S-Block Elements properties and overview

S-Block Elements

  1. S-block elements include the filling of s-subshell.
  2. S-BLOCK ELEMENTS includes group-1 and group-2
    1. Group-1 called alkali metals because they form hydroxide on reaction with water which is strongly alkaline in nature.
    1. Group-2 called alkaline earth metals except beryllium (Be). These are so called because their oxides and hydroxide are alkaline in nature and their metal oxide are found in earth crust.
  3. General configuration of s-block elements are ns1-2

They have one or two electrons in their outermost s-subshell.

Alkali and Alkaline earth metals

Physical state-

Alkali metal Alkaline earth metal
General  Configuration = ns1 One electron in their outermost shell General configuration = ns2 Two electrons in the outermost shell
Francium is the Radioactive element Radium is the radioactive element
All are silvery white solid All are greyish white
Alkali metals are paramagnetic and while their ions are diamagnetic and colorless Alkaline earth metals are diamagnetic
      5. These are malleable and ductile, metallic          lustre   and  light soft. These metals are slightly harder then alkali metals.
      6. Cesium is the softest metal in s-block  Be is the hardest metal in the s-block

Key points-

1. Alkali metals are soft because of large atomic size, BCC crystal structure (but HCP in Li), loose packing (68% packing efficiency) and weak metallic bond.

Whereas alkaline earth metals are hard because of small size, FCC crystal structure, packing capacity – 74%, and strong metallic bond.

2. Beryllium is not an alkaline earth metal because its oxide is not purely basic in nature, it is amphoteric in nature. Also its oxide is not found in earth crust.

Atomic size-

Alkali Metals-

Alkali metals have the largest atomic and ionic radii in their respective periods. Their size increases from Li to Fr due to addition of an extra shell.

   Li ˂ Na ˂ K ˂ Rb ˂ Cs ˂ Fr

Alkaline earth metals-

Alkaline radius is smaller than alkali metals because as we move from left to right, electron enters in the same shell. Nuclear charge increases, Zeff increases (extra charge on nucleus attracts the electron cloud more). Size increases gradually Be to Ra-

Be ˂ Mg ˂ Ca ˂ Sr ˂ Ba

Melting and boiling point-

Alkali Metals-

1. All these metals are soft. The alkali metals have only one valence electron. These alkali metals have large atomic radii and weak interatomic bonds. Hence, melting and boiling point is low.

2. Decreasing order of M.pt and b.pt –

   Li ˂ Na ˂ K ˂ Rb ˂ Cs

3. As the size of metals atoms increases, repulsion of non-bonding electrons (the electrons in an atom that does not participate in bonding with other atoms) increases.

Alkaline earth metals-

1. These are harder than alkali metals. So, metallic bond is stronger due to smaller atomic size and they have two electrons in the valence shell. Hence, their m.pt and b.pt increases.

2. Decreasing order of M.pt and B.pt –

 Be ˃ Ca ˃ Sr ˃ Ba ˃ Mg

3. Melting point and boiling of Ca , Sr and Ba is higher than Mg because of presence of d-orbital in the outermost shell which forms stronger metallic bond.

Ionization potential-

Alkali Metals-

Alkali metals have lowest ionization potential because as we move down the group, atomic size increases due to increasing in the no. of shells. The alkali metals have only one valence electron in their outermost shell. So, the valence electrons are loosely held by nucleus. By losing one electron, they acquire stable noble gas configuration.

  1. First ionization energy of alkali metals has lower.

                    Na     ——->         Na+  +  e

  • Second ionization energy

Secondary ionization energy of alkali metals are very high because after losing one electron, they become noble gas configuration means stable. So, it becomes very difficult to remove second electron from stable noble gas configuration. Therefore, their second ionization enthalpy (IE2) is very high.

  •  Decreasing order of ionization potential-

   Li ˃ Na ˃ K ˃ Rb ˃ Cs 

Alkaline earth metals-

Alkaline earth metals have atomic size smaller than alkali metals. So, ionization energy is more than group 1 elements.

  1. First ionization energy of alkaline earth metals is higher because of completely filled s-orbital.
  • Second ionization energy

After losing one electron from ns2 electronic configuration, it has become ns1. Then, to remove second electron, less ionization energy is required. So, second ionization energy is less than alkali metals.

