Group 2 Elements: Alkaline Earth Metals

Alkaline Earth Metals

The group 2 of the periodic table consists of six metallic elements. They are Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium (Ra). The name alkaline earth metals was given to magnesium, calcium, barium & strontium since their oxides were alkaline in nature and these oxide remained unaffected by heat or fire and existed in earth.

Occurrence

Like alkali metals, alkaline earth metals are also highly reactive and hence do not occur in the free state but are likely distributed in nature in the combined state as silicates, carbonates, sulphates and phosphates.

Minerals

Be – Beryl (Be₃Al₂Si₆O_­18) & Phenacite (Be₂SiO₄)

Mg – Magnesite MgCO₃, Dolomite CaMg(CO₃)₂, Epsomite MgSO₄.&H₂O

Ca – Limestone (CaCO₃), fluoropatite [3(Ca₃(PO₄)₃.CaF₂], Gypsum (CaSO₄.2H₂O), Anhydrite (CaSO₄)

Sr – Celestite (SrSO₄), Strontianite (SrCO₃)

Br – Barytes (BaSO₄)

Electronic Configuration

The general electronic configuration of alkaline earth metals is ns².

Be – 1s²2s² Mg – 1s²2s²sp⁶3s²

Ca – 1s²2s²2p⁶3s²3p⁶4s² Sr – [Kr]5s²

Ba – [Xe]6s² Ra – [Rn]7s²

Physical Properties of Group II elements

(i) Atomic and ionic radii

The atomic radii as well as ionic radii of the members of the family are smaller than the corresponding members of alkali metals.

(ii) Ionization energy

The alkaline earth metal owing to their large size of atoms have fairly low values of ionization energies as compared to the p – block elements. However with in the group, the ionization energy decreases as the atomic number increases. It is because of increase in atomic size due to addition of new shells and increase in the magnitude of screening effect of the electrons in inner shells. Because their (IE)₁ is larger than that of their alkali metal neighbours, the group IIA metals trend to the some what less reactive than alkali metals. The general reactivity trend is

Ba > Sr > Ca > Mg > Be.

Illustration 1. The 2^nd ionization energies of the elements of group I are higher than those of the elements of group II. Explain.

Solution: The 2^nd electron in case of alkali metal is to be removed form a cation (unipostive ion) which has already acquired a noble gas configuration whereas in case of alkaline earth metals, the second electron is to be removes fro a cation which is yet to acquire the stable noble gas configuration therefore, removal of 2^nd electron in case of alkaline earth metals requires much less energy than that in case of alkali metals.

There is sharp increase in third ionization energy due to stable inert gas configuration of m⁺² ions. This explains the upper limit of +2 oxidation state for the elements.

(iii) Oxidation state

The alkaline earth metal have two electrons in their valence shell and by losing these electrons, these atoms acquire the stable noble gas configuration. Thus, unlike alkali metals, the alkaline earth metals exhibit +2 oxidation state in their compounds.

$M\to \underset{\left[noblegas\right]}{M^{+2}}+2e^{-}$

Illustration 2. The alkaline earth metals shows +2 oxidation state i.e. they always form divalent cations (M²⁺). Explain.

Solution: If ionization energy were the only factor involved, than group II elements should have formed monovalent ions i.e. Mg⁺, Ca⁺ etc rather than Mg²⁺, Ca⁺² etc.

This can be explained as follows:

(i) The divalent cations of alkaline earth metals acquires stable inert gas configuration.

(ii) The divalent cations results in stronger lattices then monovalent cations and hence a lot of energy called lattice energy released during formation of divalent cations than monovalent cation which compensates the high second ionization energy.

(iii) The existence of divalent ions in the aqueous solution is due to greater hydration of the divalent ions which counter balance the high value of second ionization energy.

The heat of hydration (hydration energy) of alkaline earth metals are approximately four times higher than alkali metals of comparable size.

e.g.

