Inorganic Chemistry II Past Lecture Notes PDF

Inorganic chemistry II Second stage / Second semester The First lecture 2024/2025 Pro. Dr. Mohammed Hamid The References:- 1- Advanced Inorganic Chemistry by F.Albert Cotton ,Geoffrey Wilkinson , Carlos A. Murillo , Manfred Bochmann ; Sixth Edition 2009 2- Concise Inorganic Chemistry by J.D.Lee...

Inorganic chemistry II Second stage / Second semester The First lecture 2024/2025 Pro. Dr. Mohammed Hamid The References:- 1- Advanced Inorganic Chemistry by F.Albert Cotton ,Geoffrey Wilkinson , Carlos A. Murillo , Manfred Bochmann ; Sixth Edition 2009 2- Concise Inorganic Chemistry by J.D.Lee ; Fifth Edition 2011 **[Syllabus :- ]** **1-Alkaline Earth Metals and their Compounds** **2- Elements of Group IIIA or 13 \[The Boron Family (*ns^2^ np^1^*)\]** **3- Elements of Group IVA or 14 \[The Carbon Family (*ns^2^ np^2^*)\]** **4- Elements of Group VA or 15 ( Element of Nitrogen Family *ns^2^np^3^*)** **5- Elements of Group VIA or 16 ( Element of Oxygen Family *ns^2^np^4^*)** **6- Elements of Group VIIA or 17 ( Element of Halogens Family *ns^2^np^5^*)** **7- Elements of 18 or Zero Group (Inert Gases or Noble Gases , *ns^2^np^6^*)** **8- Transition Elements , *d*-block Elements** **9- Transition Elements , *f*-block Elements** **Alkaline Earth Metals and their Compounds** **(Group IIA or 2, ns²)** **[POSITION OF ALKALINE EARTH METALS IN PERIODIC TABLE]** The group IIA of the periodic table consists of six elements-beryllium, magnesium, calcium, strontium, bariums and radium. These elements are collectively called as alkaline earth metals because their earths (the old name for oxide) are basic (alkaline) and group IIA is known as alkaline earth group. The oxides of three principal members calcium strontium and barium were known much earlier than the metals themselves. These oxides were alkaline in nature and existed in the earth and were named alkaline earths. The metals when discovered were also called alkaline earths. This term is now applied to all the six elements of group IIA. Element Symbol The most important minerals ----------- -------- --------------------------------- Beryllium Be Be~3~AL~2~(SiO~3~)~6~ Magnesium Mg KCl.MgCl~2~.6H~2~O Calcium Ca MgCO~3~.CaCO~3~ ,CaSO~4~.2H~2~O Strontium Sr SrCO~3~ Barium Ba BaSO~4~ , BaCO~3~ Radium Ra It is found in uranium ores The first member beryllium is less active than other members and shows some abnormal properties like lithium in 1A group. However, it shows resemblance with aluminium (a member of Iird group). *i.e.* diagonal relationship. The last member, radium is radioactive in nature. Each member of this group occupies a place just after the members of IA group in various periods of periodic table except first period. IA Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 IIA Be 4 Mg 12 Ca 20 Sr. 38 Ba 36 Ra 88 The members of this group show a marked resemblance in their properties and possess same electronic configuration. There is gradual gradation in the properties with the increase of atomic number. This justifies their inclusion in the same group of periodic table. The main properties are discussed below for this justification. ***[Electronic Configuration]*** The valence electron configuration of the atoms of the group IIA elements is ns^2^ , where n is the period number. The arrangement or the distribution of electron on various subshells in the atoms of alkaline earth metals is given below Element At. No. Electronic Configuration Configuration of the valency shell ----------- --------- ----------------------------------------------------------------------------------------------------- ------------------------------------ Beryllium 4 1s^2^ 2s^2^ \[He\]2s^2^ Magnesium 12 1s^2^ 2s^2^ 2p^6^ 3s^2^ \[Ne\]3s^2^ Calcium 20 1s^2^ 2s^2^ 2p^6^ 3s^2^ 3p^6^ 4s^2^ \[Ar\]4s^2^ Strontium 38 1s^2^ 2s^2^ 2p^6^ 3s^2^ 3p^6^ 3d ^10^4s^2^ 4p^6^ 5s^2^ \[Kr\]5s^2^ Barium 56 1s^2^ 2s^2^ 2p^6^ 3s^2^ 3p^6^ 3d^10^4s^2^ 4p^6^ 4d^10^ 5s^2^ 5p^6^ 6s^2^ \[Xe\]6s^2^ Radium 88 1s^2^ 2s^2^ 2p^6^ 3s^2^ 3p^6^ 3d^10^ 4s^2^ 4p^6^ 4d^10^ 4f^14^ 5s^2^ 5p^6^ 5d^10^ 6s^2^ 6p^6^ 7s^2^ \[Rn\]7s^2^ The outermost shell of these elements has two electrons and the penultimate shell contains 8 electrons except the