cbse-class-11-chemistry-notes-chapter-10-the-s-block-elements.pdf

Revision Notes Chapter - 11 Chemistry Chapter – 10 The s – Block Elements Alkali Metals (Group 1) They have ns1 electronic configuration and are highly reactive metals Elements...

Revision Notes Chapter - 11 Chemistry Chapter – 10 The s – Block Elements Alkali Metals (Group 1) They have ns1 electronic configuration and are highly reactive metals Elements Atomic Number Electronic Configuration Lithium 3 [ He] 2 s1 Sodium 11 [ Ne] 3s1 Potassium 19 [ Ar ] 4s1 Rubidium 37 [ Kr ] 5s1 Cesium 55 [ Xe] 6s1 Francium 87 [ Rn] 7s1 Physical Properties 1. Atomic Size In their respective times, the atoms are at their largest. As you progress through the group, the atomic size grows larger. 2. Oxidation State Group 1 elements has +1 oxidation state. 3. Density Alkali metals have low density due to their large size. Atomic mass Density  Atomic volume Class XI Chemistry www.vedantu.com 1  The group's atomic weight increases from Li to Cs , as does its volume, but the atomic weight increase outweighs the volume increase. As a result, density rises from Li to Cs 4. Exception: Density of sodium is more than that of potassium 5. Order: Li  K  Na  Rb  Cs 6. Nature of Bonds When the values of electronegativity are low, they combine with other elements to form Ionic bonds. 7. Ionization Energy The atoms in this group have lower initial ionisation energy than any other group in the periodic table. Because atoms are so big, the outer electron is only weakly bound by the nucleus, resulting in a low ionisation energy. As you move down the group, the ionisation energy drops. 8. Flame Test When alkali metals are burned on a flame, the electrons in the valence shell migrate from a lower energy level to a higher energy level due to heat absorption from the flame. When they return to their original condition, they release the additional energy in the form of visible light, which gives the flame colour. Element Colour Li Red Na Golden yellow K Violet Rb Red Violet Cs Blue 9. Standard Oxidation Potential The electrode potential of a metal in water is a measurement of its tendency to donate electrons. The standard electrode potential is defined as the concentration of metal ions being equal to one. Lithium has the largest Class XI Chemistry www.vedantu.com 2  ionisation potential, but due to its high hydration energy, it also has the highest electrode potential. 10.Hydration of Ions The ions have a lot of water in them. The degree of hydration is proportional to the size of the ion. As a result, from Li+ to Cs+, the degree of hydration falls. As a result, electrical conductivity diminishes as hydration increases. 11.Lattice Energy Ionic solids are alkali metal salts. The lattice energy of alkali metal salts with a common anion drops as one moves down the group. 12.Solubility in Liquid Ammonia M  nNH 3 [ M ( NH 3 ) x ]  e  ( NH 3 ) y (n  x  y ) The major species found in solvated metal ions and solvated electrons in dilute alkali metal solutions in liquid ammonia are solvated metal ions and solvated electrons. The colour diminishes until it disappears if the blue solution is left to stand due to the creation of metal amide. Because of the presence of solvated electrons, metal solutions in liquid conduct electricity. Because they include free electrons, the dilute solutions are paramagnetic. 13.Electronegativity Values The electronegativity values are small which decrease from lithium to cesium. 14.Reactivity The reactivity of alkali metals goes on increasing in the following order: Li  Na  K  Rb  Cs 15.Colourless and Diamagnetic Ions The number of unpaired electrons present in an ion determines whether the ion is colourless or coloured. If an anion has unpaired electrons, these electrons can be stimulated by light energy and subsequently return to the ground state to show colour. Unpaired electron ions have magnetic properties, while paired electron ions cancel out each other's magnetic fields. Diamagnetic ions are such ions. The presence of unpaired electrons causes super oxides to be para magnetic and coloured. 16.Melting and Boiling Point Class XI Chemistry www.vedantu.com 3  The cohesive energy is the force that holds the atoms or ions in a solid together. The cohesive energy is proportional to the number of electrons capable of bonding. Alkali metals contain only one valence electron that participates in bonding, and the outer bonding electron is big and diffuse, therefore the cohesive force reduces as the group gets smaller. As the atoms get bigger as you go down the group, the bonds get weaker, the cohesive energy drops, and the metal gets softer. As a result, the melting point drops as the group progresses. The boiling point also reduces the size of the group. Chemical Properties Some common reactions of Group 1 metals Reaction Comment M  H 2O MOH  H 2 Hydroxides are strongest base known. Li  O2 Li2O Monoxide formed by lithium and to a small extent by sodium. Na  O2 Na2O2 Peroxide formed by sodium and to a small extent by lithium. K  O2 KO2 Superoxide formed by potassium, rubidium and cesium. M  H 2 MH Ionic salt like hydrides. Li  N2 Li3 N Nitride formed only by lithium M  S M 2S All metals form sulphides. M  X 2 MX All metals form halides. M  NH3 MNH 2 All the metals form amides.. Class XI Chemistry www.vedantu.com 4 Reaction with Air Group 1 elements are very reactive and tarnish quickly when exposed to air. These metals form alkaline carbonates in moist air. 2 Na  O2 2 Na2O Na2O  H 2O 2NaOH 2NaOH  CO2 Na2CO3  H 2O Reaction with O2 Lithium forms Li2O , sodium forms two types of oxide (M 2O, M 2O2 ) and potassium, rubidium and cesium form superoxide ( MO2 ). Basic Nature, Ionic Nature of the Oxides Because the size of the cation increases, the basic nature of oxides changes from lithium to cesium. (i) From lithium to cesium, the cation size increases. The ionic nature of these oxides rises from lithium to cesium, according to Fajan's Rule. (ii) As the ionic nature of these metal oxides changes, solubility in water increases from lithium to cesium oxides. Reaction with water Group 1 metals react with water and liberates hydrogen and thus hydroxides are formed. 