1. Alkali Metals

2.  Alkali metals refer to six elements belonging to the Group IA of the long form of the Modern Periodic Table, viz. Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs) and Francium (Fr). Fr is a radioactive element.  These elements are called so because they form strongly alkaline oxides and hydroxides. Introduction

3. Electronic Configuration of Alkali Metals  Each of the alkali metals has one electron in their outermost (valence) shell, which is just outside an inert gas core. Element Atomic No. Electronic StructureValence Shell Electron Li 3He core, 2s 1 i.e. 1s 2 2s 1 2s 1 Na11Ne core, 3s 1 i.e. 1s 2 2s 2 2p 6 3s 1 3s 1 K19Ar core, 4s 1 i.e. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 4s 1 Rb37 Kr core, 5s 1 i.e. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 5s 1 5s 1 Cs55 Xe core, 6s 1 i.e. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 6 6s 1 6s 1 Fr87 Rn core, 7s 1 i.e. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 6 5d 10 6s 2 6p 6 7s 1 7s 1  The solitary electron in the outermost shell can be easily removed from the atom and it is the factor that is largely responsible for the chemistry alkali metals.

4. Physical Properties 1.Physical state: Aalkali metals are extremely soft and readily fused. They are highly malleable (can be pressed out into sheets) and ductile (can be drawn into wires). When freshly cut, they have a bright luster which is quickly tarnished as soon as metal comes in contact with atmosphere. 2.Density: The densities of alkali metals are quite low. Li, Na and K are lighter than water. The densities increase with the increase in atomic number from Li to Cs. However, K is lighter than Na. 3.Melting and boiling points: The melting and boiling points of alkali metals are very low, which progressively decrease when we move downward from Li to Cs. 4.Ionization energies: Alkali metals have low ionization energies, which go on decreasing from Li to Cs. 5.Electropositive character: Alkali metals show strong electropositive or metallic character, which progressively increases from Li to Cs. 6.Photoelectric effect (effect of light): The ns 1 electron in alkali metal atoms is so loosely held with the nucleus that it is ejected from the surface of these metals by (low energy photons of) ordinary lights. The high luster of the surface of these metals is due to this photoelectric effect.

5. 7.Reducing properties: Alkali metals act as strong reducing agents. This reducing power increases on going down from Li to Cs. But in aqueous condition, Li atom has the maximum reducing power. 8.Colouration to the flame: Alkali metal impart characteristic colour to flame. e.g. Li — Crimson red, Na — Golden yellow, K, Rb and Cs — Violet. 9.Electrical conductivity: Alkali metals show high electrical conductivity. 10.Oxidation states: All the alkali metals show +1 oxidation state. 11.Formation of ionic compounds: Alkali metals form ionic compounds with highly electronegative elements. The compounds are colourless unless the anions be coloured e.g. permanganates and dichromates. 12.Hydration of ions:  Alkali metals are extensively hydrated. M + (g) + aq → [M(aq)] + Hydrated cation  The degree of hydration of M + ions decreases on moving from Li + to Cs +. 13.Water solubility of alkali metal salts: Virtually all salts of alkali metals are water soluble ---- more soluble than the salts of any other periodic group.

6. Chemical Properties 1.Reactivity: Alkali metals are highly reactive especially towards highly electronegative elements like Cl 2 and O 2 with which they give ionic compounds. The reactivity increases with the increase of their atomic number from Li to Cs. 2.Reaction with O 2 or air – formation of oxides:  All the alkali metals react with O 2 of the air to form ordinary oxides, as well as higher oxides. The affinity and tendency for higher oxide formation increases with the atomic weight.  Li forms monoxide (Li 2 O), Na forms peroxide (Na 2 O 2 ) and K, Rb and Cs form superoxides (MO 2 ) 4 Li + O 2 → 2 Li 2 O 2 Li + O 2 → Na 2 O 2 M + O 2 → MO 2 (M=K, Rb, Cs)  Freshly prepared or cut alkali metals have a silvery white surface, but on exposure to air they get tarnished because of formation of a oxide coat on their surface.