first member which contains 2 electrons. Since, the last electron enters ns orbital, these are s-block elements. Beryllium shows somewhat abnormal properties as its electronic configuration is slightly different than the rest of the members. Because of their similarity in electronic configuration \[noble gas\] ns², they are included in the same group, i.e., IIA of the periodic table and closely resemble each other in the physical and chemical properties. ***[2. Physical Properties]*** \(a) Physical state: All the group IIA elements are metals and too reactive, so that cannot occur in the uncombined state in nature. They are all silvery white metals. They have greyish white lustre when freshly cut, but tarnish soon after their exposure in air due to surface oxidation. They are soft in nature but harder than alkali metals because metallic bonding is stronger than 1A elements due to possession of 2 valency electrons. However, hardness decreases with increase in atomic number. \(b) Atomic and ionic radii: The size of the atom increases gradually from Be to Ra, on account of the presence of an extra energy shell at each step. The atoms are large but smaller han corresponding IA elements since the extra charge on the nucleus attracts the electron cloud inwards. Their ions are also large and size of the ion increases from Be^2+^ to Ra^2+^ Atomic volume also increases as the atomic number increases \(c) Density: These metals are denser than alkali metals in the same period because these can be packed more tightly due to their greater nuclear charge and smaller size. The density decreases slightly up to calcium and then increases considerably up to radium. Irregular trend is due to the difference in the crystal structure of these elements. +---------+---------+---------+---------+---------+---------+---------+ | Symbol | Abundan | Atomic | Ionic | Density | Ionizat | | | | ce | | | | ion | | | | in | radius | radius | g/cc | potenti | | | | earth^' | A^0^ | A^0^ | | als | | | | ^s | | | | | | | | crust | | | | | | | | p.p.m | | | | | | +=========+=========+=========+=========+=========+=========+=========+ | | | | | | 1^st^eV | 2^nd^ | | | | | | | | eV | +---------+---------+---------+---------+---------+---------+---------+ | Be | 6 | 0.89 | 0.31 | 1.8 | 9.3 | 18.2 | +---------+---------+---------+---------+---------+---------+---------+ | Mg | 20.9 | 1.36 | 0.65 | 1.7 | 7.6 | 15.0 | +---------+---------+---------+---------+---------+---------+---------+ | Ca | 36.3 | 1.74 | 0,99 | 1.6 | 6.1 | 11.9 | +---------+---------+---------+---------+---------+---------+---------+ | Sr | 300 | 1.91 | 1.13 | 2.6 | 5.7 | 11.0 | +---------+---------+---------+---------+---------+---------+---------+ | Ba | 250 | 1.98 | 1.35 | 3.5 | 5.2 | 10.0 | +---------+---------+---------+---------+---------+---------+---------+ | Ra | 1.3x10^ | \- | 1.50 | 5.0 | 5.3 | 10.1 | | | -4^ | | | | | | +---------+---------+---------+---------+---------+---------+---------+ \(d) Melting and boiling points: The melting and boiling points of these elements are higher than corresponding alkali metals. This is due to the presence of two electrons in the valency shell and thus, strongly bonded in the solid state. However, melting and boiling points do not show any regular trend because atoms adopt different crystal structures. \(e) Ionisation energies and electropositive character: The first and second ionisation energies of these metals decrease from Be to Ba. The second ionisation energy in each case is higher than the first, nearly double the first ionisation energy. The ionisation energy of last member, radium, is slightly higher than that of barium and it is difficult to explain this anomalous behaviour. Although, the ionisation energies of these elements are higher than those of alkali metals, yet these are sufficiently low to make these atoms to lose two electron of their valency shell to form M^2+^ ions and achieve the inert gas configuration. These metals are thus, strongly electropositive in nature but less than corresponding alkali metals. The electropositive character increases from Be to Ba. Metallic character and reactivity are directly linked with the tendency to lose electron or electrons, te with