Chapter-10 The s-block elements The s-block elements: The s-block elements of the Periodic Table are those in which the last electron enters the outermost s-orbital. The s-orbital can accommodate only two electrons, two groups (1 & 2) belong to the s- block of the Periodic Table. Group 1 elements: alkali metals Lithium Li Sodium Na Potassium K Rubidium Rb Caesium Cs Francium. Fr Why they are known as alkali metals? They are collectively known as the alkali metals. These are so called because they form hydroxides on reaction with water which are strongly alkaline in nature. Electronic Configuration The general electronic configuration of alkali elements is [noble gas] ns1 Lithium Li 1s22s1 or [He]2s1 Sodium Na 1s22s22p63s1or [Ne] 3s1 Potassium K 1s22s22p63s23p64s1 or [Ar]4s1 Rubidium Rb 1s22s22p63s23p63d104s24p65s1or [Kr]5s1 Caesium Cs 1s22s22p63s23p63d104s2 4p64d105s25p66s1 or [Xe] 6s1 Francium. Fr [Rn]7s1 Occurrence: Sodium and potassium are abundant and lithium, rubidium and caesium have much lower abundances. Francium is highly radioactive; its longest-lived isotope 223Fr has a half-life of only 21 minutes. All the alkali metals have one valence electron, ns1 outside the noble gas core. The loosely held s-electron in the outermost valence shell of these elements makes them the most electropositive metals. They readily lose electron to give monovalent M+ ions. Hence,they are never found in free state in nature. Atomic and Ionic Radii The alkali metal atoms have the largest sizes in a particular period of the periodic table. Withincrease in atomic number, the atom becomes larger. The monovalent ions (M+) are smaller than the parent atom. The atomic and ionic radii of alkali metals increase on moving down the group i.e., they increase in size while going from Li to Cs. Ionization Enthalpy The ionization enthalpies of the alkali metals are considerably low and decrease down the group from Li to Cs. Hydration Enthalpy The heat energy released when new bonds are made between the ions and water molecules is known as the hydration enthalpy of the ion. The hydration enthalpy is the enthalpy change when 1 mole of gaseous ions dissolve in sufficient water to give an infinitely dilute solution. Hydration enthalpies are always negative. The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes. Li+> Na+> K+> Rb+> Cs+ Li+ has maximum degree of hydration and for this reason lithium salts are mostly hydrated, e.g., LiCl· 2H2O Physical Properties 1. All the alkali metals are silvery white, soft and light metals. 2. Because of the large size, these elements have low density which increases down the group from Li to Cs. In alkali metals, atomic weight increases down the group. Lithium has less atomic weight than Sodium. So, density of 'Na' is greater than that of 'Li'. But in the case of 'K' and 'Na', d-orbitals present in Potassium, which increases the volume of 'K'. Thereby density of 'K' decreases. So, Potassium has a low density than Sodium. 3. The melting and boiling points of the alkali metals are low indicating weak metallic bonding due to the presence of only a single valence electron in them. 4. The alkali metals and their salts impart characteristic colour to an oxidizing flame. This is because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum. 5. Alkali metals can therefore, be detected by the respective flame tests and can be determined by flame photometry or atomic absorption spectroscopy. 6. These elements when irradiated with light, the light energy absorbed may be sufficient to make an atom lose electron This property makes caesium and potassium useful as electrodes in photoelectric cells. Chemical Properties The alkali metals are highly reactive due to their large size and low ionization enthalpy. The reactivity of these metals increases down the group. Reactivity towards air: The alkali metals tarnish in dry air due to the formation of their oxides. They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form super oxides. The superoxide O2– ion is stable only in the presence of large cations such as K, Rb, Cs.
4Li + O2--2Li2O (oxide)
2Na + O2 -- Na2O2 (peroxide)
2M + O2-- MO2 (superoxide) (M = K, Rb, Cs) In all these oxides the oxidation state of the alkali metal is +1. Lithium shows exceptional behaviour in reacting directly with nitrogen of air to form the nitride, Li3N as well. Because of their high reactivity towards air and water, alkali metals are normally kept in kerosene oil. (ii) Reactivity towards water: The alkali metals react with water to form hydroxide and dihydrogen.
