The s-Block Elements Comprehensive Study

s-Block Elements – Comprehensive Notes

The s-block elements, characterized by their valence electrons residing in the outermost s-orbital, consist of Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals). Their general valence shell electron configuration is ns1 and ns2 respectively. These elements are highly reactive metals.

1. General Characteristics of s-Block Elements

  • Electronic Configuration: Outer electronic configuration is ns1 for Group 1 and ns2 for Group 2.
  • Metallic Character: Exhibit strong metallic character. They are soft (Group 1) to relatively soft (Group 2) and good conductors of heat and electricity. This is due to their large atomic size and low ionization enthalpies, leading to easy loss of valence electrons.
  • Oxidation States:
    • Group 1 (Alkali Metals): Always show a +1 oxidation state in their compounds. They readily lose their single s-electron.
    • Group 2 (Alkaline Earth Metals): Always show a +2 oxidation state in their compounds. They readily lose both their s-electrons.
  • Ionization Enthalpy (IE):
    • Generally low compared to other elements.
    • Decreases down the group due to increasing atomic size and shielding effect.
    • Second ionization enthalpy of Group 1 elements is very high (removal of electron from stable noble gas core).
    • Second ionization enthalpy of Group 2 elements is higher than the first but still relatively low, allowing for the +2 oxidation state.
  • Electronegativity: Very low electronegativity values, indicating a strong tendency to lose electrons.
  • Electropositive Character: Highly electropositive (metallic) nature, which increases down the group.
  • Reducing Nature: Powerful reducing agents as they readily lose electrons to get oxidized. Reducing power generally increases down the group.
  • Flame Colouration: Many s-block elements, when heated in a Bunsen flame, impart characteristic colours due to the excitation of their outermost electrons to higher energy levels and subsequent de-excitation, emitting light of specific wavelengths.
    • Lithium (Li): Crimson Red
    • Sodium (Na): Golden Yellow
    • Potassium (K): Lilac (Pale Violet)
    • Rubidium (Rb): Rubidium Red
    • Caesium (Cs): Sky Blue
    • Beryllium (Be) and Magnesium (Mg): Do not show flame test (due to their small size and high ionization enthalpy, the energy of the flame is insufficient to excite their electrons).
    • Calcium (Ca): Brick Red
    • Strontium (Sr): Crimson Red
    • Barium (Ba): Apple Green
  • Hydration Enthalpy:
    • The enthalpy change when one mole of gaseous ions is hydrated.
    • Decreases with increasing ionic size (down the group) because smaller ions have a higher charge density and thus stronger attraction for water molecules.
    • Li+ has the highest hydration enthalpy among alkali metals.
    • Be2+ has the highest hydration enthalpy among alkaline earth metals.
    • Higher hydration enthalpy leads to a greater extent of hydration (e.g., LiCl.2H2O, MgCl2.6H2O).
    • Hydration also affects ionic mobility: smaller hydrated ion size (larger bare ion size) leads to greater mobility in aqueous solution. For example, Li+ is the most hydrated and thus has the lowest ionic mobility among alkali metals in aqueous solution.

2. Group 1 Elements: Alkali Metals (Li, Na, K, Rb, Cs, Fr)

2.1. General Characteristics and Trends

  • Electronic Configuration: ns1.
  • Atomic and Ionic Radii: Largest in their respective periods, increasing consistently down the group due to addition of new electron shells.
  • Density: Low densities, generally increasing down the group. Exception: Potassium (K) is lighter than Sodium (Na) due to an unusually large increase in atomic volume for K.
  • Melting and Boiling Points: Very low, decreasing down the group. This is due to weak metallic bonding resulting from only one valence electron contributing to the metallic bond.
  • Metallic Lustre: Soft, silvery-white metals that can be cut with a knife. They tarnish rapidly on exposure to air due to surface oxidation.
  • Oxidation State: Always +1.
  • Ionization Enthalpy: Lowest in their respective periods, decreasing down the group. They readily form M+ ions.
  • Electropositive Character: Highly electropositive, increasing down the group.
  • Reducing Nature: Very strong reducing agents. Their reducing power increases down the group in the gaseous state. However, in aqueous solution, Lithium (Li) is the strongest reducing agent (despite higher IE) due to its exceptionally high hydration enthalpy (Li+ gets highly hydrated, releasing a lot of energy, which compensates for the high ionization energy).