  • Decreasing order of ionization potential-

                  Be ˃ Mg ˃ Ca ˃ Sr ˃ Ba

Oxidation state-

Alkali Metals-

The alkali metals show +1 oxidation state.

Alkaline earth metals-

The alkaline earth metals show +2 oxidation state.

Electropositive and metallic character-

Alkali Metals-

Due to the large size, electron can be removed easily to form m+ ion. The electropositive character increases down the group means from Li to Cs.

Alkaline earth metals-

An alkaline earth metals have atomic size is less than alkali metals. The ionization potential of alkaline earth metals is more than alkali metals. So, the electro positivity is  less than alkali metals.

Density-

Alkali Metals-

As we move along the group, atomic volume increases along with atomic weight. But atomic weight increases more than atomic volume. So, density increases from Li to Cs.

Exception Density of Na ˃ K

Because empty d-orbital are present in potassium (K).

Potassium has 8 electrons in their M-shell (3rd shell) and 3rd shell has maximum capacity is 18 electrons which decreases the density.

So, order of density- Li ˂ K ˂ Na ˂ Rb ˂ Cs

Alkaline earth metals-

Normally, density increases from Be to Ba.

Exception– a) Density of Be is higher than Ca and Mg because of less volume in comparison to its mass.

b) Density of Mg is higher than Ca because Ca has empty d-orbital. So, Ca has less density than Mg.

c) Order of density Ca ˂ Mg ˂ Be ˂ Sr ˂ Ba

Note- Density order of Mg, Na, Ca and K- k ˂ Na ˂ Ca ˂ Mg

Conductivity-

Alkali Metals-

Due to presence of loosely held valence electrons which are free to move in metal structure, these elements are good conductor of electricity.

 Alkaline earth metals-

They are also good conductor of heat and electrvicity due to presence of two free electrons. So, the conductivity of alkaline earth metals is more than alkali metals.

Flame test-

Alkali Metals-

 All the alkali metals and their salts impart characteristics flame colorization.

The flame energy is sufficient to excite the electrons of alkali metals to higher energy levels. The excited state is quite unstable and therefore, when these excited electrons come back to their original energy, they emit extra energy which fall in visible region of electromagnetic radiation. So, appear colored.

Li – crimson red, Na – Golden yellow, k – Violet, RbRed violet, Cs – Blue

Alkaline earth metals-

The alkaline earth metals give characteristic flame because their ionization enthalpies are low.

Beryllium and magnesium atoms are smaller and their ionization energy are very high. So, due to their small size, bind their electrons more strongly, and not excited to their higher level. Hence, they don’t give flame test.

Ca – Brick red, Ba – Green, Sr – Dark red

Photoelectric effect-

Alkali Metals-

Alkali metals exhibit photoelectric effect. Alkali metals have low ionization energies. So, the electrons are easily ejected when exposed to light. Cesium has lowest ionization energy and hence, it can show photoelectric effect to maximum extent. That’s why, Cs is used in photo-cells.

Alkaline earth metals-

These elements don’t show this property as their atomic size is small. Hence, ionization potential is higher than alkali metals.

Standard oxidation potential-

Alkali Metals-

All the alkali metals have high positive value of standard oxidation potential (means tendency of releasing electrons in water or self ionic solution). So, these are good reducing agent having upper most position in the electrochemical series.

Note – Li  has highest standard oxidation potential (+ 3.05 volt) due to its high hydration energy such that it convert into Li+ ion by losing one electron.

Alkaline earth metals-

They have lower values of standard oxidation potential due to their small size. Then, increasing order of standard oxidation potential-

Be ˂ Mg ˂ Ca ˂ Sr ˂ Ba

Note- Order of standard oxidation potential of s-block elements-

Li ˃ K ˃ Ba ˃ Sr ˃ Ca ˃ Na ˃ Mg ˃ Be

Hydration energy (Heat of hydration)-

Alkali Metals-

  1. Alkali metals are generally soluble in water due to hydration of cations by water molecules.
  2. Smaller the cation, greater is the degree of its hydration.