2Li  2H2O 2LiOH  H2 2Na  2H2O 2NaOH  H2 Class XI Chemistry www.vedantu.com 5 2 K  2H2O 2KOH  H2 Reaction with Hydrogen Group 1 metals react with hydrogen and forms ionic hydrides. Thermal stability of LiH is high. Stability of hydrides is in the order: LiH  NaH  KH  RbH  CsH Reaction with Dilute acids These metals react quickly with dilute acids due to their alkaline nature, and the rate of reaction increases from lithium to cesium as the basic character increases. Compounds of Alkali Metal Hydroxides Caustic soda is another name for sodium hydroxide. Because of its corrosive qualities, potassium hydroxide is known as caustic potash. The strongest base in aqueous solution is caustic alkali. The solubility of hydroxides rises as they progress through the group. The bases react with acids to form salt and water KOH  HCl KCl  H2O NaOH  HCl NaCl  H2O 2NaOH  CO2 Na 2CO3  H2O Ammonia is liberated by the bases from ammonium salts NaOH  NH4Cl NH3  NaCl  H2O Class XI Chemistry www.vedantu.com 6 KOH  NH4Cl NH3  KCl  H2O In all of its reactions, potassium hydroxide is similar to sodium hydroxide. However, because potassium hydroxide is so much more expensive, it is rarely utilised. However, because potassium hydroxide is more soluble in alcohol, the equilibrium produces C 2 H 5O  ions. C2 H 5OH  OH  C 2 H 5O   H 2 O Oxides, Peroxides and Superoxides Normal oxides – monoxide: Ionic monoxides are present. They are very basic oxides that create strong bases when they react with water. Na 2O  H2O NaOH K2O  H3O KOH Peroxides Preparation:  2Na  O 2 (excess)  300 C Na 2O 2 7400 C 2Na 2 O  Na 2 O 2  Na(vapour ) Properties Na 2O2  H2SO4 (dil) Na 2SO4  H2O2 Na 2O2  2H2O 2NaOH  H2O2 Na 2O2 is a powerful oxidant as it reacts with carbon dioxides present in the air. Na 2O2  CO Na 2CO3 Na 2O2  2CO2 Na 2CO3  O2 Class XI Chemistry www.vedantu.com 7 Na 2 O 2  Cr 3 CrO 24 Structure Oxygen atom is sp3 hybridized. Peroxide ion has 18 electrons which occupies the molecular orbital as shown:  1s 2 ,  *1s 2 ,  2 s 2 ,  * 2 s 2 ,  2p 2z ,  2p 2y   2p 2z ,  * 2p 2y   * 2Pz1 The bond order being 1 so, it is diamagnetic. Superoxides Superoxides are ionic oxides M  O 2 Preparation M  O 2 ( excess ) MO2 (M  K, Rb,Cs) Superoxides are stronger oxidizing agents than peroxides. The stability of these superoxides is in the order: KO2  RbO2  CsO2 Reactions KO2  H2O KOH  H2O2  1/ 2O2 Because it creates O2 and eliminates CO 2 , KO 2 is utilised in space capsules, submarines, and breathing masks. 4KO2  2CO2 2 K2CO3  3O2 4KO2  4CO2  2H2O 4KHCO3  O2 Class XI Chemistry www.vedantu.com 8 Sodium superoxide can be made by reacting sodium peroxide with oxygen at high temperatures and pressures, rather than by burning metal in oxygen. Na 2O  O2 2NaO2 Structure The paramagnetic property is explained by the existence of one unpaired electron in a three-electron bond. The superoxide molecule has 17 electrons and a bond order of1.5, occupying the molecular orbitals as illustrated.  1s 2 ,  *1s 2 ,  2 s 2 ,  * 2 s 2 ,  2p 2x ,  2p 2y   2p z2 ,  * 2p 2y   * 2Pz1 The stability of oxides is given as: Normal oxide > peroxide > superoxide Carbonates and Bicarbonates Solid bicarbonates are formed by Group 1 metals  MHCO3 . M 2 CO3 carbonates are formed by all alkali metals. The carbonates and bicarbonates of alkali metals are extremely heat stable due to their electro positive character \left(Li2CO3 \right decomposes easily by heat). The unusual behaviour of Li 2 CO3 can be explained by the fact that: (a) Lithium's small size and high polarisation disrupts the electron cloud of the nearby oxygen atom of the massive CO32 , weakening the carbon-hydrogen bond. Li 2 CO3  Li 2 O  CO 2 (b) When a larger carbonate ion is replaced by a smaller carbonate ion, the lattice energy increases, favouring breakdown. M 2 CO3  M 2 O  CO 2  Class XI Chemistry www.vedantu.com 9 As a washing soda, Na 2CO3 is utilised. Baking soda is made from NaHCO3. Both NaHCO3 and KHCO3 have hydrogen bonds in their crystal structures. The HCO3 in NaHCO3 forms an endless chain, whereas KHCO3 forms a dimeric anion. Reactions 2HNO3  K2CO3 2KNO3  CO2  H2O  Na 2 CO3  H 2 O  CO 2 2NaHCO3   M 2 CO3  H 2 O 2M   HCO3  OH  Halides All of the metals in this group create MX halides. Because lithium ion is the smallest ion in the group, it is more likely than other metals to produce hydrated salts. Properties Alkali metal halides are excellent ionic compounds, as evidenced by the following features. (i) With the exception of lithium fluoride, all alkali halides are easily soluble in water (Lithium fluoride is soluble in non-polar solvents). (ii) Their melting and boiling points are extremely high. (a) The melting and boiling points of the same alkali metal drop in a predictable order. Fluoride > chloride > bromide > iodide This is described in terms of the metal halides' lattice energy*. The lattice energy of the same metal reduces when the halogen's electronegativity lowers. (b) The melting point of lithium halides is lower than that of sodium halides for the identical halide ion. However, as we travel down the group from sodium to cesium, the melting points of halides decline. Lithium halides Class XI Chemistry www.vedantu.com 10  exhibit aberrant behaviour due