2.2. Chemical Properties (Reactions)

  • Reactivity towards Air (Oxygen and Nitrogen):
    • Alkali metals react vigorously with air.
    • They burn in oxygen to form different types of oxides:
      • Lithium (Li): Forms mainly Lithium Oxide (Li2O) (normal oxide, O2- ion). 4Li+O2→2Li2O
      • Sodium (Na): Forms Sodium Peroxide (Na2O2) (peroxide, O2^2- ion). 2Na+O2→Na2O2
      • Potassium (K), Rubidium (Rb), Caesium (Cs): Form Superoxides (MO2) (superoxide, O2- ion). Example: K+O2→KO2
    • All alkali metals react with nitrogen at high temperatures to form nitrides (M3N). Lithium is unique as it forms Lithium Nitride (Li3N) directly with nitrogen at room temperature due to its small size and high charge density. 6Li+N2→2Li3N
  • Reactivity towards Water (H2O):
    • React violently and exothermically with water to form hydroxides and hydrogen gas.
    • 2M(s)+2H2O(l)→2MOH(aq)+H2(g)
    • Reactivity and vigour of reaction increase down the group.
  • Reactivity towards Hydrogen (H2):
    • Form ionic hydrides (MH) when heated with hydrogen at about 673 K (e.g., 2M+H2→2MH).
    • LiH has significant covalent character due to the small size of Li+. Other hydrides are purely ionic.
  • Reactivity towards Halogens (X2):
    • React readily and vigorously with halogens to form ionic halides (MX).
    • 2M+X2→2MX
    • Reactivity increases down the group.
    • Halides are typically high melting point solids. The ionic character of halides decreases in the order MF > MCl > MBr > MI.
  • Solutions in Liquid Ammonia:
    • Alkali metals dissolve in liquid ammonia to produce deep blue, conducting solutions. This is due to the formation of ammoniated cations and ammoniated electrons.
    • M+(x+y)NH3→[M(NH3)x]++[e(NH3)y]−
    • These solutions are paramagnetic (due to unpaired ammoniated electrons).
    • At higher concentrations, the solutions become bronze-coloured and diamagnetic (due to pairing of ammoniated electrons).
    • These solutions are strong reducing agents.

2.3. Important Compounds of Alkali Metals

  • Sodium Carbonate (Washing Soda), Na2CO3.10H2O:
    • Prepared by the Solvay Process (Ammonia-Soda Process).
    • Raw materials: Sodium Chloride (NaCl), Ammonia (NH3), Limestone (CaCO3) as a source of CO2, and water.
    • Key Reactions:
      • CaCO3 (s) heat-> CaO (s) + CO2 (g)
      • 2NH3(g)+H2O(l)+CO2(g)→(NH4)2CO3(aq)
      • (NH4)2CO3(aq)+H2O(l)+CO2(g)→2NH4HCO3(aq)
      • NH4HCO3(aq)+NaCl(aq)→NaHCO3(s)+NH4Cl(aq) (Sodium bicarbonate precipitates due to low solubility)
      • 2NaHCO3(s) heat-> Na2CO3(s)+H2O(g)+CO2(g) (CO2 is recycled)
      • 2NH4Cl(aq)+CaO(s)→CaCl2(aq)+2NH3(g)+H2O(l) (Ammonia is recovered)
    • Uses: In laundries, glass and paper manufacturing, detergents, and for manufacturing other sodium compounds like caustic soda.
  • Sodium Hydroxide (Caustic Soda), NaOH:
    • Prepared by the Castner-Kellner cell (electrolysis of concentrated aqueous sodium chloride solution, brine).
    • At Anode (Graphite): 2Cl−(aq)→Cl2(g)+2e−
    • At Cathode (Mercury): Na+(aq)+e−Hg​Na/Hg (sodium amalgam)
    • Reaction of Amalgam: 2Na/Hg+2H2O(l)→2NaOH(aq)+2Hg(l)+H2(g) (Mercury is recycled)
    • Uses: In soap, paper, and artificial silk manufacturing, petroleum refining, and in textiles.
  • Sodium Bicarbonate (Baking Soda), NaHCO3:
    • Can be prepared from sodium carbonate or as an intermediate in the Solvay process.
    • Uses: As a component of baking powder (reacts with tartaric acid to release CO2 for leavening), an antacid (neutralizes stomach acid), and in soda-acid fire extinguishers.
  • Sodium Chloride (NaCl):
    • Common salt, obtained from seawater (by evaporation) and underground deposits.
    • Uses: Common dietary salt, food preservative, in the manufacturing of sodium carbonate and sodium hydroxide.

