Overview — Module 6: Hydrogen and s-Block Elements

Hydrogen, alkali metals (sodium, potassium) and alkaline earth metals (magnesium, calcium) shape everyday life — vanaspati from hydrogenation, yellow street lights from sodium, plaster of Paris for bones, caustic soda for soap. This lesson covers occurrence, isotopes, properties and uses of hydrogen; water, heavy water, H₂O₂ and hydrides; alkali and alkaline earth metals (electronic configuration, trends, oxides, water reaction, diagonal relationships); anomalous Li and Be; biological roles of Na⁺, K⁺, Mg²⁺, Ca²⁺; and manufacture of NaOH, Na₂CO₃ and lime.

s-Block elements have outer configuration ns¹ (Group 1) or ns² (Group 2). They are highly electropositive, form ionic compounds, and show clear group trends with important exceptions for the first members.

Section 1: Hydrogen (17.1)

Hydrogen is the first element (1s¹) — one proton and one electron. Cavendish studied "inflammable air"; Lavoisier named it hydrogen. Free H₂ appears in volcanic gases and stellar atmospheres; stars fuse H into He, releasing enormous energy. On Earth, hydrogen is mainly in water, petroleum, natural gas, and biomolecules. Light H₂ escapes Earth's gravity, so free atmospheric H₂ is scarce.

17.1.2 Position in the Periodic Table

H resembles alkali metals (ns¹, can form H⁺) and halogens (gains e⁻ → H⁻ 1s²; liberated at anode in electrolysis of alkali hydrides). It also has a half-filled shell like Group 14. Therefore hydrogen is given a unique position — neither fully in Group 1 nor Group 17.

17.1.4–17.1.6 Properties and Uses

H₂ is colourless, odourless, lightest gas, sparingly soluble, adsorbed on Pt/Pd. Chemical: burns with pale blue flame (2H₂ + O₂ → 2H₂O); reduces heated metal oxides (ZnO, CuO); forms NH₃, CH₄, HCl with non-metals; forms ionic hydrides with electropositive metals (2Na + H₂ → 2NaH).

Uses: synthetic petroleum from coal; methanol; hydrogenation of oils (Ni catalyst, ~443 K → vanaspati); ammonia/fertilizers; rocket fuel; balloon filling (low density).

Isotope effects are unusually large for hydrogen because the relative mass difference between H, D and T is huge compared with isotopes of heavier elements. H₂ adsorbs and reacts with Cl₂ much faster than D₂. Tritium’s low-energy β emission and half-life of 12.33 years make it useful as a tracer but demand careful handling. Natural hydrogen is essentially protium chemistry; heavy water and deuterated reagents are specialised laboratory and nuclear materials.

Section 2: Compounds of Hydrogen (17.2)

17.2.1 Water

H₂O is covalent, bent (H–O–H 104.5°), highly polar due to high EN of oxygen. Intermolecular H-bonds dominate liquid structure. Ice floats: each O is tetrahedrally linked to four H (two covalent, two H-bonds) in an open network with voids — density of ice < liquid water. Essential for aquatic life in winter.

17.2.2 Heavy Water (D₂O)

Water with deuterium instead of protium. Prepared by prolonged electrolysis — O–H breaks faster than O–D, so residual water enriches in D₂O (~30,000 L ordinary water → 1 L D₂O). Uses: nuclear reactor moderator (slows neutrons); reaction mechanism studies; starting material for deuterium compounds (CaC₂ + 2D₂O → C₂D₂ + Ca(OD)₂).

17.2.3 Hydrogen Peroxide (H₂O₂)

Prepared by various methods (e.g. from peroxodisulphates / organic routes in industry). Acts as both oxidising and reducing agent. Decolorises acidified KMnO₄ (oxidation of H₂O₂ / reduction of MnO₄⁻ depending on conditions). Uses: bleaching, antiseptic, rocket propellant (concentrated), laboratory oxidant.

17.2.4 Hydrides

Ionic hydrides: Groups 1–2 heat metal in H₂ → white crystalline MH / MH₂. Contain H⁻; electrolysis of fused hydride liberates H₂ at anode. React violently with water: H⁻ + H₂O → OH⁻ + H₂.

