Title chemistry narayana.material.pdf ✅

SYNOPSIS Introduction  The elements in which the last electron enters the outermost s-orbital are called s-block elements  As the s-orbital can accommodate only two electrons, so s-block contains two groups (IA & IIA)  Group IA of the periodic table consists of the elements Lithium, Sodium, Potassium, Rubidium, Caesium and Francium  They are collectively known as the alkali metals because they form hydroxides on reaction with water which are strongly alkaline in nature  Among the alkali metals sodium and Potassium are most abundant while Lithium, Rubidium and Caesium have much lower abundances.  Francium is highly radioactive, it is short-lived Isotope 223Fr. Halflife of this Isotope is only 21 minutes Physical Characteristics 1.Electronic Configuration and Occurence  The outer electronic configuration of alkali metals is ns1 with one valence electron in outermost shell. Element Symbol Electronic Configuration Lithium Li(3) [He]2s 1 Sodium Na(11) [Ne]3s1 Potassium K(19) [Ar]4s1 Rubidium Rb(37) [Kr]5s1 Caesium Cs(55) [Xe]6s1 Francium Fr(87) [Rn]7s1  The loosely held s-electron in the outermost valence shell of these elements makes them most electropositive metals.  They readily lose electron to give monovalent M+ ions. They are not found in free state in nature.  The alkali metals are so reactive they cannot be displaced by another element, so they are isolated by electrolysis of their molten salts. 2. Atomic and Ionic radii The alkali metal atoms have the largest sizes in their respective periods of the periodic table.  With increase 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 increases on moving down the group. Li Na > K > Rb > Cs  The second ionization energies are very high, they cannot form divalent ions. Thus, alkali metals are univalent only and form ionic compounds. 4. Hydration Enthalpy:  The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.  The size of alkali metal ions in aqueous solution is in the following order. Li+ >Na+ > K+ > Rb+ > Cs+  Li+ has maximum degree of hydration. So, Lithium salts are mostly hydrated, eg: LiCl.2H2O  Ionic mobility and conductance in aqueous solution is in the order Cs+ > Rb+ > K+ > Na+ > Li+  Alkali metals are silvery white, soft and light metals.  These elements have low density which increases down the group from Li to Cs due to their large size.  Potassium (K) is lighter than Sodium (Na) because the increase in atomic size of Potassium is more compared to it’s mass as it contains vacant ‘d’ shell 6. Melting and Boiling Points  The melting and boiling points of the alkali metals are low indicating weak metallic bonding due to the presence of a single valence electron and large size. Decreasing order of melting and boiling points. Li > Na > K > Rb > Cs 7. Flame Colours  The alkali metals and their salts impart characteristic colour to an oxidizing flame. GROUP 1 ELEMENTS : ALKALI METALS
 Heat from the flame excites the outermost electron to a higher energy level and excited electron comes back to the ground state by the emission of radiation in the visible region. Alkali metals can be detected from flame tests and determine by flame photometry Li-Crimson red, Na-yellow, K- violet, Rb- red violet, Cs - blue  These elements when irradiated with light, the light energy absorbed may be sufficient to make an atom lose electron. This porperty makes caesium and potassium useful as electrodes in photoelectric cells. Chemical Characteristics (i) Reaction With Air  Reactivity order is Li Na K Rb Cs      The alkali metals tarnish in dry air due to formation of their oxides 2 2 4 2 Li O Li O   4 lim 2 Na O ited Na O  2    2 2Na O Excess Na O  2    2 2 The superoxide ion ( 1 O2  ) is stable only in the presence of large cations such as K, Rb and Cs M O Excess MO M K Rb Cs  2     2  , ,   In all oxides, alkali metals show +1 oxidation state  Lithium burns in air form oxide and nitride ,Li3 N.  Due to high reactivity with air and water, alkali metals are (except lithium) kept in kerosene  The solubility and basic strength of oxides increase in the order: Li O Na O K O Rb O Cs O 2     2 2 2 2  The stability of peroxides and superoxides increases in the order: Na O K O Rb O Cs O 2 2 2 2 2 2 2 2    and KO RbO CsO 2 2 2   (ii) Reaction With Water  The reactivity increases down the group Li < Na < K < Rb < Cs  The alkali metals react with water to form hydroxides and dihydrogen 2 2 2 2 2 M H O MOH H     (M=Li, Na, K, Rb, Cs)  Lithium has lowest SRP value but it is less reactive with water than that of sodium because of its small size and very high hydration energy  Other Alkalimetals reacts explosively with water and they also react with proton donors like alcohol, carboxylic acid, ammonia and 1- alkynes (iii) Reaction With Hydrogen  Reactivity of alkali metals towards hydrogen increases down the group Li < Na < K < Rb < Cs  Li reacts with hydrogen at 1073K to form covalent hydride ,remaining alkali metals form ionic hydrides at 673K 2 2 2 M H MH   (M=Li, Na, K, Rb, Cs)  Alkali metal hydrides are ionic solids with high melting points.  