Notes and Study Materials -S – Block Elements (Alkali & Alkaline Earth Metals)

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About this unit

Group I and group 2 elements:General introduction, electronic configuration, occurrence, anomalous properties of the first element of each group, diagonal relationship, trends in the variation of properties (such as ionization enthalpy, atomic and ionic radii), trends in chemical reactivity with oxygen, water, hydrogen and halogens; uses.Preparation and Properties of Some important Compounds:Sodium carbonate, sodium chloride, sodium hydroxide and sodium hydrogencarbonate, biological importance of sodium and potassium.Industrial use of lime and limestone, biological importance of Mg and Ca.

s-BLOCK ELEMENTS – ALKALI METALS

ELEMENTS OF GROUP 1

Li  – Lithium    Na – Sodium  K  – Potassium
Rb – Rubidium Cs – Caesium Fr – Francium
These elements are known as alkali metals.
Lithium is known as a bridge element and was discovered by Arfwedson.
Sodium and potassium were discovered by Davy, rubidium and caesium by Bunsen and Kirchhoff while francium by Perey.
These do not occur in the  native state (i.e.,do not occur free in nature).

GENERAL CHARACTERISTICS

Physical properties of alkali metals are:-

ELECTRONIC CONFIGURATION

These are s-block elements and have one electron in the valence shell in
s-orbital. In general  their electronic configuration may be represented as [noble gas ] ns1  where ‘n’ represents the valence shell.
Element     
Atomic no.
Electronic configuration
Valence shell configuration
Li
3
[He] 2 s1
2 s1
Na
11
[Ne] 3 s1
3 s1
K
19
[Ar] 4 s1
4 s1
Rb
37
[Kr] 5 s1
5 s1
Cs
55
[Xe] 6 s1
6 s1
Fr
87
[Rn] 7 s1
7 s1

SIZE OF THE ATOMS – ATOMIC RADII

  • The alkali metals atoms have the largest atomic radii in their respective periods.
  • Atomic radii increases as we move down the group from Li to  Cs due to the addition of a new shell at each step.

SIZE OF THE ION – IONIC RADII

  • The ions of the alkali metals are much smaller than their corresponding atomic radii due to lesser number of shells and contractive effect of the increased nuclear charge.
  • The ionic radii like atomic radii of all these alkali metal ions goes on increasing on moving down the group because of the same reason.

DENSITY

  • These are light metals having low densities. Lithium is the lightest known metal.
  • On moving down the group, both the atomic size and atomic mass increases and since the increase in latter is not compensated by increase in former,consequently density increases from Li to Cs.
  • The density of potassium is lesser than that of sodium because of the abnormal  increase in size on moving from Na to K.

MELTING AND BOILING POINTS

  • The melting and boiling points of alkali metals are quite low and decreases down the group due to weakening of metallic bond.
  • Fr is a liquid at room temperature.

SOFTNESS

These are soft,malleable and ductile solids which can be cut with  knife. They possess metallic lustre when freshly cut due to oscillation of electrons.

ATOMIC VOLUME

Atomic volume of alkali metals is the highest in each period and goes  on increasing down the group
Element
Li
Na
K
Rb
Cs
Gram atomic volume in cm3
13
24
46
56
71

IONISATION ENERGY

  • The first ionisation energy of alkali metals is the lowest amongst the elements in their respective periods and decreases on moving down the group.
Element
Li
Na
K
Rb
Cs
Fr
IE, (kJ mol–1)
520
496
419
403
376
—-
  • The second ionisation energies of all the alkali  metals are very large because on releasing an electron from the elements, the resulting ions acquire noble gas (stable) configurations.

ELECTROPOSITIVE CHARACTER

Because of their low ionisation energies, alkali metals are strongly electropositive or metallic in nature and this character increases from Li to Cs.

CRYSTAL STRUCTURE

All alkali metals possess body centred cubic structures with coordination number 8

OXIDATION STATE

  • The alkali metal atoms show only +1 oxidation state, because their unipositive ions have the stable gas electronic configuration in the valence shell.
  • Since the alkali metal ions have noble gas configuration with no unpaired electrons, they are diamagnetic and colourless but their permanganates and dichromates compounds are coloured.

HYDRATION OF IONS

  • All alkali  metal salts are ionic (except Lithium) and soluble  in water due to the fact that cations get hydrated by water molecules. The  degree of hydration depends upon the size of the cation. Smaller the size of a cation, greater is its hydration energy.

Relative ionic radii :

Relative ionic radii in water or relative degree of hydration :
  • The alkali metal ions exist as hydrated ions    in the aqueous solution.
  • Since the degree of hydration of decreases as we go down the group, the hydration  energy of alkali metal ions decreases from

FLAME COLOURATION

  • All alkali metals and their salts impart characteristic colours to the flame  because of the bonding of the outermost electron.The outer electrons of these atoms are excited to higher energy levels. On returning to the original state they give out visible light of characteristic wavelength. This gives a characteristic colour to the flame.
  • On moving down the group, the ionisation energy goes on decreasing and hence the energy or the frequency of emitted light goes on increasing in the order Li  < Na < K < Rb < Cs. As a result, the colour shows following trend-
Li
Na
K
Rb
Cs
crimson
golden
pale  
purple
sky blue
red
yellow
violet
(violet)
 

PHOTOELECTRIC EFFECT

Due to low I.E., alkali metals especially K and Cs show photoelectric effect (i.e.  eject electrons when exposed to light) and hence are used in photoelectric cells.