Covalent hydrides: Groups 13–17; mostly gases (CH₄, NH₃, H₂S, HCl). Stability decreases down a group (HI less stable than HCl).

Interstitial hydrides: Transition metals trap H in lattice (non-stoichiometric). Pd absorbs H₂ reversibly — used to purify hydrogen. Density of the hydride is lower than the metal because the lattice expands. Composition can be varied with temperature and pressure; pumping at high temperature releases H₂ again.

Ionic hydrides prove H⁻ by electrolysis: fused alkali hydride (often in alkali halide melt) evolves hydrogen at the anode — opposite to aqueous electrolysis where H₂ appears at the cathode. That single observation cements the ionic model for Group 1 and 2 hydrides and links this chapter to L13 electrochemistry.

17.2.5 Hydrogen Economy

Oceans hold vast hydrogen in water. H₂ as fuel gives only H₂O — clean. Challenges: cheap production (electrolysis is energy-costly; solar-catalysed water splitting is researched) and safe storage. Fuel cells (L13) pair H₂ with O₂ for electricity.

Section 3: Alkali Metals — Group 1 (17.3.1)

Li, Na, K, Rb, Cs (Fr radioactive). Outer config ns¹. Soft, low density, low mp, highly electropositive. Occurrence: NaCl (sea), Chile saltpetre (NaNO₃), carnallite (KCl·MgCl₂·6H₂O). Flame colours: Li crimson red, Na yellow, K pale violet, Rb/Cs violet.

Chemical properties: With O₂ — Li forms Li₂O; Na forms Na₂O₂; K, Rb, Cs form superoxides MO₂. With water — vigorous formation of hydroxide + H₂ (Li less vigorous than expected despite most negative E° — small size, high IE). Hydrides MH; halides MX. Basic character of oxides and hydroxides increases down the group.

Diagonal Relationship Li–Mg & Anomalous Lithium

First member of Group 1 resembles second of Group 2 (similar polarizing power). Li and Mg: similar EN; form nitrides and carbides; normal oxides; carbonates/nitrates/hydroxides decompose to oxides; halides soluble in organic solvents.

Anomalous Li: forms stable Li₂O (not peroxide); Li₂CO₃, LiOH, LiNO₃ decompose easily; LiF, Li₂CO₃, Li₃PO₄ less soluble; more covalent compounds; forms Li₃N with N₂; reacts slowly with water; no solid bicarbonate/superoxide.

Physical trends down Group 1: atomic/ionic radii increase; ionization enthalpy decreases; electropositive character and reactivity increase; melting and boiling points decrease (weaker metallic bonding as atoms enlarge). Density generally increases (K is a slight exception). All show characteristic flame colours because the outer electron is easily excited and emits visible light on return to the ground state — sodium’s intense yellow is the classic street-lamp colour.

Polarizing power ∝ (ionic charge)/(ionic radius)² explains diagonal relationships. From Li to Na, charge is constant but radius increases → polarizing power falls. From Li to Be, charge increases and radius decreases → polarizing power rises. From Be to Mg, polarizing power falls again and becomes comparable to Li — hence Li–Mg similarity. The same logic pairs Be with Al.

Manufacture: NaOH, Na₂CO₃, NaHCO₃

Castner–Kellner: Brine electrolysis with Hg cathode. Cl₂ at Ti anodes; Na dissolves in Hg (high H₂ overvoltage). Amalgam + water → pure NaOH + H₂; Hg recycled. Products: NaOH, Cl₂, H₂.

Solvay (ammonia–soda): NH₃-saturated brine + CO₂ (from CaCO₃) → NaHCO₃ precipitates (common ion / low solubility) → heat → Na₂CO₃. NH₃ recovered from NH₄Cl with CaO. Economical: cheap raw materials, recycle NH₃ and CO₂. Washing soda = Na₂CO₃·10H₂O.

NaHCO₃ from CO₂ through cold Na₂CO₃ solution. Uses of Na₂CO₃: glass, soap, hard-water treatment, paper, volumetric analysis.

Biological Na⁺ and K⁺: maintain osmotic pressure; nerve/muscle function; form buffers with weak acids (extracellular Na⁺, intracellular K⁺).

Alkali metal carbonates (except Li₂CO₃) are thermally very stable and melt without decomposition; stability increases down the group.

