Group 1st A
- Lithium (Z=87), Sodium (Z= 11), Potassium (Z=19), Rubidium (Z=37), Caesium (Z=55) and Francium (Z= 87) are the elements of IA group of the periodic table . These elements are known as alkali metals, because they form water-soluble hydroxides, which are strongly alkaline. Francium is the radioactive elements among IA group elements.
- Atoms of alkali metals have single valence electron and form cations with only a single positive charge when they take part in compound formation. All the alkali metals are paramagnetic because of the presence of one valence electron in the valence shell. General outer electronic configuration of alkali metals is ns1 (where n = 2 3, 4, 5, 6 and 7 ). Their penultimate quantum layers have the stable electronic configuration of noble gas s2 p6, (for lithium s2). The alkali metal cations, Li+, Na+, K+ , Rb+ and Cs+ all are diamagnetic because of absence of unpaired electrons . All these cations have stable s2 p4 noble gas configuration.
- The alkali metals are extremely reactive metals. They all readily lose one electron, the outermost s-electron, to form a single positively charged ion that has the same electronic configuration as one of the noble gases. The relatively simple chemistry of alkali metals is probably due to the ease with which their single outermost electron can be lost in order to achieve stable rare gas electronic configuration. As a group, the alkali metals have the lowest first ionization energies among all the elements in the periodic table. Because of their strong chemical activity, the alkali metals occur in nature only in the form of compounds. Alkali metals never occur in free state in nature. Compounds of sodium and potassium are the most abundant, while those of remaining alkali metals occur in nature rarely. Francium is radioactive element and short lived. These metals are readily oxidised in air. The less electropositive such as lithium is found as silicate and the most electropositive elements of IA occur as chlorides. To prevent oxidation they are stored in closed vessels of low density (0.53 g/cc). Lithium is the lightest or least dense of all metals, with a density roughly half that or water.
Lithium is kept wrapped in paraffin wax.
- They are all silvery white, soft and malleable metals. They are so soft that they can easily be cut with a knife.
- All the alkali metals show metallic luster, which can be explained by the oscillation of free electrons. Alkali metals are most electropositive in the periodic table. The metallic character increases from lithium to Caesium. The elements are all metals and are good conductors of heat and electricity.
- The atomic radius in a group increases from lithium to Caesium. In each period, the alkali metal has maximum atomic size. As the size of atom increases from Li to Cs, there is more repulsion from the non-bonding electrons, resulting in decrease in cohesive energy and an increased softness.
- Alkali metals have only one valence electron. Thus their metallic bonding is weak. They are, therefore, soft metals of low density and have low melting and boiling points. They also have low enthalpies of fusion, vaporization and atomization.
- Alkali metals have low density because of their large atomic size. Density of these elements gradually increases from lithium to Caesium. However, potassium is lighter than sodium. Due to weak intermolecular attraction, these metals have low melting points. Thus energy required for binding the atoms in the crystal lattice of the metal is low. The melting point of alkali metals also decreases from lithium to calcium. The boiling point of alkali metals also decreases from lithium to Caesium.
- The ionisation energy of these elements is low and gradually decreases from lithium to Caesium. Thus Caesium is the most electropositive element in the periodic table. Because of low ionisation energies , these elements have great tendency to lose 1st electron to form positive ions. Thus these elements are said to have strong electropositive or metallic character. Because ionisation energy decreases on moving down a group, from lithium to Caesium, the electropositive character or metallic character increases from lithium to Caesium.
- Alkali metals have very low electronegativities and this property decreases from lithium to Caesium.
- All the alkali metals exhibit +1 oxidation state. Since after losing 1 electron, the resulting cation possesses the stable noble gas configuration, other oxidation states are not possible.
- The vapours of alkali metals impart characteristic colours to the flame. For example, caramine red for lithium, yellow for Na, violet (lilac) for K, violet rose for Rb and blue for Cs. The flame test is due to the excitation and de-excitation of electrons of outer energy levels. In flame test, conc. HCl is added to convert the given salt into volatile chloride.
- The metals form compounds mainly with ionic bonds, because they readily give up the s-electron from the outermost shell. Small highly polarizing ion such as Li+ tend to form covalent compounds . Most other cations of IA group elements are large and have low charge densities. Hence they form stable compounds even with polarizing ions such as carbonate, peroxides, tri lodide etc. LiCl is soluble in pyridine because of its covalent nature. Bond present in the molecules like Li2 , Na2, K2 etc is covalent . It should be noted that alkali metals form diatomic molecules such as Li2, Na2, K2, etc to the extent of 1% in vapour state.
