s-block elements Chapter 10 Class XI

s-block elements

Chapter - 10

Representative elements = s + p block elements

s-block : Group 1 and group 2

Group 1 = alkali metals as they form water soluble bases like NaOH, KOH

     = Li, Na, K, Rb, Cs, Fr

General configuration = ns¹,   valency = valence electron = 1 (all metals)

 Group 2 = alkaline earth metals or ground metals

     = Be, Mg, Ca, Sr, Ba, Ra

General configuration = ns²,   valency = valence electron = 2 (all metals).

Occurrence – All are reactive metals and do not exist in free state but are present in combined state as halides, oxides, carbonates, silicates, nitrates and phosphates.

Diagonal Relationship – Some elements of certain groups in second period resemble with the certain elements of next higher period (third period). This is known as diagonal relationship.                                    

Reasons: Similar atomic size, electropositive character, ionization energy, polarising power.

For Eg. (i) Li is diagonally related to Mg

(ii) Be is diagonally related to Al

(iii) B is diagonally related to Si.

Anomalous behaviour of elements of second period –

Reasons:

a) small size,

b) absence of d-orbital in valence shell,

c) high ionization energy,

d) high polarisation power of its cation.

Alkali metals - Li, Na, K, Rb, Cs, Fr (ns¹)

Occurrence:

Na and K are 7^th and 8^th most abundant elements in earth’s crust.

Li= spodumene [LiAl(SiO₃)₂]; 

Na= as NaCl (rock salt), Na₂B₄O₇.10H₂O (borax), Na₂CO₃.10H₂O (washing soda), NaNO₃ (saltpetre), Na₂SO₄; 

K= KCl, K₂SO₄.

Physical properties:

1) atomic size –

a) Alkali metals have largest size among all groups.

Group 1 > group 2 elements.

Reason: atomic size decreases across the period because as atomic number increases, nuclear charge increases, force of attraction between nucleus and outermost electron increases and hence size decreases.

b) Down the group, size increases as number of shells increases.

Order of ionic and atomic radii: Li < Na < K < Rb < Cs.

2) Ionization enthalpy –

a) Alkali metals have low ionization energies.

Reason: as size decreases across the period, force of attraction increases between nucleus and outermost electrons and hence energy required to remove an electron from outermost shell. So, IE increases.

b) Down the group, IE increases as size increases. Li > Na > K > Rb > Cs.

3) Oxidation state = +1. Alkali metals are highly electropositive metals.

4) They are soft metals and can be cut by knife.

5) Metallic character increases down the group.

6) They have high melting points and boiling points as their intermetallic bonds are weak.

7) All alkali metals form ionic bonds except Li (Li forms covalent bonds) and have low densities.

8) Na and K are highly reactive and catch fire easily by reacting with air and moisture. So, they are kept immersed in kerosene oil.

9) Flame colouration – All alkali metals impart colour to flame.

Li = Crimson, Na = Yellow, K = Pale violet, Rb = Violet, Cs = Blue.

Reason: On heating the alkali metal or its salt in a flame, the electrons are excited easily to higher energy levels because of absorption of energy. When these excited electrons return to their ground state, they emit energy which falls in visible region and imparts a characteristic colour to the flame. Different colours appear due to the difference in the amount of energies emitted. Visible region wavelength = 400nm to 750nm.

Note: Be and Mg do not show flame colouration because they have small size and high ionization energy. Hence, it is difficult for them to lose electrons.

10) Photoelectric effect –

It is the phenomenon of emission of electrons from metal surface when electromagnetic radiation of sufficient energy are made to strike against it. All alkali metals exhibit photoelectric effect except Li as Li has very small size and high ionization energy, so it does not release electron easily when exposed to light.

Chemical properties:

1)     Reaction with oxygen – alkali metals form basic oxides with oxygen.                There are three forms of oxides are – oxide (O^2-), peroxide (O₂^2-) and superoxide (O₂^-).

