Alkali and Alkaline Earth Metals

• Comparative chemistry of Group IA, Group IIA, and Group 14 is examined.

Group IA (Alkali Metals)

• Includes Hydrogen (H), Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), and Francium (Fr).
• All silvery-white metals except Cesium, which is golden yellow.
• Excellent conductors of electricity.
• Soft and highly reactive.
• Have one valence electron, forming univalent ionic and colorless compounds.
• Oxides and hydroxides are strong bases.
• Oxosalts are very stable.
• Largest atomic size in their period.
• Smallest ionic size in their period after losing the single valence electron in their outermost shell.
• Lowest ionization energy in their respective period.
• Very reactive and form simple ionic compounds soluble in water.
• Stored under inert solvents (hydrocarbon solvent).
• Found combined in nature due to high reactivity.
• Readily form alloys among themselves (e.g., Na/K) and with other metals (e.g., Na/Hg).
• Tarnish rapidly in air, forming a layer of oxide, peroxide, and dioxide.
• Adopt a body-centered cubic structure at ambient temperature.
• At low temperature, lithium forms a hexagonal close-packed structure.
• Lithium is the first member and shows marked differences from other members.
• Head elements usually have a diagonal relationship with the diagonally next element.
• Sodium and potassium make up about 4% of the Earth's crust.
• Elements do not occur together primarily due to different ion sizes.
• Ores: tourmaline, spodumene, petalite, borax, mirabilite, sylvite, carnallite, lepidolite, avogadrite, londonite.
• Francium is found in trace amounts in uranium ores.
• Reactive metals cannot be extracted using furnace but can be extracted by electrolysis.
• Lithium gives crimson flame, sodium gives yellow, potassium gives lilac, rubidium gives red-violet, and caesium gives blue.

Key properties of Group IA Metals

• Atomic properties like symbol, number, mass.
• Valence electron configuration, melting/boiling point.
• Density, atomic radius, ionization energy.
• Oxidation state, ionic radius, electron affinity, electronegativity.
• Standard electrode potential.
• Products of reaction with oxygen and nitrogen, and type of oxide formed.
• *Values are cited for four-coordinate ions except for Rb+ and Cs+, whose values are given for the six-coordinate ion.

Lithium and Magnesium

• Exhibit high-level covalent character in their bonding.
• Carbonates decompose into metal oxide and carbon dioxide.
• Down the group IA, carbonates become more stable to thermal decomposition.
• Form normal oxides with oxygen (Group IA forms peroxide and dioxides).
• Peroxides form by reacting LiOH or Mg(OH)2 with H2O2.
• Lithium forms organometallic compounds similar to magnesium.
• Lithium combines with nitrogen to give Li3N; magnesium gives Mg3N2.
• LiF and MgF2 sparingly soluble in water.
• LiOH is much less soluble than other alkali metal hydroxides; Mg(OH)2 is sparingly soluble.
• LiClO4 is much more soluble in water than other alkali metal perchlorates.
• Lithium and magnesium nitrates decompose on heating to give oxide, nitrogen oxide, and oxygen.
• Sodium nitrate and later alkali metal nitrates decompose to give oxide.

Chemical Properties of Group IA Metals

• React with water to liberate hydrogen and form hydroxides.
• Tarnish rapidly in dry air; lithium forms a mixture of oxides and nitride.
• Form oxides or ozonides depending on reaction.
• All group IA metals form azides except lithium.
• React with hydrogen to form hydrides.
• React with halogens to form halides.
• React with interhalogen compounds to form ionic polyhalide compounds.
• React with liquid ammonia to form metal amide and hydrogen.
• Lithium reacts when heated with carbon to form carbide.
• Form soluble carbonates except Li2CO3.
• Form sulphates.
• Form phosphides, arsenides, and stibnides.
• Form sulfides, selenides, and tellurides.

Differences Between Lithium and Other Group IA Metals

• Li has higher melting and boiling points.
• Li is much harder.
• Reacts less readily with oxygen.
• Lithium hydroxide is less basic.
• Only Lithium forms nitride Li3N in group IA.
• Only Lithium reacts directly with carbon to form carbide.
• Lithium forms more complexes, and its salts are more heavily hydrated.
• Halides and alkyls of lithium are more covalent.