electropositive nature. Thus, these characters increase gradually from Be to Ba. Be Mg Ca Sr Ba Ra ^Electropositive\ nature\ increases\ Metallic^ ^character\ increases\ Reactivity\ of\ the\ metals\ increases^ \(f) Oxidation states: The alkaline earth metals form a basic oxide with general formula RO. All show a stable oxidation state +2 in their compounds. The second ionisation energy is nearly double the first ionisation energy for all these elements. This should cause these elements to exhibit a stable +1 oxidation state and form compounds like BaCl , SrBr , Cal etc., instead of BaCl~2~, SrBr~2~, Cal~2~, etc. However, the lattice energy increases as the charge on the ion increases. The increase in the lattice energy on account of the second electron from ns² is much more than the energy required (second ionisation energy) to remove it. Hence, the stability of +2 oxidation state is due to high lattice energy. The second factor responsible for +2 oxidation state is the hydration energy which is high for M^2+^ ions. On account of the availability of energy, the process does not stop to M^+^ state but reach to M^2+^ state readily. Since, the bivalent ions, M^2+^, have an inert gas configuration, it is very difficult to remove the third electron and hence oxidation state higher than +2 is not possible. Amongst alkaline earth metals, beryllium has the highest ionisation energy, i.e., least electropositive in nature. Thus, beryllium has the minimum tendency to form Be^2+^ ion and hence a number of compounds of beryllium are covalent in nature. \(g) Hydration of ions and hydration energy: The M^2+^ jons of alkaline earth metals are extensively hydrated to form hydrated ions, \[M(H₂O)~x~\]^2+^ and during hydration a huge amount of energy, called hydration energy, is released. M²+ + xH₂O → \[M(H₂O)~X~\]^2+^ + Energy The degree of hydration and the amount of hydration energy decreases as the size of the ion increases from Be^2+^ to Ba^2+^ The hydration energies of alkaline earth metal ions are higher than those of alkali metal ions and thus the compounds of alkaline earth metals are more extensively hydrated than alkali metals. Magnesium chloride and calcium chloride exist as MgCl~2~ 6H₂O and CaCl~2~-6H₂O, respectively, while sodium chloride and potassium chloride exist as NaCl and KCl. The ionic mobilities or ionic conductance of these ions increase from \[Be(H₂O)~x~ \]^2+^ to \[Ba(H₂O)~x~ \]²+ because \[Be(H₂O)~x~\]^2+^ becomes heavy due to high degree of hydration. \(h) Electronegativity : The tendency to attract electrons is low. The electronegativity values are thus small and decrease from Be to Ra. Symbol m.pt.(K) b.pt.(K) Oxid. Potential (Volt.) Electronegativity -------- ---------- ---------- ------------------------- ------------------- Be 1560 2745 1.97 1.5 Mg 924 1363 2.36 1.2 Ca 1124 1767 2.84 1.0 Sr 1062 1655 2.89 1.0 Ba 1002 2078 2.92 0.9 Ra 973 \- \- \- \(i) Conductivity : On account of the presence of two loosely bond valency electrons per atom which can move freely throughout the crystal lattice, the alkaline earth metals are good conductors of heat and electricity. \(j) Flame colouration: In the case of Ca, Sr. Ba and Ra, the electrons can be excited by the supply of energy to higher energy levels. When the excited electrons return to the original level, the energy is released in the form of light. In beryllium and magnesium, the electrons are tightly held and hence excitation is rather difficult, thus do not show flame colouration. Ca, Sr, Ba and Ra impart a characteristic colour to the flame. Ca-brick red; Sr-crimson; Ba-green; Ra-crimson \(k) Reducing nature: The alkaline earth metals have the tendency to lose electrons and change into bivalent cation: M M^2+^ + 2e Hence, they act as strong reducing agents. The reducing nature increases as the atomic number increases. Strength of a reducing agent is linked with the value coxidation potential. The values of the oxidation potential increases from Be to Ba, hence the strength as a reducing agent increases in the same order. The oxidation potentials are lower than those of the alkali metals, hence, the alkaline earth metals are weaker reducing agents than alkali metals. The reason for the lower values of