Comprehensive Notes on p-Block Elements (Groups 15, 16, 17, and 18)

General Introduction to p-Block Elements

• The p-block elements are located in Groups 13 to 18 of the periodic table.

• General Electronic Configuration: $ns^2 np^{1-6}$ (Except Helium, which has a $1s^2$ configuration).

• Influencing Factors: Properties are influenced by atomic sizes, ionisation enthalpy, electron gain enthalpy, and electronegativity.

• The d-orbital Factor: The absence of d-orbitals in the second period and the presence of d (or d and f) orbitals in heavier elements (third period onwards) significantly impact properties.

• Diversity: p-block contains metals, metalloids, and non-metals.

Group 15 Elements: The Nitrogen Family

• Members: Nitrogen ($N$), Phosphorus ($P$), Arsenic ($As$), Antimony ($Sb$), and Bismuth ($Bi$).

• Occurence:

  • Nitrogen: 78% of the atmosphere by volume. Found as sodium nitrate ($NaNO_3$, Chile saltpetre) and potassium nitrate ($KNO_3$, Indian saltpetre). Essential in proteins, plants, and animals.

  • Phosphorus: Found in the apatite family, $Ca_9(PO_4)_6 \cdot CaX_2$ where $X = F, Cl, OH$ (e.g., fluorapatite $Ca_9(PO_4)_6 \cdot CaF_2$). Essential in bones, living cells, milk, and eggs (phosphoproteins).

  • Others: Arsenic, Antimony, and Bismuth occur mainly as sulphide minerals.

• Electronic Configuration: Valence shell is $ns^2 np^3$. The p-orbitals are half-filled, providing extra stability.

• Atomic and Ionic Radii: Increase down the group. From $N$ to $P$, there is a large increase; from $As$ to $Bi$, the increase is small due to the presence of filled d and/or f orbitals in heavier members.

• Ionisation Enthalpy: Decreases down the group as atomic size increases. Group 15 enthalpies are much higher than Group 14 due to half-filled stable configurations. Order: \Delta_i H_1 < \Delta_i H_2 < \Delta_i H_3.

• Electronegativity: Decreases down the group; the difference is less pronounced in heavier elements.

• Physical Properties:

  • All group 15 elements are polyatomic.

  • $N_2$ is a diatomic gas; others are solids.

  • Metallic character increases down the group ($N, P$ non-metals; $As, Sb$ metalloids; $Bi$ metal).

  • Boiling points increase top to bottom; melting points increase up to $As$ then decrease to $Bi$.

  • All elements except Nitrogen show allotropy.

Chemical Properties and Reactivity Trends (Group 15)

• Common Oxidation States: $-3, +3, +5$.

  • $-3$ state tendency decreases down the group due to increased size and metallic character.

  • $+5$ state stability decreases down the group. $BiF_5$ is the only well-characterized $Bi(V)$ compound.

  • $+3$ state stability increases down the group due to the inert pair effect.

  • Nitrogen also shows $+1, +2, +4$ with oxygen. Phosphorus shows $+1, +4$ in some oxoacids.

• Disproportionation: In acid solution, Nitrogen states $+1$ to $+4$ tend to disproportionate. Example: $3HNO_2 \rightarrow HNO_3 + H_2O + 2NO$. Phosphorus intermediate states disproportionate into $+5$ and $-3$ in alkali and acid.

• Covalency: Nitrogen is restricted to a maximum covalency of 4 ($s$ and three $p$ orbitals). Heavier elements expand covalency using vacant d-orbitals (e.g., $PF_6^-$).

• Anomalous Properties of Nitrogen:

  • Due to small size, high electronegativity, high ionisation enthalpy, and non-availability of d-orbitals.

  • Forms $p\pi-p\pi$ multiple bonds (e.g., $N \equiv N$). Heavier elements cannot do this effectively because orbitals are too large/diffuse.

  • Bond enthalpy of $N \equiv N$ is very high ($941.4\,kJ\,mol^{-1}$).

  • The single $N-N$ bond is weaker than the single $P-P$ bond due to high interelectronic repulsion of non-bonding electrons (small bond length).

