Comprehensive Study Guide to P-Block Elements (Groups 13 to 18)

General Introduction to P-Block Elements

The $p$-block of the periodic table encompasses elements belonging to Groups $13$ through $18$. In these elements, the valence electrons enter the outermost $p$-orbitals. Their general valence shell electronic configuration is $ns^2\,np^{1-6}$ (with the exception of Helium, which possesses a $1s^2$ configuration).

The maximum oxidation state exhibited by a $p$-block element equals the total number of valence electrons (the sum of the $s$- and $p$-electrons). This is commonly designated as the group oxidation state. Besides the group oxidation state, $p$-block elements also exhibit lower oxidation states that typically differ from the group oxidation state by units of two. The relative stability of these lower oxidation states increases down the group for the heavier members due to the inert pair effect (the reluctance of the valence $s$-electrons to participate in bonding owing to poor shielding by intervening $d$- and $f$-electrons).

Group 15 Elements: The Nitrogen Family

Group $15$ elements, historically known as the pnictogens (meaning "suffocation-forming"), include Nitrogen ($N$), Phosphorus ($P$), Arsenic ($As$), Antimony ($Sb$), Bismuth ($Bi$), and the synthetic element Moscovium ($Mc$).

Electronic Configuration and Atomic Numbers

The atomic numbers and electronic configurations follow the standard Aufbau buildup:

• Nitrogen ($N$): $Z = 7$; Electronic Configuration: $[He]\,2s^2\,2p^3$

• Phosphorus ($P$): $Z = 15$; Electronic Configuration: $[Ne]\,3s^2\,3p^3$

• Arsenic ($As$): $Z = 33$; Electronic Configuration: $[Ar]\,3d^{10}\,4s^2\,4p^3$

• Antimony ($Sb$): $Z = 51$; Electronic Configuration: $[Kr]\,4d^{10}\,5s^2\,5p^3$

• Bismuth ($Bi$): $Z = 83$; Electronic Configuration: $[Xe]\,4f^{14}\,5d^{10}\,6s^2\,6p^3$

• Moscovium ($Mc$): $Z = 115$

The general valence shell electronic configuration is $ns^2\,np^3$. The $s$-orbital is completely filled, and the $p$-orbitals are half-filled, conferring exceptional electronic stability.

Occurrence

• Nitrogen: Molecular nitrogen ($N_2$) comprises $78\%$ by volume of the atmosphere. In the earth's crust, it occurs as sodium nitrate ($NaNO_3$, Chile saltpetre) and potassium nitrate ($KNO_3$, Indian saltpetre). It is also found in the form of proteins in plants and animals.

• Phosphorus: Occurs primarily in minerals of the apatite family, $Ca_9(PO_4)_6 \cdot CaX_2$ (where $X = F, Cl,\text{ or } OH$), such as fluorapatite ($3Ca_3(PO_4)_2 \cdot CaF_2$ or $Ca_9(PO_4)_6 \cdot CaF_2$). It is an essential constituent of animal and plant matter, bones, living cells, and phosphoproteins present in milk and eggs.

• Arsenic, Antimony, and Bismuth: Occur primarily as sulphide minerals in the earth's crust.

Atomic and Physical Properties

Quantitative Data Table (NCERT Table 7.1)

Detailed Periodic Trends

1. Atomic and Covalent Radii: Covalent radius increases down the group ($N < P < As < Sb < Bi$). A substantial increase is observed from $N$ ($70\,pm$) to $P$ ($110\,pm$). From $As$ ($121\,pm$) to $Bi$ ($148\,pm$), only a small increase is noted due to the poor shielding effect of completely filled $d$- and $f$-orbitals.

2. Ionisation Enthalpy: Decreases down the group due to the increase in atomic size ($\Delta_i H_1: N > P > As > Sb > Bi$). Because of their stable half-filled $p^3$ configurations, Group $15$ elements exhibit higher first ionisation enthalpies than the corresponding Group $14$ and Group $16$ elements. Note an anomaly in higher ionisation steps: $\Delta_i H_2$ and $\Delta_i H_3$ of $Bi$ are slightly higher than those of $Sb$ due to poor shielding by $4f$-electrons.