+ - 2M +2H2O--2M + 2OH + H2 (M = an alkali metal) Other than Li, all the metals of the group react explosively with water. Although lithium has most negative E value, its reaction with water is less vigorous than that of sodium which has the least negative Evalue among the alkali metals. This behaviour of lithium is attributed to its small size and very high hydration energy. (iii) Reactivity towards dihydrogen: The alkali metals react with dihydrogen at about 673K (lithium at 1073K) to form hydrides. All the alkali metal hydrides are ionic solids with high melting points.
+ - 2M + H2 --2M H (iv) Reactivity towards halogens: The alkali metals readily react vigorously with halogens to form ionic halides, M+ X–.
Lithium halides are somewhat covalent. The Li+ ion is very small in size and has high tendency to distort electron cloud around the negative halide ion i.e.,high polarisation capability.Since, anion with large size can be easily distorted, among halides, lithium iodide is the most covalent in nature. (v) Reducing nature: The correct order of reducing character of alkali metals is Na + + M (g)+H2O→M (aq) hydration enthalpy With the small size of its ion, lithium has the highest hydration enthalpy which accounts for its high negative E∘ value and its high reducing power. (vi) Solutions in liquid ammonia: The alkali metals dissolve in liquid ammonia giving deep blue solutions which are conducting in nature. + - M+(x + y)NH3---- [M(NH3) x] +[e(NH3) y] blue colour of the solution is due to the ammoniated electron which absorbs energy in the visible region of light and thus imparts blue colour to the solution. The solutions are paramagnetic and on standing slowly liberate hydrogen resulting in the formation of amide. + - M (am)+ e + NH3(1) --- MNH2(am) + ½H (g) (where ‘am’ denotes solution in ammonia.) In concentrated solution, the blue colour changes to bronze colour and becomes diamagnetic. Uses of alkali metals Lithium metal is used to make useful alloys, for example o with lead to make ‘white metal’ bearings for motor engines, o with aluminium to make aircraft parts, and o with magnesium to make armour plates. Lithium is used in thermonuclear reactions. Lithium is also used to make electrochemical cells. Sodium is used to make a Na/Pb alloy needed to make PbEt4 and PbMe4. These organolead compounds were earlier used as anti-knock additives to petrol, but nowadays vehicles use lead-free petrol. Liquid sodium metal is used as a coolant in fast breeder nuclear reactors. Potassium has a vital role in biological systems. Potassium chloride is used as a fertilizer. Potassium hydroxide is used in the manufacture of soft soap. Potassium is also used as an excellent absorbent of carbon dioxide. Caesium is used in devising photoelectric cells. General characteristics of the compounds of the alkali metals All the common compounds of the alkali metals are generally ionic in nature. Oxides and Hydroxides On combustion in excess of air lithium forms mainly the oxide, Li2O and some peroxide Li2O2, Li(s)+O2(g) → 2Li2O(s) sodium forms the peroxide, Na2O2 and some superoxide NaO2 while Na(s)+O2(g) → 2Na2O potassium, rubidium and caesium form the superoxide, MO2. Under appropriate conditions pure compounds M2O, M2O2 and MO2 prepared. The stability of the peroxide or superoxide increases, as the size of the metal ion increases, is due to the stabilisation of large anions by larger cations through lattice energy effects. The oxides and the peroxides are colourless when pure, but the superoxides are yellow or orange in colour. The peroxides are diamagnetic (no unpaired electrons) in nature. But superoxides are paramagnetic (unpaired electrons) in nature. Sodium peroxide is widely used as an oxidising agent in inorganic chemistry. These oxides are easily hydrolysed by water to form the hydroxides + - M2O + H2O -2M + 2OH + - M2O2+ H2O --2M + 2OH + H2O2 + - 2MO2+ H2O -2M + 2OH + H2O2 + O2 The hydroxides are white crystalline solids. The hydroxides are the strongest of all bases and dissolve freely in water with evolution of much heat on account of intense hydration. Question Name the alkali metal which forms superoxides when heated in excess of air and