Group 15 Elements

The p-block elements are located in Groups 13 to 18 of the periodic table. Their general valence shell electronic configuration is ns$^2$np$^{1-6}$ (except Helium with 1s$^2$). Properties are influenced by atomic size, ionisation enthalpy, electron gain enthalpy, electronegativity, and the presence/absence of d or f orbitals.

Group 15 includes Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi), and Moscovium (Mc).

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Occurrence

• Nitrogen: Most abundant gas in atmosphere ($\sim 78\%$ by volume). In Earth's crust, found as sodium nitrate ($\textsf{NaNO}_3$, Chile saltpetre) and potassium nitrate ($\textsf{KNO}_3$, Indian saltpetre). Present in proteins.
• Phosphorus: Found in apatite family minerals ($\textsf{Ca}_9\text{(PO}_4)_6.\textsf{CaX}_2$, X = F, Cl, or OH). Essential in bones, living cells, and phosphoproteins (milk, eggs).
• Arsenic, Antimony, Bismuth: Found mainly as sulfide minerals.
• Moscovium (Mc): Synthetic, radioactive element. Atomic number 115. Chemistry not fully established due to short half-life.

Table 7.1 summarises properties of N, P, As, Sb, Bi.

Property	N	P	As	Sb	Bi
Atomic number	7	15	33	51	83
Atomic mass/g mol$^{-1}$	14.01	30.97	74.92	121.75	208.98
Electronic configuration	[He]2s$^2$2p$^3$	[Ne]3s$^2$3p$^3$	[Ar]3d$^{10}$4s$^2$4p$^3$	[Kr]4d$^{10}$5s$^2$5p$^3$	[Xe]4f$^{14}$5d$^{10}$6s$^2$6p$^3$
Ionisation enthalpy ( $\Delta_i H$/(kJ mol$^{-1}$)	I	1402	1012	947	834	703
	II	2856	1903	1798	1595	1610
	III	4577	2910	2736	2443	2466
Electronegativity	3.0	2.1	2.0	1.9	1.9
Covalent radius/pma	70	110	121	141	148
Ionic radius/pm (E$^{3-}$)	171	212	222
Ionic radius/pm (E$^{3+}$)				76	103
Melting point/K	63*	317d	1089e	904	544
Boiling point/K	77.2*	554d	888f	1860	1837
Density/[g cm$^{-3}$(298 K)]	0.879g	1.823	5.778h	6.697	9.808

(a E$^{III}$ single bond (E = element); b E$^{3–}$; c E$^{3+}$; d White phosphorus; e Grey $\alpha$-form at 38.6 atm; f Sublimation temperature; g At 63 K; h Grey $\alpha$-form; * Molecular N$_2$).

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Electronic Configuration

General valence shell configuration: ns$^2$np$^3$. The completely filled s-orbital and half-filled p-orbitals contribute to extra stability of the electronic configuration.

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Atomic And Ionic Radii

Covalent and ionic radii increase down the group due to increased number of electron shells. Significant increase from N to P. Smaller increase from As to Bi due to presence of filled d and f orbitals in heavier elements.

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Ionisation Enthalpy

Ionisation enthalpy decreases down the group due to increasing atomic size. Group 15 elements have higher ionisation enthalpy values compared to Group 14 elements in the same period due to their extra stable half-filled p-orbitals. Order of successive ionisation enthalpies: $\Delta_i H_1 < \Delta_i H_2 < \Delta_i H_3$.

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Electronegativity

Electronegativity generally decreases down the group with increasing atomic size. Difference is less pronounced among heavier elements.

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Physical Properties

All elements are polyatomic. Nitrogen is a diatomic gas (N$_2$). Others are solids. Metallic character increases down the group (N, P are non-metals; As, Sb are metalloids; Bi is a metal), due to decreasing ionisation enthalpy and increasing atomic size. Boiling points generally increase down the group. Melting point increases up to As, then decreases to Bi. All elements except N show allotropy.

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Chemical Properties

Common oxidation states: –3, +3, +5. Stability of –3 state decreases down the group (Bi rarely shows –3). Stability of +5 state decreases down the group (due to inert pair effect), while +3 state stability increases (Bi(V) is rare). Nitrogen also shows +1, +2, +4 with oxygen. Nitrogen's maximum covalency is 4 (limited by s and p orbitals). Heavier elements can expand covalency beyond 4 due to vacant d orbitals (e.g., PF$_6^{-}$).