Group IIA (Alkaline Earth Metals)

• Includes Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra).
• Silvery, but Be and Mg are greyish.
• Highly reactive but less reactive than Group IA.
• Divalent and colorless ionic compounds.
• Oxides and hydroxides are less basic.
• Oxosalts are less stable to heat.
• Smaller atomic and ionic sizes but higher densities than Group IA counterparts.
• Two valence electrons participate in bonding, leading to harder metals with higher cohesive energy and higher melting/boiling points.
• Beryllium is the head element and shows marked differences.
• Dissolve in liquid ammonia like Group IA elements.
• Form amides slowly in liquid ammonia; Group IA directly yields metals.
• Cannot be extracted via oxidation or reduction but are purified by electrolysis.
• Ores include beryl, phenacite, kieserite, carnallite, limestone, gypsum, celestite, strontianite, baryte. Radium is radioactive and found in uranium ore.
• Flame test: beryllium (white), magnesium (brilliant white), calcium (brick red), strontium (crimson), barium (apple-green), radium (crimson red).

Key properties of Group IIA Metals

• Atomic properties like symbol, number, mass.
• Valence electron configuration, melting/boiling point.
• Density and atomic radius.
• First ionization energy, most common oxidation state, and ionic radius.
• Electron affinity, electronegativity, and standard electrode potential.

Chemical properties of Group IIA Metals

• Be reacts with steam, others react with water.
• Beryllium hydroxide is amphoteric.
• Bicarbonates are produced by bubbling excess carbon dioxide.
• All react with acids and liberate hydrogen.
• All burn in oxygen to form oxides.
• Thermal decomposition of oxosalts also gives oxide.
• All form sulfate, the solubility decreases down the group.

Chemical reactions of Group IIA Metals

• From nitrate by reaction of the nitric acid with carbonate, oxides or hydroxides.
• React with hydrogen to form hydride except beryllium.
• React with halogens to form halide.
• Burn in dinitrogen and form ionic nitride M3N2.
• All and their oxides react at high temperature with carbon to give carbide.

Relationship between beryllium and aluminum:

• Both form carbides which react with water to produce methane, while other Group IIA carbides produce ethyne.
• Form covalent hydrides, halides, and oxides.
• Chlorides fume in moist air.
• Oxides are amphoteric while magnesium oxide is basic.
• Form tetrahydroxo complexes with excess OH- ions.
• Hydrides are electron deficient and polymeric.
• Form many complexes, unlike Group IA and IIA.
• Be and Al are rendered passive by nitric acids.
• Standard electrode potentials are close.
• Salts are extensively hydrolyzed and are among the most soluble salts known.

Group 14 (Carbon Group)

• Includes Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), Lead (Pb), and Flerovium (Fl).
• Members vary from non-metal (C) to metalloid (Si, Ge) to metal (Sn, Pb).
• Carbon is abundant in biological systems (proteins, carbohydrates, fats).
• Silicon is abundant in the Earth's crust (24%).
• Lead exhibit allotropy.
• Ores include diamond, graphite, fullerene, charcoal, peat, Quartz, zeolite, Cassiterite, teallite, Galena, Anglesite, Winklers' ore.

Key properties of Group 14 elements:

• Atomic symbol and number.
• Atomic mass.
• Valence electron configuration.
• Melting/boiling point.
• Density.

Differences between carbon silicon and other elements:

• First element differ from the rest of the elements because of its smaller size and higher electronegativity, being more covalent and being less metallic than the other members of the group.
• Forms strong π – π multiple bonds.
• Exhibits catenation, or the ability to form chain-like structures.
• Contain only s and p orbital electron.

Relationship between B and Si:

• Both form acidic oxides.
• Form polymeric oxide structures.
• Form flammable gaseous hydrides.

Reactions of Group 14 Elements:

• Form tetravalent hydrides.
• Form tetrahalides with halogens.
• React with oxygen to produce oxide, dioxide.
• React with nitrogen to form stable compounds.
• Carbon forms stable sulphide with sulphur.

Atoms, molecules and structures

• Lewis octet rule: elements tend to bond to achieve eight valence electrons, resembling noble gases.
• Guidelines to obey the Lewis octet rule.

VSEPR (Valence Shell Electron-Pair Repulsion) Theory

• Each group of valence electrons around a central atom is located as far away as possible to minimize repulsions.
• Regions of high electron concentration repel one another, dictating molecular shape.
• Distinction between single and multiple bonds is disregarded.
• Lone pairs exert stronger repulsions than bonding pairs.
• AXEm notation identifies atom and lone pairs,
• Electrons repel each other in the order: lone pair-lone > lone pair-atom > atom-atom.
• Molecule shape is determined by location of the bonded atoms.

Explore properties of Group IA (alkali metals) and Group IIA (alkaline earth metals). Understand electronic configurations, ionic size, reactivity, and trends in properties. Investigate reasons for storing alkali metals under inert solvents.