ELECTRICAL CONDUCTIVITY

Due to the presence of loosely held valence electrons which are free to move throughout the metal structure, the alkali metals are good conductors of heat and electricity. Electrical conductivity increases from top to bottom in the order

REDUCING CHARACTER

  • All the alkali metals are good reducing agents
    and it is due to their low ionisation energies.
    Their reducing character, follows the order
  • Among the alkali metals Li has the highest negative electrode potential, which depends upon its (i) heat of vaporisation (ii) ionisation energy and (iii) heat of hydration and hence Li is the strongest reducing agent.
Elements
 Li
  Na
K
 Rb
Cs
Fr
(V)at at 298 K
–3.05
–2.71
–2.93
–2.99
–2.99
 —

CHEMICAL PROPERTIES

ALKALI METALS FORM IONIC COMPOUNDS

(Lithium can form covalent compounds because of its high ionisation energy) and others form ionic compounds because of their large atomic size and low I.E. 

ALKALI METALS ARE VERY REACTIVE

Due to low I.E. and high electropositive character the alkali metals are chemically very reactive.

ACTION OF AIR

On exposure to moist air, their surface is tarnished due to the formation of  their oxides, hydroxides and carbonates at the surface.
Hence they are kept under inert liquid kerosene oil but lithium is kept wrapped in paraffin wax because it floats on the surface of kerosene oil due to its very low density.

ACTION OF OXYGEN

  • All the alkali metals when  heated with oxygen form different types of oxides for example, lithium forms lithium oxide, sodium forms sodium peroxide , while K, Rb and Cs form their respective superoxides ( where M=K, Rb or Cs). The increasing stability of peroxides and superoxides of alkali metals from Li to Cs  is due to stabilisation of larger anions by larger cations through lattice energy.
  • Superoxides are coloured and paramagnetic as these possess three electron bond where one unpaired electron is present.Sodium peroxide acquires yellow colour due to the presence of traces of superoxide as an impurity.is orange, is brown and   is orange in colour.
  • All oxides, peroxides and superoxides are basic in nature.
  • The solubility and basic strength of oxides increase in the order  
  • The stability of peroxides and superoxides increases in the order
       

ACTION WITH WATER AND OTHER COMPOUNDS CONTAINING ACIDIC HYDROGEN

  • All the alkali metals readily react with water  evolving hydrogen.
(where M=Li, Na, K, Rb or Cs)
The reactivity with water increases on descending the  group from Li to as
Li < Na < K < Rb < Cs  due to increase in electropositive character in the same  order.
  • Alkali metals also react with alcohols and acetylene and liberate

ACTION OF HYDROGEN  

Alkali metals combine with hydrogen to form ionic hydrides M+H
The reactivity of alkali metals towards hydrogen decreases as we move down the group i.e. Li  > Na > K > Rb > Cs, due to the decreasing lattice energy of these hydrides with the increasing size of the metal cation. Thus the stability of  hydrides follows the order
LiH > NaH > KH > RbH > CsH

REACTION WITH HALOGENS

  • Alkali metals combine readily with halogens to form ionic halides
[where M= Li,Na, K etc. and X = F,Cl, Br,I]
  • The reactivity of alkali metals towards a particular halogen increases in the order  :
Li < Na  < K < Rb < Cs
while that of halogen  towards a particular alkali metal decreases in the order :
  • All alkali halides except LiF are freely soluble in water (LiF is soluble in non-polar solvents. Since it has strong covalent bond.)
  • The power of the cation to polarise the anion  is known as the polarising power while the tendency of the anion to get polarised is known as its polarisability. The polarising power of cation and polarisability of anion depends on the following factors (which are collectively referred to as Fajan’s rules)
    • Size of the cation – Smaller the size of cation greater is its polarising power. So LiCl is more covalent than KCl.
    • Size of the anion – Bigger the anion, larger is its polarisability. Hence the covalent character of lithium halides  is in the order –
LiI > LiBr > LiCl > LiF
    • Charge of the ion and electronic configuration – Larger the charge on the cation, greater is its polarising power
Thus the covalent character of various halides is in the order
when two cations have same charge and size, the one having 18 electrons in their outermost shell will have larger polarising power than a cation having 8 electrons in the outermost shell. For example CuCl is more covalent than NaCl.
Above rules help to predict the ionic /covalent character of metal halides.

MELTING POINTS OF ALKALI METAL HALIDES

  • For the same alkali metal, the melting points decrease in the order with the     increase in the size of halides ion. Fluorides > chlorides > bromides > iodides  
  • For the same halide ion, melting points decreases with the increasing size of the  metal but lithium halides being covalent have lower melting point than corresponding sodium halides.

REACTION WITH NITROGEN

Only lithium reacts  with nitrogen and forms lithium nitride ()

REACTION WITH SULPHUR AND PHOSPHORUS

Alkali metals  react with sulphur and phosphorus on heating and form respective sulphides and phosphides.

SOLUBILITY  IN LIQUID AMMONIA

All alkali metals dissolve in liquid ammonia giving deep blue solution, which has some characteristic properties given below due to formation of ammoniated metal cations and ammoniated electrons in the solution.
            
  • Colour – The blue colour is due to the excitation of ammoniated electron to higher energy levels and the absorption of photons occurs in the red region of the spectrum.Thus the solution appears blue.But at very high concentration  the solution attains the colour like that of metallic copper.
  • Conductivity – It is highly conducting because of the presence of ammoniated electrons and ammoniated cations.However, on cooling,the conductivity increases further.
  • Paramagnetism – It is paramagnetic due to the presence of an unpaired electrons and ammoniated  cations.However the paramagnetism decreases with increasing concentration due to the  association of ammoniated electrons to yield diamagnetic species containing electron pairs.
  • Reducing property – Due  to the presence of ammoniated electrons, solution is a very powerful reducing agent and used in organic chemistry under the name Birch reduction.