The order of ionic nature of alkali metal hydrides LiH NaH KH RbH CsH     (IV) Reaction With Halogens  Reactivity of alkali metal towards a particular halogen increases in the order Li < Na < K < Rb < Cs  Reactivity of halogens towards particular alkali metal decreases in the order F Cl Br I 2 2 2 2     The alkali metal reacts with halogens to forms ionic halides M X     Lithium halides are covalent due to the high polarisation ability of lithium ion. Li ion has high tendency to distort electron cloud of halide ion due to it’s small size (covalent nature is  polarising power).  Polarisability of anions increase with increase of size  Among the halides, Lithium iodide is most covalent in nature. Reducing Nature  Alkali metals are strong reducing agents.  Li is strong reducing agent and sodium is weak reducing agent.  Small size lithium has highest hydration enthalpy and it has high negtive SRP value so it has high reducing power. M M     s  g sublimation enthalpy M M e     g g     ionization enthalpy M H O M   g 2   aq     hydration enthalpy
Solutions in Liquid Ammonia  Alkali metals dissolve in liquid ammonia  Alkali metal atom readily lose the valence electron in ammonia to form solution.  The ammoniated electrons absorbs energy in visible region and imparts blue colour to solution.   3     3 3 + x y Ammoniated Ammoniated cation electron M x y NH M NH e NH               The blue colour solution of alkali metals in ammonia is paramagnetic and on standing slowly liberate hydrogen resulting in the formation of amide. 3 2 2 min 2 2 2 Metala e M NH MNH H     In concentrated solution, (above 3M) the blue colour changes to bronze colour on warming and becomes diamagnetic Uses : (i) White metal - Li+ Pb (ii) It is used in thermonuclear reactions (iii) Li is used in electrochemical cells. (iv) Sodium is used to make tetra ethyl lead Pb(Et)4 and tetra methyl lead Pb(Me)4 (v) Liquid sodium metal is used as coolant in fast breeder nuclear reacters. (vi) Potassium has vital role in Biological system (vii) KCl is used as fertilizer (viii) Caesium is used in photo electric cells General Characteristics of the Compounds of the Alkali Metals (i) Oxides and Hydroxides  All the alkali metal on exposure to air or oxygen burn vigorously, forming oxides on the surface of the metals.  Lithium forms mainly monoxide (Li2O) and some peroxide (Li2O2 ).  Sodium forms the peroxide Na O2 2  and some superoxide.  Other elements form superoxides MO M K 2 : , Rb,Cs  .  The stability of above oxides is based on the fact that a small cation can stabilize a small anion and a large cation can stabilize a large anion.  The oxides and the peroxides are colourless when pure, but the superoxides are yellow or organe in colour and paramagnetic due to presence of unpaired electron in superoxide ion.  Sodium peroxide (oxone) is used as oxidising agent in purifying air because it release oxygen when reacts with carbondioxide. 2 2 2 2 3 2 2 2 2 Na O CO Na CO O     The oxides are easily hydrolysed by water to form the hydroxides 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 M O H O MOH M O H O MOH H O MO H O MOH H O O           The alkali metal hydroxides are strongest of all bases and dissolve in water with evolution of heat due to high enthalpy of hydration. Halides:  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. 2 2 2 2 3 2 2 2 2 2 2 M O HX MX H O MOH HX MX H O M CO HX MX CO H O           (M=Li, Na, K, Rb or Cs) (X=F, Cl, Br or I)  For given halide: LiX NaX KX RbX CsX     (Increasing ionic character)  For given alkali metal: MF MCl MBr MI    (Increasing covalent character)  The alkali metal halides are high melting, colourless crystalline solids.  All of these halides have high negative enthalpies of formation:  The f H   values of fluorides become less negative as we go down the group.  The f H   for chlorides, bromides and iodides become more negative as we go down the group  The order of melting and boiling points of MX MF > MCl > MBr > MI (M=Alkalimetal)  All the halides are soluble in polar solvents like water.  LiF is less soluble in water due to it’s high lattice energy and CsI has low solubility due to smaller hydration enthalpy of it’s two ions.