Oxides: 4 Li + O₂ → 2Li₂O

               4 Na + O₂ → 2 Na₂O

               4 K + O₂ → 2K₂O

Peroxides: 2 Na + O₂ → Na₂O₂

                  2 K + O₂ → K₂O₂

Superoxides: K + O₂ → KO₂

Ø  Lithium can form only oxide (Li₂O) and cannot form peroxide and superoxide because of its very small size. Li⁺ has strong positive field (high charge density) around it which attracts the negative charge so strongly that it does not allow oxide anion O^2- to combine with another oxygen atom to from peroxide ion, O₂^2-.

While Na⁺ ion has larger size than Li⁺, so it has weaker positive field around it which allow oxide anion to combine with another oxygen atom to form peroxide ion.

Ø  Superoxides are paramagnetic and hence are coloured due to the presence of unpaired electrons in π^*2p_x² = π^*2p_x¹. Normal oxides are diamagnetic and re colourless.

2)     Reaction with hydrogen – These elements form hydrides.

2 M + H₂ → 2 MH

Such as 2 Na + H₂ → 2 NaH

Order of reactivity of alkali metals towards hydrogen decreases down the group because lattice energy of hydrides decreases as the size of cation increases while going down the group and hence stability of hydrides decreases.

Stability order: LiH > NaH > KH > RbH > CsH.

3)     Reactivity and electrode potential –

Electrode potential is a measure of the tendency of an element to lose electrons in the aqueous solution.

Ø  Alkali metals act as strong reducing agents because they have to lose only one electron from their valence shell and have large atomic size and low ionization energy.

Ø  Down the group, ionization energy decreases so tendency to lose electrons increases and hence, reducing character increases down the group. Order: Na < K < Rb < Cs < Li.

Ø  Lithium is the strongest reducing agent in aqueous solution.

Reason:

Reducing character depends upon ionisation energy and tendency to lose electron from isolated gaseous atom.

Factors affecting: (i) Sublimation: Na(s) → Na(g)

(ii) Ionisation Enthalpy: Na(g) → Na⁺(g) + e^-

(iii) Heat of hydration.

More the tendency of an element to lose electrons, more is the reducing character.

But in case of Lithium, it has smallest size and high heat of hydration enthalpy which compensates the ionisation energy required to remove electron. So, it can easily remove an electron and behaves as a strong reducing agent. Lithium has most negative standard electrode potential. (E⁰ = -3.05V)

4)     Reaction with water – These elements form hydroxides and react vigorously with water.

Ø          2 Na + 2 H₂O → 2 NaOH + H₂,                   2 K + 2 H₂O → 2 KOH + H₂

Ø          Na₂O + H₂O → 2 NaOH

Ø                  K₂O₂ + 2 H₂O → 2 KOH + H₂O₂

Ø                 2 KO₂ + 2 H₂O → 2 KOH + H₂O₂ + O₂.

5)     Reaction with halogens – These elements form halides.

2M + X₂ → 2MX

Reactivity order: Li < Na < K < Rb < Cs & F₂ > Cl₂ > Br₂ > I₂.

6)     Covalent character and polarisation – When a cation approaches an anion, electron cloud of anion is attracted towards cation and gets distorted. This effect is called polarisation.

The power of cation to polarise anion is called polarisation power and tendency of anion to get polarised is called its polarizability.

Greater the polarisation, greater is the covalent character.

By Fajan’s rule – Covalent character increases if:

(i)            The size of cation is small.

Order of covalent character: LiCl > NaCl > KCl > RbCl > CsCl

(ii)          Anion is larger in size.

Order of covalent character: LiI > LiBr > LiCl > LiF

(iii)         Charge of an ion is high.

Order of covalent character: NaCl < MgCl₂ < AlCl₃.

7)     Lattice energy – Amount of energy required to separate one mole of solid ionic compound into its gaseous ions or vice versa.

Greater the lattice energy is, higher is the melting point and boiling point of a compound, greater is the stability of the compound and lesser is the solubility.

Order of stability: LiF > LiCl > LiBr > LiI.

Bond strength of LiF is higher because of small sized cation and anion forming a stable lattice and has higher lattice enthalpy.

8)     Hydration energy – It is the amount of energy released when one mole of gaseous ions combined with water to form hydrated ions.