COMPLEX FORMATION

Alkali metals have a weak tendency to form complexes but polydentate ligands such as crown ethers and cryptands form highly stable complexes collectively called as Wrap Around Complexes. Cryptands are macrocyclic molecules with N and O  atoms and their complexes are called cryptates. The name cryptate came from the  fact that metal ion is hidden in the structure.

NATURE OF HYDROXIDES

Alkali metals hydroxides are very strong bases, highly soluble in water and are not decomposed on heating.However, LiOH decomposes on heating to give  because latter is more stable than former.
Their  basic strength increases from LiOH to CsOH due to a corresponding decresae in the I.E., of the metal in a group,i.e., the order:-  
LiOH < NaOH < KOH < RbOH < CsOH

NATURE OF CARBONATES AND BICARBONATES

  • is unstable towards heat and decomposes to give
The thermal stability of carbonates increases with  the increasing basic strength of metal hydroxides on moving down the group.Thus the order is
  • The bicarbonates of all the alkali metals are known. All the bicarbonates (except which exits in solution) exist as solids and on heating form carbonates.
  • The solubility of the carbonates and bicarbonates increases  on moving down the group due to lower lattice energies. Thus, order is

NATURE OF NITRATES  

On heating decomposes to give   while the nitrates of the other alkali metals decompose on  heating to form nitrites and.
All nitrates are soluble in water.

NATURE OF SULPHATES

  • Li2SO4 is insoluble in water whereas other sulphates i.e, Na2SO4, K2SO4 are soluble in water.
  • Lithium sulphate does not form alums and is also not amorphous with other sulphates.

ANOMALOUS BEHAVIOUR OF LITHIUM

Lithium, the first member of the  alkali metal family shows an anomalous behaviour because of the following main reasons:-
  • It has the smallest size in the group
  • It has very high ionization energy and highest electronegativity  in the group.
  • It has no vacant d-orbital in the valence shell.
 
As a result, it differs from the other member of the alkali metal family  in following respects:
  • Lithium is harder than other alkali metals, due to strong metallic bond.
  • Lithium combines with to form lithium monoxide, , whereas other alkali metals form Peroxides, and superoxides.
  • Lithium, unlike the other alkali metals, reacts with nitrogen to form the nitride.
  • LiOH  is a weak base and decomposes to give the corresponding oxide while the hydroxides of alkali metals are stable to heat and sublime as such.
  • , LiF and lithium phosphate are insoluble in water while the corresponding salts of other alkali metals are soluble in water.
  • LiH is the stablest among all the alkali metal hydrides.
  • decomposes on heating to evolve whereas other alkali metal carbonates do not.
  • Lithium nitrate on heating  evolves O2 and NO2 and forms Li2O while other alkali metal nitrates on heating evolve and form their respective nitrites.
  • Lithium when heated with ammonia forms lithium imide  while other alkali metals form amides of the general formula ( where M=Na,K, Rb and S).
  • Only lithium combines directly with carbon to form lithium carbide, , while other alkali metals  react with ethyne to form the corresponding metal carbides.

DIAGONAL RELATIONSHIP

Lithium shows diagonal relationship with magnesium, the element of group 2 and this resemblance is due to polarising power, i.e,  is similar for both of these elements.

Lithium resembles magnesium in the following respects:
  • The atomic radius of Lithium is  while that of magnesium is
  • The ionic radius of which is very close to that of Mg2+ ion (0.65Å).
  • Lithium (1.0) and magnesium (1.2) have almost similar electronegativities.
  • Both Li and Mg are hard metals.
  • LiF is partially soluble in water like .
  • Both decompose water only  on heating.
  • Alkyls of lithium and magnesium are soluble in organic solvents.
  • Both combine with to form  monoxides,e.g., and MgO.
  • Both LiOH and are weak bases.
  • Both LiCl and are predominantly covalent.
  • Both Li and Mg  combine with to form their respective  nitrides, and.
  • The hydroxides and carbonates of both Li and Mg  decompose on heating and form their respective oxides.
  • Both lithium and magnesium nitrates on heating  evolve and leaving behind their oxides.
Fire caused by burning of alkali metals is extinguished by sprinkling .
A mixture of and dil.HCl is commercially called Oxone and is used for bleaching delicate fibres.

METALLURGY OF SODIUM

OCCURRENCE AND MINERALS

  • Sodium does not occur in the  free state because of its high reactivity.
  • Important  minerals of sodium are –
    • Common salt or rock salt,NaCl
    • Chile saltpetre,
    • Sodium carbonate,
    • Sodium sulphate or Glauber’s salt  Na2SO4. 10 H2O
    • Cryolite,
    • Borax,

EXTRACTION  OF SODIUM

  • Sodium metal is extracted by electrolysis of fused NaCl containing a little and KF at 873 K. This process is known as Down process.
  • Reactions during electrolysis
  • Difficulties during the process
    • Sodium cannot be extracted from  aqueous NaCl because the metal liberated at the cathode reacts with to form metal hydroxide and.
    • NaCl melts at 800º C and it is difficult to attain and maintain this high temperature.
    • Molten Na forms a metallic fog (colloidal solution ) with fused NaCl.
Above difficulties were removed by adding and KF  to fused NaCl which themselves do not undergo decomposition at the voltage employed and lower the melting point of NaCl to about. The electrodes are separated by a wire gauze to prevent the reaction between Na and

METALLURGY OF POTASSIUM

OCCURRENCE AND MINERALS

  • Potassium also does not occur in free state.
  • Important minerals of potassium are
    • Sylvine, KCl
    • Carnallite,
    • Feldspar,
    • Kainite,

EXTRACTION OF POTASSIUM

Potassium is not obtained by the electrolysis of fused KCl because K has lower boiling point (1039 K)  than the melting point of KCl (1063K) and hence it get vaporises. Therefore, K metal is extracted by the following methods :-
  • By the electrolysis of fused KOH – The reaction involved are
  • Modern method – By the  reduction of molten KCl with metallic sodium in stainless steel vessel  at 1120-1150 K.

