Chapter 10 — The s-Block Elements

Fig 1 — Educational Diagram: Introduction

Overview
s‑Block elements are the chemical elements in which the outermost (valence) electron(s) occupy an s‑orbital. They comprise Group 1 (alkali metals: Li, Na, K, Rb, Cs, Fr) and Group 2 (alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra) of the periodic table.

General electronic configuration
Group 1: ns¹
Group 2: ns²

Occurrence and extraction (brief)
Most s‑block metals are highly reactive and not found free in nature; they occur as salts and minerals. Extraction methods include electrolysis (for alkali metals from molten salts) and reduction with carbon or other reagents (for some alkaline earths), depending on chemical stability.

Physical properties (general)

• Shiny, silver‑coloured metals (except some oxides/tarnished surfaces).
• Good conductors of heat and electricity, malleable and ductile.
• Low ionization energies (especially Group 1), soft (alkali metals softer than alkaline earths).
• Melting and boiling points are generally low and decrease down Group 1; Group 2 have higher melting points than Group 1.

Chemical properties (general)

• They form predominantly ionic compounds (M⁺ for Group 1, M²⁺ for Group 2).
• Oxidation states: +1 in alkali metals; +2 in alkaline earth metals (Be shows covalent character; Mg is less reactive with cold water).
• React readily with water and oxygen to form hydroxides and oxides; alkali metal hydroxides are strongly basic.
• Tendency to lose valence electron(s) — strong reducing agents.

Trends down a group (important)

• Atomic radius increases down the group (added shells).
• Ionization energy decreases down the group (outer electron more easily removed).
• Reactivity of alkali metals increases down the group (easier to lose the single valence electron).
• Hydration energy and polarising ability decrease with increasing ionic size.

Notable special features
Diagonal relationship: Li (Group 1) and Mg (Group 2) show similar properties (e.g., comparable ionic sizes, formation of nitrides, and covalent behaviour in some compounds). Many s‑block compounds are basic (oxides and hydroxides).

Simple characteristic reactions
Examples of typical reactions that illustrate s‑block behaviour are given in the formulas list below. These reactions show the strong tendency to form M⁺ or M²⁺ ions, to react with water to give hydroxides and hydrogen, and to form ionic salts with halogens.

Why s‑block is important
s‑Block elements have wide industrial and biological importance: sodium and potassium are essential electrolytes; calcium is a major component of bones and building materials; magnesium and aluminium alloys (Al is p‑block) are structural; lithium is used in batteries.

Fig 2 — Educational Diagram: Electronic Configuration

What is electronic configuration? Electronic configuration is the arrangement of electrons in the atomic orbitals of an atom. It determines chemical behaviour (valency, ion formation, bonding, reactivity) and is written using shells (n = 1,2,3…), subshells (s, p, d, f) and the number of electrons in each subshell, e.g. 1s2 2s2 2p6 3s1.

Basic rules and principles:

• Aufbau principle: Electrons fill orbitals of lower energy first. Typical order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p …
• Pauli exclusion principle: An orbital holds a maximum of two electrons with opposite spins.
• Hund’s rule: For degenerate orbitals (same energy), electrons occupy them singly with parallel spins before pairing.

Important capacities and rules: Maximum electrons in a shell = 2n^2. Subshell capacities: s = 2, p = 6, d = 10, f = 14. The n + l rule helps predict filling order (lower n+l fills first; if equal, lower n fills first).

s-Block specifics (Class 11 focus): The s-block includes Group 1 (alkali metals) and Group 2 (alkaline earth metals). Their valence configurations are:

• Alkali metals (Group 1): ns1 (e.g. Na: [Ne] 3s1)
• Alkaline earth metals (Group 2): ns2 (e.g. Ca: [Ar] 4s2)

Because the outermost electron(s) occupy an s-orbital, s-block elements readily lose 1 (Group 1) or 2 (Group 2) electrons to form stable cations (M+ or M2+). For example: M (g) → M+ (g) + e−. The ease of losing these electrons (first ionization energy) largely controls metallic character and reactivity.

Trends in s-block (down a group): atomic radius increases, first ionization energy decreases, reactivity of alkali metals increases (they become stronger reducing agents), and metallic character increases. These trends follow from increased principal quantum number (n) and greater shielding of nucleus.

Why electronic configuration matters in real life: The ns1 or ns2 configuration explains why alkali metals form ionic salts (NaCl), why Ca2+ is important in biological systems (bone hydroxyapatite), why Mg forms strong alloys and is used in construction/aircraft, and why Li with its single valence electron is used in rechargeable batteries.

Fig 3 — Educational Diagram: Electronic Configuration and Position in the Periodic Table

Overview

The s‑block of the periodic table comprises Group 1 (alkali metals) and Group 2 (alkaline earth metals), together with hydrogen and helium. The position of an element in the periodic table is determined by its electronic configuration — in particular by the number of electrons in its outermost (valence) s‑orbital.

General electronic configurations

• Alkali metals (Group 1): outer configuration = ns1 (e.g., Li: 1s2 2s1, Na: 1s2 2s2 2p6 3s1)
• Alkaline earth metals (Group 2): outer configuration = ns2 (e.g., Be: 1s2 2s2, Mg: 1s2 2s2 2p6 3s2)
• Hydrogen: 1s1 (placed in Group 1 by valence but chemically unique). Helium: 1s2 (placed with noble gases for chemical inertness, though it is an s‑block element electronically).

How configuration gives position

• Group number (for s‑block) = number of valence s electrons (1 or 2).
• Period number = principal quantum number n of the outermost s orbital (e.g., Na is in period 3 because outer electron is in 3s).

Why s‑block elements show characteristic properties

• They have 1 or 2 valence electrons which are relatively loosely bound; this makes them highly electropositive and predisposes them to form M+ (Group 1) or M2+ (Group 2) ions, achieving noble gas configurations.
• Shielding by inner electrons is effective, so ionization energy decreases and atomic radius increases down a group; effective nuclear charge (Zeff) across a period increases, pulling electrons closer and reducing atomic radius.

Key trends related to electronic configuration

• Atomic radius: increases down a group (n increases), decreases across a period (Zeff increases).
• Ionization energy (IE): decreases down a group, increases across a period.
• Electropositive/metallic character: increases down a group, decreases across a period.
• Ionic formation: Alkali metals → M+; alkaline earth metals → M2+; ionic radii typically less than corresponding atomic radii for cations.

Simple rationale using quantum numbers

The principal quantum number n of the outermost s electron identifies the period. The s subshell can hold a maximum of 2 electrons (ms = +1/2, −1/2), so groups 1 and 2 are filled by 1 or 2 s electrons respectively before p‑block filling begins.

Special notes

• Helium has an s2 configuration like Group 2 elements electronically but is placed with noble gases because of its filled shell and chemical inertness.
• Some chemical behaviour (e.g., diagonal relationships like Li–Mg) can arise from a balance of charge density, size and polarising power, despite different groups.

Fig 4 — Educational Diagram: Occurrence and Distribution

Overview
S‑block elements (Groups 1 and 2) are highly reactive metals that do not occur in nature in the free state. Instead they occur as ions in sea water, dissolved in salts and brines, and as constituents of various minerals and ores in the Earth’s crust. Their distribution depends on ionic size, charge and geochemical behaviour: lighter elements (Na, Mg, Ca, K) are common, while Li, Be, Sr, Ba, Rb and Cs are present in much smaller amounts and in specific minerals.

Main natural reservoirs

• Sea water: Major source of Na+, Mg2+, Ca2+ and K+. Sodium and magnesium are abundant in oceans as dissolved salts.
• Earth’s crust (minerals/ores): Most s‑block elements are found as carbonates, sulfates, halides, oxides or silicates (e.g., halite NaCl, calcite CaCO3, magnesite MgCO3, dolomite CaMg(CO3)2, barite BaSO4, celestite SrSO4, spodumene LiAlSi2O6, beryl Be3Al2Si6O18).
• Saline brines and evaporite deposits: Sources of K, Li, and Mg (e.g., sylvite KCl, carnallite KMgCl3·6H2O).
• Biological systems: Na+, K+ are vital electrolytes; Mg2+ is central in chlorophyll; Ca2+ in bones and shells.

Why they are not found free
Alkali and alkaline earth metals have low ionization energies and react readily with oxygen, water and other anions to form stable compounds (oxides, hydroxides, carbonates, halides). Hence they occur as compounds in minerals or dissolved ions, not as elemental metals.

Trends in distribution

• Alkali metals (Group 1): Sodium (Na) and potassium (K) are common; Li is comparatively scarce and concentrated in pegmatites and brines; Rb and Cs occur in trace amounts in certain minerals (lepidolite, pollucite).
• Alkaline earth metals (Group 2): Magnesium (Mg) and calcium (Ca) are very abundant in crust and oceans; Sr and Ba occur in specific sulfate minerals (celestite, barite); Be is rare and confined to special silicates (beryl).
• Geochemical mobility: Elements that form soluble ions (Na+, K+, Mg2+) are mobile and abundant in oceans and soils; those that form insoluble sulfates or carbonates concentrate in particular mineral deposits.

Importance for extraction and industry
Because they occur as compounds, industrial extraction uses methods tailored to the source: electrolysis of molten salts (for alkali metals), chemical precipitation from brines (for Mg, Li), thermal decomposition of carbonates/oxides, or mining of mineral ores (barite, celestite, spodumene, beryl).

Summary points

• S‑block elements are widespread but mostly as compounds, not free metals.
• Major reservoirs: ocean (Na+, Mg2+), crust (carbonates, silicates, sulfates), brines/evaporites (K, Li).
• Abundance order (qualitative): For alkali metals Na > K > Li > Rb > Cs; for alkaline earths Ca > Mg > Sr > Ba > Be.

Fig 5 — Educational Diagram: Occurrence and Isolation

Overview
The s‑block elements (Group 1: alkali metals; Group 2: alkaline earth metals) are generally highly electropositive and reactive. Because of this they are rarely found in the free state in nature and occur mostly as compounds (salts, carbonates, sulfates, silicates, etc.). Isolation therefore requires first concentrating the ore and then converting it to a suitable salt that can be reduced or electrolysed to give the metal.

Occurrence — where they are found

• Alkali metals (Na, K, Li, etc.) — abundant in Earth's crust and seawater. Common minerals: halite (NaCl), sylvite (KCl), carnallite (KMgCl3·6H2O), spodumene (LiAl(SiO3)2), lepidolite (Li‑mica). Seawater is a major reservoir of Na+ and Mg2+.
• Alkaline earth metals (Be, Mg, Ca, Sr, Ba) — occur as carbonates, sulfates and silicates. Important ores: magnesite (MgCO3), dolomite (CaMg(CO3)2), limestone (CaCO3), gypsum (CaSO4·2H2O), barite (BaSO4), celestine (SrSO4), beryl (Be3Al2(Si6O18)).
• Why not found free: these metals have very negative standard reduction potentials and strongly reduce O2 and H2O, forming stable ionic compounds. Thus they are found as salts rather than native metals.

General steps in isolation

1. Concentration of the ore: gravity separation, magnetic separation, froth flotation, leaching, etc., to remove gangue.
2. Chemical conversion: convert ore to a soluble or fusible salt (e.g., chloride, carbonate, fluoride) amenable to purification.
3. Purification: crystallisation, precipitation, ion exchange, solvent extraction.
4. Reduction/electrolytic isolation: because of high reactivity, most s‑block metals are produced by electrolysing molten salts or by strong chemical/thermal reduction (e.g., Pidgeon process for Mg).