3. Electronegativity: Decreases down the group with increasing atomic size ($N (3.0) > P (2.1) > As (2.0) > Sb (1.9) = Bi (1.9)$).

4. Boiling and Melting Points:

   • Boiling Point: Increases from Nitrogen to Antimony and then slightly decreases at Bismuth: $N < P < As < Bi < Sb$ (or decreasing: $Sb > Bi > As > P > N$). The decrease from $Sb$ to $Bi$ is attributed to the inert pair effect reducing bonding tendencies.

   • Melting Point: Increases up to Arsenic and then decreases down to Bismuth: $N < P < Bi < Sb < As$ (decreasing: $As > Sb > Bi > P > N$). Arsenic possesses a giant layered covalent lattice yielding a peak melting point of $1089\,K$.

5. Density: Increases monotonically down the group ($N < P < As < Sb < Bi$).

6. Physical Nature: Dinitrogen is a diatomic gas ($N_2$), whereas all other elements are polyatomic solids ($P_4, As_4, Sb_4$). Nitrogen and phosphorus are non-metals, arsenic and antimony are metalloids, and bismuth is a typical metal. Except for nitrogen, all members exhibit allotropy.

Chemical Properties and Oxidation States

• Common Oxidation States: $-3, +3, \text{ and } +5$.

• Stability of $-3$: The tendency to exhibit a $-3$ oxidation state decreases down the group due to increases in atomic size and metallic character. Bismuth hardly forms any compounds in the $-3$ state.

• Stability of $+5$ vs $+3$: The stability of the $+5$ oxidation state decreases down the group, while the stability of the $+3$ oxidation state increases down the group. This is caused by the inert pair effect, where the outermost $ns^2$ electrons remain unshared. For bismuth, $Bi(V)$ compounds are powerful oxidising agents; $BiF_5$ is the only stable pentahalide.

• Disproportionation Reactions:

  • For nitrogen, all oxidation states from $+1$ to $+4$ tend to disproportionate in acid solution:     $3HNO_2 \xrightarrow{H^+} HNO_3 + 2NO + H_2O$

  • For phosphorus, nearly all intermediate oxidation states ($+1$ to $+4$) disproportionate into $+5$ and $-3$ in acidic and alkaline media:     $4H_3PO_3 \xrightarrow{\Delta} 3H_3PO_4 + PH_3$

  • For heavier elements ($As, Sb, Bi$), intermediate oxidation states become increasingly stable toward disproportionation.

• Maximum Covalency: Nitrogen is restricted to a maximum covalency of $4$ ($1$ $s$ + $3$ $p$ orbitals) because it lacks vacant $d$-orbitals in its valence shell. Heavier elements can expand their covalency to $5$ or $6$ using vacant $d$-orbitals (e.g., $PF_6^-$).

Anomalous Properties of Nitrogen

Nitrogen differs from the rest of the group due to:

1. Small atomic size.

2. High electronegativity ($3.0$).

3. High ionisation enthalpy ($1402\,kJ\,mol^{-1}$).

4. Non-availability of $d$-orbitals in its valence shell.

Specific Consequences of Anomalous Nature

• Multiple Bonding ($p\pi-p\pi$): Nitrogen possesses the unique ability to form strong $p\pi-p\pi$ multiple bonds with itself ($N \equiv N$) and with other small, highly electronegative atoms ($C, O$). Heavier members cannot form $p\pi-p\pi$ bonds because their atomic orbitals are too large and diffuse to undergo effective lateral overlap.

• Bond Enthalpy of Dinitrogen: Dinitrogen exists as a diatomic molecule with a triple bond ($1\,\sigma$ and $2\,\pi$). Its bond dissociation enthalpy is exceptionally high ($941.4\,kJ\,mol^{-1}$), making it chemically inert at room temperature.

• Single Bond Enthalpies and Catenation: The single $N-N$ bond is significantly weaker than the single $P-P$ bond ($N-N: 160\,kJ\,mol^{-1}$ vs $P-P: 201\,kJ\,mol^{-1}$). This weakness arises from strong interelectronic repulsions among the non-bonding electron pairs situated close together due to the short $N-N$ bond length. Consequently, nitrogen has a much weaker catenation property than phosphorus.