Group 15 Reactivity Towards Other Elements

• Reactivity towards Hydrogen: Forms hydrides ($EH_3$). Stability decreases from $NH_3$ to $BiH_3$ (measured by bond dissociation enthalpy). Reducing character increases (Ammonia is mild; $BiH_3$ is strongest). Basicity decreases: NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3.

• Reactivity towards Oxygen: Forms $E_2O_3$ and $E_2O_5$. Oxides in higher oxidation states are more acidic. Acidic character decreases down the group. $N, P$ oxides are acidic; $As, Sb$ amphotenic; $Bi$ basic.

• Reactivity towards Halogens: Forms $EX_3$ and $EX_5$. Nitrogen does not form pentahalides. $EX_5$ are more covalent than $EX_3$.

• Reactivity towards Metals: Forms binary compounds with $-3$ oxidation state (e.g., $Ca_3N_2$ , $Ca_3P_2$, $Na_3As_2$, $Zn_3Sb_2$, $Mg_3Bi_2$).

Compounds of Nitrogen

• Dinitrogen ($N_2$):

  • Preparation: Commercially by liquefaction and fractional distillation of air ($N_2$ distils at $77.2\,K$). In lab: $NH_4Cl(aq) + NaNO_2(aq) \rightarrow N_2(g) + 2H_2O(l) + NaCl(aq)$. Pure nitrogen: thermal decomposition of sodium or barium azide ($Ba(N_3)_2 \rightarrow Ba + 3N_2$).

  • Properties: Colorless, odorless, tasteless, non-toxic. Inert at room temperature ($N \equiv N$ bond). Combines with Li ($Li_3N$) and Mg ($Mg_3N_2$) at high heat.

• Ammonia ($NH_3$):

  • Preparation: Haber's Process ($N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$; $\Delta_f H^0 = -46.1\,kJ\,mol^{-1}$). Optimum: $200\,atm$, $700\,K$, iron oxide catalyst with $K_2O$ and $Al_2O_3$.

  • Properties: Pungent odor. Trigonal pyramidal structure. Highly soluble in water, forming weakly basic solution ($OH^-$). Acts as a Lewis base (donates lone pair to metals like $Cu^{2+}$ and $Ag^+$).

• Oxides of Nitrogen: See Table 7.3 for specifics. Note: $NO_2$ dimerises because it has an odd number of electrons, forming stable $N_2O_4$.

• Nitric Acid ($HNO_3$):

  • Preparation: Ostwald's Process: (1) Catalytic oxidation of $NH_3$ to $NO$. (2) $NO$ to $NO_2$. (3) $NO_2$ in water to $HNO_3$. Concentrated to 68% by distillation and 98% with sulphuric acid.

  • Properties: Strong acid and powerful oxidiser. Attacks metals (except gold/platinum). Reaction with Copper: Dilute ($8HNO_3$) yields $NO$; Conc. ($4HNO_3$) yields $NO_2$. Reaction with Zinc: Dilute yields $N_2O$; Conc. yields $NO_2$. Passive effect on $Cr, Al$.

  • Brown Ring Test: ${NO_3}^- + 3Fe^{2+} + 4H^+ \rightarrow NO + 3Fe^{3+} + 2H_2O$, then $[Fe(H_2O)_6]^{2+} + NO \rightarrow [Fe(H_2O)_5(NO)]^{2+}$ (Brown complex).

Phosphorus Allotropes and Compounds

• Allotropes:

  • White Phosphorus: Waxy, poisonous, translucent solid. Soluble in $CS_2$, glows in dark (chemiluminescence). Reactive due to angular strain ($60^\circ$ angle). Consists of discrete $P_4$ tetrahedra.

  • Red Phosphorus: Obtained by heating white P at $573\,K$ in inert atmosphere. Polymeric structure. Less reactive than white P.

  • Black Phosphorus: $\alpha$-black and $\beta$-black forms. Formed under high pressure.

• Phosphine ($PH_3$): Prepared by reacting calcium phosphide ($Ca_3P_2$) with water/HCl, or white P with $NaOH$. Rotten fish smell, highly poisonous. Spontaneous combustion used in Holme's signals.