$\Delta$H_hyd for Na⁺ (size 102 pm) = -397 KJmol^-1

$\Delta$H_hyd for Ca⁺² (size 100 pm) = -1650 KJmol^-1

Larger hydration energy is due to the fact that the alkaline earth metals ions, because of their much larger charge to size ratio, exert a much stronger electrostatic attraction on the oxygen of water molecule.

Lattice energies decreases as atomic number increases

(v) Nature of metallic bonding in alkaline earth metals

The alkali metal two electrons are involved in the metallic bonding. Moreover, sizes of alkaline earth metal ions are smaller than those of alkali metal ions. Consequently, stronger metallic bonds are formed which result in the close packing of the atoms. Due to the presence of stronger metallic bonds, alkaline earth metals have

(a) Higher melting points (b) Higher boiling points

(c) higher densities (d) Harder than the corresponding alkali metals.

(vi) Density

The alkaline earth metals are denser and harder than the corresponding alkali metals.

The atoms of alkaline earth metals have smaller size and are hence held by stronger metallic bonds, as compared to alkali metals. Therefore, they are more closely packed in their crystal lattice which accounts for high density and increased hardness of these elements.

(vii) Characteristic flame colouration

Expect Be & Mg (due to high ionization energy), the alkaline earth metals impart characteristic colour when introduced into flame of a burner. This property is due to the ease of excitation of their valence electrons. When elements or their compounds are introduced to flame, the electron absorbs energy from the flame and gets excited to higher energy levels. When these electrons return to their ground state, they emit absorbed energy in form of visible light having characteristic wavelengths. Depending upon the wavelength of light emitted, different colours are impart to the flame. Salts (generally chlorides) impart characteristic colours to the Bunsen flame.

Illustration 3. Mg forms Mg²⁺, but Na²⁺ does not exist. Explain.

Solution: Na metal after the loss of one electron attains a noble gas configuration of neon. Therefore, the removal of second electron is energetically unfavourable. Hence, Na²⁺ does not exist.

Electropositive or Metallic Character

The alkaline earth metals are highly electropositive and hence metallic and their electropositive or metallic character increases down the group. However they are less electropositive or metallic than the alkali metals. It is due to smaller size and higher ionization energies as compared to alkali metals, hence have less tendency to loose electron than those of alkali metals (group I)

Like the alkali metals they also form predominantly ionic compounds but tendency of covalency is greater, particularly with Be and Mg because of their smaller atomic and ionic radii. Be forms compounds which are essentially covalent.

Melting and boiling points

The alkaline earth metals have higher melting and boiling points as compared to those of alkali metals which is attributed to their small size and more close packed crystal lattice as compared to alkali metals and presence of two valence electrons.

Heat of Hydration

$\bullet$The heats of hydration of M²⁺ decreases with an increase in their ionic size and their values are greater than that of alkali metal ions.

$\bullet$Alkaline earth metal ions, because of their larger charge to size ratio, exert a much stronger electrostatic attraction on the oxygen of water molecule surrounding them.

$\bullet$Since the alkaline earth metals (except Be) tend to lose their valence electrons readily, they act as strong reducing agents as indicated by E⁰red values. The particularly less negative value for Be arises from the large hydration energy associated with the small size of Be²⁺ and the relatively large value of heat of sublimation.

Solubility

$\bullet$Basic nature of oxides increases down the group but solubilities of sulphates and carbonates decrease as ionic size increases.

$\bullet$The solubility of most salts decreases with increased atomic weight, though usual trend is reversed with fluorides and hydroxides in this group.

Physical Properties of groups 2 elements (alkaline earth metals)

* ppm (parts per million) ** Percentage by weight

Reactivity and Electrode potential

All the alkaline earth metals are highly reactive elements since they have a strong tendency to lose the two valences s-electrons to form the corresponding dipositive ions having inert gas configuration. The high reactivity arises due to their low ionization energies and high negative values of their standard electrode potentials. Further, the chemical reactivity of alkaline earth metals increase on moving down the group because the I.E. decreases and electrode potentials become more and more negative with increasing atomic number from Be to Ra. Thus, beryllium is the least reactive while Ba (or Ra) is the most reactive element. Further since the ionization energies of alkaline earth metals are higher and their electrode potential is less negative than the corresponding alkali metals. They are less reactive than corresponding alkali metals.