Section 4: Alkaline Earth Metals — Group 2 (17.3.2)

Be, Mg, Ca, Sr, Ba (Ra radioactive). Outer config ns². Smaller and less reactive than neighbouring alkali metals; still strong reductants. Occurrence: carnallite, limestone/chalk/marble (CaCO₃), dolomite (CaCO₃·MgCO₃), gypsum (CaSO₄·2H₂O), beryl (Be₃Al₂(SiO₃)₆).

Trends: atomic/ionic radii ↑ down group; IE ↓; metallic character ↑; density generally ↑. Flame: Ca brick red, Sr crimson, Ba sea green; Be and Mg no colour (high IE, small size).

Chemical: Burn to MO (BeO–CaO); larger metals form peroxides more readily. Be does not react with water/steam; Mg reacts with steam; Ca, Sr, Ba with cold water → M(OH)₂ + H₂. Halides MX₂. Carbonates sparingly soluble; thermal stability BeCO₃ < … < BaCO₃. Sulphates: BeSO₄, MgSO₄ soluble; solubility decreases to BaSO₄; thermal stability of sulphates also ↑ down group.

Anomalous Be / diagonal Be–Al: High IE and small size → covalent compounds; BeO amphoteric; resistant to acid (oxide film, passive in conc. HNO₃); forms complexes [BeF₄]²⁻; does not give blue solutions in NH₃. Shares many properties with aluminium.

CaO (quicklime): roast CaCO₃. Slaked lime Ca(OH)₂; with CO₂ reforms CaCO₃. Uses of lime: steel slag, cement, glass, lime–soda, water softening, CaC₂.

Biological Mg²⁺ and Ca²⁺: Mg²⁺ intracellular, chlorophyll; Ca²⁺ extracellular fluids, bones/teeth (apatite, fluoroapatite enamel), blood clotting, heart rhythm, nerve transmission.

Group 2 metals are less reactive than Group 1 but still reduce water (except Be). E° values are strongly negative; Be’s less negative E° reflects high atomization enthalpy and very high second IE, partly offset by large hydration energy of small Be²⁺. Complexes are more common for the smaller ions — chlorophyll is a magnesium complex; [BeF₄]²⁻ shows Be’s tendency to form covalent/complex species. CaCO₃ exists as calcite (stable, 6-coordinate Ca) and aragonite (metastable, 9-coordinate); calcite is preferred thermodynamically but conversion has a high activation barrier.

Quicklime production is among the world’s largest inorganic processes. Roasting limestone drives off CO₂; the reverse reaction (slaked lime + CO₂ → CaCO₃) is used in lime mortars and water treatment. Mixing lime with silica and clay produces cement — a cornerstone of civil engineering, rooted entirely in Group 2 chemistry.

Exam Connections and Chapter Summary

Exam favourites: unique position of H; isotopes and isotope effect; ice density; D₂O uses; hydride classification; Group 1 oxide types (Li₂O / Na₂O₂ / KO₂); diagonal Li–Mg and Be–Al; anomalous Li and Be; Castner–Kellner and Solvay outlines; carbonate thermal stability and sulphate solubility trends; flame colours; bio roles of Na, K, Mg, Ca.

Connect industrial chemistry to L13 electrochemistry (Hg overvoltage in Castner–Kellner) and L12 common ion (NaHCO₃ precipitation in Solvay). Trends follow periodic size and IE logic — first members always break the pattern.

Hydrogen and s-block elements form the foundation of Module 6. From nuclear fusion in stars to soap and cement on Earth, these elements and their compounds are both cosmically abundant and industrially indispensable.

Intext checkpoints: isotopes and radioactive T; why balloons use H₂; ice density; D₂O uses; hydride types (NaH is ionic); Group 1 oxide products; diagonal Li–Mg; Castner–Kellner and Solvay; BeO amphoteric; carbonate stability order; Ca²⁺ in bones and clotting. If you can answer those without notes, the chapter is exam-ready.

Next lessons (L18–L20) move into p-block chemistry where diagonal relationships continue (B–Si) and non-metal chemistry expands. Keep s-block trends as your baseline for comparing reactivity, oxide character, and industrial processes across the periodic table.