Typical industrial methods (examples)

• Sodium (Na) — produced by electrolysis of molten NaCl in the Downs cell. CaCl2 is often mixed with NaCl to lower the melting point and electrical conductivity is ensured. Major reactions:
  At cathode: Na+ + e− → Na (liquid)
At anode: 2Cl− → Cl2 (g) + 2e−
Overall: 2NaCl (l) → 2Na (l) + Cl2 (g)
• Potassium (K) — prepared by electrolysis of fused KCl (or by displacement by sodium in some older processes). K is more volatile, so special handling is needed.
• Lithium (Li) — recovered from brines (evaporation to Li2CO3) or from minerals (spodumene). Metallic Li is obtained by electrolysis of molten LiCl (often mixed with KCl to lower the melting point) or by reduction of LiCl with sodium in some lab methods.
• Magnesium (Mg) — two main industrial routes:
  • Dow (electrolytic) process: Mg(OH)2 from seawater (precipitated by adding Ca(OH)2) → convert to MgCl2 → dehydrate and electrolyse molten MgCl2.
  • Pidgeon (thermal reduction) process: MgO (often from calcined dolomite) reduced by ferrosilicon (or Si) under vacuum at high temperature to give Mg vapour which is condensed. Representative reaction:
  2MgO + Si → 2Mg (g) + SiO2
• Calcium (Ca), Strontium (Sr), Barium (Ba) — usually obtained by electrolysis of their molten chlorides. Calcium can also be obtained by reduction of CaO/CaCl2 by aluminium in some specialized processes.
• Beryllium (Be) — extracted from beryl or bertrandite. The ore is converted to BeO and subsequently to salts (e.g., BeF2 or BeCl2); metallic Be is obtained by reduction with magnesium or by specialised electrolytic reduction of molten Be salts. (Beryllium extraction is technically demanding because of toxicity and strong chemical bonds.)

Why electrolytic methods?
Alkali and alkaline earth metals are too electropositive to be reduced by carbon or common reducing agents. Electrolysis of fused salts supplies the required electrons directly to reduce the cations to the metal. For some less reactive alkaline earth metals, high‑temperature chemical reductions (with Si or Al/FeSi) are also used.

Practical notes
Industrial extraction aims to minimize energy use (mixing salts to lower melting points), avoid contamination (use of inert cathode materials, controlled atmosphere), and handle corrosive halogen gases (Cl2, F2) safely.

Fig 6 — Educational Diagram: Anomalous Behavior of Lithium and Beryllium

What is meant by anomalous behavior? In Groups 1 and 2 most elements show a gradual trend of properties down the group. Lithium (Li) and beryllium (Be), the first members of their groups, differ markedly from the rest of their respective groups — this is called their anomalous behaviour.

Main reasons (common to both Li and Be)

• Very small atomic and ionic size (Li+ and Be2+ are unusually small).
• High ionisation enthalpy and relatively high electronegativity (compared to other group members).
• High polarising power (high charge/radius2) — strong tendency to distort anions.
• High hydration energy (for Li+) and high covalent character in compounds.
• Diagonal relationship: Li shows similarities with Mg and Be with Al because of comparable size/charge effects.

Consequences & specific features — Lithium (Li)

• Li forms covalent or covalent-like compounds more than other alkali metals. Example: LiCl has more covalent character than NaCl (polarising Li+).
• Oxides: Li forms oxide (Li2O) whereas other alkali metals form peroxides or superoxides. Example reactions:
  4 Li + O2 → 2 Li2O
2 Na + O2 → Na2O2 (peroxide)
K + O2 → KO2 (superoxide)
• Li forms nitride readily: 6 Li + N2 → 2 Li3N (other alkali metals do not form stable nitrides under mild conditions).
• Li compounds often have high lattice and hydration energies; Li2CO3 is more stable and less soluble than other group I carbonates.
• Li forms important organometallic reagents (organolithium, R–Li), widely used in organic synthesis and as strong bases:
  R–X + 2 Li → R–Li + LiX
• Applications tied to anomalous properties: Li-ion batteries (high electrochemical potential, small size), LiOH used for CO2 scrubbing in spacecraft, Li2CO3 in medicine (bipolar disorder).

Consequences & specific features — Beryllium (Be)

• Be forms predominantly covalent compounds (Be–X bonds have strong covalent character). Example: BeCl2 is covalent and polymeric in solid state; in vapour it exists as discrete linear molecules.
• BeO and Be(OH)2 are amphoteric (not strongly basic like heavier group 2 oxides). Example reactions:
  BeO + 2 HCl → BeCl2 + H2O
BeO + 2 NaOH + H2O → Na2[Be(OH)4]
• High ionisation energies and small size make Be less reactive: it does not react with water and shows reluctance to form M2+ ionic compounds like other alkaline earths.
• Be forms covalent hydride BeH2 (polymeric) rather than ionic hydride typical of heavier group 2 metals.
• Diagonal relationship: Be shows some similarities with Al (e.g., amphoteric oxides, covalent behavior).

Underlying rule: Fajan's rules (qualitative)

• Small cation + high charge + large polarizable anion → more covalent character. Li+ and Be2+ strongly polarise anions, producing covalent character in their compounds.

Summary

The anomalous behaviour of Li and Be arises from their very small size, high charge density, high ionisation energy and high polarising power. These factors make their chemistry more covalent, less metallic and cause properties (oxides, hydrates, nitrides, solubility, reactivity) that differ from the rest of Group 1 and Group 2 respectively.

Fig 7 — Educational Diagram: Isolation and Preparation of Elements

Overview: s‑Block elements (alkali and alkaline‑earth metals) are very electropositive and highly reactive. They do not occur in nature as free metals but as salts and minerals. Their isolation requires methods that overcome very negative reduction potentials. Major industrial methods are electrolytic reduction of fused salts, chemical (metallothermic) and thermal reduction of oxides.

Occurrence & concentration: Alkali metals occur as halides (NaCl, KCl) in sea water and evaporitic deposits; lithium in pegmatites or brines (LiCl, Li2CO3); alkaline‑earth metals occur as carbonates, sulphates, silicates (limestone CaCO3, dolomite CaMg(CO3)2, magnesite MgCO3, beryl for Be).

Why electrolysis? Because standard reduction potentials of M+/M are very negative (for example Li+ + e− → Li E° ≈ −3.04 V, Na+ ≈ −2.71 V, K+ ≈ −2.93 V, Ca2+ ≈ −2.87 V, Mg2+ ≈ −2.37 V). Chemical reducing agents (C, CO) cannot reduce many s‑block oxides; thus molten‑salt electrolysis is the common route.

Main methods:

• Electrolysis of molten salts (primary industrial method): Molten chloride or fluoride of the metal is electrolysed. Example: Downs cell for sodium (molten NaCl, CaCl2 added to lower melting point). Cathode reaction: Na+ + e− → Na(l). Anode reaction: 2Cl− → Cl2(g) + 2e−. Overall: 2NaCl(l) → 2Na(l) + Cl2(g).
• Electrolysis for other metals: Magnesium and calcium are also obtained by electrolysis of fused MgCl2 and CaCl2 respectively. Lithium is obtained by electrolysis of molten LiCl (often mixed with KCl to lower melting point).
• Chemical/thermochemical reduction (metallothermic): Some metals (e.g., Be, Mg in some processes) are prepared by reducing their oxides or halides with a more reactive metal (Mg or Al) or by silicothermic reduction. Example: Pidgeon process for Mg (silicothermic): MgO + Si → Mg(g) + SiO2 (Mg vapor is condensed). Another example: BeF2 reduced by Mg: BeF2 + Mg → Be + MgF2 (in practice more complex).
• Hydrometallurgical routes (for concentration): Minerals are treated (acid/alkali treatment, precipitation, ion exchange) to obtain pure salts (e.g., Li2CO3 from brines by precipitation and purification) before final electrolysis.
• Refining: Metals such as sodium and magnesium are purified by fractional distillation (low boiling metals) or zone refining when needed.

Practical/industrial notes:

• Molten‑salt electrolysis requires inert or consumable anodes (graphite anode evolves Cl2 or sometimes O2 if oxide present) and suitable cathodes to collect liquid metal.
• To lower cell voltage and melting point, binary or ternary salt mixtures (e.g., NaCl–CaCl2, LiCl–KCl) are used.
• Safety: products are highly reactive with water/air (Na, K react violently; Mg burns in air), so handling and storage are under inert atmosphere or hydrocarbons.

Summary: Isolation of s‑block elements relies mainly on electrolytic reduction of molten salts, supplemented by metallothermic and silicothermic processes for some metals. Choice depends on thermodynamics (reduction potential), cost, and physical properties (melting point, vapor pressure).

Fig 8 — Educational Diagram: Diagonal Relationship

Definition: The diagonal relationship is an observed similarity in chemical behavior between certain pairs of elements that are diagonally adjacent in the periodic table (moving one step down and one step to the right). The most commonly cited diagonal pairs are Li–Mg, Be–Al and B–Si. Although these elements belong to different groups, similarities arise because some of their important properties change in opposite directions across a period and down a group, so the net result is similar for the diagonal pair.

Why it occurs (key factors):

• Comparable ionic size and charge/radius ratio — the effective size and the charge/size (ionic potential) of a cation may be similar for diagonal neighbours.
• High polarizing power — small radius and relatively high charge (or high ionization energy) cause strong polarization of anions, increasing covalent character (Fajans’ concepts).
• Similar electronegativity and ionization enthalpy — trends across a period (increasing electronegativity/IE) and down a group (decreasing) can cancel, producing similar values diagonally.
• Similar hydration and lattice energies — comparable balance of solvation vs lattice energies leads to analogous solubility and thermal stability.

Typical similarities (illustrative):

• Li and Mg: comparable ionic potential leads to similar bonding tendencies. Both form nitrides (Li3N and Mg3N2), both give important organometallic reagents (organolithiums R–Li and Grignard reagents R–MgX) and many of their salts show enhanced covalent character compared with heavier group members.
• Be and Al: both form covalent chlorides (BeCl2 is covalent/polymeric; AlCl3 has covalent character and forms Al2Cl6 dimers in vapor), and both have amphoteric hydroxides/oxides (Be(OH)2 and Al(OH)3 / BeO and Al2O3).
• B and Si: both form predominantly covalent network oxides and hydrides (B2O3, SiO2; B2H6, SiH4) and show covalent bonding and polymeric structures rather than simple ionic solids.

Limitations: The diagonal relationship is qualitative and most pronounced for elements in the top three periods. It weakens for heavier elements because size and other trends become more dominant. Not all properties are similar—differences in oxidation states, coordination chemistry and specific reactivities remain.

Importance for students: The diagonal relationship helps explain why some light members of Group 1/2 show behaviour unexpected for their group and why some compounds of these elements resemble those of elements in the neighbouring group diagonally. It links periodic trends (size, ionization energy, electronegativity) to chemical behavior.

Fig 9 — Educational Diagram: Physical Properties and Trends

Overview: The s-block of the periodic table contains Group 1 (alkali metals, electronic configuration ns1) and Group 2 (alkaline earth metals, ns2). These elements are metals with characteristic physical properties that change in regular ways (trends) down a group and across a period. Understanding these trends uses concepts such as atomic/ionic size, nuclear charge, metallic bonding and electrostatic (Coulomb) interactions.

Common physical properties

• Metallic luster, malleability (soft for alkali metals), and good electrical and thermal conductivity.
• Low ionization energies and low electronegativities compared with p- and d-block elements.
• Form cations (M+ for Group 1, M2+ for Group 2) easily, so they tend to form ionic compounds.
• Melting and boiling points are generally lower for alkali metals than for most metals; alkaline earth metals are harder and have higher melting points than alkali metals (because they contribute two valence electrons to the metallic bond).