why? Solution: Potassium forms superoxides when heated in excess of air. This is due to the stabilization of large size cation by large size anion. K + O2 → KO2 (potassium superoxide) Question Why lithium forms only lithium oxide and not peroxide or super oxides? Solution: Due to the small size of lithium, it has a strong positive field around it. On combination with the oxide anion (O2–), the positive field of lithium ion restricts the spread of negative charge towards another oxygen atom and thus prevents the formation of higher oxides. Halides of Alkali Metals The alkali metals combine directly with halogens under appropriate conditions forming halides of general formula MX. These halides can also be prepared by the action of aqueous halogen acids (HX) on metals oxides, hydroxides or carbonate. M2O + 2HX → 2MX + H2O MOH + HX → MX + H2O M2CO3 + 2HX → 2MX + CO2 + H2O (M = Li, Na, K, Rb or Cs) (X = F, Cl, Br or I) All these halides are colourless, high melting crystalline solids having high negative enthalpies of formation. The value decreases in the order: Fluoride > Chloride > bromides > Iodide Thus, fluorides are the most stable while iodides are the least stable. The trends in melting points, boiling points and solubility of alkali metals halides can be understood in terms of polarization effects, lattice energy and hydration of ions. Lattice Energies Lattice energy is defined as the amount of energy required to separate one mole of solid ionic compound into its gaseous ions. Evidently greater the lattice energy, higher is the melting point of the alkali metals halide and lower is its solubility in water Hydration Energy It is the amount of energy released when one mole of gaseous ions combineswith water to form hydrated ions. M+ (g) + aq → M+ (aq) + hydration energy X- (g) + aq → X- (aq) + hydration energy Higher the hydration energy of the ions greater is the solubility of the compound in water. Further the extent of hydration depends upon the size of the ions. Smaller the size of the ion, more highly it is hydrated and hence greater is its hydrated ionic radius and less is its ionic mobility (Conductance). A delicate balance between lattice enthalpy and hydration enthalpy determines the ultimate solubility of a compound in water. For e.g., Low solubility of LiF (0.27 g/100 g H2O) is due to its high lattice energy -1 (- 1005 KJmol ) whereas the low solubility of CsI (44g/100g H2O) is due to smaller hydration energy of the two ions (-670 KJ/mol). The solubility of the most of alkali metal halides except those of fluorides decreases on descending the group since the decrease in hydration energy is more than the corresponding decrease in the lattice energy. Due to small size and high electronegativity, lithium halides except LiF are predominantly covalent and hence are soluble in covalent solvents such as alcohol, acetone, ethyl acetate, LiCl is also soluble in pyridine. In contrast NaCl being ionic is insoluble in organic solvents. Due to high hydration energy of Li+ ion, Lithium halides are soluble in water except LiF which is sparingly soluble due to its high lattice energy. For the same alkali metal, the melting point decreases in the order fluoride > chloride > bromide > iodide because for the same alkali metal ion, the lattice energies decrease as the size of the halide ion increases. For the same halide ion, the melting point of lithium halides are lower than those of the corresponding sodium halides and thereafter they decrease as we move down the group from Na to Cs. The low melting point of LiCl (887 K) as compared to NaCl is probably because LiCl is covalent in nature and NaCl is ionic. Salts of Oxoacids Since the alkali metals are highly electropositive, therefore their hydroxides are very strong bases and hence they form salts with all oxoacids (H2CO3, H3PO4, H2SO4, HNO3, HNO2 etc). They are generally soluble in water and stable towards heat. The carbonates (M2CO3) of alkali metals are stable up to 1273 K, above which they first melt and then eventually decompose to form oxides. Li2CO3, however is considerably less stable and decomposes readily. Li2CO3 + Δ→ Li2O + CO2 + -2 This is due to large size difference between Li and CO3 which makes the crystal lattice unstable. Being strongly basic, alkali metals also form solid bicarbonates. No other metals form solid bicarbonates though NH4HCO3 bicarbonate though it does exist in solution.Lithium, however does not form solid bicarbonate. 