Reducing Character

The alkaline earth metals are weaker reducing agents than the alkali metals. Like alkali metals, their reducing character also increases down the group. This is due to the reason that the alkaline earth metals have greater tendency to lose electrons so, they act as reducing agent but since their I.E. are higher and their electrode potentials are less negative than the corresponding alkali metals, therefore alkaline earth metals are weaker reducing agents than alkali metals. The sulphates are stable to heat whereas the carbonates decompose to give MO and CO₂, the temperature of decomposition increasing from Mg to Ba. BeCO₃ is kept in the atmosphere of CO₂ to prevent its decomposition.

BeCO₃ MgCO₃ CaCO₃ SrCO₃ BaCO₃

<100°C 540°C 900°C 1290°C 1360°C

Occurrence and uses of alkaline earth metals

Group IIA (Alkaline earth metals) and groups IIB (Zn, Cd, Hg) Mg acts as a bridge element between IIA and IIB.

Difference between alkaline earth metals and alkali metals

Both alkaline earth metals and alkali metals are s – block elements as the last electron enters the ns – orbital. They resemble with each other in some respects but still there are certain dissimilarities in their properties on account of different number of electrons in the valency shell, smaller atomic radii, high ionization potential, higher electro negativity etc.

Chemical Properties of Group – II elements

Reaction with water – (Formation of hydroxides)

The electrode potential of Be (Be²⁺/Be = -1.97 V) is least negative amongst all the alkaline earth metals. This means that Be is much less electropositive than other alkaline earth metals and hence does not react with water or steam even at red heat.

The electrode potential of Mg (Mg⁺²/Mg = -2.37 V), although more negative than that of Be yet is still less negative than those of alkali metals and hence it does not react with cold water but reacts with boiling water or steam.

$Mg+H_{2}O\to MgO+H_2$

or, $Mg+2H_{2}O\to Mg{\left(OH\right)}_2+H_2$

Mg, infact, forms a protective layer of oxide on its surface, therefore, despite its favourable electrode potential it does not react readily with water unless the oxide layer is removed by amalgamating it with mercury. In the formation of oxide film, Mg resembles Al.

Ca, Sr and Ba have more negative electrode potentials similar to those of the corresponding group I alkali metals and hence react with even with cold water, liberating H₂ and forming the corresponding metal hydroxides.

Ca + 2H₂O $\to$ Ca(OH)₂ + H₂

Reactivity of alkaline earth metals increases as we move down the group. However, the reaction of alkaline earth metals is less vigorous as compared to alkali metals.

Reaction with air (Nitrogen and Oxygen)

(a) Formation of oxides and nitrides

Be metal is relatively unreactive in the massive form and hence does not react below 873K. However, powdered Be is more reactive and burns brilliantly on ignition to give a mixture of BeO & Be₃N₂.

$2Be+O_2\left(air\right)\stackrel{\Delta }{\to }2BeO$

$3Be+N_2\left(air\right)\stackrel{\Delta }{\to }B{e}_3{N}_2$

magnesium is more electropositive than Be and hence burns with dazzling brilliance in air to form a mixture of MgO and magnesium nitride.

$Mg+air\stackrel{\Delta }{\to }MgO+M{g}_3{N}_2$

Ca, Sr and Ba being even more electropositive react with air readily to form a mixture of their respective oxides and nitrides.

The reactivity towards oxygen increases as we go down the group. Thus Ca, Ba and Sr are stored in paraffin but Be and Mg are not because they form protective oxide layer on their surface.