• Lack of $d\pi-p\pi$ and $d\pi-d\pi$ Bonding: Nitrogen cannot form $d\pi-p\pi$ bonds, unlike phosphorus and heavier elements (e.g., $R_3P=O$ or $R_3P=CH_2$). Furthermore, phosphorus and arsenic form $d\pi-d\pi$ bonds with transition metals when compounds such as $P(C_2H_5)_3$ and $As(C_6H_5)_3$ act as ligands.

Group 15 Hydrides ($EH_3$)

All group 15 elements form volatile hydrides of the formula $EH_3$ ($NH_3, PH_3, AsH_3, SbH_3, BiH_3$):

• Basic Strength: $NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3$. Basicity depends on the electron pair availability on the central atom. As size increases, electron density disperses over a larger volume, diminishing basicity.

• Thermal Stability and Bond Dissociation Enthalpy: $NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3$. As the central atom size increases, the $E-H$ bond length increases, bond strength decreases, and thermal stability falls.

• Reducing Power: $NH_3 < PH_3 < AsH_3 < SbH_3 < BiH_3$. $BiH_3$ is the strongest reducing agent among the hydrides because its $Bi-H$ bond is the weakest and cleaves most readily. $NH_3$ is only a mild reducing agent.

• Bond Angles: $NH_3\,(107.8^\circ) > PH_3\,(93.6^\circ) > AsH_3\,(91.8^\circ) > SbH_3\,(91.3^\circ) > BiH_3\,(90^\circ)$.

• Boiling Point Anomaly: $PH_3 < AsH_3 < NH_3 < SbH_3 < BiH_3$. Ammonia ($NH_3$) has an abnormally high boiling point compared to $PH_3$ and $AsH_3$ due to extensive intermolecular hydrogen bonding.

Group 15 Oxides and Halides

• Oxides: Elements form trioxides ($E_2O_3$) and pentoxides ($E_2O_5$).

  • The oxide in the higher oxidation state is more acidic than that in the lower oxidation state: $E_2O_5 > E_2O_3$.

  • Acidic character decreases down the group:

  • $N_2O_3, P_2O_3$: Purely acidic.

  • $As_2O_3, Sb_2O_3$: Amphoteric.

  • $Bi_2O_3$: Predominantly basic.

• Halides: Elements form trihalides ($EX_3$) and pentahalides ($EX_5$).

  • According to Fajans' rules, higher oxidation state halides ($EX_5$) are more covalent than lower oxidation state halides ($EX_3$) because cations with a $+5$ state have higher polarising power.

  • Nitrogen cannot form pentahalides due to the absence of vacant $d$-orbitals.

  • All trihalides of Group $15$ elements (except $BiF_3$, which is predominantly ionic) are covalent. All trihalides of nitrogen are unstable except $NF_3$.

Group 16 Elements: The Oxygen Family (Chalcogens)

Group $16$ comprises Oxygen ($O$), Sulphur ($S$), Selenium ($Se$), Tellurium ($Te$), Polonium ($Po$), and Livermorium ($Lv$). They are termed chalcogens ("ore-forming") because copper and other metal ores occur predominantly as oxides and sulphides.

Electronic Configuration and Atomic Numbers

• Oxygen ($O$): $Z = 8$; Configuration: $[He]\,2s^2\,2p^4$

• Sulphur ($S$): $Z = 16$; Configuration: $[Ne]\,3s^2\,3p^4$

• Selenium ($Se$): $Z = 34$; Configuration: $[Ar]\,3d^{10}\,4s^2\,4p^4$

• Tellurium ($Te$): $Z = 52$; Configuration: $[Kr]\,4d^{10}\,5s^2\,5p^4$

• Polonium ($Po$): $Z = 84$; Configuration: $[Xe]\,4f^{14}\,5d^{10}\,6s^2\,6p^4$

• Livermorium ($Lv$): $Z = 116$

The general valence shell electronic configuration is $ns^2\,np^4$.

Occurrence

• Oxygen: Most abundant element on Earth ($46.6\%$ by mass of the earth's crust; dry air contains $20.946\%$ oxygen by volume).

• Sulphur: Abundance in the earth's crust is $0.03 - 0.1\%$.