Anomalous properties of nitrogen:

• Small size, high electronegativity, high ionisation enthalpy, absence of d orbitals.
• Forms p$\pi$-p$\pi$ multiple bonds with itself (N$\equiv$N triple bond, very high bond enthalpy) and small electronegative atoms (C, O). Heavier elements do not form effective p$\pi$-p$\pi$ bonds.
• N-N single bond is weaker than P-P single bond due to interelectronic repulsion in small N atom. Catenation is weaker in N.
• Cannot form dp-p$\pi$ bonds unlike heavier elements (e.g., R$_3$P=O).

(i) Reactivity towards hydrogen: Form hydrides of type $\textsf{EH}_3$ ($\textsf{E} = \textsf{N}, \textsf{P}, \textsf{As}, \textsf{Sb}, \textsf{Bi}$). Properties of hydrides (Table 7.2):

Property	NH$_3$	PH$_3$	AsH$_3$	SbH$_3$	BiH$_3$
Melting point/K	195.2	139.5	156.7	185	–
Boiling point/K	238.5	185.5	210.6	254.6	290
(E–H) Distance/pm	101.7	141.9	151.9	170.7	–
HEH angle ($^\circ$)	107.8	93.6	91.8	91.3	–
$\Delta_f H$/kJ mol$^{-1}$	–46.1	13.4	66.4	145.1	278
$\Delta_{diss}H$(E–H)/kJ mol$^{-1}$	389	322	297	255	–

• Stability decreases from $\textsf{NH}_3$ to $\textsf{BiH}_3$ (decreasing bond dissociation enthalpy).
• Reducing character increases from $\textsf{NH}_3$ (mild) to $\textsf{BiH}_3$ (strongest).
• Basicity decreases from $\textsf{NH}_3$ to $\textsf{BiH}_3$.
• $\textsf{NH}_3$ has hydrogen bonding (due to high electronegativity and small size of N), giving higher m.p./b.p. than $\textsf{PH}_3$.

(ii) Reactivity towards oxygen: Form $\textsf{E}_2\text{O}_3$ and $\textsf{E}_2\text{O}_5$ oxides. Acidity decreases down the group. $\textsf{N}_2\text{O}_3$, $\textsf{P}_2\text{O}_3$ are acidic; $\textsf{As}_2\text{O}_3$, $\textsf{Sb}_2\text{O}_3$ are amphoteric; $\textsf{Bi}_2\text{O}_3$ is basic. Higher oxidation state oxide is more acidic than lower one.

(iii) Reactivity towards halogens: Form $\textsf{EX}_3$ and $\textsf{EX}_5$. N forms only $\textsf{NF}_3$ (no d orbitals). $\textsf{EX}_5$ are more covalent than $\textsf{EX}_3$ (due to higher +5 ox. state). Trihalides except $\textsf{NF}_3$ and $\textsf{BiF}_3$ are covalent and stable. $\textsf{BiF}_3$ is ionic.

(iv) Reactivity towards metals: React to form binary compounds with –3 oxidation state (e.g., $\textsf{Ca}_3\text{N}_2$, $\textsf{Ca}_3\text{P}_2$).

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Dinitrogen

Dinitrogen ($\textsf{N}_2$) is the elemental form of nitrogen, a diatomic gas with a triple bond ($\textsf{N}\equiv\textsf{N}$).

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Preparation

• Commercial production: By liquefaction and fractional distillation of air. Liquid $\textsf{N}_2$ (b.p. 77.2 K) distils first, separating from liquid $\textsf{O}_2$ (b.p. 90 K).
• Laboratory preparation: By treating aqueous ammonium chloride ($\textsf{NH}_4\text{Cl}$) with sodium nitrite ($\textsf{NaNO}_2$).
  • $\textsf{NH}_4\text{Cl(aq)} + \textsf{NaNO}_2\text{(aq)} \to \textsf{N}_2\text{(g)} + 2\textsf{H}_2\text{O(l)} + \textsf{NaCl(aq)}$
  Impurities (NO, $\textsf{HNO}_3$) can be removed. Also by thermal decomposition of ammonium dichromate ($\textsf{(NH}_4)_2\text{Cr}_2\text{O}_7$).
  • $\textsf{(NH}_4)_2\text{Cr}_2\text{O}_7 \xrightarrow{\text{Heat}} \textsf{N}_2\text{(g)} + 4\textsf{H}_2\text{O(l)} + \textsf{Cr}_2\text{O}_3\text{(s)}$
• Very pure nitrogen: By thermal decomposition of sodium or barium azide ($\textsf{Ba(N}_3)_2$).
  • $\textsf{Ba(N}_3)_2\text{(s)} \to \textsf{Ba(s)} + 3\textsf{N}_2\text{(g)}$