COMPOUNDS OF SODIUM

SODIUM CHLORIDE, COMMON SALT OR TABLE SALT, NaCl

  • It is obtained by evaporation of sea water in sun but due to presence of impurities like  it is deliquescent It is purified by passing HCl gas through the impure saturated solution of NaCl and  due to common ion effect, pure NaCl gets precipitated.
  • 28% NaCl solution is called Brine.

SODIUM HYDROXIDE, CAUSTIC SODA, NaOH

PREPARATION
  • Causticizing process ( Gossage process) – A 10% solution of  is treated with milk of lime, .
  • Electrolytic process – In this process a concentrated solution of sodium chloride is electrolysed where is evolved at the anode and at the cathode. However gas reacts with  NaOH forming NaCl and  sodium hypochlorite.

Mercury cathode process (Castner – Kellner cell)
This process is  used to avoid reaction between NaOH and .NaOH is obtained by the electrolysis of (aqueous) solution of brine. The cell has three compartments and  involves following reactions :-
 
 In outer compartment
Anode – Graphite rods
Cathode – Mercury
Electrolyte -Brine solution  
Reaction – At  Anode :
      
At Cathode :

 

In central compartment
Anode – Mercury
Cathode – Iron rods
Electrolyte – dil. solution of  NaOH
Reaction – At  Anode :
At Cathode :
  • Lowing’s Process
  • Pure Sodium Hydroxide
          
The filtration on evaporation give pure NaOH

 

PROPERTIES
  • It is a hygroscopic, deliquescent white solid, absorbs CO2 and moisture from the atmosphere.
  • Reaction with salts:- It reacts with metallic salts to form hydroxides out of which some are unstable and decompose to insoluble oxides,
    • Formation of  insoluble hydroxides, e.g.
    • Formation of unstable hydroxides, e.g.
    • Formation of insoluble hydroxides which dissolve in excess of NaOH e.g. Zn, Al, Sb, Pb, Sn and As.
    • Formation of ammonia from ammonium salts :-
  • Reaction with halogens:-
(where X= Cl,Br or I)
  • Reaction with metals :- Less electropositive metals like Zn, Al and Sn etc. give H2 gas with NaOH.
  • Reaction with sand:-
  • Reaction with CO :-
  • Reaction with non-metals e.g.with P, Si, S, F, etc.
  • It breaks down the proteins of the skin flesh to a pasty mass and hence it is commonly known as caustic soda.

SODIUM CARBONATE OR WASHING SODA  :

PREPARATION
  • Solvay or ammonia – soda process :- In this process,NaCl (brine), ammonia and are taken as raw materials. The involving reactions are
  • Electrolytic Process :- In this Nelson cell is used for the manufacture of NaOH, CO2 under pressure is blown with steam
  • Leblance Process :-This is now an absolute method.
PROPERTIES
  • Sodium Carbonate crystallizes from water as decahydrate which efflorescence on exposure to dry air forming monohydrate which on heating change to anhydrous salt (soda-ash).
  • On hydrolysis it forms an alkaline solution
  • Aqueous sodium carbonate solution react with CO2 gas and forms sodium bicarbonate.
  • It is used as fusion mixture

SODIUM BICARBONATE, BAKING SODA, NaHCO3

PREPARATION
It is obtained as an intermediate product in Solvay ammonia process.

 

PROPERTIES
  • Heating effect :- It gives
  • In aqueous medium it is alkaline due to hydrolysis:
  • It is used as a constituent of baking powder and in medicine to remove acidity of the stomach (as antacid).
  • It is present in Selidlitz powder.
  • Baking powder is a mixture of starch, sodium bicarbonate and potassium hydrogen tartarate.
  • Fire extinguishers contain
  • is called Glauber’s salt, anhydrous is called salt cake, is called chile salt-petre, is called nitre cake, mixture   of and dil. HCl is called oxone.
  • When common salt is fused with a little, 5% to 10%and some sugar, it acquires a dark purple colour and has a characteristic saline taste. It is used in medicine and is useful for digestion.It is called kala namak or black salt or sulemani namak.

COMPOUNDS OF POTASSIUM

POTASSIUM HYDROXIDE,CAUSTIC POTASH, KOH

PREPARATION
  • It is prepared in a cell similar to that used for NaOH. In this cell electrolysis of an aqueous solution of KCl takes place.
  • It is also prepared by the action of soda lime on potassium carbonate.

POTASSIUM CARBONATE, POTASH, PEARL ASH, K2CO3

PREPARATION
It is prepared by following two methods-
  • By Leblanc process
  • By Precht process (magnesia process)

POTASSIUM CYANIDE, KCN:

PREPARATION
  • By heating potassium ferrocyanide with metallic potassium
  • It is used  in electroplating and due to the formation of soluble  complexes with gold and silver, it is used in extraction of these metals.

POTASSIUM CHLORATE, KCLO

PREPARATION
  • By passing through boiling concentrated KOH solution.
  • By the action  of KCl on (obtained by electrolysis of NaCl at 345-350K).

PROPERTIES
It used as an oxidising agent and in the laboratory and preparation of .

 

is called Indian salt petre or nitre.

s-BLOCK ELEMENTS – ALKALINE EARTH METALS

ELEMENTS OF GROUP 2

Be – Beryllium Mg – Magnesium Ca – Calcium
Sr – Strontium  Ba – Barium Ra – Radium (Radioactive)
These metals are known as alkaline earth metals as their oxides are alkaline and occur in earth crust.
Radium was discovered from the ore pitchblende by Madam Curie. It is used in the treatment of cancer. These metals do not occur in the native form (i.e., do not occur in free state).