- Because of low ionisation energies and strong electropositive character, alkali metals such as K, Rb and Cs show photoelectric effect .
- In the electrochemical series all alkali metals are arranged above hydrogen. They are easily and readily oxidised by water. The reaction with water is so exothermic that the heat generated often ignites the H2 (g) produced, if air is present. These metals replace hydrogen from water. The alkali metals readily react with non-metals such as halogens and sulphur to form ionic salts. These reactions are exothermic, frequently violent and can even be explosive.
- When these metals are heated at 350-400°C in an atmosphere of hydrogen, they form hydrides, in which hydrogen is electronegative. The ease of formation ionic hydrides increases from Li to Cs.
- The alkali metals must be protected from exposure to air because they react with O2, H2O and CO2. When burnt in air, lithium forms monoxide (Li2O), sodium forms monoxide in limited O2and peroxide in excess O2. Other alkali metals form largely the super oxides. In the presence of atmospheric oxygen we will get the mixture of oxides.
The reactions of alkali metals with water and oxygen are so exothermic that they can cause a fire or even an explosion. It is therefore, dangerous to expose alkali metals to the air, which contains both oxygen and water vapour. The hydroxides of alkali metals are hygroscopic, readily absorb CO2, form air and destroy glass. They dissociate in solution almost completely.
- The alkali metals replace hydrogen from those acids which are not oxidizing agents (eg. HCI, dilute H2 SO4). While reacting with oxidizing acids such as conc. H2 SO4, conc. HNO3 and dilute HNO3 , these metals do not replace hydrogen but H2 S , N2 O and NH3 are respectively formed.
- Compounds of alkali metals are ionic crystalline solids and with very few exceptions these compounds are soluble in water. Alkali metals crystallise in body centred cubic lattices. The co-ordination number is 8. Seawater contains sodium in the form of ionic salt, NaCl. NaCl is essential for digestion in animals and humans, but too much salt can cause high blood pressure and retention of water in tissues. Potassium is an important element for plant growth. Potassium salts such as KNO3or KCl are used as fertilizers.
- Because of strong tendency to lose the s-electron , the alkali metals are powerful reducing agents. Lithium in aqueous solution is as strong a reducing agent as Caesium. This is probably due to high hydration energy of small lithium ion, which compensates for high ionisation energy. The hydration energy of alkali metal ions follows the order: Li+ > Na+ > K+ > Rb+ > Cs+ . Due to extensive hydration, Li+ ion has the highest hydration energy, as a result of which reduction potential of Li is higher than other alkali metals. Thus most powerful reducing agent in solution is lithium.
- The reactivity of alkali metals towards all reagents (except nitrogen) increases as the ionisation energy decreases from lithium to Caesium. Fox example, the vigorous nature of the reaction of water with alkali metals increases from Li to Cs. Lithium is only slowly attacked by water at 298 K, though it liberates more energy in its reaction with water. Sodium reacts more vigorously, whereas potassium catches fire and rubidium and cesium react violently or explosively.
- The fact that lithium reacts with water less vigorously may probably be due to its higher melting point. Other alkali metals have low melting points (below 373 K) and melt due to the heat liberated in their initial reaction with water, thus exposing more metal surface to attack. The reaction, therefore, proceeds much more rapidly than with lithium which remains in solid state.
- Despite its high ionisation energy, lithium is the only alkali metal, which reacts directly with nitrogen and carbon to form the nitride and carbide respectively. Other alkali metals do not react directly with either nitrogen or carbon. This is probably due to the fact that small size of the nitride and carbide ions coupled with the small size of Li+ ions results in high lattice energies for these compounds. Thus heats of formation of these compounds are also high resulting in more stability.
- All alkali metals form amalgams with mercury and the reaction is exothermic. Sodium amalgam is less reactive than the sodium metal. All alkali metals combine directly With P, As, P and S.
- When heated with NH3gas at 300-400°C in presence of Fe, alkali metals form amides of the type MNH2. The expected order of reactivity of the metal and the corresponding stability of the salt formed is Cs > Rb > K > Na > Li.
- Except lithium nitrate, the nitrates of alkali metals are thermally stable. Alkali metal sulphates form double salts called alums. Lithium does not form these alums because of its small size.
- The only monoxide obtained by direct combination of the metal with oxygen is that of lithium, Li2 O. Monoxides are ionic and colourless compounds and contain the oxide ion O2-. They are strongly basic. With the exception of lithium monoxide, all others can be produced by reducing the nitrite or nitrate with free metal.
- Lithium monoxide is slowly hydrolysed (due to its lattice energy ) and all other monoxides are hydrolysed with water to form hydroxides.