oxidation potentials is due to high heats of atomisation (sublimation) and ionisation energies. \(1) Colour and magnetic property: Since, the divalent ions have noble gas configuration with no unpaired electrons, their compounds are diamagnetic and colourless unless the anion is coloured. The metals are also diamagnetic in nature as all the orbitals are fully filled with spin paired electrons, e.g., ***[Chemical Properties]*** **(a) Occurrence:** Alkaline earth metals are reactive elements and hence do not occur free in nature. Magnesium and calcium are found in abundance in nature. Beryllium is not very abundant. Strontium and barium are much less abundant. Radium is a rare element.Calcium and magnesium are the most common and commercially useful of the alkaline earth elements. We can see in the table given below, calcium is the fifth and magnesium is the eighth most abundant element in the earth\'s crust. **Ten most Abundant Elements in the Earths Crust** No. Element Mass percentage No. Element Mass percentage ----- ---------- ----------------- ----- ----------- ----------------- 1 Oxygen 46.6 6 Sodium 2.8 2 Silicon 27.7 7 Potassium 2.6 3 Aluminum 8.3 8 Magnesium 2.1 4 Iron 5.1 9 Titanium 0.4 5 Calcium 3.6 10 Hydrogen 0.1 Like the alkali metals, the group IIA elements occur i nature as silicate rocks. They also occur as carbonates an sulphates, and many of these are commercial sources of alkalin earth metals and compounds. These metals occur in nature largely as carbonates, sulphate and silicates. **(b) Extraction:** The metals of this group are not easy to produce on account of following reasons: (i) The metals cannot be produced by chemical reduction because they are themselves strong reducing agents and they react with carbon and form carbides. \(ii) They are strongly electropositive and react with water and so aqueous solutions cannot be used for displacing them with another metal. \(iii) The electrolysis of aqueous solutions of their salts produces hydrogen at cathode rather than the metal as the metal reacts with water. Electrolysis of an aqueous solution can be carried out by using mercury as cathode, but recovery of the metal from amalgam is difficult. These metals are best isolated by electrolysis of their fused metal halides containing NaCl. NaCl lowers the fusion temperature and makes the fused mass as good conductor of electricity. **(c) Reactivity towards water:** Calcium, strontium, barium and radium decompose cold water readily with evolution of hydrogen. M + 2H₂O →M(OH)₂ + H₂ Magnesium decomposes boiling water but beryllium does not react with water, even when red hot, its protective oxide film survives even at high temperature as its oxidation potential is lower than the other members. Reactivity of alkaline earth metals increases as we move down the group as the oxidation potential increases. However, the reaction of alkaline earth metals is less vigorous than alkali metals. \(d) Reactivity towards atmosphere: Except beryllium, these metals are easily tarnished in air as a layer of oxide is formed on their surface. The effect of atmosphere increases as the atomic number increases. Barium in powdered form bursts into flame on exposure to air. M + air →MO + M3N2 ^(Ca,\ Sr\ or\ Ba)^ \(e) Reactivity towards acids: Like alkali metals, the alkaline earth metals freely react with acids and displace hydrogen. M + H₂SO4 →MSO4 + H₂ M+2HC1 → MCl₂ + H₂ Beryllium behaves differently as it dissolves in caustic alkalies also with liberation of hydrogen. It is due to diagonal relationship with aluminium. Be is thus amphoteric in nature. Be + 2NaOH → Na₂BeO₂ + H₂ ^Sodium\ beryllate^ **(f) Affinity for non-metals:** Alkaline earth metals have great affinity for non-metals. They directly react with non-metals at the appropriate temperature. **(i) Reaction with hydrogen:** Except beryllium, all combine with hydrogen directly to form hydrides of the type MH₂ when heated with hydrogen. M + H₂→ MH₂ BeH₂ and MgH₂ are covalent in nature while other hydrides are ionic in nature. Calcium, strontium and barium hydrides liberate hydrogen at anode on electrolysis in the fused state. Ionic hydrides are violently decomposed by water evolving hydrogen, CaH₂ is technically called hydrolith and used on large scale for the production of hydrogen. CaH₂ + 2H₂O→ Ca(OH)~2~ +2H₂ \[BeH₂ is not obtained by direct combination of beryllium and hydrogen. It is formed by reacting beryllium chloride with lithium aluminium hydride. 