s-Block Elements : Chemical Properties

CHEMICAL PROPERTIES: –

Due to their low ionisation energies, alkali metals are highly electropositive and chemically reactive. Further since the ionsation energies decrease with increase in atomic number, their reactivity also increases from Li to Cs.

1. Action of Air: Alkali metals are so reactive that they tarnish rapidly when exposed to air because of the formation of oxides, hydroxides and ultimately carbonates at the surface. Hence they are kept under an inert liquid like kerosene oil which prevents them from coming in contact with air and moisture.

               4 Na(s) + O2 (g)     →         2Na2 O(s)

                            Sodium                   Sodium oxide

          Na2 O(s) + H2O (l)      →        2 NaOH(s)

                        Sod. oxide                  Sod. hydroxide

      2 NaOH(s) + CO2 (g)    →          Na2 CO3 + H2O (l)

                                                            Sod. carbonate

2. Oxides: Alkali metals when heated with oxygen form oxides, the nature of which depends upon the nature of the alkali metal.Under ordinary conditions, lithium forms the monoxide (Li2O), sodium forms the peroxide.

(Na2 O2) and the other alkali metals form mainly superoxides (MO2) along with a small amount of peroxides.

                   4Li + O2      →                    2Li2O

                   Lithium monoxide              

                   2Na + O2    →                    Na2O2

                   Sodium peroxide

                    K + O2        →                    KO2­ (Potassium super oxide)

The increasing stability of peroxides and superoxides of alkali metals from Li to Cs is due to the fact that the strong positive field around the small lithium ion attracts the negative charge so strongly that it does not permit the monoxide anion O2– to combine with another oxygen to form peroxide ion O22–. On the other hand, the weak positive field around sodium and potassium ions allows the dispersal of the negative charge of the monoxide anion to form peroxide and superoxide ions.

{ O }^{ 2 }\quad \underrightarrow { \quad 1/2\quad { O }_{ 2 }\quad } \quad { O }_{ 2 }^{ 2- }\quad \underrightarrow { \quad { O }_{ 2 }\quad } \quad 2({ O }^{ 2- })                  

        Monoxide ion     Peroxide ion       Superoxide ion

Structure of the three ions can be represented as below.

{ O }^{ 2- }\quad \quad \overset { \theta }{ O } -\overset { \theta }{ O } \quad \quad \overset { \quad \quad \theta \quad \\ ------\quad }{ O-O }                                                    

          Oxide ion            Peroxide ion       Superoxide ion

Note that the superoxide ion (O2–) has a three-electron bond which makes it paramagnetic and coloured. Yellow colour of sodium peroxide is most probably due to the presence of a small amount of superoxide.

The normal oxides of alkali metals are monoxides, M2O. These are highly soluble in water giving strongly alkaline solution due to the formation of hydroxides, M+ OH–, which are very largely ionised.

                   M2O + H2 O → 2M+ + 2OH–

The higher oxides (peroxides and superoxides) are important oxidising agents. They react with water and dilute acids forming hydrogen peroxide/oxygen.

          Na2O2 + 2H2 O →  2Na+OH– + H2O2

          Na2 O2 + H2 SO4 → Na2 SO­4 + H2O2

          2KO2 + 2H2 O  →   2KOH + H2 O2 + O2

          4KO2 + 2H2 SO4  → 2K2 SO4 + 2H2 O + 3O2

 

3. Action of compounds containing acidic hydrogen, e.g. water, alcohol and acetylene:

Alkali metals react readily and vigorously with water forming hydroxides with the liberation of hydrogen.

2M + 2H2 O→   2MOH + H2

(here M = Li, Na, K, Rb or Cs)

The reactivity with water increases on descending down the group from Li to Cs. It is because of low enthalpy of fusion of heavy members of the family. Heavy Metal melts by the heat produced by the reaction and spread over larger area of water.