(b) Formation of Nitrides

All the alkaline metals burn in dinitrogen to form ionic nitrides of the formula, M₃N₂. This is in contrast to alkali metals where only Li forms Li₃N.

3M + N₂ $\stackrel{\Delta }{\to }{M}_3{N}_2$

Be₃N₂ being covalent is volatile while the nitrides of all other elements are crystalline solids.

All these nitrides decompose on heating and react with water liberating NH₃.

$B{e}_3{N}_2\stackrel{\Delta }{\to }3\mathbf{B}e+N_2$

$B{a}_3{N}_2+6H_{2}O\stackrel{\Delta }{\to }3Ba{\left(OH\right)}_2+2N{H}_3$

$C{a}_3{N}_2+6H_{2}O\stackrel{\Delta }{\to }3Ca{\left(OH\right)}_2+2N{H}_3$

(c) Formation of Peroxides

Since larger cations stabilize larger anions. Therefore, tendency to form peroxide increases as the size of the metal ion becomes larger. Thus BaO₂ is formed by passing air over heated BaO at 773K.

$2BaO+O_2\stackrel{773K}{\to }2\mathbf{B}a{O}_2$

$2SrO+O_2\underset{Highpressrue}{\overset{773K}{\to }}2Sr{O}_2$

SrO₂ is prepared in similar way but under high pressure and temperature. CaO₂ is not formed this way but can be prepared as the hydrate by treating Ca(OH)₂ with H₂O₂ and then dehydrating the product.

$Ca{\left(OH\right)}_2+H_2{O}_2\to Ca{O}_2.2H_{2}O$

Crude MgO₂ has been made using H₂O₂ but peroxide of beryllium is not known.

All peroxide are white crystalline ionic solids containing the peroxide ion$O_2^{2-}$. Treatment of peroxide with acids liberates H₂O₂.

$Ba{O}_2+2HCl\to \mathbf{B}aC{l}_2+H_2{O}_2$

Reaction with hydrogen – (Formation of hydrides)

All the alkaline earth metals except Be combine with hydrogen directly on heating to form metal hydrides of formula MH₂.

$M+H_2\stackrel{\Delta }{\to }M{H}_2$

The hydride of beryllium can also be obtained by the reduction of BeCl₂ with LiAlH₄.

$2BeC{l}_2+LiAl{H}_4\to 2Be{H}_2+LiCl+AlC{l}_3$

Both BeH₂ and MgH₂ are covalent compounds having polymeric structures in which H – atoms between beryllium atoms are held together by three centre – two electron (3C - 2e) bonds as shown below:

The hydrides of other elements of this groups i.e. CaH₂, SrH₂ and BaH₂ are ionic and contain the H^- ions.

All the hydrides of alkaline earth metals reacts with water liberating H₂ gas and thus act as reducing agents.

$M{H}_2+2H_{2}O\to M{\left(OH\right)}_2+2H_2$

CaH₂ is called Hydrolith and is used for production of H₂ by action of water on it.

Reaction with carbon – (Formation of carbides)

When BeO is heated with carbon at 2175 – 2275 K a brick red coloured carbide of the formula Be₂C is formed

$2BeO+2C\stackrel{2175-2275K}{\to }B{e}_{2}C+2CO$.

It is a covalent compound and react water forming methane.

$B{e}_{2}C+4H_{2}O\to 2Be{\left(OH\right)}_2+C{H}_4$

The rest of the alkaline earth metals (Mg, Ca, Sr & Ba) form carbides of the general formula, MC₂ either when the metal is heated with carbon in an electric furnace or when their oxides are heated with carbon.

$Ca+2C\stackrel{1375K}{\to }Ca{C}_2$

$CaO+3C\stackrel{2275K}{\to }Ca{C}_2+CO$

All these carbides react with water producing acetylene gas.

$Ca{C}_2+2H_{2}O\to HC\equiv CH+Ca{\left(OH\right)}_2$

Reaction with Halogens

The alkaline earth metals react with halogens at elevated temperature to form the halides of the types MX₂.