• Phosphorus Halides:

  • $PCl_3$: Colorless oily liquid. Pyramidal shape ($sp^3$). Hydrolyses to $H_3PO_3$.

  • $PCl_5$: Yellowish white powder. Trigonal bipyramidal structure in gas phase (axial bonds longer than equatorial). Ionic solid in solid state: $[PCl_4]^+[PCl_6]^-$. Hydrolyses to $POCl_3$, then $H_3PO_4$.

• Oxoacids of Phosphorus: All contain at least one $P=O$ and one $P-OH$ bond. $P-H$ bonds impart reducing properties (e.g., $H_3PO_2$ has two $P-H$ bonds and is a strong reducer). Basicity is determined by ionisable $H$ atoms in $P-OH$ groups (e.g., $H_3PO_4$ is tribasic, $H_3PO_3$ is dibasic).

Group 16 Elements: The Chalcogens

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• Members: Oxygen ($O$), Sulphur ($S$), Selenium ($Se$), Tellurium ($Te$), and Polonium ($Po$).

• Occurence: Oxygen is 46.6% of Earth's crust by mass. Sulphur (0.03-0.1%) found as gypsum ($CaSO_4 \cdot 2H_2O$), epsom salt ($MgSO_4 \cdot 7H_2O$ ), and sulphides (galena $PbS$, zinc blende $ZnS$).

• Trends:

  • Electronic Configuration: $ns^2 np^4$.

  • Atomic Size: Increase down the group.

  • Ionisation Enthalpy: Decreases down the group. Lower than Group 15 because Group 15 has stable half-filled orbitals.

  • Electronegativity: Oxygen is second only to Fluorine in electronegativity.

• Physical Properties: Oxygen/Sulphur (non-metals), Selenium/Tellurium (metalloids), Polonium (radioactive metal). Oxygen is diatomic; Sulphur is polyatomic ($S_8$).

• Chemical Reactivity:

  • Oxidation States: $-2$ stability decreases. Oxygen is mostly $-2$ (but $+2$ in $OF_2$). Others show $+2, +4, +6$. $+6$ stability decreases and $+4$ increases down the group (inert pair effect).

  • Hydrides ($H_2E$): Acidic character increases (H_2O < H_2S < H_2Se < H_2Te) due to bond dissociation enthalpy decrease. Thermal stability decreases.

  • Halides: Hexahalides only stable as fluorides ($SF_6$ is remarkably stable). Tetrafluorides ($SF_4$ gas, $SeF_4$ liquid, $TeF_4$ solid) have see-saw geometry.

Oxygen and Ozone

• Dioxygen ($O_2$): Prepared by heating chlorates ($KClO_3$ with $MnO_2$) or oxides like $Ag_2O, HgO$. Paramagnetic molecule.

• Simple Oxides:

  • Acidic: Non-metal oxides ($SO_2, Cl_2O_7, CO_2$).

  • Basic: Metallic oxides ($Na_2O, CaO$).

  • Amphoteric: Dual behavior ($Al_2O_3$).

  • Neutral: $CO, NO, N_2O$.

• Ozone ($O_3$): Prepared by silent electrical discharge on oxygen. Powerful oxidising agent (liberates nascent oxygen: $O_3 \rightarrow O_2 + O$). Structure is angular with 117 degree angle and equal bond lengths (resonance hybrid).

Compounds of Sulphur

• Allotropes: Rhombic Sulphur ($\alpha$, stable at room temp) and Monoclinic Sulphur ($\beta$). Transition temperature is $369\,K$.

• Sulphur Dioxide ($SO_2$): Sharp pungent smell. Prepared by burning sulphur or roasting sulphide ores. Acts as a reducing agent when moist (decolourises acidified $KMnO_4$).

• Sulphuric Acid ($H_2SO_4$): Manufactured by Contact Process: (1) Burning Sulphur to $SO_2$. (2) Catalyst oxidation of $SO_2$ to $SO_3$ over $V_2O_5$ at $2\,bar$, $720\,K$. (3) Absorption in $H_2SO_4$ to form Oleum ($H_2S_2O_7$). (4) Dilution to acid. Characteristics: low volatility, strong acid, high water affinity, strong oxidiser.