                   Cs > Rb > K > Na > Li                  (reactivity with water)

Reaction with K, Rb and Cs is so vigorous that the evolved hydrogen catches fire spontaneously. On account of high reactivity of alkali metals with water, these are not kept in water but stored under kerosene oil.

These metals also react with alcohols forming oxide with the evolution of hydrogen, e.g.

                   2Li + 2C2 H5 OH →          2C2 H5OLi + H2

             Ethyl alcohol                 Lithium ethoxide

2M + HC≡ CH  →   M—C≡C–M  +   H2

M = Na, K etc.             Alkali metal                acetylide

 

4. Reaction with nitrogen: Lithium is the only alkali metal that reacts with nitrogen at room temperature to form lithium nitride.

          6Li(s) + N2 (g)   →   2 Li3N (Lithium nitride)

 

5. Action with hydrogen and Hydrides: Alkali metals react with hydrogen forming ionic hydrides, M+H–. The reactivity of group I metals with hydrogen decreases from Li to Cs.

                   Cs > Rb > K > Na > Li   (Reactivity with hydrogen)

                   2M + H2 → 2MH              (where M = Li, Na, K, etc.)

The alkali metal hydride being highly ionic are attacked readily by water to give back hydrogen.

                   MH + H2O →  MOH + H2

The stability of hydrides being highly ionic readily decreases down the group.

 

6. Action with halogens:

The alkali metals combine readily with halogens forming halides. As the electropositive character increases from top to bottom in the group, the ease of formation of alkali metal halides increases from Li to Cs. However there are some variations.

Alkali metal halides are easily prepared by the direct combination of the elements, M and halogen. They are normally represented by the formula M+X– but Cs and Rb, being of large size, also form polyhalides, e.g., CsI3. (I3– has a linear shape)

Properties of halides of alkali metals: As evident from their following properties, alkali metal halides are ideal ionic compounds.

(i)     All the alkali halides except lithium fluoride, are freely soluble in water (LiF is soluble in non-polar solvents LiCl dissolves in pyridine).

(ii)     They have high melting and boiling points.

(iii)    They are good conductors of electricity in the fused state.

(iv)    They have ionic crystal structure. However, lithium halides have partly covalent character due to polarising power of Li+ ions.

The alkali florides are most stable and then the chlorides, bromides and iodides.

 

7. Solubility in liquid ammonia:

All alkali metals are soluble in liquid ammonia. Dilute alkali metal-ammonia solutions are blue in colour. With increasing concentration of metal in ammonia, the blue colour starts changing to that of metallic copper after which further amount of metal does not dissolve.

The blue solution of an alkali metal ammonia shows certain characteristics which are explained on the basis of formation of ammoniated (solvated) metal cations and ammoniated electrons in the metal ammonia solution in the following way.

                   M     →           M+ + e–

                   M+ + x NH3  →        [M (NH3) x] +

                   e– + y NH3 →  [e (NH3) y]–

                   M + (x + y) NH3   → [M (NH3) x] +    + [e (NH3) y]–

                                                                Ammoniated                 Ammoniated

                                                                metal cation                  electron

The important characteristics of the alkali metal-ammonia solution are as follows.

(A)    Conductivity: The blue solution has high electrical conductivity due to the presence of ammoniated electron present in the cavities formed by the electronic polarisation between the electrons and ammonia molecules. Consequently the metal solution occupies large volume and has lower density than the solvent itself.

(B)    Paramagnetism: The blue solution is paramagnetic. This is again due to the presence of an unpaired electron in the cavities in ammonical solution.

(C)    Colour:  The blue colour of the solution is due to excitation of free electrons to higher energy levels. The absorption of photons takes place in the red region of the spectrum and hence the solution appears blue in the transmitted light.

As the concentration of the alkali metal increases, the metal ion cluster formation takes place and at very high concentration the solution becomes coloured like that of metallic copper.