Key trends down a group

• Atomic and ionic radii: Increase down the group because new electron shells are added, increasing the size despite increased nuclear charge.
• Ionization energy (IE): Decreases down the group because outer electrons are farther from the nucleus and more shielded, so they are easier to remove.
• Metallic character and reactivity: Increase down the group for alkali metals (easier to lose the single valence electron). For alkaline earth metals reactivity also increases but less dramatically.
• Density: Generally increases down the group (more mass per volume), but there are exceptions (e.g., K less dense than Na).
• Melting/boiling points: For alkali metals they generally decrease down the group (weaker metallic bonding due to larger atomic size); for alkaline earth metals trends are less uniform — Be and Mg have anomalously high melting points due to strong bonding from small size and higher charge density.

Group-to-group comparisons (Group 1 vs Group 2)

• At the same period, Group 2 atoms have smaller atomic radii than Group 1 (greater nuclear charge for same shell) and higher ionization energies.
• Group 2 metals are harder, have higher melting points and are less reactive than Group 1 because they hold two valence electrons more tightly and form M2+ ions.

Reasons / physical explanations

• Metallic bonding: Strength depends on number of delocalized electrons and charge density. Group 2 atoms contribute two electrons to the electron sea, usually giving stronger bonds than Group 1.
• Coulombic interactions: Electrostatic attraction between oppositely charged ions (or between nucleus and electrons) depends on charge and distance; this underlies lattice energies, ionization energies and many trends.
• Charge density effects: Small ions with higher charge (e.g., Be2+) have high charge density — strong attraction to electrons or water molecules — causing high ionization energies, strong lattice energies and stronger covalent character in bonds.

Anomalies and important notes

• Li and Be often deviate from simple trends: their very small sizes and high polarising power lead to stronger metallic/covalent character, higher melting points (Li), and anomalous chemistry (Be forms covalent compounds).
• Trends often show smooth change but several physical properties (melting point, density) show irregularities due to changes in metallic structure, packing, or electronic effects.

Relevance / real-world impact: These physical properties explain why alkali metals are used where lightweight, soft, reactive metals are needed (e.g., Li in batteries), and why alkaline earth metals (Mg, Ca) are used for structural materials, biological roles (Ca in bones), and alloys.

Fig 10 — Educational Diagram: Ionization Enthalpy and Electron Affinity

Definitions

Ionization enthalpy (IE) — the enthalpy change required to remove an electron from one mole of gaseous atoms (or ions). First ionization enthalpy (IE1):

X(g) → X⁺(g) + e⁻ ; ΔH = IE1 (always endothermic, IE1 > 0, unit: kJ mol⁻¹)

Electron affinity (EA) — the enthalpy change when one mole of gaseous atoms accepts an electron to form one mole of gaseous anions. (Sign convention: if energy is released the enthalpy change is negative.)

X(g) + e⁻ → X⁻(g) ; ΔH = EA (often exothermic, EA < 0 for atoms that release energy)

Factors affecting IE and EA

• Effective nuclear charge (Z_eff): higher Z_eff → higher IE and more negative EA.
• Atomic size (distance of outer electron from nucleus): larger radius → lower IE and less negative EA.
• Shielding (inner electrons): more shielding → lower IE, less tendency to gain electron (less negative EA).
• Electronic configuration and subshells: removing/adding to half-filled or filled subshells can be less favorable (causing exceptions).

Trends in the periodic table (relevant for s‑block)

• Across a period (left → right): IE generally increases (nuclear charge increases, radius decreases); EA generally becomes more negative (atoms more likely to gain electrons), with some exceptions due to subshell stability.
• Down a group: IE decreases (outer electron farther from nucleus, more shielding); EA becomes less negative (atoms less likely to accept an electron).
• S‑block characteristics: alkali metals (group 1) have low IE1 → readily lose one electron → common +1 oxidation state. Alkaline earth metals (group 2) have higher IE1 and moderate IE2 → commonly form +2 ions. Successive IEs rise sharply, especially when a core electron must be removed; this explains stable oxidation states.

Successive ionization enthalpies

IEn (n = 1, 2, 3...) is the energy to remove the nth electron from the gaseous ion. Successive IEs increase (IE1 < IE2 < IE3 ...). Large jumps occur when an electron is removed from a more tightly bound inner shell (e.g., very large jump from IE1 to IE2 for alkali metals).

Common exceptions

• Across a period small local decreases: e.g., IE of B < Be (electron removed from 2p vs 2s), IE of O < N (paired electron in O’s p orbital causes increased repulsion).
• Electron affinity is small or positive for atoms with filled subshells (group 2) or half‑filled subshells (group 15) because adding an electron disrupts a stable configuration.

Importance and applications

• Explains why metals (especially s‑block) lose electrons in chemical reactions (highly reactive alkali metals).
• Helps predict ionic formation and stability (e.g., why Na → Na⁺ and Cl + e⁻ → Cl⁻ combine to form NaCl).
• Used in processes like metallurgy, electrochemistry and understanding flame reactions and ionic bonding energies.

Fig 11 — Educational Diagram: Trends in Atomic and Ionic Radii

Definition: Atomic radius is a measure of the size of an atom (commonly reported as covalent, metallic or van der Waals radius depending on context). Ionic radius is the effective radius of an ion in a crystal lattice. For s‑block elements (Groups 1 and 2) the metallic radius is usually used for neutral atoms and the ionic radius for their M+ or M2+ ions.

Main causes of trends

• Effective nuclear charge (Zeff): As the nuclear charge felt by the valence electrons increases (Zeff increases), electrons are pulled closer and the radius decreases.
• Shielding and principal quantum number (n): Adding electron shells (increasing n) increases size because outer electrons are farther from the nucleus and experience greater shielding.
• Orbital penetration: s orbitals penetrate close to the nucleus, so for s‑block elements the increase in Zeff across a period more effectively reduces radius.
• Electron–electron repulsion: Adding electrons to the same shell increases repulsion and tends to expand size; removing electrons (forming cations) reduces repulsion and contracts size.

Specific trends

• Down a group (Group 1 or Group 2): Atomic and ionic radii increase markedly. Reason: each successive element has an additional electron shell (higher n) and increased shielding outweighs increased nuclear charge.
• Across a period (left to right): Atomic radii decrease. For s‑block across a period (e.g., from Group 1 → Group 2 → nonmetals) the increased nuclear charge pulls valence electrons inwards; s electrons' penetration makes this effect pronounced.
• Cations vs parent atoms: Cations are significantly smaller than their neutral atoms because electrons are removed (sometimes removing an entire shell) and remaining electrons feel a stronger Zeff.
• Anions vs parent atoms: Anions are larger than their neutral atoms because added electron(s) increase electron–electron repulsion and expand the electron cloud.
• Charge effect on ionic size: For ions of same element, higher positive charge → smaller ionic radius (e.g., Mg2+ << Mg). For isoelectronic series (same electron count) ionic radius decreases with increasing nuclear charge (more protons pull the same electrons closer).

Why s‑block shows clear trends: s orbitals are spherical with good penetration and relatively weak shielding by inner electrons; thus changes in nuclear charge and shell number produce large, regular changes in size for alkali and alkaline earth metals.

Practical consequences: Ionic radius affects lattice energy, solubility, coordination number and reactivity. For example, smaller cations with higher charge give larger lattice energies (more stable ionic solids) and stronger hydration in water.

Fig 12 — Educational Diagram: Electronegativity and Electropositive Character

Electronegativity (χ) is the ability of an atom in a molecule to attract the shared pair of electrons towards itself. It is a relative scale; the most commonly used values are on the Pauling scale. Electronegativity depends mainly on nuclear charge, atomic radius and shielding. Higher nuclear charge and smaller radius give higher χ.

Electropositive character is the tendency of an element (typically a metal) to lose electrons and form positive ions (cations). Electropositive elements have low ionization energies and low electronegativities. Electropositive character is essentially the opposite of electronegativity.

Trends in the periodic table (relevant to s‑block elements):

• Across a period (left to right): electronegativity generally increases, electropositive character decreases (because effective nuclear charge increases and atomic radius decreases).
• Down a group: electronegativity decreases, electropositive character increases (because atomic radius and electron shielding increase, ionization energy decreases).
• For s‑block specifically (Group 1 and Group 2): elements are highly electropositive. Reactivity and electropositive character increase down the group (e.g., Li < Na < K < Rb < Cs).

Why these trends occur:

• Electronegativity falls down a group because the valence electron is farther from the nucleus and shielded by inner electrons, so the nucleus attracts bonding electrons less strongly.
• Electropositive character increases down a group because ionization energy (energy required to remove an electron) decreases, so atoms lose electrons more easily.

Consequences for chemical behavior of s‑block elements:

• Form predominantly ionic compounds (e.g., NaCl, CaF2) because of large difference in electronegativity with nonmetals.
• Oxides and hydroxides are strongly basic (e.g., Na2O, CaO, NaOH, Ca(OH)2) — basicity increases down Group 2.
• Strong reducing agents: alkali metals and alkaline earth metals readily donate electrons in redox reactions.

Common electronegativity scales and simple relations:

• Mulliken electronegativity (approximate): χ_M = (I.E. + E.A.) / 2, where I.E. is first ionization energy and E.A. is electron affinity (both in same energy units).
• Pauling interpretation: difference in electronegativity relates to extra bond energy of A–B compared with average of A–A and B–B bonds. (Used to produce Pauling χ values.)
• Empirical rule for bond type: if Δχ (difference in Pauling electronegativity) > ~1.7 the bond is largely ionic; if < ~1.0 it is largely covalent (values approximate).

Summary: s‑block elements have low electronegativity and high electropositive character, which increases down the group. This explains their strong tendency to form cations, make ionic solids, act as reducing agents, and form basic oxides/hydroxides.

Fig 13 — Educational Diagram: Ionization Enthalpy and Electronegativity

Ionization Enthalpy (IE)

Ionization enthalpy (also called ionization energy) is the minimum energy required to remove an electron from one mole of gaseous atoms or ions. It is an endothermic process and is usually expressed in kJ mol^-1.

General representation:

X(g) → X⁺(g) + e^-   ΔH = IE₁

Successive ionization energies: removing more than one electron is progressively harder. For example:

X⁺(g) → X²⁺(g) + e^-   ΔH = IE₂

Key factors affecting IE:

• Atomic radius / distance of electron from nucleus: larger distance → lower IE.
• Nuclear charge (Z): greater effective nuclear charge (Z_eff) → higher IE.
• Electron shielding: inner electrons shield valence electrons → lower IE.
• Electron configuration / stability: fully filled or half-filled subshells are relatively stable and show higher IE.

Periodic trends:

• Across a period (left → right): IE generally increases due to increasing Z_eff and decreasing atomic radius.
• Down a group: IE decreases because valence electrons are farther from the nucleus and more shielded.

Particulars for s‑block elements (Group 1 & 2):

• Alkali metals (Group 1) have the lowest first IE in each period — they lose their single valence s electron easily, explaining their high reactivity (e.g., Na, K).
• Alkaline earth metals (Group 2) have higher IE₁ than Group 1 (two valence electrons, greater nuclear charge), but their IE values are still relatively low compared to p‑block nonmetals.
• Successive IE shows a very large jump after removing all valence electrons (e.g., Na: IE₁ ≈ 496 kJ mol^-1, IE₂ ≈ 4562 kJ mol^-1).

Electronegativity (χ)

Electronegativity is a measure of the tendency of an atom in a chemical bond to attract shared electrons. It is a dimensionless relative scale (commonly the Pauling scale).

Key points:

• Higher χ → stronger attraction for bonding electrons (e.g., F ≈ 3.98, O ≈ 3.44 on the Pauling scale).
• Trend: across a period χ increases (atoms become less metallic and attract electrons more); down a group χ decreases (atoms are larger and less able to attract bonding electrons).
• Electronegativity difference (Δχ) between two bonded atoms indicates bond character: small Δχ → covalent, large Δχ → ionic. A rough guideline: Δχ > 1.7 (Pauling) indicates largely ionic character.