3. Group 2 Elements: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)

3.1. General Characteristics and Trends

  • Electronic Configuration: ns2.
  • Atomic and Ionic Radii: Smaller than corresponding alkali metals (due to higher nuclear charge), but increase down the group.
  • Density: Higher than alkali metals. Generally increase down the group, but Magnesium (Mg) is slightly an exception.
  • Melting and Boiling Points: Higher than alkali metals (due to stronger metallic bonding from two valence electrons), generally decrease down the group (irregular trend).
  • Metallic Lustre: Silvery-white, relatively soft (harder than alkali metals) and generally tarnish in air.
  • Oxidation State: Always +2.
  • Ionization Enthalpy: Higher than alkali metals (due to smaller size and higher effective nuclear charge). Decreases down the group. The second IE is higher than the first but is still low enough for the +2 state.
  • Electropositive Character: Less electropositive than alkali metals, but increases down the group.
  • Reducing Nature: Strong reducing agents, but less strong than alkali metals. Reducing power increases down the group.

3.2. Chemical Properties (Reactions)

  • Reactivity towards Air (Oxygen and Nitrogen):
    • Form normal oxides (MO) and nitrides (M3N2) when heated in air.
    • 2M+O2→2MO
    • 3M+N2→M3N2
    • Barium (Ba) and Strontium (Sr) can also form peroxides (MO2) upon heating in excess oxygen.
  • Reactivity towards Water (H2O):
    • React with water to form hydroxides (M(OH)2) and hydrogen gas.
    • M(s)+2H2O(l)→M(OH)2(aq)+H2(g)
    • Reactivity increases down the group. Beryllium (Be) does not react with water. Magnesium (Mg) reacts only with hot water or steam. Calcium (Ca), Strontium (Sr), Barium (Ba) react increasingly vigorously with cold water.
  • Reactivity towards Hydrogen (H2):
    • Form hydrides (MH2) when heated with hydrogen.
    • Beryllium Hydride (BeH2) is covalent and polymeric. Other hydrides (MgH2, CaH2, etc.) are ionic.
  • Reactivity towards Halogens (X2):
    • React readily to form halides (MX2).
    • M+X2→MX2
    • Beryllium Halides (BeX2) are predominantly covalent and polymeric in solid state. They exist as dimeric bridged structures in vapor phase. Other halides are ionic.
  • Solutions in Liquid Ammonia:
    • Like alkali metals, alkaline earth metals also dissolve in liquid ammonia to form deep blue-black conducting solutions due to the formation of ammoniated ions and electrons.

3.3. Important Compounds of Alkaline Earth Metals

  • Calcium Oxide (Quicklime / Lime), CaO:
    • Prepared by heating limestone (CaCO3) in a rotary kiln at 1070-1270 K.
    • CaCO3(s) heat-> CaO(s)+CO2(g)
    • Uses: In cement and mortar, metallurgy (flux), manufacturing of soda lime, and in the purification of sugar.
  • Calcium Hydroxide (Slaked Lime), Ca(OH)2:
    • Prepared by adding water to quicklime (a process called “slaking of lime”).
    • CaO(s)+H2O(l)→Ca(OH)2(aq) (exothermic reaction)
    • A suspension of Ca(OH)2 in water is called “milk of lime”, and an aqueous solution is “limewater”.
    • Uses: As a building material (mortar, plaster), in white wash, as an antacid, and in the preparation of bleaching powder.
  • Calcium Carbonate (Limestone), CaCO3:
    • Found naturally as limestone, marble, chalk, dolomite.
    • Uses: As a building material, a flux in metallurgy, an antacid, and in the manufacture of cement, glass, and paper.
  • Plaster of Paris (POP), (CaSO4).1/2H2O:
    • Chemically known as calcium sulfate hemihydrate.
    • Prepared by heating Gypsum (CaSO4.2H2O) at 393 K.
    • CaSO4.2H2O heat at 393 K-> (CaSO4).1/2H2O+3/2H2O
    • Note: Heating above 393 K leads to “dead burnt plaster” (anhydrous CaSO4), which loses its setting property.
    • Uses: Building material (plaster), casts for fractured bones, dental impressions, making statues, and in dentistry.
  • Cement:
    • A complex mixture of silicates (e.g., dicalcium silicate, tricalcium silicate) and aluminates (e.g., tricalcium aluminate) of calcium, along with small amounts of iron oxide.
    • Key ingredients: Limestone (CaCO3) and Clay (provides Al2O3, SiO2, Fe2O3).
    • Setting of cement: An exothermic process involving hydration and hydrolysis reactions that form a hard, solid mass.