Action of Acids

The alkaline earth metals readily react with acids liberating hydrogen.

$M+2HCl\to MC{l}_2+H_2\left(M=Be,Mg,Ca,Sr,Ba\right)$

Reaction with Ammonia

Like alkali metal, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solution from which ammoniates ${\left[M{\left(N{H}_3\right)}_6\right]}^{2+}$ can be recovered.

Illustration 4. How does the basicity of oxides of group 2 increases down the group?

Solution: The basicity increases down the group

\begin{gathered} \,\,\,\,{\text{BeO < MgO < CaO < SrO < BaO}} \hfill \\ {\text{amphoteric}}\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,{\text{strongly - basic}} \hfill \\ \end{gathered}

General characteristics of compounds of the Alkaline earth metals

(a) Oxides

The oxides MO are obtained either by heating the metals in oxygen or by thermal decomposition of their carbonates.

$2M+O_2\stackrel{\Delta }{\to }2MO\left(M=Be,Mg,Ca\right)$

$MC{O}_3\stackrel{\Delta }{\to }MO+C{O}_2\left(M=Be,Mg,Ca,Sr,Ba\right)$

Expect BeO all other oxides are extremely stable ionic solids due to their high lattice energies.

These have high melting point, have very low vapour pressure, are very good conducts of heat, are chemically inert and act as electrical insulators. Therefore, these oxides are used for lining furnaces and hence used as refractory materials.

Due to small size of beryllium ion, BeO is covalent but still has high melting point because of its polymeric nature.

(b) Hydroxides

The hydroxides of Ca, Sr & Ba are obtained either by treating the metal with cold water or by reacting the corresponding oxides with water. The reaction of these oxides with H₂O is also sometimes called as slaking.

$M+2H_{2}O\to M{\left(OH\right)}_2+H_2\left(M=Ca,Sr,Ba\right)$

$MO+H_{2}O\to M{\left(OH\right)}_2\left(M=Ca,Sr,Ba\right)$

Be(OH)₂ and Mg(OH)₂ being insoluble are obtained from suitable metal ion solutions by precipitation with OH^- ions.

$BeC{l}_2+2NaOH\to Be{\left(OH\right)}_2\downarrow +2NaCl$

$MgS{O}_4+2NaOH\to Mg{\left(OH\right)}_2\downarrow +N{a}_{2}S{O}_4$

Properties

(i) Basic Character

All the alkaline earth metal hydroxides are bases except Be(OH)₂ which is amphoteric. This basic strength increases as we move down the group. This is because of increase in size which results in decrease of ionization energy which weakens the strength of M – O bonds in MOH and thus increase the basic strength. However, these hydroxides are less basic than the corresponding alkali metal hydroxides because of higher ionization energies, smaller ionic sizes and greater lattice energies.

(ii) Solubility in Water

Alkaline earth metals hydroxides are less soluble in water as compared to alkali metals.

The solubility of the alkaline earth metal hydroxides in water increases with increase in atomic number down the group. This is due to the fact that the lattice energy decreases down the group due to increase in size of the alkaline earth metals cation whereas the hydration energy of the cation remains almost unchanged. The resultant of two effects i.e.

$\Delta H_{solution}=\Delta H_{lattice}-\Delta H_{Hydration}$

becomes more negative as we move from Be(OH)₂ to Ba(OH)₂ which accounts for increase in solubility.

Halides

The alkaline earth metals combine directly with halogen at appropriate temperature forming halides MX₂.

These halides can also be prepared by the action of halogen acids (HX) on metals, metals oxides, hydroxides and carbonates.

$M+2HX\to M{X}_2+H_2$

$MO+2HX\to M{X}_2+H_{2}O$

$M{\left(OH\right)}_2+2HX\to M{X}_2+2H_{2}O$

$MC{O}_3+2HX\to M{X}_2+H_{2}O+C{O}_2$

Properties

(1) All beryllium halides are essentially covalent and are soluble in organic solvents. They are hydroscopic and fume in air due to hydrolysis. On hydrolysis, they produce acidic solution.