Common definitions / scales:

• Pauling scale — empirical, most commonly used.
• Mulliken electronegativity: χ = (IP + EA)/2 where IP is ionization potential and EA is electron affinity (both in the same energy units).
• Allred–Rochow: based on effective nuclear charge and covalent radius (used less at Class 11 level).

Connections between IE and electronegativity:

• Both depend on effective nuclear charge and atomic size. Higher IE generally correlates with higher electronegativity.
• In s‑block elements, low IE corresponds to low electronegativity, which explains their tendency to form cations and ionic compounds.

Applications and implications (especially for s‑block):

• Low IE of alkali metals explains violent reactions with water producing H₂ and hydroxides (e.g., 2 Na + 2 H₂O → 2 NaOH + H₂).
• Electronegativity differences determine polarity of bonds and properties of compounds — e.g., ionic character in NaCl vs covalent character in H₂O.
• Successive IE jumps explain why alkali metals commonly form +1 ions while alkaline earth metals form +2 ions.

Note for students: remember units (kJ mol^-1) for IE and that electronegativity is a relative (unitless) scale. Use periodic trends and electron configurations to predict behavior.

Fig 14 — Educational Diagram: Anomalous Behavior of the First Element

Definition / Overview
In the s‑block (alkali and alkaline earth metals), the first element of each group (lithium in Group 1; beryllium in Group 2) shows properties that are different from the rest of the group. These differences are called the anomalous behavior of the first element.

Main causes

• Very small atomic and ionic radii: Small size increases polarizing power and hydration energy.
• High ionization enthalpy and relatively higher electronegativity: Makes formation of ions less favoured than for heavier congeners.
• High polarizing power (charge/radius²): Small cations strongly polarize anions, increasing covalent character (Fajans' rules).
• Absence of available d‑orbitals and compact electronic structure: Influences bonding and structures (more covalent, polymeric, or directional bonding).

Consequences — Lithium (Li) in Group 1

• Tendency to form compounds with significant covalent character (e.g., organolithium compounds R–Li, lithium halides show more covalency than heavier alkali halides).
• Forms oxide (Li₂O) on burning in oxygen, whereas heavier alkali metals form peroxides or superoxides: 4Li + O₂ → 2Li₂O; 2Na + O₂ → Na₂O₂; K + O₂ → KO₂ (superoxide).
• Forms a stable nitride (Li₃N) on direct combination with nitrogen: 6Li + N₂ → 2Li₃N (other alkali metals do not form stable nitrides under normal conditions).
• Smaller ionic radius → higher hydration energy and stronger solvation effects; Li+ often behaves differently in solution and solid state (e.g., strong Lewis acidity, forms complexes and aggregates).
• Diagonal relationship with Mg: Li shows chemical similarities with magnesium (e.g., both form nitrides, carbonates decompose on heating, tendency to form covalent compounds).

Consequences — Beryllium (Be) in Group 2

• Be forms predominantly covalent compounds (BeCl₂ is covalent and forms chains/dimers) unlike ionic character of heavier alkaline earth chlorides.
• BeO and Be(OH)₂ are amphoteric (can react with both acids and bases), whereas heavier group‐2 oxides/hydroxides are more basic.
• Be does not give typical ionic chemistry; it forms complex covalent hydrides (BeH₂ is polymeric) and forms coordinative compounds with electron donors.
• Diagonal relationship with Al: Be and Al show similar behaviour (e.g., amphoteric hydroxides, covalent halides).

Fajan's rule summary (qualitative)
Small highly charged cation + large polarizable anion → increased covalent character. Li+ and Be2+ are highly polarizing cations because of their small radii.

Typical equations (examples)
4Li + O₂ → 2Li₂O
2Na + O₂ → Na₂O₂ (peroxide)
K + O₂ → KO₂ (superoxide)
6Li + N₂ → 2Li₃N

Practical significance / Real‑life links
Lithium’s special chemistry underlies lithium‑ion batteries (Li+ insertion/extraction, organolithium reagents in synthesis) and pharmaceutical uses (lithium carbonate in mood stabilizers). Beryllium’s covalent compounds and light weight make its alloys and BeO useful in aerospace and high‑temperature ceramics (with serious toxicity considerations for Be compounds).

How to remember
Small size → high polarizing power → more covalent character and unique compounds. Also remember the diagonal relationships: Li ↔ Mg and Be ↔ Al.

Fig 15 — Educational Diagram: Polarizing Power and Covalent/Ionic Character (Fajans' Rules)

Overview: Fajans' rules give simple qualitative guidelines to predict whether a chemical bond between a cation and an anion will be largely ionic or will show significant covalent character. The central idea is polarization: a small, highly charged cation distorts (polarizes) the electron cloud of a large, easily deformable anion. The greater the distortion, the greater the covalent character.

Fajans' rules (summary):

• Smaller cations with higher positive charge have greater polarizing power (they pull electron density toward themselves).
• Larger anions with higher negative charge are more easily distorted (higher polarizability), so they develop more covalent interaction.
• Cations with an electron configuration close to a noble gas (no outer d or f electrons) tend to be less polarizing than transition metal cations of the same size/charge; cations with vacant orbitals (e.g., Al3+) can promote covalency.
• Thus, covalent character increases when: small/highly charged cation + large/highly charged anion are paired.

Physical picture: A purely ionic bond is an electrostatic attraction between point charges. In reality, the anion's electron cloud is spread out; a strong electric field from the cation pulls the electron cloud toward the cation, producing partial electron sharing (polarization) and hence covalent character.

Why size and charge matter:

• Charge: Higher cation charge → stronger electric field → stronger pull on anion electrons.
• Size of cation: Smaller radius → larger charge density → stronger field at the anion's position.
• Size of anion: Larger radius → electrons are held more loosely and are easier to distort (higher polarizability).

Consequences and examples in chemical behavior:

• Compounds predicted as ionic by charge balance may show covalent features: e.g., AlCl3 is largely covalent (forms Al2Cl6 dimers) while AlF3 is more ionic.
• Smaller alkali cations polarize halide anions more: LiI is more covalent than CsI; LiF is more ionic than LiI.
• Covalent character often reduces aqueous solubility and increases volatility or solubility in organic solvents compared with strongly ionic salts.

Relation to other measures: Polarizing power is related to cation charge density (charge divided by size). The actual electric field at the anion depends on Coulomb's law, so field strength at distance r is proportional to Z/r^2, while qualitative polarizing power is often expressed as Z/r (charge/ionic radius) for simplicity. Anion polarizability scales roughly with its volume (~r^3).

Limitations: Fajans' rules are qualitative. Electronic structure, orbital availability, and covalent orbital overlap (especially for transition metals) also strongly affect bonding. For quantitative ionic/covalent percentages use electronegativity models or spectroscopic data.

Fig 16 — Educational Diagram: Physical Properties

Overview
The s‑block elements (alkali metals, group 1; and alkaline earth metals, group 2) show characteristic physical properties that arise from their single (group 1) or two (group 2) valence s‑electrons and metallic bonding with free (delocalized) electrons. These properties include metallic lustre, malleability/ductility, good electrical and thermal conductivity, generally low ionization energies, and trends in atomic/ionic radii, density, melting/boiling points and hardness down each group.

Key features and causes

• Metallic character: Low ionization energies allow easy formation of cations; delocalized valence electrons give metallic bonding and conductivity.
• Lustre, malleability, ductility: Sea of electrons permits layers of atoms to slide past each other without fracturing.
• Electrical & thermal conductivity: High due to mobile conduction electrons; conductivity tends to be high for metals with more free electrons per atom and closely packed structures.
• Atomic and ionic radii: Increase down a group because successive shells are added. Alkali metal ionic radius (M+) is large because only one electron is removed; alkaline earth (M2+) radii are smaller relative to their neutral atoms but still increase down the group.
• Density: Alkali metals are light (very low densities; many float on water). Alkaline earth metals are generally heavier and denser than alkali metals because of higher nuclear charge and closer packing in crystal lattices.
• Melting and boiling points: Determined by metallic bonding strength (number of delocalized electrons, charge density, packing). Alkali metals have relatively low melting points and show a steady decrease down the group (Li > Na > K > Rb > Cs). Alkaline earth metals generally have higher melting points than alkali metals; Be and Mg are particularly high due to strong bonding arising from small size and high charge density.
• Hardness: Alkali metals are soft (can be cut with a knife); alkaline earth metals are harder and less reactive physically.

Trends (summary)

• Atomic/ionic radius: increases down the group.
• Ionization enthalpy: decreases down the group (explains increasing reactivity of alkali metals).
• Metallic character: increases down the group.
• Melting point: alkali metals generally decrease down the group; alkaline earth metals show higher m.p. for Be and Mg and then a decline.
• Density: alkali metals low and increase slowly down the group; alkaline earth metals are denser, generally increase down the group.

Qualitative connection to energetics
Physical properties are linked to electronic structure: low ionization energy ⇒ easy electron delocalization ⇒ metallic bonding. For ionic salts of s‑block metals, lattice energy is important; lattice energy is inversely related to ionic radii (smaller ions → larger lattice energy → higher melting point of ionic solid).

Experimental/observational notes
- Alkali metals are soft and can be cut; they have characteristic flame colours (Li: crimson, Na: yellow, K: lilac) used in qualitative identification.
- Alkaline earth metals are less reactive and harder; Mg and Al (though Al is p‑block) form protective oxide layers that affect observed physical behaviour.

Fig 17 — Educational Diagram: Bonding: Ionic vs Covalent Character

Overview: Bonds lie on a continuum from purely ionic (electrostatic attraction between ions) to purely covalent (shared electron pairs). Most real compounds have mixed character. In s‑block chemistry (alkali and alkaline earth metals) ionic bonding predominates, but significant covalent character appears when the cation is small/highly charged and the anion is large/highly polarizable.

Key concepts:

• Ionic bond: Electrostatic attraction between cation and anion. Typical properties: high melting/boiling points, crystalline solids, conduct electricity when molten/dissolved.
• Covalent bond: Electron sharing between atoms; directional bonding; lower melting/boiling points for molecular covalent substances, poor electrical conductors.
• Continuum: Ionic and covalent are end points. Electronegativity difference and electron cloud distortion determine the degree of ionic vs covalent character.

Fajans' rules (to predict covalent character in ionic compounds):

• Small cation and large charge on cation → large polarizing power → increases covalent character.
• Large anion with many electrons (more easily distorted) → high polarizability → increases covalent character.
• If cation has an inert gas configuration (closed shell) or anion is isoelectronic with noble gas, ionic character is favored.

Polarizing power and polarizability: Polarizing power of a cation is roughly proportional to its charge/size (commonly approximated ∝ Z/r²); polarizability of an anion increases with its size and number of electrons. A small, highly charged cation (e.g., Be²⁺, Li⁺) strongly distorts a large anion (I^-, Br^-), producing substantial covalent character.

S‑block examples and trends:

• Among alkali metals (Li → Cs) covalent character of their halides decreases down the group; LiX are most covalent because Li⁺ is smallest and most polarizing.
• Alkaline earth metals: BeCl₂ is largely covalent (molecular, linear monomer in gas phase; polymeric with covalent chains in solid), MgCl₂ shows partial covalency, whereas CaCl₂, SrCl₂ are more ionic.
• LiI, LiBr show more covalent character than NaI/NaBr. MgO is strongly ionic (high lattice energy) and has a very high melting point.

Consequences of increased covalent character in an ionic compound: Lower melting point (relative), decreased electrical conductivity in molten state, greater solubility in nonpolar solvents, formation of molecular species (e.g., BeCl₂ polymerizes or forms covalent chains).

Quantifying ionic character: Electronegativity difference gives a measure. Pauling proposed an empirical expression to relate difference in electronegativity to partial ionic character (see formulas).

Summary: Ionic vs covalent character depends on charge and size of ions and electronegativity differences. Use Fajans' rules qualitatively for s‑block salts: small/highly charged cations and large/polarizable anions → increased covalent character. Many s‑block compounds therefore show a range of bonding behavior rather than being purely ionic.