to their covalent character, whereas sodium and other halides are ionic in nature. As we advance along the group of ionic halides, the melting point lowers as the lattice energy decreases. NaCl  KCl  RbCl  CsCl (iii) Solubility of halides of alkali metals: Alkali metal halides have a range of solubilities. The solubility of alkali metal fluorides in water, for example, gradually increases from lithium to caesium. Lithium chloride has a far better solubility in water than sodium chloride when it comes to chlorides. This is owing to the lithium ion's tiny size and high hydration energy. However, when the lattice energy of the crystals decreases, solubility in water increases steadily from sodium chloride to cesium chloride. (iv) In the fused condition, they are good conductors of electricity. (v) They are made up of ionic crystals. Lithium halides, on the other hand, have a partially covalent character due to the polarising power of lithium ions. (vi) The lattice energy and polarising power are responsible for the structure and stability (solubility) of alkali metal halides. (a) Lattice Energy: The energy produced during the production of a crystal lattice from gaseous cations and anions, or the energy necessary to split one mole of a solid ionic compound into its gaseous ions, is known as lattice energy. As a result, lattice energy (the force of attraction between ions) is a direct measure of ionic crystal stability; the higher the lattice energy of a molecule, the lower its solubility in water. When an ionic compound crystal comes into contact with a polar solvent like water, the water molecule's hydrogen end (positive pole) is attracted to a negative ion, whereas the oxygen end (negative pole) is drawn to a positive ion. Solvation (or hydration, if the solvent is water) of the ions is the process of polar solvent molecules attaching to the ions. When ions are stabilised by solvation, a considerable amount of solvation energy (or hydration energy) is released, which, if it exceeds the crystal's lattice energy, causes the ionic compound to dissolve in the solvent. In the case of lithium fluoride, however, if the solvation energy is insufficient to oppose the lattice energy, the material stays insoluble. The combination of small lithium ions and small fluoride ions gives lithium fluoride its Class XI Chemistry www.vedantu.com 11 high lattice energy. The lattice energy of a particular ion increases as the size of the oppositely charged ion decreases. (b) Polarising power and polarisability (Fajan’s Rule): Although an ionic bond in a chemical like M X is thought to be 100 percent ionic, it is discovered to have significant covalent character in some circumstances (e.g., lithium halides). When two oppositely charged ions approach each other, the nature of the link between them, according to Fajan, is determined by the action of one ion on the other. When two oppositely charged ions come into contact, the positive ion attracts electrons from the anion's outermost shell while repelling the anion's positively charged nucleus. The anion is distorted, deformed, or polarised as a result of this. The ability of a cation to distort an anion is called polarisation power, and the anion's susceptibility to be polarised by the cation is called polarisability. If the degree of polarisation is modest, an ionic bond is established; however, if the degree of polarisation is considerable, electrons are attracted from the anion to the cation by electrostatic attraction, resulting in a higher electron density between the two ions and a covalent connection. In general, the stronger an ion's polarisation power or polarisability, the greater its proclivity to form covalent bonds. Because polarisation power grows as the size of the cation decreases and polarisability increases as the size of the anion increases, the polarisation in a compound containing big negative ions and small positive ions may be so strong that the connection becomes covalent. In nature, lithium iodide, which is made up of lithium ions (the smallest alkali metal ion) and iodide ions (the largest halide ion), is found to be highly covalent. Other examples of such ionic-covalent compounds are AlCl3 , FeCl3 ,SnCl4 , etc. Reactions Ionic polyhalides are formed when alkali metal halides react with halogen and interhalogen compounds. KI  I 2 K  I3  KBr  ICl K[BrICl] Class XI Chemistry www.vedantu.com 12 KF  BrF3 K  BrF4  Sulphates They form sulphates of type M 2SO 4 Anomalous Behaviour of Lithium Although lithium has many of the properties of the group I elements, it also differs from them in a number of ways. The incredibly small size of the lithium atom and its ion causes this unusual behaviour. The high charge density of the lithium ion is due to its small size. As a result, out of all the alkali metal ions, lithium ion has the most polarising power. As a result, it has a significant distorting impact on a negative ion. As a result, the lithium ion has a strong proclivity for solvation and the creation of covalent bonds. It's also worth noting that the polarising power of the lithium ion is similar to that of the magnesium ion, making the two elements resemble very much in their properties. (i) Lithium is substantially more difficult to work with than the other elements in Group I. (similarity with magnesium which is also a hard metal). (ii) Lithium has a high melting point and boiling point. (iii) Lithium, unlike the other elements in this group, is the least reactive, as shown by the following points. (a) It is not influenced by air, unlike others. (b) It, unlike others, decomposes water slowly (similar to magnesium). (c) It rarely reacts with bromine, unlike many others. (d) On