$BeC{l}_2+2H_{2}O\to Be{\left(OH\right)}_2+2HCl$

(2) The halides of all other alkaline earth metals are ionic. Their ionic character, however increases as the size of the metal ion increase.

(3) Except BeCl₂ all other chlorides of group 2 form hydrates but their tendency to form hydrates decreases for eg – MgCl₂.6H₂O, CaCl₂.6H₂O.

(4) The hydrated chloride, bromides and iodides of Ca, Sr and Ba can be dehydrated on heating but those of Be and Mg undergo hydrolysis.

(5) BeF₂ is very soluble in water due to the high hydration energy of the small Be⁺² ion. The other fluorides (MgF₂, CaF₂, SrF₂ and BaF₂) are almost insoluble in water. Since on descending the group lattice energy decreases more rapidly than the hydration energy. Therefore whatever little solubility these fluorides have that increase down the group.

The chlorides, bromides and iodides of all other elements i.e. Mg, Ca, Sr, Ba are ionic have much lower melting points than the fluorides and are readily soluble in water. The solubility decreases some what with increasing atomic number.

(6) Except of BeCl₂ and MgCl₂, the other chlorides of alkaline earth metals impart characteristics colour to flame.

CaCl₂ = Brick red colour

SrCl₂ = Crimson colour

BaCl₂ = Grassy green colour

Uses

(i) Calcium fluoride or fluorospar (CaF₂) is by far the most important of all the fluorides of the alkaline earth metals since it is the only large scale source of fluorine.

(ii) CaCl₂ is widely used for melting ice on roads, particularly in very cold countries because 30% eutectic mixture of CaCl₂/ice freezes at 218 K as compared to NaCl /ice at 255K.

(iii) CaCl₂ is also used as a desiccant (drying agent) in the laboratory.

(iv) Anhydrous MgCl₂ is used in the electrolytic extraction of magnesium.

Solubility and Thermal stability of oxo salts

The salts containing one or more atoms of oxygen such as oxides, hydroxides, carbonates, bicarbonates, nitrites, nitrates, sulphates, oxalates and phosphates are called oxo salts.

Sulphates

The sulphates of alkaline earth metals (MSO₄) are prepared by the action of sulphuric acid on metals, metals oxides, hydroxides and carbonates.

$M+H_{2}S{O}_4\to MS{O}_4+H_2$

$MO+H_{2}S{O}_4\to MS{O}_4+H_{2}O$

$M{\left(OH\right)}_2+H_{2}S{O}_4\to MS{O}_4+2H_{2}O$

$MC{O}_3+H_{2}S{O}_4\to MS{O}_4+C{O}_2+H_{2}O$

Properties of sulphates

The sulphates of alkaline earth metals are all white solids.

(a) Solubility

The solubility of the sulphates in water decreases down the groups i.e. Be > Mg > Ca > Sr > Ba.

Thus BeSO₄ and MgSO₄ are highly soluble, CaSO_­4 is sparingly soluble but the sulphates of Sr, Ba and Ra are virtually insoluble.

Reason

The magnitude of the lattice energy remains almost constant as the sulphate is so big that small increase in the size of the cation from Be to Ba does not make any difference. However the hydration energy decreases from Be⁺² to Ba⁺² appreciably as the size of the cation increase down the group. Hence, the solubilities of sulphates of alkaline earth metals decrease down the group mainly due to the decreasing hydration energies from Be⁺² to Ba⁺². The high solubility of BeSo₄ and MgSO₄ is due to high hydration energies due to smaller Be⁺² and Mg⁺² ions.

(b) Stability

The sulphates of alkaline earth metal decomposes on heating giving the oxides and SO₃.

