P-Block Elements: Groups 13 & 14

The p-Block Elements

• The last electron enters the outermost p orbital.

• There are six groups of p-block elements, numbered 13 to 18.

• Valence shell electronic configuration: $ns^2 np^{1-6}$ (except He, which is $1s^2$).

• The inner core electronic configuration influences physical and chemical properties.

• Maximum oxidation state = total number of valence electrons (s + p electrons).

• p-block elements can also show other oxidation states, typically differing by 2 from the group oxidation state.

• In B, C, and N families, the group oxidation state is most stable for lighter elements.

• Oxidation states 2 units less than the group oxidation state become more stable for heavier elements ('inert pair effect').

• Non-metals and metalloids are found only in the p-block.

• Non-metallic character decreases down the group; the heaviest element is most metallic.

• Non-metals have higher ionization enthalpies and electronegativities than metals.

• Non-metal oxides are acidic or neutral, while metal oxides are basic.

• The first member of a p-block group differs due to size and the absence of d-orbitals.

• Second-period elements (B, C, N, O, F) are restricted to a maximum covalence of 4 (using 2s and 2p orbitals).

• Third-period elements can expand their covalence above 4 using vacant 3d orbitals.

  • Example: Boron forms only $[BF<em>4]^-$, while aluminium forms $[AlF</em>6]^{3-}$.

• Heavier elements can form $p\pi - d\pi$ or $d\pi - d\pi$ bonds, but these are weaker than $p\pi - p\pi$ bonds.

• Coordination number can be higher for heavier elements.

  • Example: $NO<em>3^-$ (three-coordinate) and $PO</em>4^{3-}$ (four-coordinate).

Group 13 Elements: The Boron Family

• Elements: Boron (B), Aluminium (Al), Gallium (Ga), Indium (In), Thallium (Tl), and Nihonium (Nh).

• Boron is a non-metal, aluminium is a metal with similarities to boron, and Ga, In, Tl are exclusively metallic.

• Boron sources: orthoboric acid ($H<em>3BO</em>3$), borax ($Na<em>2B</em>4O<em>7 \cdot 10H</em>2O$), kernite ($Na<em>2B</em>4O<em>7 \cdot 4H</em>2O$).

• Boron isotopes: $\^{10}B$ (19%) and $\^{11}B$ (81%).

• Aluminium is the most abundant metal in the earth's crust (8.3% by mass).

• Aluminium minerals: Bauxite ($Al<em>2O</em>3 \cdot 2H<em>2O$) and cryolite ($Na</em>3AlF_6$).

• Ga, In, and Tl are less abundant.

• Nihonium (Nh): synthetic, radioactive, atomic number 113, symbol Nh, atomic mass 286 g/mol. Electronic configuration [Rn] 5f14 6d10 7s 2 7p 2

Electronic Configuration

• Outer electronic configuration: $ns^2 np^1$.

• B and Al have a noble gas core.all

• Ga and In have a noble gas plus 10 d-electrons.

• Tl has a noble gas plus 14 f-electrons plus 10 d-electrons.

Atomic Radii

• Atomic radius generally increases down the group.

• Ga has a smaller atomic radius than Al due to poor screening by 3d electrons.

Ionization Enthalpy

• Ionization enthalpy generally decreases down the group, but not smoothly.

• Discontinuities between Al and Ga and In and Tl are due to the poor shielding effect of d- and f-electrons.

• Order of ionization enthalpies: \DeltaiH1 < \DeltaiH2 < \DeltaiH3.

• The sum of the first three ionization enthalpies is very high.

Electronegativity

• Electronegativity first decreases from B to Al, then increases slightly.

• This trend is related to atomic size discrepancies.

Physical Properties

• Boron is a hard, black, non-metallic solid with many allotropic forms and a high melting point.

• Other members are soft metals with low melting points and high electrical conductivity.

• Gallium has an unusually low melting point (303 K) and a high boiling point (2676 K), making it useful for measuring high temperatures.

• Density increases down the group from B to Tl.

Chemical Properties

• Oxidation state and trends in chemical reactivity:

  • Boron forms only covalent compounds due to its small size and high ionization enthalpy.

  • Aluminium can form $Al^{3+}$ ions and is highly electropositive.

  • The stability of the +1 oxidation state increases down the group: Al < Ga < In < Tl (inert pair effect).

  • Thallium exhibits a predominant +1 oxidation state, while +3 is highly oxidizing.

  • Compounds in the +1 oxidation state are more ionic than those in +3.