- The peroxides contain [ - O - O - ]2- ion and are salts of H2 O2 . The peroxide ion has two electrons more than O2molecule. The peroxides are diamagnetic and slightly yellow in colour due to the defects in crystal structure. Peroxides are powerful oxidising agents.
- Superoxides are the direct oxidation products of K, Rb and Cs, in excess of oxygen. Sodium super oxide can also be prepared by heating sodium peroxide with oxygen under pressure in a sealed vessel.
- Super oxides contain O22- ion which has one electron more than the oxygen molecule.
- The stability of peroxides and Superoxides increases with increase with increase in size of cation and shows stabilisation of large anions by large cation through lattice energy effects. The inability of lithium to form a peroxide or superoxide may probably be due to strong polarisation power of its ion. The Li+ ion is strong enough to restrict the oxidation of an oxide ion,O2- to the peroxide ion ,O22-. Similarly Na+ prevents the oxidation of O22- to O2-.
- Nickel and cast iron are resistant to fused caustic alkalis. Elements having electronegativity values less than 1.5 are not attacked by alkali hydroxides. Those elements having electronegativity values between 1.5 and 2 .0 dissolve in hydroxides and H2 is evolved with the formation of alkali metal salt of the oxide . Elements having electronegativities greater than 2 dissolve without evolution of hydrogen forming their hydride or the alkali metal derivative of that hydride as one of the products.
- Hydrides of the type MH have been prepared for all alkali metals, except francium. They are colourless, crystalline, ionic solids. The hydrides are largely ionic with the negative hydride ion H- , present in the lattice . The alkali metal hydride posses Na Cl -type structure. The ionic character of the bonds in these hydrides increases as the atomic number of metal increases. Their stability decreases in the same order.
- Hydroxides of all alkali metals are thermally very stable except LiOH , which decomposes to the oxide when calcined.
- Lithium hydride is by fact the most stable hydride which dissociates above 1273 K. Sodium and potassium hydrides dissociate below 773K. The stability decreases from LiH to CsH.
- The Li-H bond has been fond to be only 25% ionic. As the ionic character of M-H bond increases with increase in size of the metal, it would appear that M+ H- is less stable than the covalent M-H bond . The alkali metal hydrides are strong reducing agents and reducing property increases with decrease in stability. Lithium also forms some complex of the Li AlH4 and LiBH4 , which like simple hydrides , are also good reducing agents.
- Alkali metals readily react with halogens to form MX type of halides . These are mainly ionic compounds having NaCl typoe of structure with a co-ordination number of 6, except for Cs Cl , Cs Br and CsI , whiich have a Cs Cl type structure with a co- ordination number of 6, although of 8. Due to more favourable lattice energy , lithium also exhibits a co-ordination number of 6, although the theoretically expected value is 4. Because of small size of Li+ ion, all form hydrated salts more readily than other alkali metals . In fact , LiCl, LiBr and LiI , all form trihydrates , LiX.3H2O . Other alkali metal halides are anhydrous . The melting and boiling point of halides almost metal halides comprise only singly charged ions and hence have low lattice energies . The all alkali metal halides (except LiF) are soluble in water .
- For a substance to dissolve , the hydration energy must be greater than lattice energy . Due to small size of Li+ ion , the hydration energy of LiF is considerably high , but it has low solubility in water because its lattice energy is even higher . The hydration and lattice energies of LiF are -1034 and - 1039 kJ/ mole respectively.
- The alkali metals are highly electropositive and hence their oxysalts are quite stable. Carbonates are remarkable stable . Li2CO3is less stable and decomposes readily. The strong polarising action of the small cation on the large carbonate ion presumably assists the decomposition.
Al (OH)3 + 3NaOH NaAlO3 + 3H2O
- The gain in lattice energy resulting from the substitution of a smaller oxide ion from larger CO32- ion enables the decomposition. Bicarbonates of all alkali metals , except that of Li, can be obtained in the solid form. All carbonates and bicarbonates are soluble in water . Lithium carbonate is sparingly soluble.
- Due to the presence of a single loosely held electron in the valence shell, the alkali metals are excellent conductors of heat and electricity and show photoelectric effect .
- The alkali metals can not be prepared by the reduction of their oxides, as the metals themselves are strong reducing agents. Electrolysis of the aqueous solutions of their salts produces the metal hydroxide, as the metals are very reactive towards water. For these reasons, the isolation of alkali metals is usually carried out by the electrolysis of their molten salts .