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Properties

• Colourless, odourless, tasteless, non-toxic gas.
• Low solubility in water. Low freezing and boiling points.
• Rather inert at room temperature due to high bond enthalpy of $\textsf{N}\equiv\textsf{N}$ triple bond (941.4 kJ mol$^{-1}$).
• Reactivity increases at higher temperatures. Combines directly with some metals (e.g., Li, Mg) to form nitrides (ionic). Combines with non-metals to form covalent nitrides.
  • $6\textsf{Li(s)} + \textsf{N}_2\text{(g)} \xrightarrow{\text{Heat}} 2\textsf{Li}_3\text{N(s)}$
  • $3\textsf{Mg(s)} + \textsf{N}_2\text{(g)} \xrightarrow{\text{Heat}} \textsf{Mg}_3\text{N}_2\text{(s)}$
• Combines with $\textsf{H}_2$ at high temperature (Haber's Process) to form $\textsf{NH}_3$.
• Combines with $\textsf{O}_2$ at very high temperature ($\sim 2000$ K) to form NO.

Uses: Manufacture of ammonia and other nitrogen compounds. Provides inert atmosphere (iron, steel industry, chemical reactions). Liquid $\textsf{N}_2$ as refrigerant (preserving biological materials, food, cryosurgery).

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Answer:

Ammonia

Ammonia ($\textsf{NH}_3$) is an important compound of nitrogen.

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Preparation

• Small scale: From ammonium salts by reaction with caustic soda ($\textsf{NaOH}$) or calcium hydroxide ($\textsf{Ca(OH)}_2$).
  • $2\textsf{NH}_4\text{Cl(aq)} + \textsf{Ca(OH)}_2\text{(aq)} \to 2\textsf{NH}_3\text{(g)} + 2\textsf{H}_2\text{O(l)} + \textsf{CaCl}_2\text{(aq)}$
  • $\textsf{(NH}_4)_2\text{SO}_4\text{(aq)} + 2\textsf{NaOH(aq)} \to 2\textsf{NH}_3\text{(g)} + 2\textsf{H}_2\text{O(l)} + \textsf{Na}_2\text{SO}_4\text{(aq)}$
• Large scale: By Haber’s process (synthesis from $\textsf{N}_2$ and $\textsf{H}_2$).
  • $\textsf{N}_2\text{(g)} + 3\textsf{H}_2\text{(g)} \rightleftharpoons 2\textsf{NH}_3\text{(g)}; \Delta_f H^0 = -46.1$ kJ mol$^{-1}$
  This is an exothermic and reversible reaction. Optimum conditions for high yield (according to Le Chatelier's principle): High pressure ($\sim 200$ atm), moderate temperature ($\sim 700$ K), and catalyst (iron oxide with $\textsf{K}_2\text{O}$ and $\textsf{Al}_2\text{O}_3$).

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Properties

• Colourless gas with pungent odour. Freezing point 198.4 K, boiling point 239.7 K.
• Associated through hydrogen bonding in solid and liquid states, resulting in higher m.p./b.p. than expected.
• Molecule is trigonal pyramidal with N at apex (sp$^3$ hybridisation, 3 bond pairs, 1 lone pair).
• Highly soluble in water, forms weakly basic aqueous solution ($\textsf{NH}_3 + \textsf{H}_2\text{O} \rightleftharpoons \textsf{NH}_4^{+} + \textsf{OH}^{-}$).
• Forms ammonium salts with acids (e.g., $\textsf{NH}_4\text{Cl}$).
• Acts as a weak base, precipitates metal hydroxides from salt solutions (e.g., with $\textsf{ZnSO}_4$, $\textsf{FeCl}_3$).
• Acts as a Lewis base (donates lone pair on N) forming complex compounds with metal ions (e.g., $[\textsf{Cu(NH}_3)_4]^{2+}$ (deep blue) with $\textsf{Cu}^{2+}$, $[\textsf{Ag(NH}_3)_2]\textsf{Cl}$ (colourless) with $\textsf{AgCl}$ precipitate).

Uses: Manufacture of nitrogenous fertilisers ($\textsf{NH}_4\text{NO}_3$, urea), nitric acid. Liquid $\textsf{NH}_3$ as refrigerant.

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Oxides Of Nitrogen

Nitrogen forms a variety of oxides with different oxidation states (+1 to +5). Their names, formulas, methods of preparation, and properties are summarised in Table 7.3.