From left to Right

Li+      Na+     K+    Rb+     Cs+

Degree of hydration decreasing     

                            Hydration energy decreasing

                            Hydrated ion decreasing 

                            Ionic conductance increasing                                                 

    Alkaline earth metals-

  1. Alkaline earth metal has smaller ionic size and higher charge density due to their hydration energy is high.
  2. Decreasing order of hydration energy-

     Be+2 ˃ Mg+2 ˃ Ca+2 ˃ Sr+2 ˃ Ba+2

Complex formation tendency-

Alkali Metals-

Only those elements can form complex compounds which have-

  1. Small cation size
  2. High charge density
  3. Vacant d-orbital to accept electron

Only Li+ can form complex compound, due to its small size and rest alkali metals have very less tendency to form complex compounds.

Alkaline earth metals-

Alkaline earth metals have less tendency to form complex compound, but due to small size of cation Be and Mg forms complex compounds like – [BeF4]-2

Reducing property-

Alkali Metals-

Alkali metals have high standard oxidation potential, so these are strongest reductants. Reducing character increases in down the group in gaseous or molten state-

Li ˂ Na ˂ K ˂ Rb ˂ Cs

Whereas the order of reducing character in aqueous solution-

 Li ˃ K ≈ Rb ˃ Cs ˃ Na

Alkaline earth metals-

Alkaline earth metals have less reductant than alkaline earth metals.

Order of reducing property in aqueous and gaseous medium is  –

    Be ˂ Mg ˂ Ca ˂ Sr ˂ Ba

Lattice enthalpy-

The lattice enthalpy is defined as the amount of energy which is required to break the one mole of crystal into free ions. It gives a measure of force of attraction between the ions. Larger is the force of attraction, greater will be the lattice energy.

It is mainly depends on the size of the ion and its charge.

The alkali metals salts consists of cations and anions which are held up by strong electrostatic force of attraction. The lattice enthalpy of alkali metals salts are very high. As we move from down the group, lattice enthalpy decreases.

LiCl < NaCl < KCl < RbCl < CsCl

Key point –For the cation of same valency, lattice enthalpy of ionic solid decreases with increase in the size of cation due to decrease in the force of attraction between them.

Tendency to form Bivalent ions

The alkaline earth metals exhibit a valency +2 because they can lose two electrons from their outermost shell easily. Therefore, they have tendency to form bivalent ions.

We all know that, the alkaline earth metals – I.E> I.E1

Explanation –

  • After losing the two electrons from their outermost shell, they can have noble gas configuration.
  • Bivalent ions of the alkaline earth metals are hydrated in solution.

M+2  ions of alkaline earth metals have high hydration energy which makes them stable than M+1 ion. Because it is observed that, when M+2 dissolved in water, it release much more energy than M+1 ion.

The amount of energy released in the hydration of M+2 ions is more than 2nd ionization potential which is required for the formation of such ions.

  • In solid,  +2 form stronger lattice than +1.

Greater lattice energy of M+2 ion which compensates for high second ionization potential which is responsible for its stability as compared to M+1.

Solubility –

Solubility depends on the two factors – Lattice energy and hydration energy

Net enthalpy = Lattice enthalpy – Hydration enthalpy

For the feasible reaction, net enrgy should be –ve means energy should be released.

  • If hydration energy is more than lattice energy i.e. – then salt will dissolve.
  • If hydration energy is less than lattice energy i.e. – then salt will not dissolve.

As the size of cation increases, lattice enthalpy decreases. So, solubility increases from Be to Ba.

Chemical properties of s – block elements

Alkali metals –

The alkali metals exhibit high chemical reactivity due to –

  • Low ionization energy
  • Low heat of atomization

Due to their high reactivity, they do not find free in nature.

Reactivity α 1 / I.P

Order of reactivity – Li <  Na < K <  Rb < Cs

Li is stable so reacts slowly with steam while Rb and Cs reacts even with cold water.

Alkaline earth metals –

The alkaline earth metals are less reactive than alkali metals.

Order of reactivity Be < Mg  < Ca < Sr < Ba

Be   =>  No reaction even with hot water.

Mg  => React with hot water.

Ca, Sr, Ba => React with cold water.

Reaction with Air –

Alkali metals –

  • Alkali metals get tarnished (discoloration of metal surface caused by oxidation or make dirty or spotty by exposure to air) due to the formation of oxide at their surface. Hence, they are kept in kerosene or paraffin oil.
  • They burn vigorously in air or oxygen forming oxides.
  • These elements react with moist air to form carbonates as final product.