2BeCl₂ + LiAIH~4~→ 2BeH₂ + LiCl + AlC13\] It is polymeric. (BeH₂), possesses hydrogen bridges. Three centre bonds are present in which a banana shaped molecular orbital covers three atoms Be\-\--H\-\--Be and contains two electrons. Hydrogen atoms lie in the plane perpendicular to the plane of molecule containing beryllium atoms.  The stability of the hydrides decreases with increasing atomic number because the metallic nature of the elements increases. **(ii) Reaction with oxygen** (Oxides and Hydroxides): Except Ba and Ra, these elements when burnt in oxygen form oxides of the type MO. 2M + O₂→ 2MO Beryllium metal is relatively unreactive and does not react below 600 ^0^C. but the powder form is much more reactive and burns brilliant. The element. Mg burns with dazzling brilliance evolving a lot of heat. Barium and radium, being highly electropositive, form peroxides. Thus, the affinity for oxygen increases on moving down the group. BeO is usually formed by ignition of the metal, but the other metal oxides (*MO* type) are usually obtained by thermal decomposition of the carbonates, *M*CO~3~. *M*CO~3~ ^Heat^ *M*O + CO₂ The oxides are very stable compounds (BeO and MgO are used as refractory materials) and white crystalline solids. Except BeO (predominantly covalent), all the other oxides are ionic and possess NaCl structure (face centred cubic). The reason for high stability is due to high lattice energy values which, however, decrease as the size of the metal ion increases.  Except BeO, which is amphoteric in nature, other *MO* oxides are basic in nature as they combine with water to form basic hydroxides. This reaction is highly exothermic. MO + H₂O →*M*(OH)₂ + Heat (where, *M* = Ca²+, Sr^2+^ or Ba²+) Basic nature of the oxides increases gradually from BeO to BaO. (\*The amphoteric nature is supported by its reaction with acids as well BeO + 2HCl →BeCl₂ + H₂O. BeO+ 2NaOH → Na~2~BeO~2~ + H~2~O.) \[BeO and MgO are insoluble in water as these are tightly held together in the solid state.\] Be(OH)₂ is amphoteric, but the hydroxides of other alkaline earth metals are basic. The basic strength increases gradually. Be(OH)₂ + 2HCl → BeCl₂ + 2H₂O Be(OH)2 + 2NaOH → Na₂BeO₂ + 2H₂O ^Sod.\ beryllate^ Be(OH)~2~ + 2OH^-^ → \[Be(OH)~4~1^2-^ ^Beryllate\ ion^ The solubility of the hydroxides increases with increase of atomic number of the alkaline earth metals. This is due to the fact that decrease in lattice energy is more than decrease in hydration energy on moving down the group. The increasing solubility can also be explained on the basis of values of their solubility products which increase from Be(OH)₂ to Ba(OH)₂. Be(OH)~2~ and Mg(OH)~2~ are almost insoluble in water. Metal hydroxide Be(OH)₂ Mg(OH)~2~ Ca(OH)~2~ Sr(OH)~2~ Ba(OH)~2~ Solubility product 1.6x10^-26^ 8.9x10^-12^ 1.3x10^-4^ 3.2x10^-4^ 5.4x10^-3^ (K~sp~) The hydroxides decompose on heating. The thermal stability increases from Be(OH)₂ to Ba(OH)~2~. Mg(OH)₂ → MgO + H₂O Ca(OH)~2~ → CaO + H₂O **(iii) Reaction with halogens (Halides):** The alkaline earth metals directly combine with halogens, when heated with them. M + X₂ ^Heated^ MX₂ (X₂ = F₂, Cl₂, B~2~, or I₂) The alkaline earth metal halides can be obtained by the action of halogen acids on metals, their oxides, hydroxides and carbonates. Beryllium halides are covalent in nature. This is due to small size and high charge of Be²+ ion, i.e., it has high polarising power. The glassy forms of halides are known to have chains of \-\-\-- X₂Be X₂Be\-\-\--. The halides of the type MX₂ (fluorides, chlorides, bromides and iodides) of other metals are ionic solids. The solubility of these halides decreases with increasing atomic number of the metal as there is decrease in hydration energy with the increase in the size of the metal ion. Solubility of BeF₂ will therefore be greater than BaF₂. As the ionic character increases on moving down the group. the melting points and their conductivity increase from magnesium halides to barium halides. They