- Lithium is extracted by the electrolysis of a fused mixture of LiCl and KCl ( 1:1) . KCl reduces the m.p of of LiCl from 883K to 673 K . Sodium is extracted by the electrolysis of a mixture of NaCl and CaCl2(Down's process) . CaCl2decreases the m.p of NaCl from 1076K to 778K. Potassium cannot be prepared by the electrolysis due to its low melting point and ready vapourisation. It can be prepared by passing sodium vapour over molten KCl in a counter current fractionating tower. Rubedium and caesium are made by similar methods.
- The alkali metals are best purified by distillation .
- The alkali metals are soluble in liquid ammonia giving a solution which is paramagnetic , highly conducting, highly reducing and deep blue in colour . The solution is blue when dilute and acquires and intense blue colour as more metal is added . Colour of these solutions is independent of the metal involved. Dilute solutions are paramagnetic and blue solutions have a broad absorption band at 1450nm. The blue solution has remarkably high electrical conductivity, which varies with concentration in an anomalous manner. Properties of these solutions strongly suggest that alkali metal atoms ionise in liquid ammonia forming solvated cations and solvated electrons. The solutions are strong reducing agents due to the presence of free electrons. It reduces (a) Metal halides to metals (b) N2O to N2(c) O2to O2- and then to O22- . The solution of a metal in liquid ammonia causes a large increase in volume (decrease in density). This solution of a metal in liquid ammonia causes a radius 300-400 pm. As a result , the solutions occupy far greater volume than expected from the sum of volumes of metal and solvent. This causes a decrease in density .
- Most abundant alkali metal in earth's crust is sodium . It was first isolated by Davy in 1807 . It is one of the most reactive metals and so it does not occur in free combines vigorously with water and many other elements and compounds, where it acts as a strong reducing agent . It occurs in combined form as rock salt (NaCl) , chile saltpetre or caliche (NaNO3 ), borax (Na2B4O7, 10H2O), cryolite (Na3AlF6) and felspar (NaAiSi3 O8). Sodium is extracted by Castener's process and Down's process. Sodium is soft silvery white metal, lighter than water .Its density is 0.972 g./cc. It is malleable as well as ductile.
- There are number of important sodium compounds such as NaOH (caustic soda), sodium carbonate (Na2CO3. 10 H2 O ) , Na2 CO3 , (soda ash) and sodium bicarbonate , NaHCO3 (baking soda). Sodium hydroxide is hygroscopic. When it is dissolved in water, a large amount of heat is evolved (exothermic). NaOH is prepared by the electrolysis of NaCl solution (brine). Industrially it is prepared by mixing lime with sodium carbonate. Sodium carbonate is obtained from mines or by Solvay process. Sodium chloride can be obtained from sea or underground deposits. Sodium carbonate is widely used in the manufacture of soaps, detergents and glass etc. Sodium bicarbonate or baking soda is used to make breads and cakes light and fluffy because when it is heated, CO2 is evolved. 2NaHCO3 NaCO3+ CO2 + H2O , Sodium bicarbonate is also used in fire extinguishers . Zinc uranyl acetate is one of the few fairly insoluble compounds of sodium. Its formula is NaZn (UO2)3 (CH3 COO)9 . It is pale yellow in colour . Sodium metal is used in Wurtz reaction. Sodium is also used as a coolant in atomic reactors.
- Electrolysis of aqueous NaCl gives H2 gas at cathode . Metallic sodium is released at cathode, if aqueous NaCl is electolysed-using mercury as cathode.
- Soda lye is solution of NaOH in water. Common name for KOH is caustic potash.
- NaOH can be prepared by (a) Causticizing process or lime soda process (b) Nelson's method (c) Castner Kellner methid . NaOH is deliquescent and so a standard solution ofNaOH cannot be prepared simply by weighing.
- Al , Zn, Sn , Pb, Cr and Sb are the metals which dissolve in NaOH solution and liberate H2 .
- The amphoteric substances such as Al2O3 , ZnO, Al (OH)2 dissolve in excess of NaOH solution forming aluminates and zincates.
Zn(OH)2 + 2 NaOH NaZnO2 + H2O
- Aqueous solution of NaOH and KOH can CO2 from the atmosphere.
- Alcoholic KOH is used as dehydrohalogenating agent.
- Na2 CO3 is called soda ash and Na2 CO3. 10H2O is called washing soda.
- Na2 CO3 can be prepared by Le Blanc process and Solvay process or ammonia soda process.
- Solvay process is based on low solubility of NaHCO3 . The end product of Solvay process in NaHCO3 . In soda process, the raw materials usedare Ca CO3 , NH3 and NaCl .
- K2 CO3 , known as pearl ash , can not be prepared by Solvay process due to high solubility of KHCO3 .