2MHCO3 → M2CO3 + CO2 + H2O All the carbonates and bicarbonates are soluble in water and their solubilities increase rapidly on descending the group. This is due to the reason that lattice energies decrease more rapidly than their hydration energies on moving down the group. Anomalous properties of lithium The anomalous behaviour of lithium is due to the: (i) exceptionally small size of its atom and ion, and (ii) high polarising power Difference between Lithium and other Alkali Metals Lithium is harder than other alkali metals. Melting and boiling point is higher than other alkali metals. Out of all the other alkali metals, it is the least reactive metal. It is a strong reducing agent compared to other alkali metal. It is the only alkali metal that forms its monoxide. Compared to other alkali metals, it is not capable of forming solid hydrogen carbonates. It does not react with ethyne to form ethynide. On the other hand, all other alkali metals form ethynide. It reacts slowly with bromine as compared to other alkali metals. Similarities between Lithium and Magnesium The similarity between lithium and magnesium is particularly striking and arises because of their similar sizes. Both lithium and magnesium are harder and lighter than other elements in the respective groups. Lithium and magnesium both form monoxides. 2Mg + O2 → 2MgO They both react with nitrogen to form their nitrides respectively. Li(s) + N2(g) → 2Li3N(s) Mg(s) + N2(g) -→Mg3N2 They both react slowly with water. They form oxides and hydroxides which decompose on heating. Mg(s) + 2H2O(g) → Mg (OH)2(aq) + H2(g) The oxides, Li2O and MgO do not combine with excess oxygen to give any superoxide. The carbonates of both decompose on heating to form the oxides and carbon- dioxide. 2 MgCO3 → 2MgO(s) + CO2(g) Li2CO3 + Δ→ Li2O + CO2 Lithium chloride (LiCl) and magnesium chloride (MgCl2) are both soluble in ethanol. Both LiCl and MgCl2 are deliquescent and crystallise from aqueous solution as hydrates, LiCl·2H2O and MgCl2·8H2O. Both the elements are less stable towards heat. Group 2 elements: alkaline earth metals Beryllium Be Magnesium Mg Calcium Ca Strontium Sr Barium Ba Radium Ra Why they are known as alkaline earth metals? These elements with the exception of beryllium are commonly known as the alkaline earth metals. These are so called because their oxides and hydroxides are alkaline in nature and these metal oxides are found in the earth’s crust. Electronic Configuration The general electronic configuration of alkaline earth elements is [noble gas] ns2 Beryllium Be 1s22s2 or [He]2s2 Magnesium Mg 1s22s22p63s2or [Ne] 3s2 Calcium Ca 1s22s22p63s23p64s2 or [Ar]4s2 Strontium Sr 1s22s22p63s23p63d104s24p65s2or [Kr]5s2 Barium Ba 1s22s22p63s23p63d104s2 4p64d105s25p66s2 or [Xe] 6s2 Radium Ra [Rn]7s2 Atomic and Ionic Radii The atomic and ionic radii of the alkaline earth metals are smaller than those of the corresponding alkali metals in the same periods. This is due to the increased nuclear charge in these elements. Within the group, the atomic and ionic radii increase with increase in atomic number. Ionization Enthalpies: The alkaline earth metals have low ionization enthalpies due to fairly large size of the atoms. Since the atomic size increases down the group, their ionization enthalpy decreases. The first ionisation enthalpies of the alkaline earth metals are higher than those of the corresponding Group 1 metals. This is due to their small size as compared to the corresponding alkali metals. The second ionisation enthalpies of the alkaline earth metals are smaller than those of the corresponding alkali metals. Hydration Enthalpies Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease withincrease in ionic size down the group. Be2+> Mg2+> Ca2+> Sr2+> Ba2+ The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions. Physical Properties The alkaline earth metals, in general, are silvery white, lustrous and