burning in oxygen, it forms only the monoxide Li 2 O , while the others form peroxides too. (iv) Unlike other elements, it forms nitride when it comes into contact with nitrogen, Li3 N (similarity with Mg ) (v) Lithium is much less electropositive and, therefore, several of its compounds  Li 2 CO3 and LiOH  are less stable (similarity with Mg). For example, 2LiOH Li2O  H2O Mg(OH)2 MgO  H2O Class XI Chemistry www.vedantu.com 13  (vi) When heated, lithium nitrate produces nitrogen dioxide and oxygen, leaving lithium oxide behind (similar to magnesium nitrate), but sodium and potassium nitrates produce only oxygen, leaving nitrites. (vii) The majority of lithium salts (for example, hydroxide, carbonate, oxalate, phosphate, and fluoride) are water insoluble (similarity with magnesium.) The sodium and potassium salts that correspond to them are both water soluble. (viii) Lithium halides and lithium alkyls are soluble in organic solvents, but sodium and potassium halides and alkyls are not; MgCl2 is soluble in alcohol as well. (ix) Lithium chloride, like magnesium chloride, undergoes some hydrolysis in hot water, but only to a little level; sodium chloride and potassium chloride do not. (x) Lithium sulphate, unlike sulphates of other alkali metals, does not create alums. (xi) Partially covalent lithium compounds, particularly lithium halides, exist in nature. This is owing to lithium ions' proclivity for attracting electrons (polarising power). This explains why lithium compounds have a smaller dipole moment than expected. (xii) The ions and compounds of alkali metals are more hydrated than those of other alkali metals (similarity with magnesium). Extraction of Sodium Sodium is obtained on large scale by two process: Castner’s process The electrolysis of fused sodium hydroxide takes place at 330 C , with iron as the cathode and nickel as the anode. 2NaOH 2Na   2OH At cathode: 2 Na   2e 2Na At anode: 4OH  2H 2O  O 2  4e  Class XI Chemistry www.vedantu.com 14 Oxygen and water are created during electrolysis. At the cathode, water generated at the anode is partially evaporated and partially broken down, resulting in hydrogen discharge. H 2O H   OH  At cathode: 2H   2e 2H H 2  Down’s Process Nowadays, the metal is produced using Down's technique. It includes employing iron as a cathode and graphite as an anode to electrolyze fused sodium chloride containing calcium chloride and potassium fluoride at roughly 600 degrees Celsius. The cell is made up of a steel tank with heat-resistant bricks lining it. In the centre of the cell, a circular graphite anode is installed, which is encircled by a cylindrical iron cathode. A steel gauze cylinder separates the anode and cathode, allowing fused charge to pass through. A dome-shaped steel hood protects the anode and provides an outlet for chlorine gas to escape. The molten metal that has been freed at the cathode rises and flows into the kerosene receiver. Reactions: NaCl Na   Cl At Cathode: 2Na   2e 2Na At Anode: 2Cl Cl2  2e  Class XI Chemistry www.vedantu.com 15 The sodium chloride when obtained from this method is 99.5% pure. The electrolysis of pure sodium chloride has a number of drawbacks: (i) Sodium chloride has a high fusion temperature of 803 degree celsius, which is difficult to sustain. (ii) Because sodium is volatile at this temperature, some of it vapourizes, forming a metallic fog. (iii) The electrolysis products, salt and chlorine, are corrosive at this temperature and may harm the cell's substance. Pure sodium chloride is blended with calcium chloride and potassium fluoride to overcome the problems mentioned above. At the voltage used, calcium chloride and potassium fluoride do not breakdown, but they do lower the fusion temperature. A combination containing 40% sodium chloride, 60% calcium chloride, and a trace amount of potassium fluoride reaches a fusion temperature of around 600 degrees Celsius. In the electrolytic cell, this combination is electrolyzed at 600 degrees Celsius. Example 1: Alkali metals are paramagnetic but their salts are diamagnetic. Explain. Solution: The outermost energy shell in metals is solely occupied, whereas in cations, all orbitals are double occupied (inert gas configuration). e.g., Na,1s2 , 2 s2 2p6 ,3s2 3p6 , 4s 1 paramagnetic   Na  1s 2 , 2 s 2 2p 6 , 3 s 2 3p 6 Diamagnetic Example 2: Alkali metals are good reducing agents. Explain. Solution: Alkali metals are strong reducing agents because their low ionisation enthalpy values and high oxidation potential allow them to easily lose valence electrons. Example 3: Which alkali metal ion has the maximum polarising power and why? Solution: Among the alkali metal ions, the lithium ion has the most polarising power. This is owing to the lithium ion's small size. Class XI Chemistry www.vedantu.com 16 Example 4: Lithium ion is far smaller than other alkali metal ions but it moves through a solution less rapidly that the others. Explain. OR The conductance of lithium salts is less in comparison to the salts of other alkali metals. Explain. Solution: Lithium ion has the highest degree of hydration because of its strong charge, which pulls numerous water molecules surrounding it. As a result, the size of the hydrated lithium ion is larger than that of the other alkali metal ions, affecting its mobility in solution and lowering conductance. Size: [Li(aq)]  [Na(aq)]  [K(aq)] Example 5: Sodium salts in aqueous solutions are either neutral or alkaline in nature. Explain. Solution: Strong acids or weak acids produce the anions in sodium salts. There is no hydrolysis when anions come from strong acids, and aqueous solutions are neutral. When anions come from weak acids, however, they undergo hydrolysis, resulting