M⁺ (g) + aq → M⁺ (aq) + Hydration energy

Higher the hydration energy of ions, greater is the solubility of compound in water.

Smaller the size of an ion, more it can be hydrated and greater is its hydration radius and has less ionic mobility.

Hence, Li salts are most hydrated.

Ionic Radius: Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

Hydrated Radius: Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

Ionic Mobility: Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺.

9)     Solutions in liquid ammonia – All alkali metals dissolve in liquid ammonia giving highly conducting deep blue solutions.

As these solutions contain ammoniated cations and ammoniated electrons, they are good conductors of electricity.

M + (x+y) NH₃ → M⁺(NH₃)_x + e^-(NH₃)_y

When light falls on these electrons, they get excited to higher energy levels by absorbing energy and when they come back to ground state, they impart blue color. In concentrated solution, blue color changes to bronze and becomes diamagnetic.

10)  Salts of oxoacids – Oxoacids are those in which acidic proton is on a hydroxyl group with an oxo group attached to the same atom.

For e.g. H₂SO₄, H₂SO₃, H₂CO₃

Alkali metals form oxoacid salts which are soluble in water and thermally stable.

Stability order: LiCO₃ < Na₂CO₃ < K₂CO₃ < Rb₂CO₃ < Cs₂CO₃.

As Li⁺ is a smaller cation and CO₃^2- is a larger anion. So, Li₂CO₃ has less lattice energy and is stable.

Li₂CO₃ → Li₂O + CO₂

Note: Larger anions = CO₃^2-, HCO₃^-, SO₄^2-, SO₃^2-, I^-, HSO₃^-‑ .

Uses:

Li metal is use to make alloys with Al to make aircraft parts and with Mg to make armour plates. Li is used in electrochemical cells.

Liquid sodium metal is used as a coolant in nuclear reactors.

Potassium is used as an important constituent of fertilizers. KOH is used for absorbing CO₂.

Biological importance of Na & K:

Na⁺ and K⁺ ions develop a sodium-potassium pump which participate in the transmission of nerve signals by developing a potential across cell membranes.

These ions also help in regulating the flow of water, sugars and amino acids across cell membranes. Na⁺ ions are present outside the cells while K⁺ ions are present within the cells.

K⁺ ions activate many enzymes, participate in oxidation of glucose to produce ATP.

Anomalous properties of Lithium: Different properties than other alkali metals,

(i)            Li is much harder and has high melting and boiling point.

(ii)          Li is least reactive but it is the strongest reducing agent among all alkali metals.

(iii)         LiCl is deliquescent and forms LiCl.2H₂O.

(iv)         Li forms normal oxide only and cannot form peroxide & superoxide.

(v)          Li carbonates, hydroxides and nitrates on heating decomposes to give oxide while other metal carbonates & hydroxides do not decompose at all whereas alkali metals nitrates give nitrites on heating.

Li₂CO₃ → Li₂O + CO₂

2LiNO₃ → Li₂O + 2NO₂

NaNO₃ → NaNO₂ + ½ O₂.

Diagonal relationship: Similarities between Li and Mg:

(i)            Both Li and Mg are harder and lighter in their respective groups.

(ii)          Both of them form nitride by direct combination of Li and Mg with N₂ [ Li₃N, Mg₃N₂]

(iii)         Both form normal oxide.

(iv)         LiCl and MgCl₂ are deliquescent and form hydrates. LiCl.2H₂O, MgCl₂.8H₂O.

(v)          Both react slowly with water.

(vi)         Their hydroxides and oxides are less soluble in water.

Note: Learn preparation, physical properties and chemical properties of some important compounds of Na. (Given in NCERT)

Group II: Be, Mg, Ca, Sr, Ba, Ra

General configuration = ns²

General Characteristics:

Physical properties:

(i)            Density – Harder & denser than alkali metals due to smaller size and strong intermetallic bonds.

(ii)          Melting point and boiling point – Higher than alkali metals but no regular trend.

(iii)         Ionization energy and electropositive character – Lower I.E. as compared to p-block elements but it decreases down the group. They are less electropositive than alkali metals. 1^st I.E. is lower than 2^nd I.E.