$MS{O}_4\stackrel{\Delta }{\to }MO+S{O}_3$

The temperature of decomposition of these sulpahtes increases as the basicity of the hydroxide of the corresponding metal increase down the group

Carbonates and Bicarbonates

Alkaline earth metal carbonates are obtained as white precipitates when.

(i) Calculated amount of carbon dioxide is passed through the solution of the alkaline metal hydroxides.

$M{\left(OH\right)}_2\left(aq\right)+C{O}_2\left(g\right)\to MC{O}_3\left(s\right)+H_{2}O\left(\ell \right)$

(ii) Sodium or ammonium carbonate is added to the solution of the alkaline earth metal salt such as CaCl₂.

$CaC{l}_2\left(aq\right)+N{a}_{2}C{O}_3\left(aq\right)\to CaC{O}_3\left(s\right)+2NaCl\left(aq\right)$

Properties

All carbonates are stable but beryllium carbonate is prone to hydrolysis. It contains the hydrated ion ${\left[Be{\left(H_{2}O\right)}_4\right]}^{2+}$ rather than Be⁺² and hence is precipitated only in an atmosphere of CO₂.

$BeC{O}_3+4H_{2}O\underset{presenceofC{o}_2}{\stackrel{AbsenceofC{O}_2}{\rightleftharpoons }}{\left[Be{\left(H_{2}O\right)}_4\right]}^{2+}+C{O}_3^{-2}$

Solubility

The carbonates of magnesium and other alkali earth metals are sparingly soluble is water and their solubility decreases down the group from Be to Ba. For e.g MgCO₃ is slightly soluble in water but BaCO₃ is almost insoluble. The solubility can be explained by the reason same as for sulphates.

All the carbonates of alkaline earth metals are however, more soluble in the presence of CO₂ due to the formation of corresponding bicarbonates.

For e.g.

$CaC{O}_3\left(s\right)+C{O}_2\left(g\right)+H_{2}O\left(g\right)\to Ca{\left(HC{O}_3\right)}_2\left(aq\right)$

Stability

The carbonate of all alkaline earth metals decompose on heating to form the corresponding metal oxide and CO₂.

$MC{O}_3\stackrel{\Delta }{\to }MO+C{O}_2$

The temperature of decomposition, i.e thermal stability of these carbonates, however increase down the group from Be to Ba as the basicity of metal hydroxide increases from Be(OH)₂ to Ba(OH)₂.

The sulphates are stable to heat whereas the carbonates decompose to give MO and CO₂. Thus BeCO₃ is unstable and kept in the atmosphere of CO₂ to prevent its decomposition.

BeCO₃ MgCO₃ CaCO₃ SrCO₃ BaCO₃

<100°C 540°C 900°C 1290°C 1360°C

Bicarbonates

The bicarbonates of alkaline earth metals are prepared by passing CO₂ through a suspension of metal carbonates in water.

$\underset{\left(Inso\mathrm{lub}le\right)}{M_{2}C{O}_3}+H_{2}O+C{O}_2\to \underset{\left(so\mathrm{lub}le\right)}{M{\left(HC{O}_3\right)}_2}$

All the bicarbonates of alkaline earth metals are stable only in solution and have not been isolated in the pure state.

Nitrates

Alkaline earth metals nitrates are prepared in solution and can be crystallized as hydrated salts by the action of HNO₃ on oxides hydroxides and carbonates.

$MO+2HN{O}_3\to M{\left(N{O}_3\right)}_2+H_{2}O$

$M{\left(OH\right)}_2+2HN{O}_3\to M{\left(N{O}_3\right)}_2+2H_{2}O$

$MC{O}_3+2HN{O}_3\to M{\left(N{O}_3\right)}_2+C{O}_2+H_{2}O$

(M = Be, Mg, Ca, Sr or Ba)

Magnesium nitrate crystallizes as Mg(NO₃)₂.6H₂O

White Ba(NO₃)₂ crystallises as unhydrous salt.