in alkaline solutions. Solns. of sodium carbonate or bicarbonate, for example, are alkaline. CO32  H 2 O HCO3  OH  HCO3  H 2 O H 2CO3  OH  Example 6: Why do potassium, rubidium and cesium form superoxides in preference to oxides and peroxides on being heated in excess supply of air? Solution: K  , Rb  and Cs are large cations in size and superoxide ion  O 2  is larger in size in comparison to oxide  O 2  and peroxide  O 22  ion. These metals form superoxides rather than oxides or peroxides because a larger cation can stabilise a large anion. Example 7: Why is KO 2 paramagnetic? Solution: The superoxide O2 is paramagnetic because of one unpaired electron in  *2p molecular orbital. Class XI Chemistry www.vedantu.com 17 KK (2 s) 2  * (2 s) 2   2p x    2p x   2p y   *  2p x   *  2p v  2 2 2 2 1 Example 8: Among the alkali metals which element has (i) Highest melting point (ii) Highest size of hydrated ion in solution (iii) Strongest reducing agent in solution (iv) Least electronegative Solution: The elements which have: (i) Highest melting point: Lithium (ii) Highest size of hydrated ion in solution: [Li(aq)] (iii) Strongest reducing agent in solution: Lithium (iv) Least electronegative: Cesium Example 9: What happens when following compounds are heated? (a) Li2CO3 Li2O  CO2 (b) Na2CO3.10H2O Na2CO3  10H2O (c) 4LiNO3 2Li2O  4NO2  O2 (d) 2 NaNO3 2NaNO2  O2 Example 10: (a) Arrange LiF, NaF, KF, RbF and CsF in order of increasing lattice energy. (b) Arrange the following in order of the increasing covalent character. MCl, MBr, MF, MI (where M  alkali metal) Solution: (a) CsF  RbF  KF  NaF  LiF (b) MF  MCl  MBr  MI Class XI Chemistry www.vedantu.com 18  With increasing size of the anion, covalent character increases. Example 11: Why a standard solution of sodium hydroxide cannot be prepared by weighing ? Solution: The material sodium hydroxide is deliquescent. It absorbs moisture and reacts with atmospheric carbon dioxide, both of which increase its mass. As a result, precise weighing is difficult. Example 12: What happens when: (a) Fused sodium reacts with dry ammonia. (b) Sodium hydrogen carbonate is heated. (c) Sodium hydroxide is heated with sulphur? Solution: (a) 2Na  2NH3  2NaNH2  H2 (b) Sodium carbonate is formed. 2NaHCO3 Na 2CO3  H2O  CO2 (c) Sodium thiosulphate is formed. 4 s  6NaOH Na 2 S2O3  2Na 2 S  3H2O Example 13: Give reasons for the following: (i) LiCl is more covalent than NaCl (ii) Lithium Iodide has lower melting point then LiCl (iii) MgCl2 is more covalent than NaCl (iv) CuCl is more covalent than NaCl Solution: (i) Lithium ion is more polarising than sodium ion due to its smaller size, and so lithium chloride is more covalent than sodium chloride. (ii) Due to bigger size, I  is more polarisable than Cl and hence lithium iodide is more covalent than lithium chloride. Therefore, lithium iodide has lower melting point than LiCl. Class XI Chemistry www.vedantu.com 19  (iii) Magnesium ion is more polarising than sodium ion due to its greater charge, and so magnesium chloride is more covalent than sodium chloride. (iv) Copper ion is more polarising than sodium ion because to the pseudo inert gas configuration, and so copper chloride is more covalent than sodium chloride. Example 14: Identify ( A),( B),(C) and ( D) and give their chemical formulae. (a) (A)  NaOH  Heat NaCl  NH 3  H 2O (b) NH3  CO2  H2O (B) (c) (B)  NaCl (C)  NH4Cl (d) (C)  Heat Na 2 CO3  H 2 O  (D) Identify (A), (B), (C) and (D) and give their chemical formulae. Solution: (a) NH 4 Cl  NaOH  Heat NH 3  NaCl  H 2 O Compound A is ammonium chloride  NH 4 Cl . (b) NH3  CO 2  H 2O  NH 4 HCO 3 Compound B is ammonium bicarbonate  NH 4 HCO3 . (c) NH 4 HCO3  NaCl  NaHCO3  NH 4 Cl Compound (C) is sodium bicarbonate  NaHCO3 . (d) 2NaHCO3  Na 2CO3  H 2O  CO 2 Compound (D) is carbon dioxide  CO 2 . Class XI Chemistry www.vedantu.com 20 Example 15: Arrange the following as specified: (i) MgO,SrO, K 2 O and Cs2O (increasing order of basic character) (ii) LiCl , LiBr , LiI (decreasing order of covalent character) (iii) NaHCO3 , KHCO3 , Mg  HCO3 2 , Ca  HCO 3 2 (decreasing solubility in water) (iv) LiF, NaF, RbF, KF and CsF (in order of increasing lattice energy) (v) Li, NaK (in order to decreasing reducing nature in solution) Solution: (i) MgO  SrO  K 2O  Cs2O } (ii) LiI  LiBr  LiCl (iii) NaHCO3  KHCO3  Mg  HCO3 2  Ca  HCO3 2 (iv) CsF  RbF  KF  NaF  LiF (v) Li  K  Na Example 16: (a) What happens when KO 2 reacts with water? Give the balanced chemical equation. (b) Predict giving reason the outcome of the reaction : LiI  KF  Solution: (a) When KO 2 reacts with water, oxygen is evolved and an alkaline solution containing potassium hydroxide and H2O2 is formed: 2KO 2  2H 2 O  2KOH  H 2 O 2  O 2 (b) Lithium iodide reacts with potassium fluoride and anions are exchanged in this process: Class XI Chemistry www.vedantu.com 21  LiI  KF  LiF  KI As stable compounds form, the exchange takes place, with the bigger cation stabilising the larger anion and the smaller cation stabilising the smaller anion. Alkaline Earth Metals Introduction Group-II of the periodic table contains the elements beryllium, magnesium, calcium, strontium, barium, and radium. All of these substances are metals. Calcium, strontium, and barium oxides were discovered far before the metals themselves, and they were dubbed alkaline earths because they were alkaline and found in the earth. Alkaline earth metals were given to the elements once they were found. Radium shares chemical properties with alkaline earth metals, but because it is a radioactive element, it is researched independently from the other radioactive elements. Physical Properties Atomic Size On going down the group, atomic size of elements