are good conductors in molten state. The halides are hygroscopic in nature and readily form hydrates, e.g., MgCl₂.6H₂O, CaCl~2~.6H₂O, BaCl₂ 2H₂O, etc. Calcium chloride has a strong affinity for water and is used as a dehydrating agent. However, BeCl₂ fumes in moist air due to its hydrolysis. BeCl₂ + H₂O Be(OH)₂ + 2HCl **(iv) Reaction with nitrogen :** All the alkaline earth metals burn in nitrogen to form nitrides of the type M~3~N₂. 3M + N₂ M~3~N₂ The ease of formation of nitrides decreases from Be to Ba. This is in contrast to alkali metals where only Li3N is formed. Because the N₂ molecule is very stable, it requires very high energy to form N³ ions. The large amount of energy comes from the very large amount of lattice energy evolved when the crystalline solid is formed. The lattice energy is particularly high because of the high charges on the ions M^2+^ and N^3-^ Be~3~N₂ is volatile (covalent character) while other nitrides are not volatile as they are ionic crystalline solids. The nitrides are hydrolysed with water liberating ammonia. M~3~N₂ + 6H₂O 3M(OH)₂ + 2NH~3~ **(v) Reaction with carbon (Carbides):** With the exception of Be, other metals when heated with carbon in an electric furnace or when their oxides are heated with carbon form carbides of the type MC₂. These carbides are called acetylides as on hydrolysis they evolve acetylene. M + 2C MC~2~ MO +3C MC~2~ + CO MC~2~ + 2H~2~O M(OH)~2~ + C~2~H~2~ MC₂ carbides, all have a distorted sodium chloride type of structure, M^2+^ replaces Na^+^ and \[ - C≡C- \]^2-^ replaces Cl^-^. MgC₂, on heating, changes into Mg₂C~3~. Mg~2~C~3~ on hydrolysis evolves propyne, CH~3~-C≡CH methyl acetylene). Mg2C3+ 4H₂O 2Mg(OH)₂ + C3H4 When BeO is heated with carbon at about 2000°C, a brick red coloured carbide of formula, Be₂C, is formed. This on hydrolysis evolves methane and is, thus, called methanide. Be₂C+4H₂O 2Be(OH)₂ + CH₂ It is also ionic but possesses an antifluorite structure. **(vi) Reaction with sulphur and phosphorus :** Alkaline earth metals directly combine with sulphur and phosphorus when heated with them to form sulphides of the type MS and phosphides of the type M~3~P~2~, respectively. M+S MS 3M+2P M~3~P~2~ Sulphides on hydrolysis liberate H₂S, while phosphides on hydrolysis evolve phosphine. MS+ dil.acid → H₂S M~3~P₂ + dil.acid → PH~3~ Sulphides are phosphorescent. They cannot be precipitated by passing H₂S through their salts solutions as they are decomposed by water. 2MS + 2H₂O M(OH)₂ + M(HS)2 **(g) Nature of oxy salts:** **(i) Bicarbonates an carbonates:** Bicarbonates of alkaline earth metals do no exist in solid state but are known in solutions only. When such solutions are heated, bicarbonates are decomposed with evolution of carbon dioxide. M(HCO3)2 ^Heated^ MCO3 + CO₂ + H₂O ^(Solution)^ Carbonates of alkaline earth metals (MCO3) are insoluble water. These dissolve in water in presence of carbon dioxide MCO3 + H₂O + CO₂→M (HCO3)2 Solubility of carbonates decreases on moving down t group, while stability increases. This is evident from the values of decomposition temperatures of various carbonates which increase gradually. MCO~3~ MO+CO~2~ Decomposition BeCO~3~ MgCO~3~ CaCO~3~ SrCO~3~ BaCO3 temp. (°C) 100 540 900 1290 1360 Increasing stability can be explained on the basis of polarisation and covalent character. Be^2+^ is smallest in size hence show high polarising power. BeCO~3~ is least ionic and has least stability. BeCO3 \< MgCO3 \< CaCO3 \< SrCO3 \< BaCO3 Increasing ionic character and stability The instability of BeCO3 is due to small size of Be^2+^ ion which is unable to stabilise the bigger CO~3~^2-^ ion. However, it can stabilise the smaller O^2-^ ion. The stability of other carbonates increases as the size of other cations increases gradually. The carbonates are all ionic, but BeCO3 is unusual because it contains hydrated ion \[Be(H₂O)~4~\]^2+^ rather than Be²+. **(ii) Sulphates:** Alkaline earth metals form sulphates of the type MSO4. These are prepared by the action of sulphuric acid on oxides, hydroxides or carbonates. MO + H₂SO~4~ → MSO~4~ + H₂O M(OH)₂ + H₂SO~4~ → MSO~4~ + 2H₂O MCO3 + H₂SO4→ MSO~4~ + H₂O + CO₂ The solubility of sulphates decreases on moving down the group. CaSO4 is sparingly soluble, while SrSO4, BaSO4 and RaSO4 are almost insoluble. The solubilities of BeSO4 and MgSO4 are due to high energy of solvation of smaller Be²+ and Mg²+ ions. The values of solubility products which decrease gradually also explain the decrease in solubility on moving down the group. Metal sulphate BeSO4 MgSO4 CaSO4 SrSO4 BaSO4 Solubility product very high 10 2.4×10^-5^ 7.6x10^-7^ 1.5x10^-9^ The sulphates decompose on heating to give the corresponding oxide (MO). 