In dry air only lithium gives nitride gives and oxide both while other elements give only oxide.

  • Lithium forms monoxide (Li2O), sodium forms peroxide (Na2O2) and other elements form superoxide.

Alkaline earth metals –

Alkaline earth metals have strong tendency to lose valence electrons. Therefore, they are very reactive.

  • Except Be, alkaline earth metals are easily tarnished in air as a layer of oxide is formed on the surface.
  • Barium is present in powdered form, so when it is expose to air, it get burst into flame.
  • All elements except Be, converts into carbonates in presence of moist air.
  • In dry air, Be and Mg gives nitride and oxide both while other give only oxide.

The reactivity increases down the group. Alkaline metals are less reactive than alkali metals.

Reaction with oxygen

Alkali metals –

  • Lithium forms monoxide (Li2O), because the size of lithium ion is very small and has strong +ve field around it. So, it combines with small anion like O2-.
  • Sodium reacts with O2 to form peroxide (Na2O2), because Na+ is larger cation and has weak +ve field around it. It can stablise a bigger peroxide ion.
  • K, Rb, Cs forms MO2 type oxides (superoxide) in excess of O2. Therefore, superoxides are paramagnetic and colored.

Key point – Small cation can stablise small anion and large cation can stablise large anion.

Stability Order – Normal oxide > Peroxide > Superoxide

Peroxide gives oxygen (O2) while superoxide gives H2O2.

Alkaline earth metals –

  • Alkaline earth metals react with O2 and forms ‘MO’ type oxides.
  • Ca, Sr, Ba has more reactivity, therefore froms MO2 (peroxides) at low temperature. For example – CaO2, SrO2, BaO2 etc.
  • Due to lattice defect, peroxides are colored.
  • BeO – Amphoteric

MgO – Weak base

CaO, SrO, BaO – strong base

It means basic properties are increasing from top to bottom.

Reaction with Hydrogen –

Alkali metals –

As the size of alkali metals increases from Li to Cs, the M-H bond becomes weak and stability of hydride decreases.

  • Alkali metals combine with H2, and forms ionic hydrides.

         LiH – covalent Hydride and other alkali metal hydrides are ionic.

  • Alkali metals hydrides react with H2O, gives H2 which acts as reducing agent.
  • Order of basic nature – LiH < NaH < KH < RbH < CsH

Thermal stability decreases, basic nature increases.

Alkaline earth metals

  • Except Be, all the alkaline earth metals from MH2 type hydrides (MgH2, CaH2, SrH2, BaH2) or directly with H2.
  • BeH2 and MgH2 are covalent whereas others are ionic.
  • Be and Mg have tendency of polymerization.

Reaction with water

Alkali metals

  • Alkali metals react with water, forms hydroxide and hydrogen gas is evolved.
  • Alkali metals reactivity with water increases from Li to Cs.

Li – Least reactivity towards water, Na – react vigorously, K – react with water and produces flame, Rb and Cs – react explosively.

  • These metals also react with alcohol and form alkoxide and H2.
  • Monoxides give strongly alkaline solution with water.
  • These metals react slowly with water to form H2 and metal hydroxide.
  • Be – not react with water

Mg – reacts only with hot water

Ca, Sr, Ba – reacts with cold water.

Halides –

Alkali metals –

  • Alkali metals react directly with halogen to form MX.
  • Ionic properties of MX increases from LiCl to CsCl.

LiCl – covalent (due to polarization of chloride ion by small lithium ion. Hence, it hydrolyses with water while rests are ionic, so do not hydrolyse.

  • K, Rb, Cs halides react with more halogens to give polyhalides.

Alkaline earth metals

  • Alkaline earth metals react with halogen to form MX2 e.g.- BeCl2, MgCl2, CaCl2.
  • Ionic nature of MX2 increases from BeCl2 to BaCl2. BeCl2 and BaCl2 are covalent in nature.
  • Ba burns in contact with Cl2.
  • Hydrolytic nature of these halides decreases from BeCl2 to BaCl2.