All nitrates on heating give the corresponding oxides.

$2M{\left(N{O}_3\right)}_2\to 2MO+2N{H}_3+O_2$

(M = Be, Mg, Ca, Sr or Ba)

Illustration 5 Arrange the following in order of the property indicated:

(i) carbonates of group 2 $\to$order of increasing thermal stability

(ii) sulphates of group 2 $\to$order of decreasing solubility

Solution. (i) BeCO₃ < MgCO₃ < CaCO₃ < SrCO₃ < BrCO₃

(ii) \begin{align} \,\,\,\,BeS{{O}_{4}}<MgS{{O}_{4}}<CaS{{O}_{4}}<SrS{{O}_{4}}<BaS{{O}_{4}}\\ \text{most}\,\,\text{soluble}\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\text{insoluble} \\ \end{align}

Anamolous Behaviour of Beryllium

The properties of berrylium the first member of the alkaline earth metal, differ from the rest of the member. Its is mainly because of

(i) Its small size and high polarizing power.

(ii) Relatively high electronegativity and ionization energy as compared to other members.

(iii) Absence of vacant d – orbitals in its valence shell.

Some important points of difference between beryllium and other members (especially magnesium) are given below:

(i) Be is harder than other members of its group.

(ii) Be is lighter than Mg.

(iii) Its melting and boiling points are higher than those of Mg & other members.

(iv) Be does not react with water while Mg reacts with boiling water.

(v) BeO is amphoteric while MgO is weakly basic.

(vi) Be forms covalent compounds whereas other members form ionic compounds.

(vii) Beryllium carbide reacts with water to give methane whereas carbides of other alkaline earth metals gives acetylene gas.

$B{e}_{2}C+4H_{2}O\to 2Be{\left(OH\right)}_2+C{H}_4$

$M{g}_2{C}_2+2H_{2}O\to Mg{\left(OH\right)}_2+C_2{H}_2$

$Ca{C}_2+2H_{2}O\to Ca{\left(OH\right)}_2+C_2{H}_2$

(viii) Beryllium does not exhibit coordination number more than four as it has four orbitals in the valence shell. The other members of this group has coordination number 6.

Resemblance of Beryllium with Aluminium (Diagonal relationship)

The following points illustrate the anomalous behaviour of Be and its resemblance with Al.

(i) Unlike groups – 2 elements but like aluminium, beryllium forms covalent compounds.

(ii) the hydroxides of Be, [Be(OH)₂] and aluminium [Al(OH)₃] are amphoteric in nature, whereas those of other elements of group – 2 are basic in nature.

(iii) the oxides of both Be and Al i.e. BeO and Al₂O₃ are high melting insoluble solids.

(iv) BeCl₂ and AlCl₃ have bridged chloride polymeric structure.

(v) The salts of beryllium as well as aluminium are extensively hydrolysed.

(vi) Carbides of both the metal reacts with water liberating methane gas.

$B{e}_{2}C+4H_{2}O\to 2Be{\left(OH\right)}_2+C{H}_4$

$A{l}_4{C}_3+12H_{2}O\to 4Al{\left(OH\right)}_3+3C{H}_4$

(vii) The oxides and hydroxides of both Be and Al are amphoteric and dissolve in sodium hydroxide as well as in hydrochloric acid.

$BeO+2HCl\to BeC{l}_2+H_{2}O$

$BeO+2NaOH\to N{a}_{2}Be{O}_2+H_{2}O$

$A{l}_2{O}_3+6HCl\to 2AlC{l}_3+H_{2}O$

$A{l}_2{O}_3+2NaOH\to 2NaAl{O}_2+H_{2}O$

(viii) Like Al, Be is not readily attacked by acids because of the presence of an oxide film.

Illustration 6. Give the structure of BeCl₂ in the

(i) vapour state (ii) solid state

Solution: (i) Linear molecule

(ii) Polymeric structure with bridged chlorine atom