increases. Oxidation State +2 oxidation state is exhibited by group II elements. Density Group II elements are smaller in size than group I elements, hence they have a higher density than group I elements. From beryllium to radon, density rises. Exception: Calcium has a lower density than magnesium, whereas magnesium has a lower density than beryllium. Nature of Bonds Class XI Chemistry www.vedantu.com 22 Beryllium forms mainly covalent compound. The rest of the elements in group II form ionic bonds. Hydration energy Because of their smaller size and higher charge, the hydration energies of group 2 ions are four to five times higher than those of group 1 ions. As the size of the ions grows larger, the hydration enthalpy falls. Lattice Energy The lattice energy of alkali metal salts with a common anion drops as one moves down the group. Ionization Energy Because the atoms in group 2 are smaller, the electrons are more firmly bound, requiring more energy to remove the initial electron (first ionisation energy) than in group 1. The amount of energy necessary to remove the second electron is approximately double that required to remove the first. As a result, the energy necessary to make divalent ions from group 2 elements is four times that required to make M from group 1 metals. Flame Test These elements' electrons are stimulated to higher energy levels when energy is delivered to them in a flame, as is the case with alkali metals under identical conditions. The excess energy generated by the electrons when they return to their original energy level produces visible light with distinct colours, as shown below: Element Colour Calcium Brick red Strontium Crimson red Barium Grassy green Radium Crimson Class XI Chemistry www.vedantu.com 23 The atoms of beryllium and magnesium are smaller. As a result, the electrons in these atoms are more tightly bound. As a result, the energy of the flame does not stimulate them to higher energy levels. As a result, these ingredients do not produce any colour in the bunsen flame. Standard Oxidation Potential Standard Oxidation Potential of Alkaline Earth Metals Element Oxidation Reaction Standard Oxidation Potential (volt) Be Be Be2  2e 1.85 Mg Mg Mg2  2e 2.37 Ca Ca Ca 2  2e 2.87 Sr Sr Sr 2  2e 2.89 Ba Ba Ba 2  2e 2.90 Solubility in Liquid ammonia The metals, like group 1 metals, dissolve in liquid ammonia. Due to the creation of solvated electrons, dilute solutions are blue in colour. As the solution decomposes, amides form and hydrogen gas is released. 2NH 3  2e  2NH 2  H 2 Electronegative Values Class XI Chemistry www.vedantu.com 24 Group II element’s electronegativity values are low, although they are higher than group I element’s. The value of electronegativity diminishes as you progress through the group. Colourless and Diamagnetism The elements of the alkaline earth metal group generate M2 ions, which are diamagnetic and colourless due to the lack of an unpaired electron. Melting and Boiling Point The melting point of elements in group II lowers as the cohesive force diminishes as the group progresses. Exception: Magnesium has the lowest melting point Metallic Properties Group-II elements have typical metallic characteristics. They have a nice metallic sheen and excellent electrical and thermal conductivity. GROUP – I and II Oxides Sodium Oxide Preparation (i) It's made by heating sodium to 180 degrees Celsius in a small amount of air or oxygen and then distilling the surplus sodium away. 180 2Na  1/ 2O 2  Na 2 O (ii) By heating sodium peroxide, nitrate or nitrate with sodium. Na2O2  2 Na 2Na2O Class XI Chemistry www.vedantu.com 25  2 NaNO3  10 Na 6 Na2O  N2 2 NaNO2  6 Na 4 Na2O  N2 Properties (i) It is a white amorphous mass. (ii) It gets decomposed at 400 degree Celsius into sodium peroxide and sodium. 400 2Na 2 O  Na 2 O 2  2Na (iii) It gets dissolved violently in water and yields caustic soda. Na 2 O  H 2 O  2NaOH Sodium Peroxides Preparation: It is formed by heating the metal in excess of air or oxygen at 300 degrees Celsius in a dry, carbon dioxide-free environment. 2Na  O 2  Na 2 O 2 Properties (i) It is a pale yellow solid, becoming white in air from the formation of a film of NaOH and Na 2 CO3. (ii) In cold water, sodium peroxide produces H2O2 but at room temperature it produces oxygen. Sodium peroxide produces hydrogen peroxide in ice- cold mineral acids. 0 C Na 2 O 2  2H 2 O   2NaOH  H 2 O 2 5 C 2Na 2 O 2  2H 2O 2  4NaOH  O 2 C Na 2 O 2  H 2SO 4 0 Na 2SO 4  H 2O 2 Class XI Chemistry www.vedantu.com 26  (iii) It combines with carbon dioxide to produce sodium carbonate and oxygen, which is why it's used to purify air in small spaces. Example: submarine, ill-ventilated room. 2Na 2 O 2  2CO 2  2Na 2CO3  O 2 (iv) It is an oxidising agent and oxidises charcoal, CO, NH 3 , SO 2 3Na 2 O 2  2C  2Na 2CO3  2Na Na 2 O 2  CO  Na 2CO3 Na 2 O 2  SO 2  Na 2SO 4 3Na 2 O 2  2NH 3   6NaOH  N 2 (v) It contains peroxide ion Uses: (i) For preparing H2O2 ,O2 (ii) Oxygenating the air in submarines (iii) Oxidising agent in the laboratory. Oxides of Potassium: Column K 2O White K 2 O2 White K 2O3 Red KO 2 Bright Yellow KO3 Reddish Brown Needles Preparation: Class XI Chemistry www.vedantu.com 27  (i) 2 KNO3  10 K  heating 6 K 2O  N 2 K 2 O  heating K 2O K2O  H 2O 2KOH Controlled (ii) 2 K  O 2 air at 300 C K 2O 2 (iii) Passage of O2 through a blue solution of K in liquid NH 3 yields oxides K 2O2 (white), K 2 O 3 ( red $)$ and KO 2 (deep yellow) K in liq. NH 3  K 2 O 2  K 2 O3  KO 2 2KO 2  2H 2 O  2KOH  H 2 O 2  O 2 Magnesium Oxide: It's also known as magnesia, and it's made from natural magnesite that's been heated. MgCO3  MgO  CO 2 Properties (i) It is present as a white powder (ii) The melting point of magnesium