Carbonates –

Alkali metals –

  • All alkali metals form M2CO3 type carbonates. Except Li2CO3, all carbonates are stable towards heat.
  • Thermal stability of carbonates α 1/Ionization potential
  • Order of stability of carbonates –

  Cs2CO3 > Rb2CO3 > K2CO3 > Na2CO3 > Li2CO3

Alkaline earth metals –

  • Alkaline earth metals forms MCO3 type carbonates. All the carbonates decomposes on heating.
  • Order of decreasing stability –

   BaCO3   > SrCO3  > CaCO3 > MgCO3 > BeCO3

Nitrates –

Alkali metals

  • All alkali metals forms MNO3 type nitrates where M = alkali metals. Stability increases from LiNO3 to CsCO3. Lithium decomposes into its oxide and NO2 on heating whereas other nitrates, on heating, give nitrite and oxygen.

Alkaline earth metals

  • Alkaline earth metal forms M(NO3)2 type nitrates where M is alkaline earth metals. Stability increases from Be(NO3)2 to Ba(NO3)2. They decompose on heating, gives oxide and NO2 + O2.

Be(NO3)2  forms a layer of BeO on its surface, so reaction stops.

Nitrides –

Alkali metals –

  • Lithium reacts directly with N2 to form nitride which gives NH3 on reaction with water.

Alkaline earth metals –

  • Only Be and Mg burns in N2 to give M3N2.

Formation of amalgam

Alkali metals –

Alkali metals gives amalgam with Hg. These metals react with other metals, gives mixed metals (alloys).

Alkaline earth metals –

Alkaline earth metals shows same property.

Sulphates –

Alkali metals –

  • Alkali metal forms M2SO4 type sulphates. These are ionic in nature. Their ionic properties increase from from Li to Cs.

                            Li2SO<  Na2SO4  <    K2SO<  Rb2SO<  Cs2SO4

          Li2SO4 least soluble in water.

  • These sulphides on burning with C, form sulphide.
  • Sulphates of alkali metals reacts with sulphates of trivalent metals like Fe+3, Cr+3, Al+3 to form double salts called alum i.e. – M2SO4 . M2(SO4)3 . 24H2O

Alkaline earth metals

  • Alkaline earth metals form MSO4 types sulphates. The ionic nature of their sulphates increases from Be to Ba.

                   BeSO4 < MgSO4 <  CaSO4   <  SrSO< BaSO4

And solubility of these sulphates decreases from BeSO4 to BaSO4 because of small size of Be+2 and Mg+2, so their hydration energy is more than lattice energy.

  • Order of solubility –

                  BeSO4 > MgSO4 >  CaSO4   >  SrSO> BaSO4

  • Order of thermal stability –

BeSO4 < MgSO4 < CaSO4   < SrSO4 < BaSO4

Ionic nature increases, thermal stability increases.

Reaction with acids –

Alkali metals –

  • React vigorously with acids.

Alkaline earth metals

  • Freely react with acids and displace hydrogen.

In alkaline earth metals, Be react slowly with acids like HNO3 because HNO3 is an strong oxidising agent and it forms a thin layer of oxide on the surface of metal, which protect it further attack of acid.

Key point – Be is Amphoteric in nature because it also reacts with NaOH and giving H2 whereas Mg, Ca, Sr, Ba do not react with NaOH and are purely basic.

Solubility in liquid ammonia

Alkali metal –

All the alkali metals are dissolve in liquid ammonia, produces blue color solution which are conductive in nature. As the concentration increases, color changes into bronze, paramagnetic solution becomes diamagnetic and conductivity decreases.

This ammonical electron is responsible for the blue color of the solution and its strong reducing power. The electrical conductivity is due to ammoniated cation and ammonicated electrons.

Alkaline earth metal

Out of all alkaline earth metals, Be and Mg do not dissolve in liquid ammonia because of their small size and high ionization potential whereas Ca, Sr, Ba gives blue solution due to ammoniated electrons.

But if this blue color of the solution is allowed to stand for a long time, color become fade because of metal amide formation. When ammonium salt is added, blue color of the solution disappears due to ammonia formation.

On increasing metal ion concentration, solution converts into bronze color due to the cluster formation of metal ions.

Diagonal Relationship –

Similarties between lithium and magnesium-

  1. Both lithium and magnesium are quite harder and lighter than other elements in their respective groups.
  2. Both react slowly with cold water. Their oxides and hydroxide are much less soluble and their hydroxide decomposes on hearting.
  • Both lithium and magnesium combine with oxygen and form monoxide whereas other members give peroxide and superoxide.