 Lithium halides are soluble in ethanol, acetone, ethylacetate & pyridine. Salts of Oxo-Acids Alkali metals form strong basic hydroxides and oxides. So they form salts with all oxoacids and hydroacids.  They are generally soluble in water. The oxides solubility increases from top to bottom.  They are thermally stable. Thermal stability increases from top to bottom.  Li2 CO3 decomposes readily. Li CO Li O CO 2 3 2 2     Li2 CO3 decomposes readily because Li+ has greater polarising power or polarising ability.  Lithium does not form solid bicar bonates. It exist in liqued state.  Aqueous solutions of carbonates and bicarbonates are basic in nature due to anionic hydrolysis 2 3 2 2 3 3 2 2 3 CO H O H CO OH 2 2 HCO H O H CO OH            The solubility of carbonates, nitrates and bicarbonates increase in the order: 2 3 2 3 2 3 2 3 2 3 Li CO Na CO K CO Rb CO Cs CO     , LiNO NaNO KNO RbNO CsNO 3     3 3 3 3 and LiHCO NaHCO KHCO RbHCO CsHCO 3     3 3 3 3 Anomalous Properties of Li :  Lithium is much harder. Its m.p. and b.p. are higher than the other alkali metals.  Lithium is least reactive but the strongest reducing agent among all the alkali metals.  On combustion in air it forms mainly monoxide, Li2O and the nitride, Li3 N unlike other alkali metals.  LiCl is deliquescent and crystallises as a hydrate, LiCl.2H2O whereas other alkali metal chlorides do not form such a hydrates.  Lithium hydrogencabonate is not obtained in the solid form while all other elements form solid hydrogencarbonates.  Lithium does not form ethynides.  Lithium nitrate when heated gives lithium oxide, Li2O, whereas other alkali metal nitrates decompose to give the corresponding nitrite. 3 2 2 2 3 2 2 4 2 4 2 2 LiNO Li O NO O NaNO NaNO O       LiF and Li2O are comparatively much less soluble in water than the corresponding compounds of other alkali metals.  Alkali metal carbonates do not decompose on heating except Lithium. Diagonal Relationship Between Lithium and Magnesium  Similarity between lithium and magnesium arises because of their (a) Similar sizes (b) Similar electro negativity (c) Similar polarising power. The similarities are  Both lithium and magnesium are harder and lighter than other elements in the respective groups.  Lithium and magnesium react slowly with water.  Their oxides and hydroxides are much less soluble and their hydroxides decompose on heating. Both form a nitrides, Li3 N and Mg3 N2 , by direct combination with nitrogen.  Both Li and Mg give only monoxides Li2O, MgO and they do not combine with excess oxygen to give any superoxide.  The carbonates of lithium and magnesium decompose easily on heating to form the oxides and CO2 . Solid hydrogencarbonates are not formed by lithium and magnesium  Both LiCl and MgCl2 are soluble in ethanol  Both LiCl and MgCl2 are deliquescent and crystallise from aqueous solution as hydrates. LiCl.2H2O and MgCl2 .8H2O Sodium Carbonate Na2 CO3 : Soda ash Na2 CO3 . 10H2O : Salt soda or washing soda Preparation  Sodium carbonate is prepared commonly by ammonia-soda process (or) Solvay process.  Raw materials: Brine solution, lime stone and ammonia.  Principle: Low solubility of sodium hydrogencarbonate  The raw materials reacts in the following manner     3 2 2 4 3 2 4 3 2 2 2 4 3 4 3 4 3 2 2 NH H O CO NH CO NH CO H O CO NH HCO NH HCO NaCl NH Cl NaHCO           Sodium hydrogencarbonate crystal separates. These are heated to given sodium carbonate 3 2 3 2 2 2NaHCO Na CO CO H O   