Fig 18 — Educational Diagram: Chemical Reactivity and General Reactions

Overview

S-block elements = Group 1 (alkali metals: Li, Na, K, Rb, Cs) and Group 2 (alkaline earth metals: Be, Mg, Ca, Sr, Ba). They have ns1 (Group 1) or ns2 (Group 2) valence configurations, form predominantly +1 and +2 cations respectively, and show characteristic reactivity patterns controlled by ionization energy, atomic/ionic size, lattice energy and hydration energy.

Trends in reactivity

• Ionization energy decreases down each group → metals lose electrons more easily → chemical reactivity increases down the group (Group 1 & Group 2).
• Hydration energy and lattice energy determine solubility and thermal stability of salts. Small cations (Li+, Be2+) have high polarising power → greater covalent character and special behaviour (e.g., Li and Be differ from their groups).

General types of reactions

• Reaction with oxygen (air):
  • Alkali metals form oxides, peroxides or superoxides depending on cation size: Li → oxide (Li2O); Na → peroxide (Na2O2); K, Rb, Cs → superoxides (KO2, RbO2, CsO2). The choice is governed by ability of large cations to stabilise larger oxygen anions.
  • Alkaline earth metals generally form normal oxides (MO) on burning (e.g., MgO, CaO). Heavy members may form peroxides in some conditions.
• Reaction with water:
  • Alkali metals react vigorously with water to give hydroxides and hydrogen: 2M + 2H2O → 2MOH + H2 (M = alkali metal). Reactivity increases down the group (K, Rb, Cs more vigorous than Na, Li).
  • Alkaline earth metals: Be does not react with water; Mg reacts very slowly with cold water but reacts with steam to give oxide and H2; Ca, Sr, Ba react with cold water to give M(OH)2 and H2 (less vigorous than alkali metals).
• Reaction with hydrogen:
  • Alkali metals form ionic hydrides: 2M + H2 → 2MH (e.g., NaH). These hydrides are strong reducing agents and react with water to release H2.
  • Alkaline earth metals form MH2 (e.g., CaH2) on direct combination with hydrogen under suitable conditions.
• Reaction with halogens (X2):
  • 2M + X2 → 2MX (MX are ionic halides; very exothermic). Stability increases with ionic character.
• Reaction with nitrogen:
  • Lithium uniquely reacts with nitrogen to form nitride: 6Li + N2 → 2Li3N. Other alkali metals do not form nitrides directly.
  • Group 2 metals do not normally form binary nitrides by direct reaction with N2 under mild conditions, though Ca3N2 can be prepared under forcing conditions.
• Reaction with acids:
  • All s-block metals react with dilute acids to give corresponding salt and H2: M + nH+ → M^{n+} + (n/2)H2. E.g., Mg + 2HCl → MgCl2 + H2; 2Na + 2HCl → 2NaCl + H2.
• Thermal stability of carbonates and nitrates:
  • Carbonates: Group 1 carbonates are generally stable on heating (exception: Li2CO3 which decomposes on strong heating). Group 2 carbonates decompose on heating to give oxides and CO2: MCO3 → MO + CO2 (thermal stability increases down the group).
  • Nitrates: Most Group 1 nitrates decompose on heating to give nitrite + O2: 2MNO3 → 2MNO2 + O2 (LiNO3 behaves differently). Group 2 nitrates decompose to oxide, NO2 and O2: 2M(NO3)2 → 2MO + 4NO2 + O2.
• Acid-base character:
  • S-block oxides and hydroxides are largely basic. Basic strength increases down the group (e.g., LiOH & NaOH strong; BeO and Be(OH)2 are amphoteric; Mg(OH)2 sparingly soluble).

Why these trends? (Brief rationale)

• Lower ionization energy down the group → easier electron loss → greater metallic character and reactivity.
• Cation size influences which oxygen species are stabilised (oxide vs. peroxide vs. superoxide): larger cations stabilise larger oxygen anions due to lower lattice energy penalty.
• Polarising power of small, highly charged cations (Li+, Be2+) gives covalent character and anomalous behaviour (e.g., Be forms covalent compounds and is amphoteric).

Important safety note: Alkali metals react violently with water and must be handled carefully in small pieces under oil in lab demonstrations; many reactions are exothermic and can ignite hydrogen.

Fig 19 — Educational Diagram: Hydrides

Definition
Hydrides are binary compounds of hydrogen with other elements. In hydrides hydrogen can occur as H− (hydridic), H+ (protonic) or as part of covalent bonds. Hydrides are classified into three main types: ionic (saline), covalent (molecular or polymeric) and metallic (interstitial).

Types of hydrides

• Ionic (Saline) Hydrides: Formed by very electropositive metals (group 1 and 2). They contain H− (hydride ion). Example: NaH, CaH2. Characteristics: high melting points, crystalline, react vigorously with water to give H2, strong bases and reducing agents. Typical formation (alkali): 2M + H2 → 2MH.
• Covalent Hydrides: Formed with p‑block elements (and some small cations like Be2+ giving polymeric covalent hydrides). Examples: CH4, NH3, H2O (molecular covalent); BeH2 is polymeric (linear chains). Characteristics: lower melting/boiling points (molecular ones), solubility and reactivity depend on polarity and hydrogen bonding.
• Metallic (Interstitial) Hydrides: Formed by transition metals and some lanthanides where H atoms occupy interstitial sites in metal lattices (non‑stoichiometric often). Examples: PdHx, LaNi5Hx. Characteristics: metallic conductivity retained, used for reversible hydrogen storage, often non‑stoichiometric (x variable).

Bonding and oxidation state
In ionic hydrides H has oxidation state −1 (hydridic). In hydrides with more electronegative elements (O, N, F) hydrogen is +1 (protonic). With elements of similar electronegativity (C, Si) bonding is covalent and H is formally 0 to +1 depending on context.

Preparation (typical methods)

• Direct combination: alkali/alkaline-earth metals with H2. Example: 2Na + H2 → 2NaH; Ca + H2 → CaH2.
• Reduction of halides or oxides by hydride donors (lab scale): e.g., formation of covalent hydrides by LiAlH4 or NaBH4 reductions (organic/analytical use).
• Absorption of H2 by transition metals at suitable temperatures/pressures to give interstitial hydrides (e.g., Pd + x/2 H2 → PdHx).

Chemical properties

• Ionic hydrides react with water to produce hydrogen and the metal hydroxide: 2NaH + 2H2O → 2NaOH + 2H2↑; CaH2 + 2H2O → Ca(OH)2 + 2H2↑. This makes them useful as drying agents and laboratory H2 sources.
• Ionic hydrides are strong bases (H− is a powerful proton acceptor) and strong reducing agents (reducing metals, organic functional groups in presence of suitable reagents).
• Covalent hydrides show varied chemistry: NH3 is a base and forms hydrogen bonds; CH4 is chemically inert under normal conditions; BH3 (as diborane) is a strong Lewis acid.
• Metallic hydrides can reversibly absorb/release H2 and are important for hydrogen storage technologies.

Trends and important points (s‑block focus)

• Alkali and alkaline earth metals form predominantly ionic hydrides. Smaller, highly polarizing cations (Be2+, Mg2+) can impart covalent character—BeH2 is polymeric and covalent.
• Hydride stability: formation of ionic hydrides is favoured for more electropositive metals (group 1 & 2). Hydride ion (H−) is a strong base—its reactivity with protic solvents increases practical handling hazards.
• In s‑block hydrides hydrogen has oxidation state −1; in hydrides with more electronegative elements (O, N, F) H is +1.

Uses / real‑life importance

• CaH2 and NaH — drying agents and laboratory sources of hydrogen.
• Metal hydrides (LaNi5Hx, PdHx) — hydrogen storage, hydrogen purification and sensors.
• CH4 (methane) — main component of natural gas (fuel).
• NH3 — fertilizer production and also a hydrogen carrier in some technologies.

Safety
Ionic hydrides react violently with water and moist air to release flammable H2; they must be handled in dry inert atmospheres.

Fig 20 — Educational Diagram: Reactions with Oxygen: Oxides, Peroxides and Superoxides

Overview: s‑Block metals (alkali and alkaline earth metals) react with oxygen to give three types of binary oxygen-containing salts depending on the metal cation size and electrostatic stabilization: oxides (containing O2−), peroxides (containing O22−) and superoxides (containing O2−). The type formed changes down the group as the cation becomes larger and can better stabilize larger anions.

Definitions and electronic structure:

• Oxide: contains O2− (oxide ion). O–O bond is absent (oxygen as single atoms bonded to metal).
• Peroxide: contains O22− with an O–O single bond (bond order = 1).
• Superoxide: contains O2− (one extra electron on O2), bond order = 1.5; the species is paramagnetic (has an unpaired electron).

Which metals form which species?:

• Alkali metals: Li → oxide (Li2O); Na → mainly peroxide (Na2O2); K, Rb, Cs → superoxides (KO2, RbO2, CsO2) (larger cations stabilize the larger O2− anion).
• Alkaline earth metals: Be, Mg → oxides (BeO, MgO). Heavier ones (Ca, Sr, Ba) can form peroxides (CaO2, SrO2, BaO2); superoxides are not common for divalent cations.

Representative formation reactions with oxygen:

• Oxide formation (small cations): 4Li + O2 → 2Li2O.
• Peroxide formation (intermediate cations): 2Na + O2 → Na2O2.
• Superoxide formation (large cations): K + O2 → KO2.

Chemical behaviour with water (important practical difference):

• Oxide (basic): M2O + H2O → 2MOH (e.g., Li2O + H2O → 2LiOH).
• Peroxide (gives hydrogen peroxide on hydrolysis): M2O2 + 2H2O → 2MOH + H2O2 (e.g., Na2O2 + 2H2O → 2NaOH + H2O2).
• Superoxide (gives O2 as well as H2O2): 2MO2 + 2H2O → 2MOH + H2O2 + O2 (e.g., 2KO2 + 2H2O → 2KOH + H2O2 + O2).

Trends and stability:

• Basic strength of oxides increases down the group (Li2O < Na2O < K2O etc. as M–O ionic character increases).
• Stability of peroxides and superoxides increases with increasing cation size because the larger cation better stabilizes the large O2n− anions electrostatically.
• Peroxides and superoxides are strong oxidizing agents (they can release O2 and form OH− or generate H2O2).

Bonding and properties:

• Peroxide O–O single bond length ~1.49 Å; superoxide O–O bond length shorter (~1.33 Å) due to higher bond order; oxide has no O–O bond.
• Superoxide ions are paramagnetic (one unpaired electron), peroxides are diamagnetic.

Decomposition / thermal behaviour:

• Peroxides can decompose on heating giving oxide and oxygen (simplified): 2M2O2 → 2M2O + O2.
• Superoxides may decompose to oxide + O2 on strong heating or reduction.

Uses and real‑life importance:

• KO2 (potassium superoxide) is used in closed‑circuit breathing apparatus (submarines, spacecraft) because it absorbs CO2 and releases O2.
• Na2O2 and BaO2 are used as sources of oxygen and as bleaching/oxidizing agents in industry and laboratories.
• CaO2 is used in environmental applications (slow oxygen release in soil/aquifers) and in some detergents/bleaches.
• Peroxides are precursors or sources of hydrogen peroxide (industrial and lab uses).

Safety notes: Peroxides and superoxides are strong oxidizers and may react violently with organics or moisture; they must be handled with care.

Summary in one paragraph: Reaction of s‑block metals with O2 yields oxides, peroxides or superoxides depending mainly on cation size and charge density. Small, highly polarizing cations (Li+, Be2+, Mg2+) form simple oxides; intermediate cations stabilize peroxides; large cations (K+, Rb+, Cs+) stabilize superoxide ions. These compounds differ in O–O bonding, magnetic properties, hydrolysis products (hydroxide vs H2O2 and O2) and uses in oxygen generation and oxidation chemistry.