The oxide of Li2O and MgO don’t combine with excess oxygen to form peroxide or superoxide.

  • The carbonates of lithium, magnesium decomposes easily on heating form oxide and CO2 solid bicarbonates are not formed by Li and Mg.
  • Both LiCl and MgCl2 are deliquescent and crystalline form aqueous solution as hydrates e.g – LiCl.2H2O , MgCl2.8H2O
  • Both LiCl and MgCl2 soluble in ethanol.
  • Both LiOH and Mg(OH)2 are weak bases.
  • Both form nitride by direct combination with nitrogen Li3N and Mg3N2.
  • Both have similar electro negativities, similar atomic size and similar ionic radius.

Anomalous behavior of beryllium

  1. It is the hardest of all alkaline earth metals.
  2. The melting and boiling point of the beryllium are the highest.
  3. It is least reactive due to highest ionization potential.
  4. Due to high charge density, its polarizing effect is highest and it forms covalent bond.
  5. Beryllium dissolves in alkali with evolution of hydrogen whereas other alkali metals do not react with alkali.
  • Oxides and Hydroxides of beryllium are amphoteric in nature.

The hydroxide is unstable in water and covalent in nature.

  • Like Al, Its carbide (Be2C) on hydrolysis evolves methane.
  • Due to its small size, it has strong tendency to form complex.
  • It shows diagonal relationship with Al.

Diagonal Relationship between Beryllium and aluminum –

In many of its properties, beryllium resembles aluminum. Thus –

  1. The two elements have same electro negativity and their charge/radius ratios.
  2. Both metals are fairly resistant to action of acids due to protective film of oxide on surface.
  3. Both metals are acted upon by strong alkali to form soluble complexes, beryllates [Be(OH)2]2- and aluminates [Al(OH)4].
  4. The chlorides of both beryllium and aluminum  have chloride structure in vapor phase.

  • Salt of these metals from hydrates ions e.g. – [Be(OH)4]2+ and [Al(OH)6]3+ in aqueous solutions. Due to similar charge/radius ratio of beryllium and aluminum ions have strong tendency of complexes.

For example – Beryllium forms tetrahedral complexes such as BeF42- and [Be(C2H4)2]2- and  aluminum forms octahedral complexes likes AlF63- and [Al(C2O4)3]3- .

Beryllium chloride –       

It is prepared by heating beryllium oxide with chlorine vapors in the presence of carbon.

Structure –     

In solid state, beryllium chloride has polymeric chain structure.

In this case, each Be atom is tetrahedral surrounded by four Cl-atoms in which two chlorine atoms are bonded by covalent bonds while other two chlorine atoms by coordinate bonds.

The polymeric structure of BeCl2 is due to its electron deficient nature. It has four electrons in the valence shell and can accept two electron pairs (from neighbouring Cl atoms forming coordinate bond) to complete their octet.

In vapour phase – BeCl2 exists as dimer which dissociates into linear monomer at 1200k .          

Cement –

It is light grey, heavy fine powder. It is homogenous mixture of silicates and aluminates of calcium which form more than 90% of cement are-

  1. Tricalcium silicate – 3CaO.SiO2
  2. Dicalcium Silicates (slowest setting element) – 2CaO.SiO2
  3. Tricalcium aluminates (fastest setting component )- 3CaO.AlO2
  4. Tetracalcium alumino ferrite – 4CaO.Al2O3.Fe2O3
  5. Lime, silica and alumina are essential constituents of cement. If lime is present in excess, then cement cracks during setting. If lime is less than the required, the cement is weak in strength.

Raw  materials –   

  • Limestone – Provides CaO
  • Clay         –        Provides Al2CO3 and silica ( SiO2).
  • Gyspum    – CaSO4 . 2H2O

In cement, CaO (61.5%), SiO2 (22.5%), Al2O3 (7.5%), MgO (2.5%), Na2O (1.5%), SO3 (1.5%), K2O (1.5%) and Fe2O3 (2.0%).

Setting of cement –

When water is mixed to the cement and the mixture is left, it becomes very hard. This property is called cement.

Mortar –

It is a mixture of cement,sand gravel and water.

Reinforced concrete cement –

when concrete is filled in brams , made up of iron bars is known as RCC. In this, iron imparts extra strength to structure.

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