relatively soft but harder than the alkali metals. The melting and boiling points of these metals are higher than the corresponding alkali metals due to smaller sizes. The electropositive character increases down the group from Be to Ba Calcium, strontium and barium impart characteristic brick red, crimson and apple green colours respectively to the flame. In flame the electrons are excited to higher energy levels and when they drop back to the ground state, energy is emitted in the form of visible light. The electrons in beryllium and magnesium are too strongly bound to get excited by flame. Hence, these elements do not impart any colour to the flame. The flame test for Ca, Sr and Ba is helpful in their detection in qualitative analysis and estimation by flame photometry. Chemical Properties The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group. Reactivity towards air and water: Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface. Magnesium is more electropositive and burns with dazzling brilliance in air to give MgO and Mg3N2. Calcium, strontium and barium are readily attacked by air to form the oxide and nitride. They also react with water with increasing vigour even in cold to form hydroxides. Reactivity towards the halogens: All the alkaline earth metals combine with halogen at elevated temperatures forming their halides. M+ X2 --> MX2[X= F, Cl, Br, I] Reactivity towards hydrogen: All the elements except beryllium combine with hydrogen upon heating to form their hydrides, MH2. BeH2, however, can be prepared by the reaction of BeCl2 with LiAlH4. Reactivity towards acids: The alkaline earth metals readily react with acids liberating dihydrogen. M + 2HCl → MCl2 + H2↑ Reducing nature: Like alkali metals, the alkaline earth metals are strong reducing agents. However, their reducing power is less than those of their corresponding alkali metals. Beryllium has less negative valueof reduction potential compared to other alkaline earth metals. Its reducing nature is less due to large hydration energy associated with the small size of Be2+ ion and relatively large value of the atomization enthalpy of the metal. Solutions in liquid ammonia: Like alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue- black solutions forming ammoniated ions. 2+ - M + (x+y) NH3 --> [M (NH3) x] + 2[e (NH3) y] Uses of alkaline earth metals: Beryllium is used for creating alloys. Beryllium in combination with copper is sued to create strong springs. Beryllium is used for making windows ofX-ray tubes. Magnesium in combination with aluminium is used to make aeroplanes. Magnesium is also used in flash powders and bulbs. Antacid used to get relief from acidity is a suspension of magnesium hydroxide in water. Calcium is used in the extraction of metals from their respective Radium being a radioactive metal is used in radiotherapy. General characteristics of compounds of the alkaline earth metals Oxides of Alkali Earth Metals The oxides of alkali earth metals (MO) are obtained either by heating the metals in oxygen or by thermal decomposition of their carbonates. 2M + O2 2MO (M = Be, Mg, Ca) MCO3 MO + CO2 (M = Be, Mg, Ca, Sr, Ba) 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. Hydroxides of Alkali Earth Metals 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 + 2H2O → M (OH)2 + H2 (M = Ca, Sr, Ba) MO + H2O → M (OH)2 (M = Ca, Sr, Ba) Be(OH)2 and Mg(OH)2 being insoluble are obtained from suitable metal ion solutions by precipitation with OH- ions. BeCI2 + 2NaOH → Be (OH)2 ↓ + 2NaCI MgSO4 + 2NaOH → Mg (OH)2 ↓ + Na2SO4 Properties of Hydroxides of Alkali Earth Metals (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 increases 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. Becomes more negative as we move from Be(OH)2 to Ba(OH)2 which accounts for increase in solubility. Halides of Alkaline Earth Metals 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 → MX2 + H2 MO + 2HX → MX2 + H2O M (OH)2 + 2HX → MX2 + 2H2O MCO3 + 2HX → MX2 + H2O + CO2 Except for beryllium halides all other halides of alkaline earth metals are ionic in nature. Beryllium halides are essentially in the vapour phase BeCl2 tends to form a chloro-bridged dimer which