(iv)         Atomic size and ionic size – They are smaller than alkali metals due to greater nuclear charge.

(v)          Oxidation state – only +2 (high lattice or hydration energy compensates for 2^nd I.E.)

(vi)         Reducing properties – They are strong reducing agents but weaker than alkali metals and reducing character increases down the group.

(vii)        Flame coloration – Be & Mg do not impart colour to the flame as flame energy is insufficient for excitation, the colours of the flame are as follows: Ca – Brick red, Sr – Crimsom, Ba – Apple green, Ra – Crimson.

Chemical Properties:

(i)            Reaction with air and water – They are less reactive then alkali metals. Be & Mg are kinetically inert towards air & water because of the formation of a film of oxide on their surface. Be does not react with steam even at red heat & does not gets oxidized in air below 873K though powdered Be reacts to form BeO & Be₃N₂. The reactivity towards O₂ increases down the group. BeO is amphoteric whereas other oxides are basic. The basic character increases down the group. BeO and MgO are almost insoluble in water due to high lattice energy.

(ii)          Solubility in water – In case of carbonates and sulphates (large anions), stability increase and solubility decreases down the group.

(iii)         Halides – They react directly with halogens to give MX₂.

BeCl₂ : It is prepared indirectly from its oxide.

BeO + C + Cl₂      →  BeCl₂ + CO

Ø  In solid state, structure of BeCl₂ is polymeric with two chlorine atoms are covalently bonded. In vapour phase below 1200K, it exists as dimer and above 1200K, it is a linear monomer.

Ø  Beryllium halides are covalent and all other metal halides are ionic in nature. Ionic character increases down the group.

Ø  Beryllium halides are Lewis acids as the octet of central atom is incomplete.

Ø  Beryllium halides are soluble in organic solvents.

Ø  Fluorides of other alkaline earth metals are insoluble in water due to high lattice energy but chlorides, bromides and iodides are soluble in water.

Ø  Anhydrous halides are hygroscopic and form hydrates such as MgCl₂.6H₂O, CaCl₂.6H₂O, etc.

Ø  Beryllium chloride fumes in moist air due to hydrolysis.

Anomalous behaviour of Be:

It shows anomalous behaviour due to small size, high charge over mass ratio (high polarization power), strong intermetallic bonding, high hydration energy and high ionization energy.

(i)            Beryllium salts are less stable.

(ii)          Be does not react with water.

(iii)         Be is an amphoteric metal.

(iv)         Be does not impart colour to the flame.

(v)          Beryllium carbide react with water to give methane & other carbides give acetylene.

Diagonal Relationship between Be & Al:

(i)            Be & Al are not attacked by acids because of an oxide layer on them.

(ii)          Be & Al form floro complex anions [BeF₄]^2- & [AlF₆]^3- and other alkaline earth metals d not form floro complex.

(iii)         Oxide and hydroxide of Be and Al are amphoteric in nature.

(iv)         BeCl₂ is covalent, polymeric and bridged structure, AlCl₃ is also covalent, bridged and dimer. Both are soluble in organic solvents and are lewis acids.

Important reactions:

(i)            2 BeCl₂ + LiAlH₄ → 2 BeH₂ + LiCl + AlCl₃

(ii)          BeO and Be(OH)₂ are amphoteric.

BeO + H₂O → Be(OH)₂

Be(OH)₂ + 2 OH^- → [Be(OH)_4]^2-

Be(OH)₂ + 2 HCl + 2 H₂O → [Be(OH)₄]Cl₂

(iii)         2 Be(NO₃)₂ → 2 BeO + 4 NO₂ + O₂

2 Mg(NO₃)₂ → 2 MgO + 4 NO₂ + O₂

Biological importance of Mg & Ca:

(i)            Mg – main component of chlorophyll and ATP.

(ii)          Ca – helps in blood coagulation, helps in muscle contraction & relaxation, in nerve transmission, main constituent of bones & teeth.

Uses: Mg is used in making alloys (being lighter in weight).

Mg(OH)₂ is used as an antacid (milk of magnesia).

Note: Learn preparation, physical properties and chemical properties of some important compounds of Ca. (Given in NCERT)