Group 17 Elements: The Halogens

• Members: Fluorine ($F$), Chlorine ($Cl$), Bromine ($Br$), Iodine ($I$), Astatine ($At$).

• General Properties: Electronic configuration $ns^2 np^5$. Smallest atomic radii in periods. Highest negative electron gain enthalpy. Electronegativity decreases down group.

• Physical Properties: $F, Cl$ are gases, $Br$ liquid, $I$ solid. All are coloured ($F_2$ yellow, $Cl_2$ greenish-yellow, $Br_2$ red, $I_2$ violet). Bond dissociation enthalpy order: Cl-Cl > Br-Br > F-F > I-I.

• Chemical Properties:

  • Oxidation States: Fluorine is always $-1$. Others show $-1, +1, +3, +5, +7$.

  • Oxidising Power: Fluorine is the strongest oxidiser. Reacts with water to release Oxygen.

  • Anomalous behavior of Fluorine: Due to small size, highest electronegativity, low $F-F$ bond enthalpy, and no d-orbitals.

  • Hydrides (HX): Acidic strength: HF < HCl < HBr < HI (due to decreasing bond enthalpy).

Chlorine and Interhalogens

• Chlorine ($Cl_2$): Discovered by Scheele. Prepared by $MnO_2 + 4HCl$ or by Deacon's Process ($HCl + O_2$ over $CuCl_2$ catalyst at $723\,K$). Bleaching action is permanent and occurs via oxidation ($Cl_2 + H_2O \rightarrow 2HCl + O$).

• Hydrochloric Acid ($HCl$): Colorless gas. Dissolves in water to form a strong acid. Aqua Regia: 3 parts Conc. $HCl$ and 1 part Conc. $HNO_3$; dissolves Gold ($Au$) and Platinum ($Pt$).

• Interhalogens: Compounds between two different halogens ($XX', XX'_3, XX'_5, XX'_7$). More reactive than individual halogens (except $F_2$). Structures based on VSEPR: Bent-T ($XX'_3$), Square Pyramidal ($XX'_5$), Pentagonal Bipyramidal ($IF_7$).

Group 18 Elements: The Noble Gases

• Members: Helium ($He$), Neon ($Ne$), Argon ($Ar$), Krypton ($Kr$), Xenon ($Xe$), Radon ($Rn$).

• General Properties: Closed shell configuration ($ns^2 np^6$). Monoatomic. Very low boiling points (Helium's is $4.2\,K$, lowest of any substance).

• Reactivity: Bartlett prepared first compound ($Xe^+ PtF_6^-$) in 1962. Xenon reacts with Fluorine and Oxygen.

• Xenon Compounds:

  • Fluorides: $XeF_2$ (linear), $XeF_4$ (square planar), $XeF_6$ (distorted octahedral).

  • Oxides: Partial/complete hydrolysis of fluorides yields $XeOF_4$ (square pyramidal) and $XeO_3$ (pyramidal).

• Uses:

  • Helium: Filling balloons, cryogenic agent, diving apparatus (low blood solubility).

  • Neon: Discharge tubes for advertisements.

  • Argon: Inert atmosphere for welding, filling bulbs.

Questions & Discussion

• Example 7.1: Why does nitrogen not form pentahalides? Response: It lacks d-orbitals to expand covalency beyond four.

• Example 7.2: Why does $PH_3$ have a lower boiling point than $NH_3$? Response: Unlike Ammonia, Phosphine does not form hydrogen bonds.

• Intext Request 7.1: Why are pentahalides more covalent than trihalides? Response: Central atom in higher oxidation state has higher polarising power.

• Example 7.5: Why does $NO_2$ dimerise? Response: It has an odd number of electrons; dimerisation creates the stable $N_2O_4$ with even electrons.

• Example 7.9: How does $H_3PO_2$ act as a reducer? Response: It possesses two $P-H$ bonds which impart reducing character.

• Example 7.22: Does the hydrolysis of $XeF_6$ lead to a redox reaction? Response: No, oxidation states of all elements remain unchanged in resultant products ($XeOF_4, XeO_2F_2$).