  • Trivalent compounds are electron-deficient Lewis acids (e.g., $BF_3$).

• Lewis acidity decreases down the group.

• Standard electrode potential values suggest that aluminium readily forms $Al^{3+}$ ions, while $Tl^{3+}$ is unstable and a strong oxidizing agent.

Reactivity towards air

• Boron is unreactive in crystalline form.

• Aluminium forms a protective oxide layer.

• Amorphous boron and aluminium react with oxygen to form $B<em>2O</em>3$ and $Al<em>2O</em>3$, respectively.

• They form nitrides with dinitrogen at high temperatures: $2E(s) + N_2(g) \rightarrow 2EN(s)$.

• $B<em>2O</em>3$ is acidic, $Al<em>2O</em>3$ and $Ga<em>2O</em>3$ are amphoteric, and $In<em>2O</em>3$ and $Tl<em>2O</em>3$ are basic.

Reactivity towards acids and alkalies

• Boron does not react with acids and alkalies.

• Aluminium dissolves in mineral acids and aqueous alkalies, showing amphoteric character.

• $2Al(s) + 6HCl(aq) \rightarrow 2Al^{3+}(aq) + 6Cl^-(aq) + 3H_2(g)$.

• Concentrated nitric acid makes aluminium passive due to oxide layer formation.

• $2Al(s) + 2NaOH(aq) + 6H<em>2O(l) \rightarrow 2Na^+[Al(OH)</em>4]^-(aq) + 3H_2(g)$.

• Trichlorides hydrolyze in water to form tetrahedral $[M(OH)_4]^−$ species.

• Aluminium chloride forms octahedral $[Al(H<em>2O)</em>6]^{3+}$ ion in acidified aqueous solution.

Reactivity towards halogens

• These elements react with halogens to form trihalides (except $TlI<em>3$): $2E(s) + 3X</em>2(g) \rightarrow 2EX_3(s)$.

• Anhydrous aluminium chloride hydrolyzes with atmospheric moisture, releasing HCl gas.

Important Trends and Anomalous Properties of Boron

• Tri-chlorides, bromides, and iodides are hydrolyzed in water.

• Tetrahedral $[M(OH)<em>4]^−$ and octahedral $[M(H</em>2O)_6]^{3+}$ species exist in aqueous medium (except boron).

• Monomeric trihalides are strong Lewis acids.

  • $BF<em>3 + NH</em>3 \rightarrow F<em>3B \cdot NH</em>3$.

• Boron's maximum covalence is 4 due to the absence of d orbitals.

• Other metal halides dimerize through halogen bridging (e.g., $Al<em>2Cl</em>6$).

Some Important Compounds of Boron

• Borax, orthoboric acid, and diborane are important compounds.

Borax

• Formula: $Na<em>2B</em>4O<em>7 \cdot 10H</em>2O$ or $Na<em>2[B</em>4O<em>5(OH)</em>4] \cdot 8H_2O$.

• Dissolves in water to give an alkaline solution: $Na<em>2B</em>4O<em>7 + 7H</em>2O \rightarrow 2NaOH + 4H<em>3BO</em>3$.

• On heating, borax forms a transparent liquid that solidifies into a glass-like borax bead.

• $Na<em>2B</em>4O<em>7 \cdot 10H</em>2O \xrightarrow{\Delta} Na<em>2B</em>4O<em>7 \xrightarrow{\Delta} 2NaBO</em>2 + B<em>2O</em>3$.

• Borax bead test identifies transition metals based on the color of metaborates.

Orthoboric acid

• Formula: $H<em>3BO</em>3$.

• White crystalline solid with a soapy touch, sparingly soluble in cold water, highly soluble in hot water.

• Prepared by acidifying borax: $Na<em>2B</em>4O<em>7 + 2HCl + 5H</em>2O \rightarrow 2NaCl + 4B(OH)_3$.

• Formed by hydrolysis of boron compounds.

• Layer structure with planar $BO_3$ units joined by hydrogen bonds.

• Weak monobasic Lewis acid: $B(OH)<em>3 + 2HOH \rightarrow [B(OH)</em>4]^- + H_3O^+$.

• On heating, forms metaboric acid and boric oxide: $H<em>3BO</em>3 \xrightarrow{\Delta} HBO<em>2 \xrightarrow{\Delta} B</em>2O_3$.

Diborane, $B<em>2H</em>6$

• Prepared by treating boron trifluoride with $LiAlH<em>4$: $4BF</em>3 + 3LiAlH<em>4 \rightarrow 2B</em>2H<em>6 + 3LiF + 3AlF</em>3$.