Fig 21 — Educational Diagram: Oxides, Peroxides and Superoxides

Overview
Alkali and alkaline earth metals form three important types of oxygen-containing species: oxides (O2−), peroxides (O22−) and superoxides (O2−). Which species is formed depends mainly on the size and charge density of the metal cation: small, highly polarizing cations stabilize oxide; larger cations stabilise peroxide and superoxide.

Definitions and electronic structure

• Oxide anion: O2− (formally two electrons added to O atom). Oxide in ionic compounds usually written as O2− (example: Li2O, MgO).
• Peroxide ion: O22− (two oxygen atoms joined by a single O–O bond; each O formally −1). Bond order = 1. Example: Na2O2, H2O2.
• Superoxide ion: O2− (one extra electron on O2 molecule; one oxygen unit with overall −1 charge). Bond order = 1.5. Example: KO2, RbO2, CsO2.

Bonding and molecular properties

• Using molecular orbital (MO) ideas: adding electrons to antibonding π* orbitals reduces bond order. Thus bond orders: O2 (neutral) = 2; O2− (superoxide) = 1.5; O22− (peroxide) = 1. Bond lengths increase as bond order decreases.
• Typical O–O bond lengths (approx.): O2 = 1.21 Å; O2− ≈ 1.33 Å; O22− ≈ 1.49 Å.

How they form (general reactions)

• Direct oxidation of metal: 4Li + O2 → 2Li2O (oxide); 2Na + O2 → Na2O2 (peroxide, often mixed); K + O2 → KO2 (superoxide) under suitable conditions.
• Peroxides and superoxides often form when oxygen reacts with larger cations because the larger cation better stabilizes the larger anion by lattice energy.

Chemical properties

• Basicity: Oxides of s-block metals are strongly basic (M2O + H2O → 2MOH). Peroxides and superoxides produce hydroxide plus additional oxygen species on hydrolysis:
• Peroxide hydrolysis: Na2O2 + 2H2O → 2NaOH + H2O2 (hydrogen peroxide formed).
• Superoxide hydrolysis (balanced overall): 2KO2 + 2H2O → 2KOH + H2O2 + O2 (or in practical devices the reaction is often written to show release of O2 and formation of hydroxide).
• Oxidizing behaviour: Peroxides and superoxides are good oxidizing agents (H2O2 is a common oxidizer and bleach). Superoxides liberate O2 on reaction and are used as oxygen sources.
• Reaction with acids: Oxides and peroxides give O2 or H2O2 depending on conditions; superoxides give O2 and can produce singlet oxygen or reactive O species.

Thermal stability and trends

• General trend in Group 1 (increasing atomic size Li → Cs): oxide (Li2O) → peroxide (Na2O2 often) → superoxide (K, Rb, Cs produce KO2 etc.).
• Small cations (Li+, Mg2+) polarize and stabilize oxide (O2−). Large cations (K+, Rb+, Cs+) stabilize larger anions (O2− and O22−) by lattice energy.
• Peroxide and superoxide salts are less thermally stable than simple oxides; decomposition typically releases O2.

Uses and hazards

• H2O2 (hydrogen peroxide): disinfectant, bleaching agent, laboratory reagent, rocket propellant (high concentrations).
• Na2O2 and KO2: oxygen sources in rebreathers, emergency oxygen generators, CO2 absorbents. KO2 reacts with CO2 to give carbonate and O2 (useful in closed atmospheres).
• Peroxides and superoxides are strong oxidizers and can be explosive when contaminated or heated; handle with care.

Representative reactions (summary)

• Oxide formation: 4Li + O2 → 2Li2O
• Peroxide formation (sodium): 2Na + O2 → Na2O2 (or mixture)
• Superoxide formation (potassium): K + O2 → KO2
• Oxide hydrolysis: M2O + H2O → 2MOH
• Peroxide hydrolysis: Na2O2 + 2H2O → 2NaOH + H2O2
• Superoxide hydrolysis (overall): 2KO2 + 2H2O → 2KOH + H2O2 + O2
• KO2 with CO2 (air-regeneration): 4KO2 + 2CO2 → 2K2CO3 + 3O2

Practical points for students

• Remember: oxide (O2−) = bond order 2 (shortest O–O if present as O2 fragments), peroxide (O22−) = single O–O bond, superoxide (O2−) = intermediate bond order 1.5.
• Trend by cation size: smaller cations → oxide; larger cations → peroxide/superoxide.
• Recognize formulas: Li2O (oxide), Na2O2 or Na2O/Na2O2 mixtures (sodium gives peroxides), KO2 (superoxide).

Tip: To visualise bonding changes, compare MO diagrams of O2, O2− and O22−; to see macroscopic trends, plot cation radius versus type of oxide formed or versus stability (enthalpy) of peroxide/superoxide formation.

Fig 22 — Educational Diagram: Reactions with Water and Steam

Overview
s‑Block metals (alkali metals, Group 1; alkaline earth metals, Group 2) react with water or steam to produce metal hydroxides/oxides and hydrogen. Reactivity generally increases down each group because ionization energy decreases and the metal more readily loses electrons.

General behavior

• Alkali metals (Group 1): react vigorously with cold water to give metal hydroxide and hydrogen; reaction is strongly exothermic and hydrogen can ignite for the heavier metals.
• Alkaline earth metals (Group 2): lighter members (Be, Mg) do not react (Be inert, Mg unreactive with cold water) because of high ionization energy and protective oxide layer. Mg and heavier members react with steam to give metal oxide and hydrogen; Ca, Sr, Ba react with cold water to give hydroxides and hydrogen.

Mechanism (qualitative)
The metal atom loses electrons to form M+ or M2+. Water is reduced to produce H2 while hydroxide (OH−) or oxide (O2−, when steam oxidizes metal surface) is formed. Reactions are often surface-controlled: an oxide/hydroxide layer can slow or prevent further reaction.

Key observations

• Bubbling of H2, solution becomes basic (pH increases).
• Exothermic heating; for K, Rb, Cs the hydrogen can ignite giving flames (K — lilac/violet, Na — yellow, Li — crimson).
• White solid precipitates for sparingly soluble hydroxides (e.g., Ca(OH)2).

Why some metals don't react with cold water
Be and Mg: very strong Be–O/Mg–O bonds, high ionisation energies and protective oxide films make reaction with cold water negligible. Mg reacts readily with steam because higher temperature removes the oxide film and provides energy to form MgO.

Practical/industrial significance

• Preparation of hydrogen in laboratory: Mg + H2O (steam) → MgO + H2.
• Formation of metal hydroxides used in industry: Na + H2O → NaOH + 1/2 H2 (in principle; industrial NaOH is made by electrolysis, but NaOH is the hydroxide product when sodium reacts with water).
• Slaking of lime: CaO + H2O → Ca(OH)2 (exothermic) — important in construction (mortar, plaster).

Fig 23 — Educational Diagram: Hydroxides

What are hydroxides?
Hydroxides are compounds that contain the hydroxide ion (OH−) paired with a metal cation: general formulas are MO (for M+ species in solution as M+ + OH−) or M(OH)2 (for M2+). In s‑block chemistry (Group 1 and Group 2) hydroxides are important basic compounds whose properties are largely determined by the cation size, charge and polarising power.

Preparation (typical reactions)

• Reaction of metal with water (alkali and alkaline earth metals): 2M + 2H2O → 2M+ + 2OH− + H2 (Group 1, e.g. Na)
• For alkaline earth metals: M + 2H2O → M(OH)2 + H2 (e.g. Ca → Ca(OH)2)
• From oxide + water: MO + H2O → M(OH)2 (alkaline earth oxides) or M2O + H2O → 2M+ + 2OH− (alkali oxides)
• Neutralisation: Metal hydroxide + acid → salt + H2O (e.g. NaOH + HCl → NaCl + H2O)

Key physical and chemical properties

• Basicity: Alkali metal hydroxides (LiOH, NaOH, KOH, …) are strong bases and practically fully dissociate in water. Alkaline earth hydroxides show increasing basicity down the group: Be(OH)2 is amphoteric (reacts with acids and bases), Mg(OH)2 is sparingly soluble and weakly basic, Ca(OH)2 (slaked lime) is moderately soluble, Sr(OH)2 and Ba(OH)2 are stronger bases in solution.
• Solubility trend: For Group 1 hydroxides solubility is very high and generally increases down the group. For Group 2 hydroxides solubility increases down the group (Mg(OH)2 sparingly soluble → Ca(OH)2 slightly soluble → Sr/Bar more soluble). The trend arises from the balance between lattice enthalpy (decreases down group) and hydration enthalpy (also decreases but less rapidly), favouring solubility for larger cations.
• Thermal stability: Generally increases down the group. Group 2 hydroxides decompose on heating to give metal oxide and water: M(OH)2 → MO + H2O; decomposition temperature increases down the group (Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2). LiOH is less thermally stable among Group 1 hydroxides (Li+ polarising effect).
• Chemical reactions: Neutralisation with acids, reaction with CO2 (e.g. Ca(OH)2 + CO2 → CaCO3 + H2O — important in construction and limewater tests), reaction with ammonium salts (caustic extraction), and metal hydroxides can precipitate from solutions of metal ions by adding OH−.

Special points

• Be(OH)2 is amphoteric (reacts with acids to give Be2+ salts and with bases to form beryllate complex ions).
• Hydroxides of group 1 are used as strong bases in industry (NaOH, KOH). Mg(OH)2 and Al(OH)3 (not s‑block) are used as antacids; Ca(OH)2 is widely used in construction and water treatment.

Why these trends? (brief energetic view)
Solubility and stability trends result from competition between lattice enthalpy (energy holding the solid together) and hydration enthalpy (energy released when ions are solvated). As cation size increases down a group lattice enthalpy falls faster than hydration enthalpy, so solubility and basic strength in water generally increase down the group.

Fig 24 — Educational Diagram: Reactions with Halogens and Hydrogen

Overview
s‑Block metals (alkali metals, Group 1; and alkaline earth metals, Group 2) react readily with halogens and hydrogen to give predominantly ionic binary compounds: metal halides (MX or MX2) and metal hydrides (MH or MH2). These reactions are generally exothermic and often vigorous because s‑block metals have low ionization energies and form stable M+ / M2+ ions.

Reaction with halogens (X2)
General behaviour:
- Alkali metals (M, Group 1): combine with halogens to give ionic halides. Example reaction: 2M + X2 → 2MX (e.g., 2Na + Cl2 → 2NaCl).
- Alkaline earth metals (M, Group 2): form dihalides: M + X2 → MX2 (e.g., Mg + Cl2 → MgCl2).
Key points:
• The reactions are usually vigorous (sometimes explosive), producing crystalline ionic halides with high lattice energy and high melting points.
• Fajans' rule: small, highly charged cations (Li+, Be2+) strongly polarize anions → increased covalent character. Example: BeCl2 is covalent (forms polymeric chains), LiCl shows more covalent character than other alkali chlorides. Larger cations (K+, Ca2+) produce more ionic halides.
• Reactivity trends: reactivity of metals toward halogens increases down Group 1 and Group 2 (easier to oxidize). Reactivity of halogens decreases down Group 17 (F2 > Cl2 > Br2 > I2).

Reaction with hydrogen
General behaviour:
- Alkali metals: form ionic (saline) hydrides of formula MH via combination with hydrogen: M + 1/2 H2 → MH (stoichiometrically often written as 2M + H2 → 2MH). Example: 2Na + H2 → 2NaH.
- Alkaline earth metals: form MH2 (e.g., Ca + H2 → CaH2).
Key properties of s‑block hydrides:
• They are ionic (hydride ion H−) for metals of low ionization energy; hydrides are strong reducing agents and react with water to liberate H2 (e.g., NaH + H2O → NaOH + H2; CaH2 + 2H2O → Ca(OH)2 + 2H2).
• Thermal stability and hydrogen content vary: MgH2 and other light metal hydrides are researched for hydrogen storage because they store hydrogen at high gravimetric capacity but often require heating to release H2.
• Small, highly charged cations (Be2+) do not form simple ionic hydrides; for Be, chemistry is more covalent and complex.