GENERAL CHARACTERISTICS

Physical properties of alkaline earth metals are:-

ELECTRONIC CONFIGURATION

Like alkali metals, these are s-block elements, and have two electrons in the valence shell in s-orbital. Hence their electronic configuration may be represented as [noble gas] ns2 where ‘n’ represents the valence shell.
Element
Atomic No.
Electronic Configuration
Valence Shell configuration
Be
4
[He] 2s2
2s2
Mg
12
[Ne] 3s2
3s2
Ca
20
[Ar] 4s2
4s2
Sr
38
[Kr] 5s2
5s2
Ba
56
[Xe] 6s2
6s2
Ra
88
[Rn] 7s2
7s2

SIZE OF ATOMS AND IONS (ATOMIC RADII AND IONIC RADII)

  • The atomic radii of these elements are quite large but smaller than those of the corresponding elements of group 1, due to increased nuclear charge of these elements which tends to draw the orbital electrons inwards.
  • The ionic radii are also large but smaller than those of the alkali metals.
  • The atomic as well as ionic radii go on increasing down the group due to the gradual addition of extra energy level and also because of the screening effect.

DENSITY

  • These are much denser than alkali metals because of their smaller size and greater nuclear charge.
  • The density, however, first decreases from Be to Ca and then steadily increases from Ca to Ra due to difference in crystal structure

MELTING AND BOILING POINTS

  • These have higher melting and boiling points than those of alkali metals because the number of bonding electrons in alkaline earth metals is two.
  • The melting and boiling points decrease down the group with the exception of magnesium.
  • Melting points of halides decrease as the size of the halogen increases. The correct order is
MF2 > MCl2 > MBr2 > MI2

METALLIC PROPERTIES:

They are silvery white metals, soft in nature but harder than alkali metals due to stronger metallic bonding.

ATOMIC VOLUME

Atomic volume of these metals increases considerably on moving from Be to Ra as the atomic radius increases.

IONIZATION ENERGY

  • The first I.E. of alkaline earth metals are higher than those of the corresponding alkali metals due to smaller size and higher nuclear charge.
  • The second I.E. values are higher than their first I.E. values but much lower than the second I.E. values of alkali metals.
  • On moving down the group due to increase in atomic size the magnitude of I.E. decreases.
  • The ionization potential of radium is higher than that of barium.

ELECTROPOSITIVE CHARACTER

  • These are strong electropositive elements due to their large size and comparatively low ionisation energies.
  • On moving down the group, the electropositive character increases due to increase in atomic radii.

OXIDATION STATE

  • Alkaline earth metals uniformly show an oxidation state of +2 despite the presence of high ionisation energy because
    • In the solid state, the dipositive ions M2+ form strong lattices due to their small size and high charge (i.e., high lattice energy).
    • In the aqueous solution, the M2+ cations are strongly hydrated due to their small size and high charge. The hydration energy released by the M2+ cation is very large.
  • The divalent ions are diamagnetic and colourless due to the absence of unpaired electron.

CONDUCTIVITY

These are good conductors of heat and electricity due to the presence of two loosely held valence electrons.

FLAME COLOURATION

  • Like alkali metal salts, alkaline earth metal salts also impart characteristic flame colouration.
  • As we move down the group from Ca to Ba, the ionisation energy decreases, hence the energy or the frequency of the emitted light increases. Thus,
Ca    Sr  Ba            Ra
Brick red Crimson red Apple green Crimson
  • Be and Mg because of their high ionization energies, however, do not impart any characteristic colour to the bunsen flame.

CHEMICAL PROPERTIES

  • Alkaline earth elements are quite reactive due to their low ionisation energies but are found to be less reactive than alkali metals because the alkaline earth metals have comparatively higher ionisation energy.
  • Reactivity of the group 2 elements increases on moving down the group because their ionisation energy decreases.

REACTION WITH WATER

  • Group 2 elements are less reactive with water as compared to alkali metals. They react with H2O evolving H2 gas.
  • The chemical reactivity of the metal with H2O, however increases as we move from Mg to Ba, i.e., Be does not react even with boiling water and Ba react vigorously even with cold water. Thus, increasing order of reactivity with water is
Mg < Ca < Sr < Ba

REACTION WITH OXYGEN

The affinity for oxygen increases down the group. Thus, Be, Mg and Ca when heated with O2 form monoxides while Sr, Ba and Ra form peroxides.

REACTION WITH ACIDS

  • Alkaline earth metals except Be, displace H2 from acids.
(where M = Mg, Ca, Sr or Ba)
  • Reactivity, however, increases down the group from Mg to Ba i.e.,Mg < Ca < Sr < Ba
  • Only Mg displaces H2 from a very dilute HNO3.

REACTION WITH HYDROGEN

  • Except Be, all other elements combine with hydrogen on heating to form hydride (MH2).
  • The hydride of beryllium can be prepared indirectly by reducing beryllium chloride with lithium aluminium hydride.
  • BeH2 and MgH2 are covalent and polymeric whereas the hydrides of Ca, Sr and Ba are ionic and monomeric in nature.
  • CaH2 is also called hydrolith.
  • All the hydrides react with water to evolve H2 and thus behave as strong reducing agents.

REACTION WITH HALOGENS

  • All the elements of group 2 combine with halogens at high temperature and form their halides (MX2).
  • Beryllium halides (BeF2, BeCl2 etc.) are covalent, hygroscopic and fume in air due to hydrolysis. The halides of other alkaline earth metals are fairly ionic and this character increases as the size of the metal increases.
  • The halides are soluble in water and their solubility decreases in the order:
MgX2 > CaX2 > SrX2 > BaX2
  • BeF2 is very soluble in water due to the high solvation energy of Be2+ in forming    but the fluorides of other alkaline earth metals have high melting point and they are insoluble in water.
  • BeCl2 has a polymeric structure in the solid state but exists as a dimer in the vapour state and as a monomer at 1200 K.