  • Combined as sulphates: Gypsum ($CaSO_4 \cdot 2H_2O$), Epsom salt ($MgSO_4 \cdot 7H_2O$), Baryte ($BaSO_4$).

  • Combined as sulphides: Galena ($PbS$), Zinc blende ($ZnS$), Copper pyrites ($CuFeS_2$).

  • Traces of $H_2S$ occur in volcanic gases. Organic materials (eggs, proteins, garlic, onion, mustard, hair, wool) contain sulphur.

• Selenium & Tellurium: Found as metal selenides and tellurides in sulphide ores.

• Polonium: Occurs in nature as a radioactive decay product of thorium and uranium minerals; half-life of $^{210}Po$ is $13.8\text{ days}$.

Periodic Trends and Atomic Properties

1. Atomic and Covalent Radii: Increase regularly down the group: $O < S < Se < Te < Po$.

2. Ionisation Enthalpy: Decreases down the group: $O > S > Se > Te > Po$. Group $16$ elements have lower first ionisation enthalpies than corresponding Group $15$ elements because removing one electron yields a stable half-filled $p^3$ configuration.

3. Electronegativity: Decreases down the group ($O\,(3.50) > S\,(2.58) > Se\,(2.55) > Te\,(2.01) > Po\,(1.76)$). Oxygen is the second most electronegative element after fluorine.

4. Electron Gain Enthalpy ($\Delta_{eg} H$):

   • Experimental values: $S (-200) > Se (-195) > Te (-190) > Po (-174) > O (-141\,kJ\,mol^{-1})$.

   • Crucial Anomaly: Oxygen has a less negative electron gain enthalpy than sulphur. This is due to the very small size of the oxygen atom, which causes high interelectronic repulsions in the compact $2p$-subshell upon electron addition.

5. Melting and Boiling Points: Increase down the group from $O$ to $Te$. The large difference between the melting and boiling points of oxygen and sulphur is explained on the basis of their atomicity: oxygen exists as a discrete diatomic molecule ($O_2$), whereas sulphur exists as an octatomic polyatomic molecule ($S_8$).

6. Allotropy: All elements of Group $16$ exhibit allotropy.

   • Sulphur Allotropes: Rhombic ($\alpha$-sulphur) and Monoclinic ($\beta$-sulphur). The transition temperature is $369\,K$, below which $\alpha$ is stable and above which $\beta$ is stable. Both contain $S_8$ puckered crown rings.

   • At elevated temperatures ($\approx 1000\,K$), sulphur exists primarily as $S_2$ vapour, which is paramagnetic like $O_2$ because it possesses two unpaired electrons in antibonding $\pi^*$ orbitals.

Chemical Properties

• Oxidation States: Common states are $-2, +2, +4, +6$.

  • Oxygen shows $-2$ predominantly, but exhibits $+2$ in $OF_2$, $+1$ in $O_2F_2$, $-1$ in peroxides ($H_2O_2$), and $-1/2$ in superoxides ($KO_2$).

  • The stability of the $+6$ oxidation state decreases down the group, while the stability of the $+4$ oxidation state increases down the group due to the inert pair effect.

• Hydrides ($H_2E$):

  • Acidic Strength: $H_2O < H_2S < H_2Se < H_2Te$. (Due to a decrease in $E-H$ bond dissociation enthalpy).

  • Thermal Stability: $H_2O > H_2S > H_2Se > H_2Te$.

  • Reducing Power: All hydrides except water are reducing agents. Reducing strength increases down the group: $H_2S < H_2Se < H_2Te$.

  • Boiling Point: $H_2S < H_2Se < H_2Te < H_2O$. Water has an abnormally high boiling point due to extensive intermolecular hydrogen bonding.

  • Drago's Rule: For $H_2S, H_2Se,$ and $H_2Te$, the central atom has an electronegativity $\le 2.5$ and bonded to hydrogen; hybridization does not effectively take place. The bond angles remain close to $90^\circ$ ($H_2S \approx 92^\circ, H_2Se \approx 91^\circ, H_2Te \approx 90^\circ$), in contrast to $H_2O$ ($104.5^\circ$, $sp^3$).

• Oxides ($EO_2$ and $EO_3$):

  • $SO_2$ is a gas; $SeO_2$ and $TeO_2$ are crystalline solids.