• Prepared by oxidation of sodium borohydride with iodine: $2NaBH<em>4 + I</em>2 \rightarrow B<em>2H</em>6 + 2NaI + H_2$.

• Industrially produced by reacting $BF<em>3$ with sodium hydride: $2BF</em>3 + 6NaH \xrightarrow{450K} B<em>2H</em>6 + 6NaF$.

• Colorless, highly toxic gas with a boiling point of 180 K.

• Spontaneously flammable in air, releasing a large amount of energy: $B<em>2H</em>6 + 3O<em>2 \rightarrow B</em>2O<em>3 + 3H</em>2O; \Delta_cH^\Theta = -1976 kJ \cdot mol^{-1}$.

• Hydrolyzed by water to give boric acid: $B<em>2H</em>6(g) + 6H<em>2O(l) \rightarrow 2B(OH)</em>3(aq) + 6H_2(g)$.

• Undergoes cleavage reactions with Lewis bases (L) to form borane adducts: $B<em>2H</em>6 + 2NMe<em>3 \rightarrow 2BH</em>3 \cdot NMe_3$.

• With ammonia, forms $[BH<em>2(NH</em>3)<em>2]^+ [BH</em>4]^−$ initially, then borazine ($B<em>3N</em>3H_6$) on heating.

• Diborane structure: Four terminal H atoms and two B atoms in one plane, with two bridging H atoms above and below the plane.

• Terminal B-H bonds are regular 2-center-2-electron bonds; bridging (B-H-B) bonds are 3-center-2-electron bonds.

Uses of Boron and Aluminium and Their Compounds

• Boron fibers are used in bullet-proof vests and lightweight composite materials.

• Boron-10 isotope absorbs neutrons, so metal borides are used in nuclear industry.

• Borax and boric acid are used to manufacture heat-resistant glasses (Pyrex), glass wool, and fiberglass.

• Borax is used as a flux for soldering, in glazed coatings for earthenware, and in medicinal soaps.

• Aqueous solution of orthoboric acid is a mild antiseptic.

• Aluminium is a bright silvery-white metal with high tensile strength, electrical and thermal conductivity.

• Aluminium is used in industry and everyday life (packing, utensils, construction, airplanes).

• Aluminium forms alloys with Cu, Mn, Mg, Si, and Zn.

Group 14 Elements: The Carbon Family

• Elements: Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), Lead (Pb), and Flerovium (Fl).

• Carbon is the 17th most abundant element by mass in the earth's crust.

• Carbon is found as coal, graphite, diamond, metal carbonates, hydrocarbons, and carbon dioxide gas (0.03%).

• Carbon combines with H, O, Cl, and S to form diverse materials.

• Organic chemistry is devoted to carbon-containing compounds.

• Carbon is an essential constituent of all living organisms.

• Naturally occurring carbon isotopes: $\^{12}C$ and $\^{13}C$.

• Radioactive isotope: $\^{14}C$ (half-life 5770 years), used for radiocarbon dating.

• Silicon is the second most abundant element on earth (27.7% by mass), present as silica and silicates.

• Silicon is important in ceramics, glass, and cement.

• Germanium exists in traces.

• Tin occurs as cassiterite ($SnO_2$) and lead as galena (PbS).

• Flerovium (Fl) is a synthetic radioactive element with atomic number 114.

• Ultrapure germanium and silicon are used to make transistors and semiconductor devices.

Electronic Configuration

• Valence shell electronic configuration: $ns^2np^2$.

Covalent Radius

• Significant increase from C to Si, then a smaller increase from Si to Pb due to filled d and f orbitals.

Ionization Enthalpy

• First ionization enthalpy is higher than corresponding group 13 members.

• Ionization enthalpy decreases down the group.

• Small decrease from Si to Ge to Sn, and slight increase from Sn to Pb due to poor shielding and size increase.

Electronegativity

• Group 14 elements are slightly more electronegative than group 13 elements due to smaller size.

• Electronegativity values are almost the same from Si to Pb.

Physical Properties

• All members are solids.

• Carbon and silicon are non-metals, germanium is a metalloid, and tin and lead are soft metals.

• Melting and boiling points are much higher than corresponding group 13 elements.

Chemical Properties

• Oxidation states and trends in chemical reactivity:

  • Common oxidation states are +4 and +2. Carbon also exhibits negative oxidation states.