Chemical and practical consequences
- Ionic halides from s‑block metals (e.g., NaCl, KCl, CaCl2, MgCl2) are widely used in everyday life: NaCl is common salt; KCl is a fertilizer; CaCl2 is a drying agent and used for de‑icing; MgCl2 is used in chlorides industry and as coagulant.
- Metal hydrides: NaH is used as a drying agent and reducing agent in organic synthesis; CaH2 is a desiccant; MgH2 and complex hydrides (not purely s‑block but related) are investigated for hydrogen storage for fuel applications.

Exceptions and finer points
- Li+ and Be2+ have high polarising power; their halides and hydrides show greater covalent character (LiH less ionic character than heavier MH? LiH is ionic but shows some covalent aspects compared to larger cations).
- BeCl2 is covalent and exists as a polymeric chain or dimeric species in vapour; MgCl2 has significant ionic character but some covalency.

Fig 25 — Educational Diagram: Halides

Definition: Halides of s‑block elements are binary compounds formed between alkali/alkaline‑earth metals (Group 1 and 2) and halogens (F, Cl, Br, I). General formulas are MX (alkali metal halides) and MX2 (alkaline‑earth metal halides).

Bonding and structure: Most s‑block halides are predominantly ionic because they pair small, low‑ionization‑energy metal cations (M+) with electronegative halide anions (X−). Crystal structures commonly observed are the NaCl (6:6 coordination) and CsCl (8:8 coordination) types. Smaller cations and larger anions favor higher covalent character by polarization (Fajans' rules).

Fajans' rules (summary): A cation with high charge and small radius (high polarizing power) and an anion with large size and high polarizability favor covalent character. Thus Be2+ and Li+ (small cations) make more covalent halides than larger alkali/alkaline earth cations.

Special/anomalous behaviour: Be and Li show diagonal relationship to B and Mg respectively. BeCl2 and LiCl have appreciable covalent character; BeCl2 is polymeric in the solid state (chain) and volatile as monomer/dimer on heating. BeF2 has a network structure analogous to silica.

Physical properties & trends: Alkali halides are crystalline, generally high‑melting and soluble in water (solubility depends on balance of lattice energy and hydration energy). Melting points of alkali metal chlorides (for example) decrease down the group as ionic radii increase and lattice energy decreases. Alkaline earth halides (MX2) often have higher lattice energies and higher melting points than corresponding MX; however BeX2 and sometimes MgX2 show more covalent/volatile behavior.

Chemical reactivity: Formation by direct combination: 2M + X2 → 2MX (or M + X2 → MX for monovalent M). Metal oxides/hydroxides react with hydrogen halides: M(OH)2 + 2HX → MX2 + 2H2O. Hydrolysis: strongly polarizing small cations (Be2+) hydrolyze water producing acidic/basic salts (e.g., partial hydrolysis of BeCl2).

Solubility and ionic character: Solubility depends on lattice energy vs hydration energy. Highly ionic salts (large hydration energy relative to lattice energy) are soluble. Sparingly soluble halides are described by their solubility product, Ksp, via the equilibrium MX(s) ⇌ M+ + X−.

Uses & importance: s‑Block halides are widely used: NaCl (table salt), CaF2 (fluorite, source of HF and used in optics), MgCl2 and CaCl2 (de‑icing, dust control, desiccants), MgF2 (optical coatings), LiF (specialized ceramics and fluxes). Some are important reagents in industry and synthesis.

Safety note: Many halides (especially soluble heavy‑metal halides) can be toxic; Be compounds are highly toxic and must be handled with care.

Fig 26 — Educational Diagram: Oxides and Hydroxides: Basicity and Amphoterism

Overview
Oxides and hydroxides of s‑block elements (alkali and alkaline earth metals) are mostly ionic and show basic behaviour: oxides react with water to give hydroxides, and hydroxides neutralise acids. Basicity generally increases down a group because ionic character and solubility increase. A notable exception in the s‑block is beryllium: BeO and Be(OH)₂ are amphoteric (react with both acids and bases).

Definitions
Basic oxide: an oxide that reacts with acids to form salt and water (e.g., CaO + 2HCl → CaCl₂ + H₂O).
Amphoteric oxide/hydroxide: a substance that reacts both with acids and with bases to give salts/complexes (e.g., BeO, Be(OH)₂).

Why basicity changes down the group
- Ionic character increases down the group (larger metal cation, lower polarising power), so M–O or M–OH bonds become more ionic and oxides/hydroxides behave more like bases.
- Solubility of hydroxides (especially in group 2) generally increases down the group (Be and Mg hydroxides are sparingly soluble; Ca, Sr, Ba hydroxides are more soluble), so basic behaviour becomes more pronounced down the group.
- Lattice energy vs hydration energy: for small cations (Be²⁺, Mg²⁺) high lattice energy and high covalent character reduce basicity/solubility; for larger cations hydration energy dominates less and ionic, soluble hydroxides form readily.

Amphoterism in s‑block
- Beryllium (Be) is diagonal to aluminium and shows covalent character. BeO and Be(OH)₂ are amphoteric:
BeO + 2HCl → BeCl₂ + H₂O (reaction with acid)
BeO + 2NaOH + H₂O → Na₂[Be(OH)₄] (reaction with base forming beryllate)
- Other s‑block oxides/hydroxides (Li, Na, K, Rb, Cs; Ca, Sr, Ba) are predominantly basic. Mg(OH)₂ is only sparingly soluble and behaves mainly as a weak base (used as an antacid), not strongly amphoteric.

Special oxide types in alkali metals
Alkali metals form different oxidised species depending on metal size:
- Oxide: M₂O (Li forms Li₂O)
- Peroxide: M₂O₂ (Na forms Na₂O₂)
- Superoxide: MO₂ (K, Rb, Cs form KO₂, etc.)
These differ in O–O bonding and reactivity with water (superoxides give O₂ and H₂O₂ products often).

Practical significance / real‑life uses
- CaO (quicklime) and Ca(OH)₂ (slaked lime): construction, pH control.
- NaOH (caustic soda): cleaning, soap manufacture; formed from Na₂O hydration in principle.
- Mg(OH)₂ (milk of magnesia): antacid and laxative—basic property neutralises stomach acid.
- KO₂ (potassium superoxide): used in closed breathing systems to generate O₂ and absorb CO₂ in respirators (reacts with CO₂ and H₂O).

Key teaching points to remember
- Basicity of oxides/hydroxides increases down Group 1 and Group 2 (except Be).
- Amphoterism appears when the metal has high polarising power and significant covalent character (Be in s‑block; many p‑block metals like Al, Zn, Pb are also amphoteric).
- Alkali metals can form oxide, peroxide, superoxide depending on their size; these show different reactions with water and acids.

Fig 27 — Educational Diagram: Carbonates, Bicarbonates and Nitrates

Basic definitions
Carbonate: salt of CO3²⁻. Bicarbonate (hydrogen carbonate): salt of HCO3⁻. Nitrate: salt of NO3⁻.

Formation and equilibria involving CO2
CO2 dissolves in water to form carbonic acid which is in equilibrium with bicarbonate and carbonate:

• CO2 + H2O ⇌ H2CO3 ⇌ HCO3⁻ + H⁺ ⇌ CO3²⁻ + 2H⁺
• With base: CO2 + OH⁻ → HCO3⁻; CO2 + 2 OH⁻ → CO3²⁻ + H2O

Bicarbonates
Bicarbonates (e.g., NaHCO3) are formed when CO2 reacts with hydroxides (CO2 + NaOH → NaHCO3). Bicarbonates are generally soluble (alkali metal bicarbonates are stable solids); many alkaline earth bicarbonates (e.g., Ca(HCO3)2) exist only in aqueous solution (formed by CO2 reacting with CaCO3 in water) and decompose on heating to give the carbonate:

• 2 NaHCO3 (heat) → Na2CO3 + CO2↑ + H2O
• Ca(HCO3)2 (on heating) → CaCO3↓ + CO2 + H2O (this is why scale forms on boiling hard water)

Carbonates
Carbonates of alkali metals (except Li2CO3) are generally stable to heat and are soluble (Na2CO3, K2CO3). Carbonates of alkaline earth metals (MCO3, M = Mg, Ca, Sr, Ba) are usually insoluble solids and decompose on heating to give the oxide and CO2:

• MCO3 (heat) → MO + CO2 (M = Mg, Ca, Sr, Ba)
• Exception / diagonal relation: Li2CO3 behaves like MgCO3 and can decompose on strong heating.

Thermal stability trend (carbonates and nitrates)
Thermal stability of carbonates and nitrates increases down a group. Reason: larger cations have lower polarizing power, so they distort the anion (CO3²⁻ or NO3⁻) less and the salt is more stable to heat. Small, highly charged cations (e.g., Li⁺, Mg²⁺, Be²⁺) polarize the anion strongly and cause decomposition at lower temperatures.

Nitrates
Alkali metal nitrates (NaNO3, KNO3) decompose on strong heating to nitrite and oxygen:

• 2 KNO3 (heat) → 2 KNO2 + O2↑

Alkaline earth metal nitrates (M(NO3)2) decompose to the oxide, NO2 and O2:

• 2 M(NO3)2 (heat) → 2 MO + 4 NO2↑ + O2↑ (M = Mg, Ca, Sr, Ba)
• LiNO3 behaves like the alkaline earth nitrates on heating: 4 LiNO3 → 2 Li2O + 4 NO2 + O2

Chemical behaviour with acids and water
Carbonates react with acids to give CO2:

• Na2CO3 + 2 HCl → 2 NaCl + H2O + CO2↑

Carbonate ion is basic in solution:

• CO3²⁻ + H2O ⇌ HCO3⁻ + OH⁻

Applications & environmental relevance
- NaHCO3 (baking soda): baking, antacid, fire extinguisher cartridges.
- Na2CO3 (soda ash): glass manufacture, detergents, chemical industry.
- CaCO3 (limestone): building material, lime (CaO) produced by calcination: CaCO3 → CaO + CO2; important in cement and steel industries.
- KNO3, NaNO3: fertilizers, KNO3 in gunpowder and pyrotechnics; nitrates important in nitrogen cycle and water pollution (eutrophication).

Key points to remember

• Bicarbonates are intermediate species between CO2/H2CO3 and carbonate; heating shifts bicarbonate → carbonate + CO2 + H2O.
• Thermal stability of carbonates/nitrates increases down the group due to decreasing polarizing power of cations.
• Alkali metal carbonates (except Li) are thermally stable; alkaline earth carbonates decompose to oxides and CO2.
• Alkali metal nitrates decompose to nitrites + O2; alkaline earth nitrates decompose to oxides + NO2 + O2.

Fig 28 — Educational Diagram: Carbonates, Bicarbonates and Thermal Stability

Overview
Carbonates contain the carbonate ion CO3(2-) combined with metal cations (M2+ or 2M+). Bicarbonates (hydrogen carbonates) contain HCO3(-). Thermal stability refers to how readily a carbonate decomposes on heating to give oxide and carbon dioxide (CO2) or how a bicarbonate decomposes to give carbonate, CO2 and H2O.

Formation and interconversion

• CO2 dissolves in water to give carbonic acid: CO2 + H2O → H2CO3 (in solution H2CO3 is in equilibrium with CO2 + H2O).
• Carbonate <-> bicarbonate equilibrium: CO3(2-) + H2O <-> HCO3(-) + OH(-) and H2CO3 <-> H+ + HCO3(-).
• Metal bicarbonates are typically formed by dissolving CO2 in a solution of the metal carbonate: M2CO3 + CO2 + H2O → 2MHCO3.