  • Reducing/Oxidising property: $SO_2$ is reducing ($+4$ to $+6$), whereas $TeO_2$ is an oxidising agent ($+4$ to $0$ or $+2$).

  • Acidic strength: $EO_3 > EO_2$. Acidic nature decreases down the group ($SO_2 > SeO_2 > TeO_2$).

• Halides:

  • Hexahalides ($EX_6$): Only hexafluorides are stable; all are gaseous with octahedral geometry ($sp^3d^2$). $SF_6$ is exceptionally stable due to steric protection around sulphur.

  • Tetrahalides ($EX_4$): $SF_4$ is a gas, $SeF_4$ is a liquid, $TeF_4$ is a solid. They have a see-saw shape ($sp^3d$ hybridisation with one equatorial lone pair).

  • Monohalides: Exist as dimers ($S_2F_2, S_2Cl_2, S_2Br_2, Se_2Cl_2, Se_2Br_2$). Dimeric monohalides undergo disproportionation:     $2Se_2Cl_2 \rightarrow SeCl_4 + 3Se$

Group 17 Elements: The Halogens

Group $17$ contains Fluorine ($F$), Chlorine ($Cl$), Bromine ($Br$), Iodine ($I$), Astatine ($At$, radioactive), and Tennessine ($Ts$). The name halogens is derived from Greek words meaning "salt producers".

Electronic Configuration and Atomic Numbers

• Fluorine ($F$): $Z = 9$; Configuration: $[He]\,2s^2\,2p^5$

• Chlorine ($Cl$): $Z = 17$; Configuration: $[Ne]\,3s^2\,3p^5$

• Bromine ($Br$): $Z = 35$; Configuration: $[Ar]\,3d^{10}\,4s^2\,4p^5$

• Iodine ($I$): $Z = 53$; Configuration: $[Kr]\,4d^{10}\,5s^2\,5p^5$

• Astatine ($At$): $Z = 85$; Configuration: $[Xe]\,4f^{14}\,5d^{10}\,6s^2\,6p^5$

• Tennessine ($Ts$): $Z = 117$

General valence shell electronic configuration: $ns^2\,np^5$.

Occurrence

• Fluorine: Occurs mainly as insoluble fluorides: Fluorspar ($CaF_2$), Cryolite ($Na_3AlF_6$), and Fluoroapatite ($3Ca_3(PO_4)_2 \cdot CaF_2$). Present in small quantities in soil, river plants, bones, and animal teeth.

• Chlorine, Bromine, Iodine: Found in seawater as chlorides, bromides, and iodides ($2.5\%$ $NaCl$ by mass). Dry sea beds contain deposits of carnallite ($KCl \cdot MgCl_2 \cdot 6H_2O$). Marine life accumulates iodine; seaweeds contain up to $0.5\%$ iodine, and Chile saltpetre contains up to $0.2\%$ sodium iodate ($NaIO_3$).

Atomic and Physical Properties

1. Atomic and Ionic Radii: Halogens possess the smallest atomic radii in their respective periods. Radii increase down the group: $F < Cl < Br < I$.

2. Ionisation Enthalpy: Very high, decreasing down the group: $F > Cl > Br > I$.

3. Electronegativity: Highest in their respective periods: $F\,(4.0) > Cl\,(3.2) > Br\,(3.0) > I\,(2.7)$. Fluorine is the most electronegative element in the periodic table.

4. Electron Gain Enthalpy ($\Delta_{eg} H$):

   • Order: $Cl > F > Br > I$.

   • Chlorine has a more negative electron gain enthalpy than fluorine ($Cl: -349\,kJ\,mol^{-1}$ vs $F: -333\,kJ\,mol^{-1}$) because the small size of the fluorine atom results in strong interelectronic repulsions in the compact $2p$-subshell.

5. Bond Dissociation Enthalpy ($X_2$):

   • Order: $Cl_2 > Br_2 > F_2 > I_2$.

   • $F_2$ has an anomalously low bond dissociation enthalpy ($158.8\,kJ\,mol^{-1}$ compared to $Cl_2: 242.6\,kJ\,mol^{-1}$ and $Br_2: 192.8\,kJ\,mol^{-1}$) due to large lone pair-lone pair repulsions in the compact $F_2$ molecule.