7. 3.Reaction with H 2 O – formation of hydroxide:  All the alkali metals react vigorously with H 2 O to form hydrogen gas and the metallic hydroxide and, for this reason, must be stored under kerosene, coated with paraffin, or protected in some other fashion. The activity increases as the atomic weight increases. 2M + 2H 2 O → 2MOH + H 2  All the alkali metal hydroxides are alkaline in aqueous solution with the alkalinity naturally increasing on going down the group. 4.Reaction with H 2 – formation of hydride: The alkali metals, on heating, react quite readily with hydrogen to give white crystalline salt-like hydrides. The tendency to form the hydrides decreases from Li to Cs. 2M + H 2 → 2MH 5.Reaction with halogens – formation of halides: Alkali metals combine directly with halogens to form halides. The ease with which the alkali metals form halides increases from Li to Cs. 2M + X 2 → 2MXX = F, Cl, Br, I

8. 6.Reaction with N 2 :  Only Li reacts readily with N 2 at room temperature to form lithium nitride (Li 3 N), which decomposes in water to form ammonia and lithium hydroxide. 6Li + N 2 → 2Li 3 N  Other alkali metals also give the reaction on heating especially in a discharge tube. 7.Reaction with acids: Alkali metals react violently with dilute acids. 2M + 2H + → 2M + + H 2 8.Action of Liquid ammonia: Alkali metals dissolve in liquid ammonia to form a blue solution (without the evolution of H 2 gas). But in presence of catalytic amounts of oxidizing agents (e.g. as impurity), ammonia will rapidly react with the metals to form alkali amides (e.g., NaNH 2 ) Catalyst 2M + 2NH 3 2MNH 2 + H 2 ↑ 9.Action of mercury – formation of amalgams: Alkali metals dissolve readily in mercury and form amalgams. This reaction is very exothermic. 10.Formation of complexes: Generally speaking, alkali metals do not form complexes except that the ions form solvates with H 2 O and NH 3.

9. Some Important Compounds of Alkali Metals Lithium (Li): Lithium fluoride (LiF), Lithium chloride (LiCl), Lithium nitride (Li 3 N), Lithium carbonate (Li 2 CO 3 ), Lithium- aluminium hydride (LiAlH 4 ). Sodium (Na): Sodium peroxide (Na 2 O 2 ), Sodium hydroxide or Caustic soda (NaOH), Anhydrous sodium carbonate or Soda ash (Na 2 CO 3 ), Hydrated sodium carbonate or Washing soda (Na 2 CO 3.10H 2 O), Sodium bicarbonate or Baking soda (NaHCO 3 ), Sodium chloride or Common salt (NaCl), Sodium hypochlorite (NaOCl), Sodium nitrite (NaNO 2 ), Sodamide (NaNH 2 ), Sodium cyanide (NaCN). Potassium (K): Potassium hydroxide or Caustic potash (KOH), Potassium chloride (KCl), Potassium bromide (KBr), Potassium iodide (KI), Potassium cyanide (KCN)

10. Uses & Applications of Alkali Metals Lithium  Lithium carbonate USP and Lithium citrate USP are used in the treatment of hypomanic and manic states.  LiHCO 3 (Lithia water) is used treat gout. It expels uric acid from human system. Uric acid + LiHCO3 → Li-urate (soluble in water) Sodium  Sodium is the principal cation of extracellular fluids.  Sodium compounds are widely used in pharmacy and medicine.  Following are some important sodium compounds and their uses:  Sodium acetate NF: Diuretic, Systemic alkalizer, Systemic antacid.  Sodium bicarbonate USP: Antacid, Buffering agent, For carbonation to (increase palatability)  Sodium bisulfite USP: Reducing agent  Sodium borate USP: Used externally for eye-washes and as wet dressing for wounds.  Sodium chloride USP: Electrolyte therapy, Preparation of isotonic medium.

11. Potassium Following are some official compounds of Potassium and their uses:  Potassium acetate NF: Diuretic and urinary alkalizer.  Potassium bitartrate NF: Mental excitement (insomnia, nervousness etc.)  Potassium chloride USP: Electrolyte replenisher  Potassium citrate NF: Diuretic, expectorant. Rubidium & Cesium Neither Rubidium nor Cesium finds application in pharmacy and medicine at the moment.

12. References  Madan  Haider  Block et. al.

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