(D)    Stability: The solution is quite stable and can be considered as dilute metal or an alloy in which the electrons behave essentially as in a metal with the alkali metals slightly apart due to the presence of ammonia molecules. However, in the presence of a catalyst like plantinum black, iron oxide, etc, the solution decomposes to form amide and hydrogen.

                 solution decomposes to form amide and hydrogen.

                  [e (NH3)y]–  →NH2 – + H2 + (y—1) NH3

          Or Simply, 2M + 2NH3 →  2MNH2 + H2

                                                           Metal amide     

(E)     Reducing property: The free ammoniated electrons make the solution a very powerful reducing agent. The ammonical solution of an alkali metal is rather favoured as a reducing agent than its aqueous solution because in aqueous solution the alkali metal being highly electropositive evolves hydrogen from water (thus H2O acts as an oxidising agent) while its solution in ammonia is quite stable, provided no catalyst (transition metal) is present. Some important examples of their reducing action are :

                 (i)      Reduction of metal halides to free metals

                 (ii)     Reduction of sodium nitrite to sodium hydronitrite.

                 (iii)    Removal of halogen atom from alkyl halides.

                          C2 H5 Cl + 2e– → C2H5– + Cl–

    (iv)     Removal of acidic hydrogen atom from acetylenic hydrocarbons.

                         CH3 C≡CH + e–  → CH3 C ≡ C– + H2

(F)    It has very low density

(G)    Its magnetic susceptibility is similar to that of pure metal.

 

8. Oxidation Potential:

Oxidation is a process in which electrons are lost while reduction is a process in which electrons are gained. Hence an oxidising agent is a substance which can accept electrons while reducing agent is a substance which can lose electrons. Evidently, a reducing agent must have low ionisation energy while an oxidising agent must have high ionistion energy (or high electron affinity). Now since ionsation energy decreases on moving down from Li to Cs, the reducing property should increase in the same order, except Li which is found to be the strongest reducing agent.

The tendency of an element to lose an electron is measured by its standard oxidation potential (E0), more the value of E0 of an element stronger will be its reducing character. The given as below.

                 Li/Li+         Na/Na+                K/K+           Rb/Rb+                Cs/Cs+

                 3.05V         2.71V                  2.93V         2 .99 V                3.02 V

The high values of E0 of alkali metals indicate that these are powerful reducing agent and further lithium having the highest value is the strongest of them.

However, it is interesting to note that among the alkali metals, lithium, although, has the highest ionisation energy (i.e., it holds its electron most tightly), yet it is the strongest reducing agent (i.e., loses electron easily). This discrepancy can be explained by recalling and analysing the definition of ionisation energy oxidation potential and high hydration energy of lithium ion (Li+).

Ionisation energy is the energy required to remove an electron from the isolated atom in its gaseous state, i.e.

                   M(g) → M+ (g) + e–

It short, it is the property of an isolated atom in the gaseous state.

Oxidation potential describes the change from the metal in its standard state to the ion in solution, i.e.

                   M(s) → M+ (aq) + e–

The overall process involves three consecutive steps.

          (A)   M(s) →    M(g)            ΔH1 = Sublimation energy

          (B)   M(g)  →   M+ (g) + e–  ΔH2 = Ionisation energy

          (C)  M+(g) + H2O → M+ (aq)      ΔH3 = Hydration energy

     

The first step involves change from solid to gaseous state and energy required to do so is called sublimation energy which is almost equal for all alkali metals. The second step involves

ionisation of the atom in its gaseous state and the energy required for this change is called ionsation energy which is mentioned above, is highest for lithium. The third step involves hydration of gaseous ion which is accompanied by the liberation of energy known as hydration energy. Among alkali metal ions, Li+ is hydrated to the maximum extent, and hence energy released during hydration of Li+ is maximum among alkali metal ions.

Now since oxidation potential is the net effect of all the above three steps, the ionisation energy required in step (B)  is more than compensated by the energy released in the step (C) . This explains the higher oxidation potential of lithium, i.e., the greater ease with which the following overall change takes place.