- NaHCO3, known as baking soda, is less soluble in water than Na2 CO3 because of polymeric anions formed by hydrogen bonding .
- All baking powders contain NaHCO3 and acid salt such as potassium hydrogen tartrate or calcium dihydrogen phosphate .
- An equimolar mixture of Na2 CO3 and K2 CO3 is called fusion mixture .
- Sodium bicarbonate is used to make breads and cakes light and fluffy because when it is heated CO2 is evolved. Sodium bicarbonate is also used in cold drinks, as antacid and in baking powders.
- Na2CO3 is widely used in the manufacture of soaps, detergents and glass. Both NaOH and Na2 CO3 are among the top 20 chemicals used in industries .
- A deliquescent compound such as NaOH is that which absorbs water readily from the atmosphere and thereby form concentrated solutions containing the alkali metal cation (Na+) and OH- ions.
- Na2O is used as a dehydrating agents as well as polymerising agent in organic chemistry.
- Anhydrous sodium sulphate is called salt cake and hydrated sodium sulphate is called Glauber's salt (Na2SO4. 10 H2 O). Glauber's salt shows efflorescence in dry air and it is used as purgative in medicine. Na2 SO4. 10H2O becomes anhydrous (Na2SO4), above 32.5°C (efflorescence).
- Na2 O2 is a yellow powder (Na2 O2. 8H2O). It is capable of absorbing gases like CO, CO2, SO2, NO2, etc. Hence it is used to purify air in submarines.
- Under trade name oxone, Na2 O2 is used as a bleaching agent.
- K2O2 is a white hygroscopic solid . KO2 is orange paramagnetic powder . K2O2 is a yellow hygroscopic solid and highly poisonous substance. It is more poisonous than NaCN.
- Potassium was first isolated by Davy in 1808 . It is soft, low malting (mp = 63°C) and more reactive than sodium . Carnallite , KCl . MgCl2. 6H2O is the important mineral source of potassium . In India potassium is found largely as saltpetre (KNO3 ). The chemical properties of potassium closely resemble those of sodium, but its reactions are more vigorous. Potassium is one of the few elements that form a super oxide (O2-) on reaction with air.
- An alloy of potassium with sodium (Na = 78% , K = 22%) is a liquid resembling mercury, It is low melting (mp = 13°C ) and high boiling (b.p 1450° C ) . It is used as a high temperature heat exchange fluid for use in power plants. At 100°C , it is less corrosive than steam and has a much higher thermal conductivity . Since potassium attack the glass electrodes, the same can not prepare potassium those of sodium.
- The potassium ion (K+ ) is chemically inert . The compounds of potassium are, in general more soluble than those of sodium. The main uses of potassium compounds are as fertilizers, where KCl and KNO3 are commonly used.
- Potassium salts of fatty acids are used in making soft soaps, because they are more soluble than those of sodium . Potassium is mainly obtained from deposits of NaCl and form brines in various lakes.
- Potassium salts are occasionally called potash, independent of the associated anion.
- Potassium ion gives a characteristic purple colour in a Bunsen flame.
- Yellow potassium sodium hexa nitro cobaltate (III) , K2 Na [ Co(NO2)6 ] and yellow potassium hexachloroplatinate , K2 Pt Cl6 are the less soluble potassium salts.
- Black ash is impure Na2 CO3 produced in Leblanc process when salt cake is reduced with coke
- Hypo is Na2 S2O3 . 5H2O.
- Fire extingushers contain H2 SO4 , Na2 CO3 and NaHCO3.
- Soda lime is NaOH + CaO.
(a) They are less reactive than alkali metals
- They belong to s-block and have general electronic configuration ns2 (where n 1). The atoms of IIA group elements have two electrons in their outermost shell. The penultimate layer has, as in case of alkali metals, 8 electrons (s2 p6 ) , except beryllium , which contains 2 electrons in the penultimate shell , that is , in the s-subshell Thus they possess an inert gas core and two electrons in the s-orbital of the valence shell (ns2) Radium is the radioactive element in IIA group.
- When compared with alkali metals
(b) They are less electropositive than alkali metals. Hence they are less metallic than alkali metals
(c) Their reducing power is much less than those of alkali metals
(d) They are less basic than alkali metals.
(a) The first hydration energy of the alkaline earth metal ion is about 5 times that of the alkali metal ions of the same period,
- Covalent character is predominant in the compounds of beryllium, because of its small size and high nuclear charge. Beryllium and magnesium form many covalent compounds whereas the compounds of other metals of the group are essentially those of dipositive (M2+ ) ions.