  • Compounds in +4 oxidation state are generally covalent due to high ionization enthalpies.

  • The tendency to show +2 oxidation state increases down the group: Ge < Sn < Pb (inert pair effect).

  • Carbon and silicon mostly show +4 oxidation state.

  • Germanium forms stable compounds in +4 state.

  • Tin forms compounds in both +2 and +4 states (Sn in +2 state is a reducing agent).

  • Lead compounds in +2 state are stable, and in +4 state are strong oxidizing agents.

• Tetravalent compounds have eight electrons around the central atom and are not typically electron acceptors or donors.

• Carbon cannot exceed covalence of 4, but other elements can due to d orbitals.

• Halides undergo hydrolysis and form complexes by accepting electron pairs.

  • Examples: $SiF<em>6^{2-}$, $[GeCl</em>6]^{2-}$, $[Sn(OH)_6]^{2-}$.

Reactivity towards oxygen

• All members form oxides when heated in oxygen: MO and $MO_2$.

• SiO only exists at high temperatures.

• Oxides in higher oxidation states are more acidic.

  • $CO<em>2$, $SiO</em>2$, and $GeO<em>2$ are acidic, while $SnO</em>2$ and $PbO_2$ are amphoteric.

• Among monoxides, CO is neutral, GeO is acidic, and SnO and PbO are amphoteric.

Reactivity towards water

• Carbon, silicon, and germanium are not affected by water.

• Tin decomposes steam to form dioxide and hydrogen gas: $Sn + 2H<em>2O \rightarrow SnO</em>2 + 2H_2$.

• Lead is unaffected due to a protective oxide film.

Reactivity towards halogens

• Form halides of formula $MX<em>2$ and $MX</em>4$.

• All except carbon react directly with halogens.

• Most $MX_4$ are covalent, undergo $sp^3$ hybridization, and are tetrahedral.

• Exceptions are $SnF<em>4$ and $PbF</em>4$, which are ionic.

• $PbI_4$ does not exist.

• Stability of dihalides increases down the group.

• $GeX<em>4$ is more stable than $GeX</em>2$, whereas $PbX<em>2$ is more stable than $PbX</em>4$.

• Except $CCl_4$, tetrachlorides are easily hydrolyzed due to the availability of d orbitals.
  Carbon forms $p\pi - p\pi$ multiple bonds (C=C, C≡C, C=O, C=S, C≡N).

• Heavier elements do not form $p\pi - p\pi$ bonds due to larger and more diffuse atomic orbitals.

• Carbon atoms link with one another through covalent bonds to form chains and rings (catenation).

• The order of catenation is C >> Si > Ge ≈ Sn.

• Lead does not show catenation.

Allotropes of Carbon

• Carbon exhibits crystalline (diamond, graphite) and amorphous allotropic forms.

• Fullerenes were discovered in 1985.

Diamond

• Crystalline lattice with each carbon atom undergoing $sp^3$ hybridization.

• Each C atom is linked to four other C atoms in a tetrahedral fashion.

• The C–C bond length is 154 pm.

• Rigid three-dimensional network of carbon atoms.

• Very hard substance due to extended covalent bonding.

• Used as an abrasive and in making dyes and tungsten filaments.

Graphite

• Layered structure with layers held by van der Waals forces (340 pm distance).

• Each layer is composed of planar hexagonal rings of carbon atoms.

• C–C bond length within the layer is 141.5 pm.

• Each carbon atom undergoes $sp^2$ hybridization and forms three sigma bonds.

• The fourth electron forms a π bond, delocalized over the sheet.

• Electrons are mobile, so graphite conducts electricity.

• Cleaves easily between layers, so it is soft and slippery.

• Used as a dry lubricant.

Fullerenes

• Made by heating graphite in an electric arc in the presence of inert gases.

• The main product is $C_{60}$ (Buckminsterfullerene).

• Cage-like molecules with smooth structure and no dangling bonds.

• $C_{60}$ has a soccer ball shape with twenty six-membered rings and twelve five-membered rings.

• All carbon atoms are equal and undergo $sp^2$ hybridization.

• Each forms three sigma bonds, and the remaining electron is delocalized.

Some Important Compounds of Carbon and Silicon

• Oxides of Carbon: carbon monoxide (CO) and carbon dioxide (CO2).

Carbon Monoxide

• Direct oxidation of C in limited oxygen yields carbon monoxide: $2C(s) + O_2(g) \rightarrow 2CO(g)$.