2 MSO~4~ ^Heat^ 2 MO+2SO₂ +O₂ The stability increases as the basic nature of the meta increases. This is evident from the decomposition temperatures Metal sulphate BeSO4 MgSO4 CaSO4 SrSO~4~ Decomposition temp. (°C) 500 895 1149 1374 Sulphates are reduced into sulphides on heating wit carbon. **(iii) Nitrates:** Alkaline earth metals form nitrates of the type M(NO~3~)~2~. These are prepared by the action of nitric acid with oxides, hydroxides and carbonates. 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. 2M(NO3)~2~ 2MO+4NO₂+O₂ Beryllium also forms a basic nitrate in addition to the norm salt. Basic nitrate is a covalent compound. Be(NO~3~)~2~ ^125°C^ \[Be O(NO₂).\] ^Basic\ beryllium\ nitrate^ **(h) Solutions of metals in liquid ammonia:** Like alkali metals, alkaline earth metals also dissolve in liquid ammonia to for coloured solutions. Dilute solutions are bright blue in colour due to solvated electrons. These solutions decompose very slowly forming amides and evolving hydrogen. M M^2+^ + 2e 2NH3 + 2e → 2NH₂^1-^ + H₂ M^2+^ + 2NH₂^1-^ → M(NH₂)~2~ When the solution is evaporated, hexammoniate, M(NH3)6 is formed. These slowly decompose to give amides. M(NH3)6 → M(NH₂)2 + 4NH3 ↑+H₂ ↑ Concentrated solutions of the metals in ammonia are bronze coloured. **(i) Formation of amalgams:** Alkaline earth metals combine with mercury to form amalgams. **(j) Complex formation:** Generally, the alkaline earth metals do not form complexes. However, the smaller ions have some tendency to form complexes. Beryllium forms stable complexes such as \[BeF~3~\]^-^,\[BeF~4~\]^2-^ and \[Be(H₂O)~4~\]^2+^Complexes of the type BeCl₂*R~2~* are formed where *R* is an ether, aldehyde or ketone with an oxygen as a donor atom. Beryllium is unique in forming a series of stable complexes of formula \[Be~4~O(R)~6~\],where R may be NO~3~^-^, HCOO^-^,CH~3~COO^-^, C~6~H~5~COO^-^, etc. The most important complex formed by magnesium is chlorophyll in which magnesium is bonded to the four heterocyclic nitrogen atoms. Calcium, strontium and barium form complexes only with strong complexing agents like acetylacetone, EDTA, etc. **(k) Organo-metallic compounds :** Both Be and Mg form an appreciable number of compounds with M-C bonds but only a few are known for Ca, Sr and Ba. Grignard reagents are very important in organic chemistry which can be used to form a wide variety of organic compounds. Mg + RBr ^Dry\ ether^ RMgBr (R = alkyl or aryl) ^Grignard\ reagents^ BeCl₂ reacts with Grignard compounds forming reactive dialkyls and diaryls. 2RMgCl + BeCl₂ ^Ether^ BeR₂ + 2MgCl₂ Dialkyls and diaryls of Mg, Ca, Sr and Ba can also be obtained by similar reactions. ***[SOLUBILITY OF COMPOUNDS OF ALKALINE EARTH METALS]*** In the case of the compounds of Ca, Sr and Ba the following facts are observed: \(i) The solubility of hydroxides, fluorides and oxalates increases from calcium to barium. \(ii) The solubility of carbonates, sulphates and chromates decreases from calcium to barium. The solubility of an ionic compound depends on two factors: (i) lattice energy and (ii) hydration energy. These two factors oppose each other. If lattice energy is high, the ions will be tightly packed in the crystal and, therefore, solubility will be low. If hydration energy is high, the ions will have greater ten dency to be hydrated and, therefore, the solubility will be high. In the case of hydroxides, fluorides and oxalates the lattice energies are different, i.e., lattice energy decreases as the size of the cation increases. This tends to increase the solubility as it overcomes the counter effect of decrease in hydration energy. Hence, the solubility of the hydroxides, fluorides and oxalates increases from Ca to Ba. In the case of carbonates, sulphates and