dissociates into the linear monomer at high temperatures of the order of 1200 K covalent and soluble in organic solvents. Beryllium chloride has a chain structure in the solid state Salts of Oxoacids: The alkaline earth metals also form salts of oxoacids. Sulphates of Alkaline Earth Metals The sulphates of alkaline earth metals (MSO4) are prepared by the action of sulphuric acid on metals, metals oxides, hydroxides and carbonates. M + H2SO4 → MSO4 + H2 MO + H2SO4 → MSO4 + H2O M (OH)2 + H2SO4 → MSO4 + 2H2O MCO3 + H2SO4 → MSO4 + CO2 + H2O 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, CaSO4 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 +2 +2 Be to Ba . The high solubility of BeSO4 and MgSO4 is due to high hydration energies due to smaller Be+2 and Mg+2 ions. Stability: The sulphates of alkaline earth metal decompose on heating giving the oxides and SO3. MSO4 MO + SO3 Carbonates The carbonates of magnesium and other alkaline earth metals are sparingly soluble in water and their solubility decreases down the group from beryllium to barium. As we move down the group, the lattice enthalpies of carbonates remain approximately the same. The reason being that the carbonate ion is so large that relatively small changes in the sizes of the cations from Be2+ to Ba2+. Hence, the solubilities of carbonates of the alkaline earth metals decrease down the group mainly due to decreasing hydration enthalpies of the cations from Be2+ to Ba2+. All the carbonates of alkaline earth metals are, however, more soluble in the presence of CO2 due to the formation of corresponding bicarbonates. All the carbonates in this group undergo thermal decomposition to the metal oxide and carbon dioxide gas. The term "thermal decomposition" describes splitting up a compound by heating it. MgCO3 → MgO + CO2 Nitrates: The nitrates are made by dissolution of the carbonates in dilute nitric acid. Magnesium nitrate crystallises with six molecules of water, whereas barium nitrate crystallises as the anhydrous salt. This again shows a decreasing tendency to form hydrates with increasing size and decreasing hydration enthalpy. All of them decompose on heating to give the oxide like lithium nitrate. Anomalous behaviour of beryllium The anomalous behaviour is due to the following reasons: The small size of Be atom and Be++ ion. High polarising power of Be2+ ion. Absence of vacant d-orbital in its valence shell. Higher ionisation potential and electronegativity compared to other alkaline earth metals. Beryllium is the first element of the alkaline earth metals. It shows anomalous behaviour in many of the cases like: Be is harder while other alkaline earth metals are soft. High melting and boiling point. Does not react with water even at high temperature. Form covalent compounds rather than ionic. Oxides and hydroxides show amphoteric nature. Beryllium carbide is covalent in nature and reacts with water to produce methane. Due to the absence of d-orbital, show maximum coordination no of 4 rather than 6. Diagonal Relationship between Beryllium and Aluminium (1) Both these form covalent compounds having low melting point and soluble in organic solvent. (2) Both have same value of electronegativity (which is 1.5). - - (3) Both form many stable complexes for example (BeF3) , (AlH4) . (4) Like BeO, Al2O3 is amphoteric in nature. Also, both are high melting point solids. Al2O3 + 2NaOH → 2NaAlO2 + H2O Al2O3 + 6HCl → 2AlCl3 + 3H2O (5) Be and Al both react with NaOH to liberate H2 forming beryllates and aluminates. Be + 2NaOH → Na2BeO2+H2 2Al + 6NaOH → 2Na3AlO3 + 3H2 (6) Be2C and Al4C3 both provides CH4 on treating with the water. Be2C+ 2H2O → CH4 + 2BeO Al4C3 + 6H2O → 3CH4 + 2Al2O3 (7) Both Be and Al react very slowly with dilute HCl to liberate H2. (8) Both Be and Al form polymeric covalent hydrides while hydrides of other alkaline earth are ionic. (9) Both BeCl2 and AlCl3 are soluble in organic solvents and act as catalyst in Friedel - Crafts reaction. (10) Be (OH)2 and Al (OH)3is both amphoteric whereas hydroxides of other alkaline earths are strong alkali. (11) The salts of Be and Al are extensively hydrated. (12) BeCl2 and AlCl3 both have a bridged polymeric structure.