REACTION WITH NITROGEN

These metals burn in nitrogen to form nitrides of the types M3N2 which are hydrolysed with water to evolve NH3.
  • The ease of formation of nitrides increases from Be to Ba. (Be3N2) is volatile in nature.
  • Anhydrous CaCl2 is a good drying agent due to hygroscopic nature (CaCl2.2H2O) and cannot be used to dry alcohol or ammonia as it forms addition products with them.

REACTION WITH CARBON

  • When heated with carbon, these form their respective carbides of the general formula MC2 (except Be) and are called acetylides containing the discrete anion.
  • Under these conditions beryllium, however, forms Be2C called methanide containing the discrete C4– anion.
  • All these carbides are ionic in nature and react with H2O to form acetylene (except Be2C which gives methane).
or 
  • On heating  MgC2 gives Mg2C3 called allylide which contains the discrete anion and gives allylene (methyl acetylene) on hydrolysis.

REDUCING CHARACTER

  • All the alkaline earth metals, because of their low electrode potentials, are strong reducing agents but these are weaker than the corresponding alkali metals.
  • As we move down the group from Be to Ra, the reducing character increases due to decreasing I.E. from Be to Ra.

SOLUBILITY IN LIQUID AMMONIA

  • Like alkali metals, these dissolve in liquid ammonia giving coloured solutions.
  • The tendency to form ammoniates decreases with increase in size of the metal atom (i.e., on moving down the group).

COMPLEX FORMATION

  • Complex formation is favoured in case of alkaline earth metals because of their small sizes as compared to the alkali metals.
  • Both Mg2+ and Ca2+ form six coordinate complexes with EDTA (ethylenedi-aminetetracetic acid) which are used to determine the hardness of water.
  • Beryllium due to small size forms complexes of type [BeF3], [BeF4]–2 [Be (H2O)4]2+.
  • Mg exists as a natural complex, chlorophyll where it is complexed with pyrole rings of porphyrin.

BASIC STRENGTH OF OXIDES AND HYDROXIDES

  • BeO and Be(OH)2 are amphoteric while the oxides and hydroxides of other alkaline earth metals are basic. The basic strength, however, increases from Be to Ba as the ionisation energy of metal decreases down the group thus the order:
BeO < MgO < CaO < SrO < BaO and
Be(OH)2 < Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2
  • The basic character of hydroxides of group-2 elements is lesser than those of group-1 hydroxides because of the larger size elements of latter than former group.
  • Aq. Ba(OH)2 is known as baryta water.

THERMAL STABILITIES AND NATURE OF BICARBONATES AND CARBONATES

  • Bicarbonates of these metals do not exist in solid state but are known in solution only, when these solutions are heated, these get decomposed to evolve CO2.
  • The carbonates of alkaline earth metals can be regarded as salts of weak carbonic acid (H2CO3) and metal hydroxide, M(OH)2. The carbonates decompose on heating form metal oxide and CO2.
  • The stability of carbonates and bicarbonates increases down the group.
  • Carbonates and sulphates of Ca and Mg are responsible for permanent hardness of water while their bicarbonates cause temporary hardness.

SOLUBILITY OF THE SALTS

The solubility of a salt in water depends upon two factors.
  • Lattice energy: Higher the magnitude of lattice energy, lesser will be the solubility of the salt in the given solvent.
  • Hydration energy: Higher the magnitude of hydration energy, higher will be the solubility of the salt in water (solvent).
  • Solubility of hydroxides: As the ionic size of group 2 metals increases from Be to Ba, the lattice energy decreases from Be to Ba, as follows:
  • Solubility of sulphates: The solubility of sulphates of alkaline earth metals decreases as we move down the group from Be to Ba due to the reason that ionic size increases down the group. The lattice energy remains constant because sulphate ion is so large, so that small change in cationic sizes do not make any difference. Thus the order:
The negligible solubility of BaSO4 in water is used in both qualitative and quantitative analysis.
  • Solubility of carbonates: The solubility of the carbonates in water decreases down the group due to the decrease in the magnitude of hydration energy. However these insoluble metal carbonates are dissolved in water having CO2 as shown:
Mg (HCO3)2  Mg2+ (aq) + 2HCO3 (aq.)
The order of solubility of carbonates:
BeCO3 > MgCO3 > CaCO3 > SrCO3 > BaCO3

ANOMALOUS BEHAVIOUR OF BERYLLIUM

Beryllium, the first member of group 2 differs from the rest of the members of its group due to the following reasons:
  • It has a small atomic size as well as small ionic size.
  • It has no vacant d-orbitals in valence shell.
  • It has high electronegativity value.
  • It has a high charge density.
  • Its hydration energy is high.

 

Some points of difference are:
  • Be is harder and denser than other members of the group.
  • The m.p., b.p., and ionisation energy of Be are the highest of all the alkaline earth metals.
  • Be does not react with water even at higher temperature where as other metals do.
  • BeO and Be(OH)2 are amphoteric in character whereas oxides and hydroxides of the group 2 metals are basic.
  • Be is least metallic of all the alkaline earth metals and forms covalent compounds.
  • Be forms nitride Be3N2 with nitrogen which is volatile while nitrides of others are non-volatile.
  • Be does not liberate H2 from acids (HCl, H2SO4) where as other metals do.
  • Be forms Be2C with carbon while the other members of the group form ionic carbide MC2.