chromates the anions are large in size and small changes in cation size do not alter the lattice energies, i.e., lattice energies are about the same. However, the hydration energies decrease from Ca^2+^ to Ba^2+^. Hence, the solubility of carbonates, sulphates and chromates decreases from calcium to barium. ***[DIFFERENCE BETWEEN ALKALINE EARTH METALS AND ALKALI METALS]*** Both alkaline earth metals and alkali metals are s-block elements as the last differentiating electron enters the *ns*-orbital. They resemble with each other in many 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 ionisation potential, higher electronegativity, etc. The man points of difference between alkaline earth metals and alkali metals are given below: +-----------------------+-----------------------+-----------------------+ | **Properties** | **Alkaline earth | **Alkali metals** | | | metals** | | +=======================+=======================+=======================+ | **(i) Electronic | **Two electrons are | **One electron is | | configuration** | present in the | present in the | | | valency shell. The | valency shell. The | | | con figurations is | configuration is | | | ns².** | ns¹.** | +-----------------------+-----------------------+-----------------------+ | **(ii) Valency** | **Bivalent.** | **Monovalent.** | +-----------------------+-----------------------+-----------------------+ | **(iii) | **Less | **More | | Electropositive | electropositive.** | electropositive.** | | nature** | | | +-----------------------+-----------------------+-----------------------+ | **(iv) Hydroxides** | **Weak bases, less | **Strong bases, | | | soluble and decompose | highly soluble and | | | on heating.** | stable towards | | | | heat.** | +-----------------------+-----------------------+-----------------------+ | **(v) Bicarbonates** | **These are not known | **These are known in | | | in free state. Exist | solid state.** | | | only in** | | | | | | | | **solution.** | | +-----------------------+-----------------------+-----------------------+ | **(vi) Carbonates** | **Insoluble in water. | **Soluble in water. | | | Decompose on | Do not decompose on | | | heating.** | heating (Li₂CO3 is an | | | | exception).** | +-----------------------+-----------------------+-----------------------+ | **(vii) Action of | **Directly combine | **Do not directly | | nitrogen** | with nitrogen and | combine with | | | form nitrides.** | nitrogen.** | +-----------------------+-----------------------+-----------------------+ | **(viii) Action of | **Directly combine | **Do not directly | | carbon** | with carbon and form | combine with | | | carbides.** | carbon.** | +-----------------------+-----------------------+-----------------------+ | **(ix) Nitrates** | **Decompose on | **Decompose on | | | heating evolving a | heating evolving only | | | mixture of NO₂ and | oxygen.** | | | oxygen.** | | +-----------------------+-----------------------+-----------------------+ | **(x) Solubility of | **Sulphates, | **Sulphates, | | salts** | phosphates, | phosphates, | | | fluorides, chromates, | fluorides, chromates, | | | oxalates, etc., are | oxalates, etc., are | | | insoluble in water.** | soluble in water.** | +-----------------------+-----------------------+-----------------------+ | **(xi) Physical | **Are less reactive | **Soft, low melting | | properties** | and comparatively | points. | | | harder metals. High | Paramagnetic.** | | | melting points. | | | | Diamagnetic.** | | +-----------------------+-----------------------+-----------------------+ | **(xii) Hydration of | **The compounds are | **The compounds are | | compounds** | extensively hydrated. | less hydrated. NaCl, | | | MgCl₂-6H₂O, | KCl and RbCl form | | | CaCl2-6H₂0 and BaCl₂ | non-hydrated | | | 2H₂O are hydrated | chlorides** | | | chlorides.** | | +-----------------------+-----------------------+-----------------------+ | **(xiii) Reducing | **Weaker, as | **Stronger, as | | power** | ionisation potential | ionisation potential | | | values are high and | values are low and | | | oxidation potential | oxidation potential | | | values are low.** | values are high.** | +-----------------------+-----------------------+-----------------------+