- The gradation of the properties of the IIA group elements is not as regular as in case of alkali metals, because of different structures of their crystal lattices.
- As compared to alkali metals , the alkaline earth metals have :
- Small size of atoms and ions
- Stronger metallic bonds
- Higher melting and boiling points
- Are denser and harder.
- These atoms are large in size but smaller than the corresponding alkali metals, since the increased nuclear charge draws the orbital electrons in. The second group metallic ions are also large but smaller than those of alkali metals are. The metal atom changes into a bipositive ion by the loss of two electrons present in the outermost shell.
- Because their atoms are smaller, the metals are denser and have high melting and boiling points than the alkali metals. Density also decreases on moving down the group upto calcium but increases considerably thereafter. Density of Ca is less than that of Mg, because of the presence of vacant 3d- orbital leading to much increase in atomic volume.
- The melting and boiling points are not regular in the group, mainly due to different crystal structures of the metals. However, these are higher than those of alkali metals.
- Solid Be and Mg has hexagonal colsed packed lattice . Ba has body centred cubic lattice. Ca and Sr have cubic close packed lattice at ordinary temperature, which changes to hexagonal close packed lattice at 500K and to body centred cube near their melting points.
- Since they have two valence electrons, hence they have are stronger metallic bonding and higher cohesive energy than those of alkali metals. Thus they are harder than alkali metals but become increasingly soft as the atomic number increases. They can be cut by a sharp knife.
- Due to larger nuclear charge and smaller atomic radius of IIA group elements (as compared to alkali metals), the valence electrons are more tightly held. Thus the first ionization energies of these elements are higher than those of alkali metals. The elements form colorless, largely ionic compounds in + 2 oxidation state, for example Mg Cl2.
- The ionization energy decreases on moving down from Be to Ra. Alkaline earth metals are not as electropositive as the corresponding alkali metals. The electropositive character increases on moving down the group.
- The heat of hydration of these elements decrease from Be to Ba . All these elements have hydrated salts. Ca, Sr, Ba and Ra impart flame colors of brick red, crimson red, apple green and crimson respectively. The Be and Mg do not impart color to the flame because the electrons in smaller Be and Mg atoms (2s and 3s orbitals) are more strongly bound and are not excited by the energy of the flame.
- The high negative values of standard reduction potentials indicate that metals of IIA group are good reducing agents comparable to alkali metals. Although the energies required to vaporize and ionize the atoms of IIA group elements to M2+ ions are high, the high lattice energies in the solid salts and high hydration energies of M2+ (Aq) ions compensate for this.
- Calcium , which is a constituent part of a large number of compounds , is the most abundant element in the IIA group.
- Each of these elements has a lot of isotopes which are natural .
- The IIA metal ions have considerably higher ionic charge to ionic radius ratio than group IA elements. As a consequence:
(b) Since the polarizing power of the cations is much greater, the compounds of Be and Mg have appreciable covalent character,
(c) Their compounds with large polarizing anions such as carbonates, bicarbonates, triiodides etc. are less stable.
- The extremely small size of Be2+ makes it so strongly polarizing that Be (II) compounds are almost covalent in nature. The number of covalent bonds can even be increased to four by donation and the use of sp3 hybrid orbitals, provided electron rich ligands are available. For the elements Ca to Ra, although M2+ ions are smaller than the corresponding M+ ions of alkali metals, they have low polarizing power and compounds are largely ionic. Magnesium as M2+ forms compounds, which are ionic as well as covalent in nature.
- The high hydration energy of IIA ions shows that their salts should be more soluble than the corresponding alkali metal salts. But the lattice energy is also greater and the solubility is a delicate balance between hydration energy and lattice energy. Actually, the alkaline earth metal salts are generally less soluble than the corresponding alkali metal salts, because lattice energy, in general, decreases less rapidly than the hydration energy. The hydration energy of the alkaline earth metal ions follows the order Be2+ > Mg2+ > Ca2+ > Ba2+. The solubility of the alkaline earth metal sulphates decreases with increasing cation size. BeSO4 > MgSO4 > CaSO4 > SrSO4 is fairly soluble, but BaSO4 is insoluble. The trend in the solubility of hydroxides is reverse of that of sulphates. Be (OH)2 is least soluble and solubility increases down the group as Be (OH)2 < Mg (OH)2 < Ca (OH)2 < Sr (OH)2 < Ba (OH)2 . The solubility of some halides follows the order:
- BeF2 > MgF2 > CaF2 > SrF2 > BaF2
- BeCl2 < MgCl2 < Ca Cl2 < Sr Cl2 < BaCl2
- CaF2 < CaCl2 < CaBr2 < Cal2 .