• Pure CO is prepared by dehydrating formic acid with concentrated $H<em>2SO</em>4$: $HCOOH \xrightarrow[conc.H<em>2SO</em>4]{373K} H_2O + CO$.

• Industrially prepared by passing steam over hot coke: $C(s) + H<em>2O(g) \xrightarrow{473-1273K} CO(g) + H</em>2(g)$.

• The mixture of CO and $H_2$ is water gas or synthesis gas.

• Using air instead of steam produces producer gas: $2C(s) + O<em>2(g) + 4N</em>2(g) \xrightarrow{1273K} 2CO(g) + 4N_2(g)$.

• Colorless, odorless, and almost water-insoluble gas.

• Powerful reducing agent, reduces metal oxides.

• One sigma and two π bonds between carbon and oxygen: :C≡O:.

• Acts as a donor and forms metal carbonyls.

• Poisonous due to its ability to form a stable complex with haemoglobin.

Carbon Dioxide

• Prepared by complete combustion of carbon: $C(s) + O<em>2(g) \rightarrow CO</em>2(g)$.

• In the laboratory, prepared by reacting dilute HCl on calcium carbonate: $CaCO<em>3(s) + 2HCl(aq) \rightarrow CaCl</em>2(aq) + CO<em>2(g) + H</em>2O(l)$.

• On a commercial scale, it is obtained by heating limestone.

• Colorless and odorless gas.

• Forms carbonic acid with water: $H<em>2CO</em>3(aq) + H<em>2O(l) \rightleftharpoons HCO</em>3^-(aq) + H_3O^+(aq)$.

• $HCO<em>3$-(aq) + $H</em>2O$(l) $\rightleftharpoons CO<em>3^{2-}$(aq) + $H</em>3O^+$.

• $H<em>2CO</em>3$/$HCO_3$ buffer system maintains blood pH.

• Combines with alkalies to form metal carbonates.

• Removed from the atmosphere by photosynthesis.

  • $6CO<em>2 + 12H</em>2O \xrightarrow[Chlorophyll]{h\nu} C<em>6H</em>{12}O<em>6 + 6O</em>2 + 6H_2O$ .

• Increased $CO_2$ leads to the greenhouse effect.

• Solid $CO_2$ (dry ice) is used as a refrigerant.

• Used to carbonate soft drinks and as a fire extinguisher.

• Carbon atom undergoes sp hybridization.

Silicon Dioxide, $SiO_2$

• 95% of the earth’s crust is made up of silica and silicates.

• Occurs in several forms: quartz, cristobalite, and tridymite.

• Covalent, three-dimensional network solid.

• Each silicon atom is tetrahedrally bonded to four oxygen atoms.

• Each oxygen atom is bonded to another silicon atom.

• Used in accurate clocks, radio, and television broadcasting.

• Silica gel is used as a drying agent, chromatographic support, and catalyst.

• Kieselghur is used in filtration plants.

Silicones

• Organosilicon polymers with $(R_2SiO)$ as a repeating unit.

• Prepared from alkyl or aryl substituted silicon chlorides, $R<em>nSiCl</em>{(4-n)}$ where R is alkyl or aryl group.

• Reaction of methyl chloride with silicon in the presence of copper yields $MeSiCl<em>3$, $Me</em>2SiCl<em>2$, $Me</em>3SiCl$.

• Hydrolysis of $(CH<em>3)</em>2SiCl_2$ yields straight chain polymers.

• Chain length is controlled by adding $(CH<em>3)</em>3SiCl$.

• Water-repelling due to non-polar alkyl groups.

• High thermal stability, dielectric strength, and resistance to oxidation and chemicals.

• Used as sealants, greases, electrical insulators, and for waterproofing fabrics.

• Biocompatible, so used in surgical and cosmetic plants.

Silicates

• Various silicate minerals exist: feldspar, zeolites, mica, and asbestos.

• The basic structural unit is $SiO_4^{4-}$, with silicon bonded to four oxygen atoms in a tetrahedral fashion.

• Units are joined via corners, sharing 1, 2, 3, or 4 oxygen atoms.

• Forms chain, ring, sheet, or three-dimensional structures.

• Negative charge is neutralized by metal ions.

• Man-made silicates: glass and cement.

Zeolites

• Aluminosilicates with a three-dimensional network structure.

• Cations such as $Na^+$, $K^+$, or $Ca^{2+}$ balance the negative charge.

• Used as catalysts in petrochemical industries (e.g., ZSM-5 converts alcohols to gasoline).

• Used as ion exchangers in softening hard water.