more amount of heat evolution and thus MgCl2 is more stable.
(ii) The hydration energies of M2+ ion decreases with increase in ionic radii.
Be2+ Mg2+ Ca2+ Sr2+ Ba2+
Heat of hydration kJ mol–1 2382 1906 1651 1484 1275
(iii) Heat of hydration are larger than alkali metals ions and thus alkaline earth metals compounds are more extensively hydrated than those of alkali metals e.g MgCl2 and CaCl2 exists as Mg Cl2 .6H2O and CaCl2. 6H2O which NaCl and KCl do not form such hydrates.
(iv) The ionic mobility, therefore, increases from Be2+ to Ba2+, as the size of hydrated ion decreases.
(i) The electronegativities of alkaline earth metals are also small but are higher than alkali metals.
(ii) Electronegativity decreases from Be to Ba as shown below,
Be Mg Ca Sr Ba
Electronegativity 1.57 1.31 1.00 0.95 0.89
(9) Conduction power : Good conductor of heat and electricity.
(10) Standard oxidation potential and reducing properties
(i) The standard oxidation potential (in volts) are,
Be Mg Ca Sr Ba
1.69 2.35 2.87 2.89 2.90
(ii) All these metals possess tendency to lose two electrons to give M2+ 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 H2 from acids.
(11) Characteristic flame colours
The characteristic flame colour shown are : Ca ‐ brick red; Sr –crimson ; Ba‐apple green and Racrimson.
(1) Formation of oxides and hydroxides
(i) The elements (except Ba and Ra) when burnt in air give oxides of ionic nature M2+O2‐ which are crystalline in nature. Ba and Ra however give peroxide. The tendency to form higher oxides increases from Be to Ra.
2M + O2 → 2MO (M is Be, Mg or Ca )
2M + O2 → MO2 (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 → Na2BeO2 + H2O ; BeO + 2HCl → BeCl2 + H2O
Sod. beryllate Beryllium chloride
(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
ΔH solution down the group.
ΔH solution = ΔH lattice energy + ΔH hydration energy
More negative is ΔH 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 CO2 due to formation of insoluble CaCO3 or BaCO3
Ca(OH)2 + CO2 → CaCO3 + H2O ; Ba(OH)2 + CO2 → BaCO3 + H2O
(white ppt) (white ppt)
SO2 also give white ppt of CaSO3 and BaSO3 on passing through lime water or baryta water. However on passing CO2 in excess, the white turbidity of insoluble carbonates dissolve to give a clear solution again due to the formation of soluble bicarbonates,
CaCO3 → H2O + CO2 → Ca(HCO3)2
(i) Except Be, all alkaline earth metals form hydrides (MH2) on heating directly with H2.
M+ H2 → MH2.
(ii) BeH2 is prepared by the action of LiAlH4 On BeCl2
2BeCl2 + LiAlH4 → 2BeH2 + LiCl + AlCl3.
(iii) BeH2 and MgH2 are covalent while other hydrides are ionic.
(iv) The ionic hydrides of Ca, Sr, Ba liberate H2 at anode and metal at cathode.
Anode : 2H– → H2 + 2e– Cathode : Ca2+ + 2e– →Ca
(v) The stability of hydrides decreases from Be to Ba.
(vi) The hydrides having higher reactivity for water, dissolves readily and produce hydrogen gas.
CaH2(s) + 2H2O → Ca(OH)2 + 2H2↑
(3) Carbonates and Bicarbonates
(i) All these metal carbonates (MCO3) 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.
(NH4)2 CO3 + CaCl2 → 2NH4Cl + CaCO3
Na2CO3 + BaCl2 → 2NaCl + BaCO3
(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) + CO2 (g) → MCO3(s) + H2O(l)
and sodium or ammonium carbonate is added to the solution of the alkaline earth metal salt such as CaCl2.
CaCl2 (aq) + Na2CO3 (aq) → CaCO3(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.
(iv) The carbonates of these metals decompose on heating to give the oxides, the temperature of decomposition increasing from Be to Ba. Beryllium carbonate is unstable.