6. Standard Reduction Potential and Oxidising Power:

   • Standard electrode potentials: $E^\circ(F_2/F^-) = +2.87\,V$, $E^\circ(Cl_2/Cl^-) = +1.36\,V$, $E^\circ(Br_2/Br^-) = +1.09\,V$, $E^\circ(I_2/I^-) = +0.54\,V$.

   • Oxidising power decreases down the group: $F_2 > Cl_2 > Br_2 > I_2$.

   • $F_2$ is the strongest oxidising agent and displaces all other halide ions from aqueous solution:      $F_2 + 2X^- \rightarrow 2F^- + X_2 \quad (X = Cl, Br, I)$      $Cl_2 + 2X^- \rightarrow 2Cl^- + X_2 \quad (X = Br, I)$      $Br_2 + 2I^- \rightarrow 2Br^- + I_2$

7. Physical State and Colour:

   • $F_2$: Yellow gas.

   • $Cl_2$: Greenish-yellow gas.

   • $Br_2$: Red-brown liquid.

   • $I_2$: Violet solid.

   • Colour arises from the absorption of visible light, causing excitation of outer electrons to higher energy levels.

Chemical Properties

• Oxidation States:

  • Fluorine shows strictly a $-1$ oxidation state (and lacks variable oxidation states or disproportionation reactions due to the absence of $d$-orbitals and highest electronegativity).

  • Chlorine, bromine, and iodine show $-1, +1, +3, +5, \text{ and } +7$ oxidation states. Compounds in intermediate oxidation states undergo disproportionation (e.g., $3Cl_2 + 6OH^- \xrightarrow{\Delta} 5Cl^- + ClO_3^- + 3H_2O$).

• Hydrogen Halides ($HX$):

  • Acidic Strength & Reducing Power: $HF < HCl < HBr < HI$.

  • Thermal Stability: $HF > HCl > HBr > HI$.

  • Boiling Point: $HCl < HBr < HI < HF$. $HF$ has an exceptionally high boiling point due to strong intermolecular hydrogen bonding.

  • Melting Point: $HCl < HBr < HF < HI$.

• Oxides of Halogens:

  • Fluorine: Forms $OF_2$ and $O_2F_2$. $O_2F_2$ oxidises plutonium to $PuF_6$ in spent nuclear fuel processing.

  • Chlorine: Forms $Cl_2O, ClO_2, Cl_2O_6, Cl_2O_7$. Chlorine oxides tend to explode; $ClO_2$ is used as a bleaching agent for paper pulp, textiles, and water purification.

  • Iodine: $I_2O_5$ is a very good oxidising agent and is used in the quantitative estimation of carbon monoxide ($CO$).

  • Stability Order of Oxides: $I > Cl > Br$.

Interhalogen Compounds

Interhalogens are formed when two different halogens react. They have the general formula $XX'_n$ ($n = 1, 3, 5, 7$), where $X$ is the larger, less electronegative halogen.

Preparation Conditions

• $\text{Equal volume } Cl_2 + F_2 \xrightarrow{437\,K} 2ClF$

• $\text{Excess } F_2 + Cl_2 \xrightarrow{573\,K} 2ClF_3$

• $\text{Equimolar } I_2 + Cl_2 \rightarrow 2ICl$

• $\text{Excess } Cl_2 + I_2 \rightarrow 2ICl_3$

• $Br_2 + 3F_2 \,(\text{diluted with water}) \rightarrow 2BrF_3$

• $\text{Excess } F_2 + Br_2 \rightarrow 2BrF_5$

Hydrolysis and Shapes (NCERT Table 7.11)

• Hydrolysis Rule: Gives halide of the smaller halogen and an oxoacid of the larger halogen:   $XX' + H_2O \rightarrow HX' + HOX$   $BrF_5 + 3H_2O \rightarrow 5HF + HBrO_3 \quad (\text{bromic acid / halate})$

• Geometries:

  • $XX'_1$ ($ClF, BrF, ICl, IBr$): Linear.

  • $XX'_3$ ($ClF_3, BrF_3, IF_3, ICl_3$): Bent T-shaped ($sp^3d$ hybridisation with $2$ lone pairs).