Thermal decomposition reactions (general)

• Bicarbonate decomposition on heating: 2NaHCO3 (s) → Na2CO3 (s) + CO2 (g) + H2O (g).
• Carbonate decomposition (common for M2+ cations): MCO3 (s) → MO (s) + CO2 (g). Example: CaCO3 (s) → CaO (s) + CO2 (g).
• Alkali metal carbonates (M+): most (Na, K, Rb, Cs) are thermally stable and do not decompose easily; Li2CO3 behaves anomalously and decomposes on strong heating: Li2CO3 → Li2O + CO2.

Trends and cause of thermal stability

• Thermal stability of carbonates increases down a group in the s-block. Reason: larger cations have lower polarizing (charge-dense) power and thus distort the CO3(2-) ion less, making the carbonate anion less susceptible to decomposition to CO2.
• For cations of higher charge (e.g., M2+ vs M+), polarizing power is greater, so M2+ carbonates are less stable than M+ carbonates of comparable size. Thus, most Group 2 carbonates decompose on heating, whereas Group 1 (except Li) carbonates are stable.
• Qualitative energetic view: decomposition involves breaking C–O bonds in carbonate. A small, highly polarizing cation lowers the energy required for this (by distorting the anion), so such carbonates decompose at lower temperatures. Lattice enthalpy and stability of the resulting oxide also play roles.

Group-specific notes

• Group 1 (alkali metals): Na2CO3, K2CO3, Rb2CO3, Cs2CO3 are thermally stable. Li2CO3 is less stable (Li+ is small and highly polarizing) and decomposes on strong heating.
• Group 2 (alkaline earth metals): carbonates decompose to oxides + CO2. Thermal stability increases down the group: MgCO3 < CaCO3 < SrCO3 < BaCO3. BeCO3 is unstable and tends to hydrolyze to basic carbonates.
• Bicarbonates are stable in solution for alkali metals (NaHCO3, KHCO3). Solid bicarbonates generally decompose on heating to carbonate, CO2 and H2O. Alkaline earth bicarbonates (e.g., Ca(HCO3)2) exist only in aqueous solution (they precipitate as CaCO3 on boiling).

Practical/industrial significance

• Thermal decomposition of CaCO3 in lime kilns: CaCO3 (limestone) → CaO (quicklime) + CO2 — fundamental in cement and lime industry.
• NaHCO3 (baking soda) releases CO2 on heating or when reacting with acid — used in baking, fire extinguishers, and as an antacid.
• Temporary hardness of water is due to Ca(HCO3)2 and Mg(HCO3)2; boiling decomposes bicarbonates to insoluble CaCO3 (scale) and removes hardness.

Summary of key points

• Carbonates of small, highly charged cations are less thermally stable because of strong polarization of CO3(2-).
• Thermal stability of carbonates increases down the group for both Group 1 (with Li exception) and Group 2.
• Bicarbonates decompose on heating to carbonates, CO2 and H2O; some bicarbonates exist only in solution.

Fig 29 — Educational Diagram: Important Compounds: Preparation, Properties and Uses

Overview: The s‑Block elements (alkali and alkaline earth metals) form a set of commercially and biologically important ionic compounds. Key compounds include sodium carbonate (washing soda), sodium bicarbonate (baking soda), sodium hydroxide (caustic soda), sodium chloride (common salt), calcium carbonate (limestone), calcium oxide (quicklime), calcium hydroxide (slaked lime), magnesium oxide and magnesium hydroxide. Below are concise, exam‑relevant preparations, properties and uses with main chemical equations.

1. Sodium carbonate, Na2CO3 (washing soda, often as decahydrate Na2CO3·10H2O)

• Preparation (Solvay process, overall): CaCO3 + 2 NaCl → Na2CO3 + CaCl2. Key steps:
• NH3 + CO2 + H2O + NaCl → NaHCO3 ↓ + NH4Cl (NaHCO3 precipitates)
• 2 NaHCO3 → Na2CO3 + CO2 + H2O (by heating)
• Properties: white crystalline solid, strongly basic (carbonate), soluble in water, decomposes on strong heating to oxide + CO2.
• Uses: glass and detergent manufacture, water softening, laboratory reagent.

2. Sodium bicarbonate, NaHCO3 (baking soda)

• Preparation: obtained during Solvay process as precipitate; can also be made by passing CO2 into Na2CO3 solution: Na2CO3 + CO2 + H2O → 2 NaHCO3.
• Properties: white crystalline powder, mildly alkaline, decomposes on heating: 2 NaHCO3 → Na2CO3 + CO2 + H2O.
• Uses: baking (releases CO2), antacid, fire extinguishers, cleaning agent.

3. Sodium hydroxide, NaOH (caustic soda)

• Preparation (industrial): Electrolysis of brine (chlor‑alkali process): 2 NaCl + 2 H2O → 2 NaOH + Cl2 + H2 (at appropriate electrodes).
• Properties: strong base, highly caustic, hygroscopic, soluble in water with large heat release; reacts with acids to form salts.
• Important reactions: saponification of fats: triglyceride + 3 NaOH → glycerol + 3 RCOONa.
• Uses: soap and detergent manufacture, paper/pulp, alumina extraction (Bayer process), drain cleaners, chemical industry.

4. Sodium chloride, NaCl (common salt)

• Preparation: obtained by evaporation of seawater or mining of rock salt.
• Properties: white crystalline solid, highly soluble in water, strong electrolyte (Na+ and Cl− ions).
• Uses: food seasoning and preservation, raw material for chlorine and NaOH production, de‑icing roads.

5. Calcium carbonate, CaCO3 (limestone, marble)

• Occurrence: abundant in nature (limestone, shells).
• Preparation (laboratory): CaCl2 + Na2CO3 → CaCO3 ↓ + 2 NaCl.
• Properties: white, insoluble in water, decomposes on heating: CaCO3 → CaO + CO2 (calcination).
• Uses: building stone, cement manufacture, antacid, filler in paints and plastics.

6. Calcium oxide, CaO (quicklime) and Calcium hydroxide, Ca(OH)2 (slaked lime)

• Preparation: CaO is made by calcining CaCO3: CaCO3 → CaO + CO2.
• Slaking: CaO + H2O → Ca(OH)2 (exothermic).
• Properties: CaO is basic and reacts vigorously with water; Ca(OH)2 is moderately soluble (gives lime water), basic and reacts with CO2: Ca(OH)2 + CO2 → CaCO3 ↓ + H2O.
• Uses: cement and mortar, steelmaking (flux), water treatment, soil stabilization, CO2 scrubbing.

7. Magnesium oxide, MgO and Magnesium hydroxide, Mg(OH)2

• Preparation: MgO by calcination of MgCO3 or by burning Mg metal; Mg(OH)2 by precipitation: MgCl2 + 2 NaOH → Mg(OH)2 ↓ + 2 NaCl.
• Properties: MgO is a refractory, basic oxide; Mg(OH)2 is sparingly soluble and is a mild base.
• Uses: refractory linings, antacid and laxative (milk of magnesia), filler in plastics and rubbers.

8. Special note on beryllium: Beryllium and its compounds show covalent character and amphoteric behaviour (e.g., Be(OH)2 reacts with acids and bases). This is a deviation from typical s‑block ionic behaviour.

General chemical behaviours and trends:

• Carbonates and hydroxides of group 2 become more thermally stable down the group (decomposition temperature increases down the group).
• Solubility trends: group 1 carbonates are generally soluble; group 2 carbonates are less soluble (BaCO3 least soluble among common ones). Hydroxide solubility increases down group 2 (Be(OH)2 amphoteric, Mg(OH)2 sparingly soluble, Ca(OH)2 slightly soluble, Ba(OH)2 more soluble).

Safety and environmental points: NaOH is highly corrosive; CO2 released in decomposition and calcination; correct handling, neutralization and waste treatment are required in industry and labs.

Summary: Memorize key preparations (Solvay for Na2CO3, chlor‑alkali for NaOH, calcination for CaO), characteristic reactions (decomposition of bicarbonate/carbonate, slaking of lime), and main uses (glass, soap, cement, antacids, water treatment).

Fig 30 — Educational Diagram: Nitrates and Sulfates

Overview
Nitrates (NO3−) and sulfates (SO4 2−) are important oxyanions whose salts with s‑block metals (alkali and alkaline earth metals) are studied in Class 11. Key points are their structures, bonding, solubility trends, thermal stability and common reactions.

Structures & bonding
• Nitrate ion, NO3−: trigonal planar, three equivalent N–O bonds due to resonance; overall charge −1.
• Sulfate ion, SO4 2−: tetrahedral, four equivalent S–O bonds (resonance), charge −2.
• Small, highly charged cations (Li+, Be2+, Mg2+) have large polarising power and induce partial covalent character in salts, affecting properties (solubility, thermal stability, decomposition products).

Nitrates of s‑block elements
Solubility: Most alkali and alkaline earth nitrates are soluble in water (all alkali nitrates are soluble).
Thermal behaviour:
• Alkali metal nitrates (except Li): on strong heating they give nitrites and oxygen:
2 MNO3 → 2 MNO2 + O2 (M = Na, K, etc.).
• Lithium nitrate behaves differently (Li+ is small, more covalent): on heating it decomposes to oxide, NO2 and O2. Example (balanced):
4 LiNO3 → 2 Li2O + 4 NO2 + O2.
• Alkaline earth (M2+) nitrates decompose to metal oxide, NO2 and O2:
2 M(NO3)2 → 2 MO + 4 NO2 + O2 (M = Mg, Ca, Sr, Ba).
Reason: Thermal stability of nitrates increases down the group because larger cations polarise the anion less, so nitrates of heavier metals are more thermally stable.

Sulfates of s‑block elements
Solubility trends:
• Alkali metal sulfates (M2SO4): generally soluble (Na2SO4, K2SO4); Li2SO4 is less soluble than other alkali sulfates because of lithium's small size and higher polarising power.
• Alkaline earth sulfates (M SO4): solubility decreases down the group — MgSO4 (quite soluble), CaSO4 (sparingly soluble), SrSO4 and BaSO4 (practically insoluble). BaSO4 is used for qualitative tests and medical X‑ray contrast because it is very insoluble and opaque to X‑rays.
Chemical tests: Acidify solution (to remove CO3 2−), then add BaCl2 — a white precipitate of BaSO4 confirms sulfate.

Thermal stability of sulfates
• Sulfates are generally more thermally stable than nitrates and carbonates. For alkaline earths, thermal stability tends to increase down the group (BaSO4 is most stable). Very strong heating may decompose some sulfates to metal oxide and SO3 (or to SO2 + O2 under reducing conditions):
M SO4 → MO + SO3 (requires very high temperature and depends on M).

Applications & environmental aspects
• Nitrates: KNO3 and NaNO3 — fertilizers and oxidisers (fireworks, explosives); nitrates in drinking water are a contamination concern (methemoglobinemia in infants). Brown ring test (FeSO4 + concentrated H2SO4) detects NO3−.
• Sulfates: CaSO4·2H2O (gypsum) → plaster of Paris (construction/orthopaedics); MgSO4 (Epsom salt) — medicinal/bath salt; BaSO4 — X‑ray contrast agent; sulfate ions also occur naturally in minerals and can cause scale (industrial contexts).

Why trends occur (concise rationale)
Polarising power ∝ (charge)/(ionic radius). Small high‑charge cations (Li+, Be2+, Mg2+) polarise anion electron clouds more, increasing covalency, lowering solubility and altering decomposition products. Therefore thermal stability of ionic salts (nitrates, carbonates) generally increases down a group as cation size increases and polarising power decreases.