Gp I – Alkali Metals

 

  • Lithium shows different properties due to its extremely small size and lesser metallic character so as it shows significance similarity to Mg and other group II , which is a general trend and can be assumed to be a general perception.
  • Due to high reactivity they do not occur in free state in nature the different sources Lithium occurs a silicates [K2Al2O3.6SiO2] where as sodium is found as salt in sea water NaNO3, Borax [Na2B4O7. 10H20] and Mirabilite [Na2SO4] etc. The potassium is in form of sylvine KCl, carnalities [KCl.MgCl2.6H2O] and Felspar etc.
  • Rb and Cs are at low abundance, while Fr is rare as radioactive.

The PHYSICAL PROPERTIES of the group are as under

Softness increases down the group in this group, the reason can be given by a logic the kernel size get bigger and bigger down the group and so the softness increases, this also gives a vibe that the density increases  the trend is important because of the single exception which is

      Li < K < Na < Rb < Cs < Fr (density) 

Melting point decreases down the group this is also because the metallic bonding is getting weaker and weaker as we move down the group. Now the next is the Hydration enthalpy, since charge is same on all ions the inverse relation of size is the deciding factor so,

Hydration enthalpy decreases down the group. Now some properties which are decided by the enthalpy itself is hydrated ionic radius follows,

Li+ > K+ > Na+ > Rb+ > Cs+ > Fr(hydrated ionic radius)  now since the hydrated size decreases down the group the mobility is more of a smaller ion so  mobility increases down the group.

Oxidation potential most of the redox reactions take place in aqueous medium and gp I metal are good reducing agents, while Li is strongest reducing agent in aq. Medium while we don’t have a very accurate order are trend is irregular. Keeping this point in mind will help in resolving reactions with lithium compounds.

Flame Test –                  Li                            Na                          K                       Rb                        Cs                               Crimson red          golden yellow        pale violet        red violet            sky blue

The flame test is mostly asked either directly or to identify the element present in compound, while in case of heating the Na the electron do a transaction of  3s à 2s and releases photon whose  wavelength lies in the visible range causing the metal to turn pink.

CHEMICAL PROPERTIES OF ALKALI METALS are under,

Hydroxides –  all hydroxides are ionic in nature.  Since they are basic being metal oxides we can conclude that more the metallic character more is basic so basicity increases down the group. Same is logic being ionic, more metallic more ionic so solubility increases down the group.

The nature is colourless crystalline solid and now considering the thermal stability all except LiOH are stable to heat and stability increases down the group.

 LiOH on heating decomposes to give Li2O +H2O the reason here is that the oxide have a more stable lattice than that of hydroxide of lithium

 

Reaction with ammonia –

M + NH3 —- >M+(NH3)X + ammoniated electrons

The solution of ammonia due to the presence of ammoniated electrons appears to be blue. If left for long the colour fades because of the following reaction

              M+ + NH3 + e —- > H2 + MNH2 the electrons get consumed and colour fades.

The solution when diluted is paramagnetic and is a good conductor of electricity for obvious reason. But when the conc. Is high the characteristics of the solutions change-

  • The solution turns copper bronze due to the formation the metallic luster.
  • The opposite spin electron attract and solution turns diamagnetic.
  • In presence of a catalyst like Fe we faster the reaction which forms MNH
  • While all other metals are forming MNH2, here also Li shows a unique behavior and reacts as a group II metal to form Li2

Reaction with Air –

The different metals form different oxides with oxygen

Li + O2 — > Li2O           (colorless oxide)

Na + O2 — > Na2O2     (colorless peroxide) sometimes the formation of super-oxide gives a colour.

M + O2  — > MO2         (colored super oxide) where M= K, Rb, Cs

The color of the super-oxide’s is due to the para-magnetic behavior, the O2is having two covalent bonds and a single electron, which when move from one to other atom releases photon of visible range giving the compounds color, and also the paramagnetic behavior

  • The stability of peroxides and super oxides increases moving down the group.
  • Li here also shows an anomalous behavior, when react with air it is the only metal to react with N2

Li + Air — > Li2O + Li3N the here also driving force is high lattice energy of product.

Li3N + H2O —> LiOH + NH3 the production of ammonia makes this an important reaction.

These are reaction which are given by the second group metals and so are important.

  • The Uses are important of the potassium super-oxide- it is used as a absorbed of the carbon dioxide and releasing di-oxygen gas, which makes it helpful for use in space capsules, submarines to function. Also is used to know the presence the carbon dioxide both qualitatively and quantitatively.

Reaction with Water –

Reaction is simple to give the hydroxides of respective metals

M + H2O — > MOH + H2  

  • The reaction is highly exothermic in nature
  • The reactivity is measure of metallic nature so more metallic nature and so the reactivity increases going down the group.

Li is gentle, Na melts on the surface and on exposure to sufficient water and air catches fire,

And potassium onward catches fire and explodes.

 

Due to smallest size, only Li form complexes.

Reaction with Sulphur and Phosphorus–

            M + S —> M2S              and all metals reacts with sulphur and may forms polysulphide in presence of excess sulphur like M2Sn where n= 2, 3,4,5,6.

               M + P —> M3P              the compound are according to the valences of the atom and there for easy to remember while the metal phosphide reacts with water to give phosgene and hydroxide.

Reaction with Carbon-

Only Li reacts with carbon in alkali metals to give Li2C2 Lithium Carbide

None of the other metal directly reacts but can form carbide when heated with ethylene.

Heating Effect –

  • Bicarbonates – All bicarbonates gives the same reactions on heating to decompose, into their respective carbonates. General reaction,

MHCO3        heat—->      M2CO3 + H2O + CO2

Now here we must remember that there is no difference in the product But LiHCO3 is the only which do not exist in solid form and is aq. Form with carbon dioxide atmosphere to prevent the reaction.

  • Carbonates – Now the carbonates are product on heating Bicarbonates this states that they are thermodynamic ally more stable. In Fact, all carbonates are stable to heat, except Li CO

M2CO3         heat —>     no reaction

Li2CO3         heat —->    Li2O + CO2   the reaction is analogous to the decomposition of group II carbonates decomposition, i.e. all the carbonates of Gp II decompose to give oxide and none of them are stable to heat.

Stability order —    Li2CO< Na2CO3 < K2CO< Rb2CO< Cs2CO3

The difference is in polarizing power of the atoms of Gp II and Li behaves as because of size.

Li+         O     C=O      the small size of Li atom increases its polarizing power as Gp II

Li+         O–                             having extra charge, more polarizing power { Fajan’s Rule }.

  • Nitrates – With nitrates as well Li shows a different behavior, all of them decompose but Li decompose to oxides-

MNO3       heat  —>   MNO2 + O2     where M= Na, K, Rb, Cs  

LiNO3       heat  —->    Li2O + NO2 + O2    

As we expect, the same type of decomposition is shown by Gp II Nitrates.

Abhishek kumar jha

(Chemistry at Utkarshini)

 

 

 

 

notes//inorganic-chemistry-gp-i-alkali-metals

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

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