  • $XX'_5$ ($ClF_5, BrF_5, IF_5$): Square pyramidal ($sp^3d^2$ hybridisation with $1$ lone pair). Note: $ClF_5$ is a colourless liquid at room temperature.

  • $XX'_7$ ($IF_7$): Pentagonal bipyramidal ($sp^3d^3$ hybridisation with $0$ lone pairs).

Group 18 Elements: The Noble Gases

Group $18$ comprises Helium ($He$), Neon ($Ne$), Argon ($Ar$), Krypton ($Kr$), Xenon ($Xe$), Radon ($Rn$, radioactive), and Oganesson ($Og$).

General Characteristics

• Electronic Configuration: $ns^2\,np^6$ (Helium is $1s^2$). Possess completely filled shells and high chemical inertness.

• Atomic Radii: Noble gases exhibit the largest atomic radii in their periods because their radii are expressed as van der Waals radii.

• Ionisation Enthalpy: Extremely high; decreases down the group from $He$ to $Rn$.

• Electron Gain Enthalpy: Have large positive values because added electrons must enter the next higher principal quantum shell.

• Melting and Boiling Points: Exceptionally low because monoatomic noble gas atoms are held only by weak dispersion forces. Helium has the lowest boiling point of any known substance ($4.2\,K$).

• Occurrence: All noble gases except radon occur in the atmosphere ($\approx 1\%$ by volume, of which argon is the major constituent). Radon is obtained as a radioactive decay product of Radium:   $^{226}_{88}Ra \rightarrow ^{222}_{86}Rn + ^4_2He$

Bartlett's Discovery of Xenon Compounds

Neil Bartlett observed that platinum hexafluoride ($PtF_6$) reacts with molecular oxygen ($O_2$) to form an ionic solid, dioxygenyl hexafluoroplatinate(V): $O_2 + PtF_6 \rightarrow O_2^+ [PtF_6]^-$

Recognizing that the first ionisation enthalpy of molecular oxygen ($1175\,kJ\,mol^{-1}$) is almost identical to that of xenon ($1170\,kJ\,mol^{-1}$), Bartlett reacted xenon with $PtF_6$, successfully isolating the first noble gas compound: $Xe^+ [PtF_6]^-$ (a red solid).

Xenon Fluorides and Oxides

Synthesis Conditions

• $Xe + F_2 \,(\text{excess } Xe) \xrightarrow{673\,K,\,1\,bar} XeF_2(s)$

• $Xe + 2F_2 \,(1:5\text{ ratio}) \xrightarrow{873\,K,\,7\,bar} XeF_4(s)$

• $Xe + 3F_2 \,(1:20\text{ ratio}) \xrightarrow{573\,K,\,60-70\,bar} XeF_6(s)$

• $XeF_4 + O_2F_2 \xrightarrow{143\,K} XeF_6 + O_2$

Hydrolysis Reactions

• $XeF_2$: Undergoes redox reaction with water:   $2XeF_2(s) + 2H_2O(l) \rightarrow 2Xe(g) + 4HF(aq) + O_2(g)$

• $XeF_4$: Disproportionates in water:   $6XeF_4 + 12H_2O \rightarrow 4Xe + 2XeO_3 + 24HF + 3O_2$

• Complete Hydrolysis of $XeF_6$: Forms xenon trioxide ($XeO_3$):   $XeF_6 + 3H_2O \rightarrow XeO_3 + 6HF$

• Partial Hydrolysis of $XeF_6$:   $XeF_6 + H_2O \rightarrow XeOF_4 + 2HF$   $XeF_6 + 2H_2O \rightarrow XeO_2F_2 + 4HF$

Fluoride Ion Donor and Acceptor Reactions

• Xenon fluorides react with fluoride ion acceptors to form cationic species:   $XeF_2 + PF_5 \rightarrow [XeF]^+ [PF_6]^-$   $XeF_4 + SbF_5 \rightarrow [XeF_3]^+ [SbF_6]^-$

• $XeF_6$ acts as a fluoride acceptor with alkali metal fluorides to form fluoroxenate anions:   $XeF_6 + MF \rightarrow M^+ [XeF_7]^- \quad (M = Na, K, Rb, Cs)$

Geometry and Structures of Xenon Compounds