DIAGONAL RELATIONSHIP OR RESEMBLANCE BETWEEN Be AND Al

The first member of group-2, Beryllium, shows similarities in the properties with its diagonally opposite member aluminium of the next group 13 of the next higher period, due to the similar polarizing power. i.e., ionic charge/(ionic radius)2 of Be and Al.
  • Both the metals are stable in air.
  • Both have a strong tendency to form covalent compounds.
  • Both form fluoro complex anions BeF42– and AlF63– in solution.
  • With conc. HNO3, both are rendered passive due to the formation of a thin film of their respective oxides on the metal surface.
  • Both react with conc. NaOH liberating H2.
  • Oxides and hydroxides of both are amphoteric in nature.
  • Carbides of both liberate methane on hydrolysis.
  • Anhydrous chlorides of both i.e., BeCl2 and AlCl3 act as Lewis acids and dissolve in organic solvents.
  • Both do not impart any colour to the flame.

 

Magnesium along with KClO3 or BaO2 is used in photography flash bulbs, fireworks and as a deoxidiser in metallurgical process.
MgCl2.5MgO xH2O is called Sorel’s cement or Magnesia cement and used to fill the cavities of teeth.
Mg(OH)2 in water is used in medicine as an antacid under the name ‘Milk of Magnesia’, while 12 gm of MgCO3 per 100 c.c. of H2O containing CO2 is known as ‘Fluid Magnesia’.
The finely divided BaSO4 is called Blanc fire and used in paints.
Suspension of slaked lime in water is called white wash (milk of lime).
A solution of MgCl2 + NH4Cl in ammonia is known as Magnesia Mixture.
Plaster of paris CaSO4.1/2 H2O is used in surgery for setting broken bones.
Pure Ca(H2PO4)2 is used as American baking powder.
Gypsum gives different products on heating as
Most abundant alkaline earth metal in the earth’s crust is Ca.
Be and Mg crystallize in hcp, Ca and Sr in ccp and Ba in bcc structures.
Because of comparatively higher electronegativity both Be and Mg form a large number of organometallic compounds.
CaCl2.6H2O is widely used for melting ice on roads, particularly in very cold countries, because a 30% eutectic mixture of CaCl2/H2O freezes at –55ºC as compared with NaCl/H2O at –18ºC.
Magnesium perchlorate, Mg(ClO4)2 is known as anhydrone and used as drying agent.
Mostly kidney stones containing calcium oxalate, CaC2O4.H2O which dissolves in dil. strong acids but remains insoluble in bases.

METALLURGY OF MAGNESIUM

OCCURRENCE AND IMPORTANT MINERALS

  • Magnesium occurs in the combined state in nature and it is the essential constituent of chlorophyll, the green colouring matter of the plants.
  • The important minerals of magnesium are.
    • Magnesite, MgCO3
    • Dolomite, MgCO3.CaCO3
    • Carnallite, KCl.MgCl2.6H2O
    • Epsom salt, MgSO4.7H2O
    • Asbestos, CaMg3 (SiO3)4
    • Talc, Mg2(Si2O5)2.Mg(OH)2

EXTRACTION

It is extracted by the electrolysis of fused mixture of magnesium chloride (which is obtained from carnallite and magnesite), NaCl and CaCl2 (added to provide conductivity to the electrolyte and to lower the fusion temperature of anhydrous MgCl2) at 700ºC in Dow’s process. In Dow’s process, MgCl2 is obtained from sea water as MgCl2.6H2O which can be changed to anhydrous MgCl2 only by passing dry HCl gas through it because even by strong heating it gets hydrolysed by its own water of crystallisation.
Anhydrous MgCl2 is fused with anhydrous NaCl and CaCl2 and electrolysed at 700ºC.
At Cathode : 
At anode: 

USES OF MAGNESIUM

  • Mg being a light metal forms alloys with Al and Zn which are used in aircraft construction. e.g., elektron (95% Mg + 5% Zn) used in construction of aircraft, magnalium (1-15% Mg + 85-99% Al) used in construction of aircraft and light instruments.
  • Magnesium powder is used in flash bulbs used in photography.

COMPOUNDS OF MAGNESIUM

MAGNESIUM OXIDE, MAGNESIA, MgO

With MgCl2, it forms a mixture of composition MgCl2.5MgO.xH2O which is known as Sorel’s cement or magnesia cement.

MAGNESIUM HYDROXIDE, MILK OF MAGNESIA, Mg(OH)2

Its aqueous suspension is used in medicine as an antacid.

MAGNESIUM SULPHATE OR EPSOM SALT, MgSO4.7H2O

It shows isomorphous nature with ZnSO4.7H2O, deliquescence and efflorescence. It is used as a purgative in medicine and as a stimulant to increase the secretion of bile.

MAGNESIUM CARBONATE, MAGNESITE, MgCO3

  • It dissolves in water in the presence of CO2.
  • Its 12% aqueous solution is known as fluid magnesia and is used as an antacid, laxative and in toothpastes.

MAGNESIUM CHLORIDE, MgCl2.6H2O

It is a deliquescent, white crystalline solid.

METALLURGY OF CALCIUM

OCCURRENCE AND IMPORTANT MINERALS

  • It is an important constituent of bones and teeth (as calcium phosphate), sea shells and corals (as calcium carbonate).
  • The important minerals are
    • Limestone, marble, chalk or calcite, CaCO3
    • Dolomite, MgCO3.CaCO3
    • Gypsum, CaSO4.2H2O
    • Fluorspar, CaF2
    • Anhydrite, CaSO4
    • Hydroxyapatite, 3Ca3(PO4)2.Ca(OH)2
    • Phosphorite, Ca3(PO4)2.

EXTRACTION

It is extracted by the electrolysis of a fused mixture of calcium chloride and calcium fluoride (lowers the fusion temperature of the electrolyte).