          Li(s) → Li+ (aq) + e–

Thus the greater reducing power of lithium is due to its large heat of hydration which is due to its small size of its ion and high charge densities (polarizability) on its surface.

 9. Formation of alloys:

The alkali metals form alloys amongst themselves as well as with other metals. The alkali metals dissolve readily in mercury due to the formation of amalgam, the process is highly exothermic.

 

10. Complex formation:

In order to form complex compounds, a metal should have following characteristics.

           (A)      Small size.

           (B)      High nuclear charge.

           (C)      Tendency to accept electrons (i.e. presence of empty orbitals).

Now since the alkali metals have none of these characteristics, they have little tendency to form complexes. However, since lithium has a small size, it forms certain complexes and complex forming tendency falls markedly down the group as the atomic size increases.

 

11. Hydroxides:

Alkali metal hydroxides, MOH, are prepared by dissolving the corresponding oxide in water.

 

Properties of metal hydroxides:  

(i)      These are white crystalline solids, highly soluble in water and alcohols.

(ii)     Since alkali metals are highly electropositive, their hydroxides form the strongest bases known. These dissolve freely in water with the evolution of much heat to give strongly alkaline solution. The strength of a base depends on the ionisation of the hydroxide. As the ionic character increases down the group of hydroxides increases down the group. Hence the basic character of the alkali metal hydroxides increases from LiOH to CsOH.

(iii)    They melt without decomposition and are good conductors of electricity in the fused state.

(iv)    These are stable to heat and do not lose water even at red heat except lithium hydroxide which decomposes on heating.

2LiOH   →  Li2 O + H2 O

(v)     The trends in various properties are shown at below :-

          Solubility \quad \underrightarrow { \begin{matrix} LiOH\quad \quad NaOH\quad \quad KOH\quad \quad RbOH\quad \quad CsOH \\ Increase \end{matrix} }

          Basic Strength \overrightarrow { \quad \quad \quad increase\quad \quad \quad }  

          Thermal Stability \overrightarrow { \quad \quad \quad increase\quad \quad \quad }       

 

12. Carbonates and Bicarbonates:

The 1st group elements form carbonates and bicarbonates of formula M2CO3 and MHCO3 respectively.

All carbonates are highly stable towards heat except Li2 CO3 which decomposes to form Li2O and CO2 . The bicarbonates on the other hand decompose to form carbonates carbondioxide and water

            2NaHCO3  \underrightarrow { \quad \quad \triangle \quad \quad }      Na2 CO3 + H2O + CO2

            2LiHCO3   \underrightarrow { \quad \quad \triangle \quad \quad }    Li2CO3 + H2O + CO2

The carbonates and bicarbonates are soluble in water.

The different trends in the properties are shown below :-

         Li2CO3          Na2CO3  K2CO3,       Rb2CO3,     Cs2CO3.

        Stability (thermal)    \underrightarrow { \quad \quad increase\quad \quad }   

        Solubility   \underrightarrow { \quad \quad increase\quad \quad }                   

        Solubility & Stability   \underrightarrow { \quad \quad LiHC{ O }_{ 3 }\quad \quad NaHC{ O }_{ 3 }\quad KHC{ O }_{ 3 }\quad RbHC{ O }_{ 3 }\quad CsHC{ O }_{ 3 }\quad \quad }

 

13. Nitrates:

The nitrates can all be prepared by action of HNO3 on the corresponding carbonates or hydroxides. All the nitrates are highly soluble in water LiNO3 and NaNO3 are highly deliquescent. Hence KNO3 is used in gun powder. The nitrates generally do not decompose on heating. They are highly stable. However NaNO3 and KNO3 decompose at 250ºC to form NaNO2 + O2 or KNO2 + O2. However at very high temperatures NaNO3 decomposes to form Na2O + NO2.