- On descending the group, the metal ions increase in size and both the lattice energy as well as hydration energy decrease. If, on descending the group the lattice energy decreases more rapidly than the hydration energy, the compounds become soluble. In case of hydroxides, the lattice energy depends on size of the cation. Be (OH)2 has the largest lattice energy of any group IIA hydroxide. The hydroxides become more soluble because of rapid decrease in lattice energy down the group.
- The lattice energies of the sulphates do not change greatly from BeSO4 to BaSO4, because the anion is such larger than any group IIA cation. The trend in the solubilities of the sulphates is perhaps similar to the trend in hydration energy of the ions. The hydration of the small Be2+ ion is by far the most exothermic of any ion of the group. Thus BeSO4 is the most stable sulphate.
- The thermal stabilities of the carbonates varies directly with the size of the cation. The increasing order of thermal stabilities of the carbonates is BeCO3 < Mg CO3 < Ca CO3 < Sr CO3 < BaCO3. BeCO3 is very unstable mainly due to the enhanced stability of BeO over BeCO3. Cations with large ionic potential have large power of polarization.
- All the alkaline earth metals are highly electropositive (but less electropositive than IA group elements ) as shown by their respective electrode potentials , high chemical reactivities and ionic nature of their compounds . The chemistry of the elements is demonstrated by their tendency to lose their two valence electrons to form cation with two positive charges. These cations Be2+, Mg2+, Ca2+, Sr2+ and Ba2+ are all diamagnetic and have noble gas configuration. The metals are strong reducing agents and combine to form ionic compounds (except Be, which has remarkable polarizing power due to its small size and high charge density).
- The reactivity of the elements increases with increase with increase in atomic number.
- In general, metallic nature, reactivity and reducing power increases from Be to Ba. Tendency to form peroxides increases from Be to Ba. Oxides of these metals become more basic and soluble as we pass from Be to Ba. Hydroxides become more basic and soluble as we go from Be to Ba. Hydration energy increases from Be2+ to Ba2+.
- All the alkaline earth metals react with water, but these reactions are much less violent than the corresponding reaction of an alkali metal. With beryllium and magnesium the reaction with water is slow. In fact, beryllium does not react with water, magnesium decomposes hot water, whereas other elements react with cold water liberating H2 gas and forming the corresponding hydroxide.
- Alkaline earth metals become more reactive as the outer electrons are further remote from the nucleus. Thus barium is the most reactive member of the family, excluding radium, which is radioactive. Be is less electropositive and less and less likely to lose its two outer electrons than the other elements because of its small size. Many of the compounds of Be are not ionic, but polar covalent.
- All the elements react with oxygen when heated in the gas forming the monoxide of the type MO. Be and Mg react less readily because of the formation of a protective oxide film on their surface. Barium also forms some peroxide alongwith monoxide. If the reaction is carried out under pressure, peroxides are formed both by Sr and Ba. On heating in air the elements form a mixture of both oxides and nitrides.
- On heating with nitrogen all the elements of the group form nitrides of the type M3 N2. The nitrides contain N3- ion and are ionic nature.
- With the exception of Be, all alkaline earth metals react directly with hydrogen to form ionic hydrides of the type MH2.
- Since the alkaline earth metal ions have rare gas configuration , they are colorless and diamagnetic.
- All alkaline earth metals form halides of the type MX2. All halides of Be are covalent and electron deficient. Anhydrous beryllium halides are polymeric and contain three central bonds. MgBr2 and MgI2 are soluble in acetone because of their covalent nature.
- The hydroxides may be obtained by slaking the oxides with water. MgO slakes very slowly, but the oxides of Ca, Sr and Ba slake readily.
- BeO as well as Be (OH)2 are amphoteric, Mg(OH) 2 is mild base and its aqueous suspension, known as milk of magnesia is used as an antacid. The Ca(OH) 2 and Sr(OH)2 are moderately strong bases, while Be(OH)2 is almost as strong as the alkali hydroxides. All the oxides, BeO, MgO, CaO, SrO and BaO have NaCl - type structure (4: 4 Coordination).
- Beryllium halide are covalent, hygroscopic and fumes in air because of hydrolysis. BeF2 is one of the few metal fluorides, which does not ionize completely in solution. BeF2 is soluble in water, MgF2 is sparingly soluble, while CaF2 , SrF2 and BaF2 are insoluble . All other halides are soluble. The solubility decreases from BeF2 to BaF2 due to difference in their structures. BeF2 has a linear polymeric chain structure, while others have crystal lattice structures. In the vapor state BeCl2 molecule is linear and has no dipole moment. This corresponds to Be in sp-valenece state. Other halides are hygroscopic and form hydrates of the type MgC2. 6H2O, BaCl2. 2H2O, CaCl2. 6H2O etc.