4. Anomalous Properties of Lithium and Beryllium

These elements, being the first members of their respective groups, show properties different from other members. This is primarily due to:

  • Very small atomic and ionic size.
  • High ionization enthalpy.
  • High electronegativity.
  • High polarizing power (charge/radius ratio).
  • Absence of d-orbitals in their valence shell.

4.1. Anomalous Behaviour of Lithium (Li)

  • Hardness: Much harder than other alkali metals.
  • Melting/Boiling Points: Higher than other alkali metals.
  • Reactivity with Water: Least reactive towards water among alkali metals.
  • Oxides: Forms only normal oxide (Li2O) on burning in oxygen. Others form peroxides/superoxides.
  • Nitride Formation: Unique in forming stable nitride (Li3N) directly with atmospheric nitrogen at room temperature.
  • Superoxides: Does not form superoxides.
  • Hydrates: Forms hydrated salt (LiCl.2H2O) due to high hydration enthalpy. Others mostly do not.
  • Solubility: Its salts (e.g., carbonates, fluorides, phosphates) are generally less soluble than those of other alkali metals.

4.2. Anomalous Behaviour of Beryllium (Be)

  • Hardness: Much harder than other alkaline earth metals.
  • Melting/Boiling Points: Higher than other alkaline earth metals.
  • Reactivity with Water: Does not react with water, even at high temperatures.
  • Bonding: Forms predominantly covalent compounds (e.g., BeCl2, BeH2). Other alkaline earth metals form ionic compounds.
  • Nature of Oxides/Hydroxides: Oxides (BeO) and hydroxides (Be(OH)2) are amphoteric (react with both acids and bases), unlike others which are basic.
  • Coordination Number: Does not show coordination number more than 4 (due to absence of d-orbitals). Other alkaline earth metals can show higher coordination numbers.
  • Polymeric Structures: Forms polymeric and covalent hydrides and halides (e.g., polymeric BeH2, bridged dimeric BeCl2 in vapor phase, polymeric in solid).

5. Diagonal Relationship

Similarity in properties between an element and the element diagonally opposite to it in the next group. This occurs because the effect of increasing nuclear charge across a period is somewhat compensated by the increasing atomic size down a group, leading to similar ionic potentials (charge/radius ratio) and electronegativity.

  • Lithium (Li) and Magnesium (Mg):
    • Both are hard metals.
    • Both react slowly with water (Li is less reactive than Na, Mg is less reactive than Ca).
    • Both form nitrides directly with N2 (Li3N, Mg3N2).
    • Both form normal oxides on burning in oxygen.
    • Hydroxides (LiOH, Mg(OH)2) are weak bases and sparingly soluble.
    • Carbonates (Li2CO3, MgCO3) decompose easily on heating to form oxides and CO2.
    • Chlorides (LiCl, MgCl2) are deliquescent and form hydrates (LiCl.2H2O, MgCl2.6H2O).
    • Both do not form superoxides.
  • Beryllium (Be) and Aluminium (Al):
    • Both form predominantly covalent compounds.
    • Both have a strong tendency to form complexes.
    • Oxides (BeO, Al2O3) and hydroxides (Be(OH)2, Al(OH)3) are amphoteric.
    • Both are attacked by strong acids and strong bases.
    • Carbides (Be2C, Al4C3) give methane on hydrolysis (Be2C+4H2O→2Be(OH)2+CH4; Al4C3+12H2O→4Al(OH)3+3CH4).
    • Chlorides (BeCl2, AlCl3) are Lewis acids and form bridged polymeric structures in the solid state and dimeric bridged structures in the vapor phase.

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