oxide is 2850 degree Celsius. So, it is used in the manufacturing of refractory bricks for furnances. (iii) It imparts alkaline reaction and it is very slightly soluble in water. Calcium Oxide It is manufactured by dissolving lime stone at a high temperature of roughly 1000 degrees Celsius and is known as fast lime or lime. CaCO3 CaO  CO2  4200cal Properties Class XI Chemistry www.vedantu.com 28  (i) It is a white amorphous powder having a melting point of 2570 degree Celsius. (ii) When heated in an oxygen-hydrogen flame, it produces strong light (lime light). (iii) It is a basic oxide that reacts with acidic oxides, such as sulphur dioxide. CaO  SiO 2  CaSiO3 CaO  CO 2  CaCO3 (iv) On combination with water, it produces slaked lime CaO  H 2 O  Ca(OH) 2 Magnesium Peroxide and Calcium Peroxide: These are obtained by passing H2O2 in a suspension of Mg(OH)2 and Ca(OH)2 Uses: Magnesium peroxide is a whitening agent and an antimicrobial in tooth paste. HYDROXIDES Sodium Hydroxides Preparation (i) Electrolysis of Brine: NaCl Na   Cl At Anode: 2Cl  Cl 2  2e At Cathode: Na  e  Na 2Na  2H 2 O  2NaOH  H 2 Class XI Chemistry www.vedantu.com 29  (ii) Caustication of Na 2 CO3 (Gossage's method): Na 2CO3  Ca(OH)2 2NaOH  CaCO3 Since the K sp  CaCO3   K sp  Ca(OH) 2  , the reaction shifts towards right. Properties (i) It's a crystalline white solid that's highly corrosive and deliquescent. (ii) It is resistant to heat. (iii) Its aqueous solution has an alkaline pH and feels soapy to the touch. (iv) FeCl3  3NaOH  Fe(OH)3  3NaCl NH4Cl  NaOH  NaCl  NH3  H2O ZnCl2  2NaOH  Zn(OH),  2NaCl Zn(OH),  2NaOHExcess  Na1ZnO 2  2H, O (v) Acidic and amphoteric oxidise gets dissolved easily, e.g: CO 2  2NaOH  Na 2CO3  H 2O Al2 O3  2NaOH  2NaAlO 2  H 2O (vi) Aluminium and Zn metal gives H 2 from NaOH 2Al  2NaOH  2H 2O  3H 2  2NaAlO 2 (vii) Several non metals such as phosphorous, sulphur, calcium etc. yield a hydride instead of hydrogen e.g. 4P  3NaOH  3H 2O  PH 3  3NaH 2 PO 2 Potassium Hydroxide Class XI Chemistry www.vedantu.com 30 Preparation: It is prepared by the electrolysis of aqueous solution of potassium chloride. Properties (a) It is stronger base compared to sodium hydroxide (b) Potassium hydroxide is more soluble in water as compared to sodium hydroxide. (c) In alcohol, NaOH is sparingly soluble but KOH is highly soluble. (d) As a reagent KOH is less frequently used but in absorption of CO2 , KOH is preferably used compared to NaOH. Because KHCO3 formed is soluble whereas NaHCO3 is insoluble. Magnesium Hydroxide It is found as the mineral brucite in nature. Preparation: It's made by combining a caustic soda solution with a magnesium sulphate or chloride solution. MgSO 4  2NaOH  Na 2SO 4  Mg(OH) 2 Properties (i) It can only be dried at temperatures up to 100 degrees Celsius; otherwise, at greater temperatures, it will break down into its oxide. Mg(OH) 2  MgO  H 2O (ii) It is alkalinizing because it is mildly soluble in water. (iii) It dissolves in NH4Cl solution Mg(OH) 2  2NH 4Cl  MgCl 2  2NH 4OH Thus, Mg(OH)2 is not therefore precipitated from a solution of Mg2 ions by NH4OH in presence of excess of NH4Cl Class XI Chemistry www.vedantu.com 31 Calcium Hydroxide Preparation: It can be prepared easily by spraying water on quicklime. CaO  H 2 O  Ca(OH) 2 Properties (i) The solubility in water of calcium hydroxide is very less. (ii) It has a lower solubility in hot water than in cold water. When a result, as the temperature rises, so does the solubility. (iii) Carbon dioxide is readily absorbed by calcium hydroxide and is used as a test for the gas. CARBONATES Preparation (i) Leblanc Process: NaCl  H2SO4 (conc.)  mild heating NaHSO 4  HCl Strongly NaCl  NaHSO 4 heated   Na 2SO 4  HCl Na 2SO 4  4C  Na 2 S  4CO  Na 2 S  CaCO3  Na 2CO3  CaS (ii) Solvay Process: NH 3  H 2 O  CO 2  NH 4 HCO3 NaCl  NH 4 HCO3  NaHCO3  NH 4 Cl 150 C 2NaHCO3   Na 2 CO3  H 2 O  CO 2 Properties Class XI Chemistry www.vedantu.com 32  (i) Soda ash is anhydrous sodium carbonate that does not disintegrate when heated but melts at 852 degrees Celsius. (ii) Sodium carbonate forms a number of hydrates. (iii) Hydrated Na 2 CO3 is called washing soda  Na 2CO3 10 H 2 O  and is prepared by Le Blanc process solvay process and electrolytic process. (iv) Sodium carbonate absorbs carbon dioxide and produces sodium bicarbonate, which can be calcined to obtain pure sodium carbonate at 250 degrees Celsius. (Tex translation failed) (v) It was causticized by lime after being dissolved in acid with effervescence of carbon dioxide. Na 2 CO3  HCl  2NaCl  H 2O  CO 2 Na 2 CO3  Ca(OH) 2  2NaOH  CaCO3 Uses: Sodium carbonate is widely used as a smelter in glass making Potassium Carbonate It can be made by the leblanc method, but not by the solvay process since potassium carbonate is water soluble. Properties: It resembles with Na 2 CO3 , its melting point is 900 C but a mixture of Na 2 CO 3 and K 2 CO 3 melts at 712 C. Uses: Potassium carbonate is used in glass manufacturing. Calcium Carbonate Marble, limestone, chalk, and calcite are examples of natural calcite. It's made by dissolving marble or limestone in HCl , eliminating any iron or aluminium, precipitating with NH 3 , and adding  NH 4 2 CO3 to the solution. CaCl2   NH 4 2 CO3  CaCO3  2NH 4Cl Class XI Chemistry www.vedantu.com 33 Properties (i) It is dissociates above 1000 degree Celsius as follows: CaCO3  CaO  CO 2 (ii) It dissolves in carbon dioxide-containing water to produce calcium bicarbonate, but boiling separates it from the solution. H 2 O  CO 2 Ca  HCO3 2 CaCO3  boiling Magnesium Carbonate It is found in nature as magnesite, which is isomorphic to calcite. It is precipitated as a white by adding sodium bicarbonate to a solution of a