- The alkaline earth metals generally dissolve in dilute mineral acids, but Be is rendered passive by nitric acid. With acid solutions all of them react to yield the respective salts. Most salts of alkaline earth metals (Ca, Sr, and Ba) such as fluorides, sulphates and phosphates are sparingly soluble in water.
- Beryllium dissolves in cold conc. aqueous alkali evolving hydrogen and forming alkaline earth metal beryllate.
Be + OH- + H2O HBeO3- + H2
- Other elements of the group do not dissolve in alkali.
- Sulphates, carbonates and hydroxides of these metals decompose on heating to yield oxides.
- Nitrates of these metals are water soluble and decompose to oxide, NO2 and O2 gases.
- When Ca, Sr and Ba are dissolved in liquid ammonia, a paramagnetic, highly reducing and highly conducting deep blue colored solution is obtained . These properties are due to the presence of solvated electrons in solution.
- Marie and Pierre Curie discovered radium in 1898. They produced pure radium in 1911. It was the first radioactive element to be discovered. Radium is extracted from pitchblende, a uranium ore. About 7 tonnes of pitchblende are required to produce 1g of pure radium. All isotopes of radium decay spontaneously to other elements, emitting dangerous radiations. The final product of the nuclear disintegration is lead.
- Barium and strontium are found frequently as sulphates, BaSO4 and SrSO4 or the carbonates, BaCO3 and SrCO3. Magnesium occurs in the earth crust as magnesite (MgCO3), dolomite (CaCO3), carnallite (KCl, MgCl2. 6H2O), Epsom salt (MgSO4. 7H2O) and asbestos CaMg3 (SiO3) 4. Magnesium ions are present in seawater to the extent of 2%. Magnesium is also present in chlorophyll.
- Beryllium forms the carbide Be2C while other metals give ionic carbides of the MC2 type. When reacted with water Be2C liberates methane, while other carbides liberate acetylene.Mg2C3 liberates propyne.
- The nitride of Be, (Be3N2) is volatile, while nitrides of other metals are non-volatile.
- The tendency of common salt (NaCl) to get sticky in humid summer is partly due to the presence of small amounts of MgCl2 in it.
- Heating gypsum at 120° C 5°C gives Plaster of Paris (CaSO4. 1/2H2O) which combines vigorously with water and sets to a hard mass. Setting of plaster of Paris is accompanied by evolution of heat (exothermic reaction), catalyzed by common salt (NaCl) and retarded by borax and alum. Setting of plaster of Paris is due to hydration as well as transition. Hardening of plaster of Paris is because of transition of orthorhombic gypsum to monoclinic gypsum. Strength of plaster of Paris on hardening is because of interlacing of needles of monoclinic gypsum On setting, plaster of Paris undergoes expansion .
- Gypsum or alabaster or selenite is calcium sulphate dihydrate, CaSO4. 2H2O. Heating gypsum at 200°C gives dead burnt (CaSO4). Heating gypsum at 400°C gives CaO, O2 and SO2.
- When Ca CO3 is heated to high temperature, it decomposes to form CO2 and CaO, commonly known as lime or quick lime. Quick lime reacts vigorously with water to form a strong base, Ca (OH) 2, which is much less soluble in water than Ba (OH) 2.The formation of Ca (OH)2 is an exothermic reaction . Ca (OH) 2 is known as slaked lime because it is formed when CaO has slaked its thirst for water. When mixed with sand, Ca (OH) 2 hardens as mortar (1 part lime +3parts sand) and cement, by absorbing CO2 from air.
- Barium is not found free in nature. In combined state it occurs as Barytes or heavy spar, BaSO4 and BaCO3. Barium burns in air producing mostly BaSO4. It also burns in CO2 to form BaO and carbon or CO.Barium is strong reducing agent (SRP Ba2+ Ba = -2.90 V ) . It is silvery white metal and m.p. is 850° C . Barium is good conductor of heat and electricity. Barium forms alloys with other metals and these alloys are used in vacuum tubes and sparking plugs, where the low ionization energy of Ba facilitates of good sparks. Barium forms two oxides, BaO and BaO2. BaO2 is peroxide, containing O22- ion.
- Magnesium is used as a de-oxidizer in metallurgy and as a fuse in aluminothermite process. It is also used in flash bulbs, pyrotechnics and in fireworks. Magnalium is an alloy of Al and Mg. Elecktron is an alloy of Mg and Zn.