USES OF CALCIUM

  • It is used to remove air from vacuum tubes, sulphur from petroleum and oxygen from molten steel.
  • It is used as a reducing agent in the extraction of such metals from their oxides where carbon is ineffective.

COMPOUNDS OF CALCIUM

CALCIUM OXIDE, QUICK LIME, BURNT LIME, LIME, CaO

PREPARATION
By the thermal decomposition of calcium carbonate.

 

PROPERTIES
  • It is a basic oxide.
  • Its aqueous suspension is known as slaked lime Ca(OH)2.
  • On heating with ammonium salts it gives ammonia
  • It reacts with carbon to form calcium carbide.
  • It is used as basic flux, for removing hardness of water for preparing mortar (CaO + Sand + Water).

CALCIUM HYDROXIDE, SLAKED LIME, LIME WATER, Ca(OH)2

PREPARATION
By dissolving quick lime in water.

 

PROPERTIES
  • Its suspension in water is known as milk of lime.
  • It gives CaCO3 (milky) and then Ca(HCO3)2 with CO2.
  • It reacts with Cl2 to give bleaching powder CaOCl2.

CALCIUM CHLORIDE, CaCl2.6H2O

  • It is a deliquescent solid which is a by-product of Solvay’s process.
  • Fused Calcium chloride is a good dessicant (drying agent), but it can not be used to dry alcohol or ammonia as it forms addition product with them.

CALCIUM CARBONATE, LIMESTONE, MARBLE, CHALK, SLATE, CALCITE, CaCO3

PREPARATION
By passing CO2 through lime water.

 

PROPERTIES
It is insoluble in H2O but dissolves in the presence of CO2, due to the formation of calcium bicarbonate.

GYPSUM, CALCIUM SULPHATE DIHYDRATE, CaSO4.2H2O

  • It is naturally occurring calcium sulphate and also known as alabaster.
  • On heating at 390K, it gives plaster of paris.
  • It is added to cement to slow down its rate of setting.

PLASTER OF PARIS, CALCIUM SULPHATE HEMIHYDRATE, CaSO4.1/2 H2O

  • When it is mixed with water, it forms first a plastic mass which sets into a solid mass with slight expansion due to rehydration and its reconversion into gypsum.
  • On heating at about 200ºC, it also forms dead burnt plaster of paris (it has no tendency to set).

CALCIUM CARBIDE OR CALCIUM ACETYLIDE, CaC2

PREPARATION
By heating a mixture of quick lime (CaO) and powdered coke in an electric furnace at 3300K.

 

PROPERTIES
  • It reacts with water to form acetylene.
  • When heated with nitrogen, it forms calcium cyanamide which on reaction with steam under pressure gives NH3.
  • Nitrolim (a mixture of calcium cyanamide and carbon) is used as a fertilizer.

BLEACHING POWDER, CALCIUM HYPOCHLORITE, CHLORIDE OF LIME, CaOCl2

PREPARATION
By passing a current of chlorine over dry slaked lime.
MANUFACTURE
The manufacture of bleaching powder is carried out in (i) Hasenclever plant or (ii) Bachmann’s plant

 

PROPERTIES
  • It is a mixture (mixed salt) of calcium hypochlorite (Ca.(OCl)2.4H2O) and basic calcium chloride (CaCl2.Ca(OH)2.H2O).
  • Its aqueous solution gives Ca2+, Cl and OCl ions.
  • With dil. H2SO4, it gives nascent oxygen which causes its oxidising and bleaching power.
  • With excess of dil. H2SO4 (or CO2), it forms Cl2 known as available chlorine.
The average percentage of available chlorine is 35 – 40%. Theoretically it should be 49%, which diminishes on keeping the powder due to following change 

 

Available chlorine is estimated by
  • Arsenite method (Penot’s method)
  • Iodometric method (Bunsen and Wagner’s method)
  • It gives O2 in presence of catalyst COCl2.

 

USES
It is used for bleaching, as disinfectant and germicide in sterilization of water, for making wool unshrinkable and in the manufacture of Chloroform.

CEMENT

Cement is essentially a mixture of complex silicates and aluminates of Ca containing less than 1.0% free lime and some gypsum (CaSO4.2H2O)

 

COMPOSITION
An approximate composition is as follows :
Lime
CaO
60-69%
62%
Silica
SiO2
17-25%
22%
Alumina
Al2O3
3-8%
7.5%
Magnesia
MgO
1-5%
2.5%
Iron oxide
Fe2O3
0.5-5%
2.5%
Sulphur trioxide
SO3
1-3%
1.5%
Sodium oxide
Na2O
0.3-1.5%
1.0%
Potassium oxide
K2O
0.3-1.5%
1.0%

 

Ratio of Silica and alumina

 

Ratio of CaO and 6(SiO2 + Al2O3 + Fe2O3)

 

White Cement : It does not contain ferric oxide

 

PROCESS
Two processes are employed (i) Wet process (ii) Dry process
Raw material : Lime and Clay

 

MANUFACTURE
Clay + Lime ➝ Cement clinker ➝ Cement
Gypsum regulates the setting time

 

SETTING OF CEMENT
When mixed with water, the cement forms a gelatinous mass sets to hard mass when three dimensional cross links are formed between … Si-O-Si—and —Si-O-Al— chains.
The reactions involved in the setting of cement are :
  • Hydration : Hydration of 3CaO.Al2O3 and 2CaOSiO2 forming colloidal gel.
  • Hydrolysis : Hydrolysis of 3CaOAl2O3 and 3CaO.SiO2 forming precipitates of Ca(OH)2 and Al(OH)3

 

Fly ash : A waste product of steel industry possess properties similar to cement. It is added to cement to reduce its cost.
Rice Husk : It has high silica content and employed to make cement.

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