magnesium salt, although only basic carbonate, known as magnesia alba, with the approximate composition MgCO3.Mg(OH)2.3H2O , is precipitated. Properties: The properties of magnesium carbonate is same as calcium carbonate. BICARBONATES Sodium Bicarbonates Preparation: By absorption of CO 2 in Na 2 CO3 solution. Na 2CO3  H2O  CO2 2NaHCO3 Uses: It is used in pharmaceutical industries and also as baking powder. Potassium bicarbonates Preparation: Same as NaHCO3 Properties: Same as NaHCO3 But it is more alkaline and more soluble in water compared to NaHCO3. Class XI Chemistry www.vedantu.com 34 Magnesium Bicarbonate boiling MgCO3  CO 2  H 2O Mg  HCO3 2 Calcium Bicarbonate CaCO3  CO 2  H 2O Ca  HCO3 2 CHLORIDES Sodium Chloride It is prepared by the method of brine which contains 25% sodium chloride. Properties (i) It is nonhygroscopic, whereas ordinary salt contains magnesium chloride, which makes it hygroscopic. (ii) It's used to make freezing mixture in the lab [freezing mixture is ice- common salt mixture with a temperature of-25 degrees Celsius]. (iii) To melt ice and snow from the road. Potassium Chloride It also occurs in nature as sylvite (KCl) or carnallite KCl.MgCl2.6H2O. Uses: Potassium chloride is used as a fertiliser. Magnesium Chloride Preparation: It is prepared by dissolving MgCO3 in dilute hydrochloric acid. MgCO3  2HCl  MgCl 2  H 2O  CO 2 Properties (i) It crystallises as hexahydrate. MgCl2  6H 2O. Class XI Chemistry www.vedantu.com 35  (ii) It is deliquescent solid. (iii) This hydrate undergoes hydrolysis as follows : MgCl2  H 2 O  Mg(OH)Cl  HCl Mg(OH)Cl  MgO  HCl (iv) Anhydrous MgCl2 can be prepared by heating a double salt like. MgClO2.NH4Cl.6H2O as follows:  H 2O MgCl2.NH 4Cl2.6 H 2O  MgCl2 ; NH 4Cl  strong  MgCl2  NH 3  HCl Sorel Cement: It's a paste-like mixture of magnesium oxide and magnesium chloride that hardens when left to stand. This is utilised in dental fillings, flooring, and other applications. Calcium Chloride (i) In the solvay process, it is a by-product. (ii) The carbonate can alternatively be made by dissolving it with hydrochloric acid. CaCO3  2HCl  CaCl 2  H 2O  CO 2 Properties (i) It's made out of deliquescent crystals. (ii) It gets hydrolysed like MgCl2 hence anhydrous CaCl2 cannot be prepared. CaCl2  H2O CaO  2HCl (iii) Anhydrous CaCl2 is used in drying gases and organic compounds but not NH or alcohol due to the formation of CaCl2.8NH3 and CaCl2  4C2H5OH SULPHATES Class XI Chemistry www.vedantu.com 36 Sodium Sulphate Preparation It is created by heating common salt with sulphuric acid in the first step of the leblanc process. 2NaCl  H 2SO 4  Na 2SO 4  2HCl Thus the salt cake formed is crystallised out from its aqueous solution as Na 2SO4.10H2O. This is called as Glauber's salt. Properties Sodium sulphate is reduced to Na 2S when it is fused with carbon. Na 2SO 4  4C  Na 2S  4CO Uses: It is used in pharmaceutical industries. Potassium Sulphate It's found in stassfurt potash deposits as schonite and kainite, and it's made by dissolving it in water and crystallising it. It separates as crystals from the solution, whereas Na 2SO 4 is a decahydrate. Magnesium Sulphate Preparation (i) It's made by dissolving kieserite MgSO4.H2O in boiling water and then crystallising the resulting hepta hydrate solution. Epsom salt is what it's called. (ii) It is also obtained by dissolving magnesite in hot dilute H2SO4 MgCO3  H 2SO 4  MgSO 4  H 2O  CO 2 (iii) It is isomorphous with FeSO4  7H2O, ZnSO4  7H2O Class XI Chemistry www.vedantu.com 37  Calcium Sulphate It occurs as anhydrite CaSO4 and as the dihydrate CaSO 4.2H 2O , gypsum, alabaster or satinspar. Properties (i) Calcium sulphate solubility increases till a certain point and then declines as temperature rises. (ii) Because of its porous nature, plaster of Paris is utilised in the construction of wood. Example 17: (a) Mg 3 N 2 when reacted with water, gives off NH 3. but HCl is not obtained from MgCl on reaction with water at room temperature. Why? (b) The crystalline salts of alkaline earth metals contain more water of crystallization than corresponding alkali metal salts. Why? Solution: (a) Because Mg 3 N 2 is a salt of a strong base and a weak acid  NH 3  , it can be hydrolyzed. Mg 3 N 2  6H 2 O  3Mg(OH) 2  2NH 3 (b) In comparison to alkali metal ions, alkaline earth metal ions have a stronger tendency to hydrate due to their small size and high nuclear charge. As a result, alkaline earth metal salts contain more crystallisation water than alkali metal salts. Example 18: What happens when: (i) Beryllium carbide reacts with water. (ii) Magnesium nitrate is heated. (iii) Quick lime is heated in electric furnace with powdered coke. (iv) Sodium chloride solution is added to zinc chloride solution. Class XI Chemistry www.vedantu.com 38 Solution: (i) A gas named methane is evolved Be 2 C  4H 2 O  2Be(OH) 2  CH 4 (ii) A gas which is brown in colour, NO2 , is evolved 2MgNO3  2MgO  4NO 2  O 2 (iii) Calcium carbide is formed CaO  3C  Electric Arc CaC2  CO (iv) A white precipitate of zinc hydroxide is evolved which forms sodium zincate on dissolving with excess of sodium hydroxide. ZnCl2  2NaOH  Zn(OH) 2  2NaCl Zn(OH) 2  2NaOH  Na 2 ZnO 2  2H 2 O Example 19: Draw a structure of: (i) BeCl2 in vapour state (ii) BeCl2 in solid state Solution: It features a chlorobridged dimer structural bond in the vapour state. At 1000 degrees Celsius, it dissociates into a linear monomer. In the solid state, it has a polymeric structure with chlorobridges, in which a halogen atom connected to one beryllium atom forms a coordinate bond with a lone pair of electrons and a covalent binding with another atom. Class XI Chemistry www.vedantu.com 39 Class XI Chemistry www.vedantu.com 40