Oxidation state of nitrogen	Name	Formula	Common methods of preparation	Physical appearance and chemical nature
+1	Dinitrogen oxide [Nitrogen(I) oxide]	N$_2$O	NH$_4$NO$_3$ $\xrightarrow{\text{Heat}}$ N$_2$O + 2H$_2$O	colourless gas, neutral
+2	Nitrogen monoxide [Nitrogen(II) oxide]	NO	2NaNO$_2$ + 2FeSO$_4$ + 3H$_2$SO$_4$ $\to$ Fe$_2$(SO$_4$)$_3$ + 2NaHSO$_4$ + 2H$_2$O + 2NO	colourless gas, neutral
+3	Dinitrogen trioxide [Nitrogen(III) oxide]	N$_2$O$_3$	2NO + N$_2$O$_4$ $\xrightarrow{\text{250K}}$ 2N$_2$O$_3$	blue solid, acidic
+4	Nitrogen dioxide [Nitrogen(IV) oxide]	NO$_2$	2Pb(NO$_3$)$_2$ $\xrightarrow{\text{Heat}}$ 4NO$_2$ + 2PbO + O$_2$	brown gas, acidic
+4	Dinitrogen tetroxide [Nitrogen(IV) oxide]	N$_2$O$_4$	2NO$_2$ $\xrightarrow{\text{Cool}}$ N$_2$O$_4$ $\xrightarrow{\text{Heat}}$ 2NO$_2$	colourless solid/liquid, acidic
+5	Dinitrogen pentoxide [Nitrogen(V) oxide]	N$_2$O$_5$	4HNO$_3$ + P$_4$O$_{10}$ $\to$ 4HPO$_3$ + 2N$_2$O$_5$	colourless solid, acidic

Nitrogen oxides often exhibit resonance structures and multiple bonds. Their Lewis structures and bond parameters are provided (Table 7.4).

Name	Lewis dot structure	Main resonance structures	Bond parameters
Dinitrogen oxide			N-N bond length 112 pm N-O bond length 119 pm Linear shape
Nitrogen monoxide			N-O bond length 115 pm Linear shape
Dinitrogen trioxide			N-N bond length 186 pm N-O bond lengths 114, 121 pm Planar shape
Nitrogen dioxide			N-O bond length 120 pm Bond angle 134$^\circ$ Bent shape
Dinitrogen tetroxide			N-N bond length 175 pm N-O bond length 121 pm Planar shape
Dinitrogen pentoxide			N-O bond lengths 118, 132 pm Bond angles 114$^\circ$, 134$^\circ$ Non-planar shape

NO$_2$ contains an odd number of valence electrons (17). It can dimerise to form $\textsf{N}_2\text{O}_4$ which has an even number of valence electrons (34), making it more stable.

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Nitric Acid

Nitric acid ($\textsf{HNO}_3$) is the most important oxoacid of nitrogen.

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Preparation

• Laboratory: By heating $\textsf{KNO}_3$ or $\textsf{NaNO}_3$ with concentrated $\textsf{H}_2\text{SO}_4$.
  • $\textsf{NaNO}_3 + \textsf{H}_2\text{SO}_4 \xrightarrow{\text{Heat}} \textsf{NaHSO}_4 + \textsf{HNO}_3$
• Large scale: By Ostwald’s process (catalytic oxidation of $\textsf{NH}_3$).
  • Step 1: Catalytic oxidation of $\textsf{NH}_3$ by air $\textsf{O}_2$ over Pt/Rh gauge catalyst at 500 K and 9 bar.
    • $4\textsf{NH}_3\text{(g)} + 5\textsf{O}_2\text{(g)} \xrightarrow{\text{Pt/Rh gauge catalyst}} 4\textsf{NO(g)} + 6\textsf{H}_2\text{O(g)}$
  • Step 2: NO combines with $\textsf{O}_2$ to form $\textsf{NO}_2$.
    • $2\textsf{NO(g)} + \textsf{O}_2\text{(g)} \to 2\textsf{NO}_2\text{(g)}$
  • Step 3: $\textsf{NO}_2$ dissolves in water to give $\textsf{HNO}_3$. NO is recycled.
    • $3\textsf{NO}_2\text{(g)} + \textsf{H}_2\text{O(l)} \to 2\textsf{HNO}_3\text{(aq)} + \textsf{NO(g)}$
  Aqueous $\textsf{HNO}_3$ is concentrated by distillation ($\sim 68\%$) and further by dehydration with concentrated $\textsf{H}_2\text{SO}_4$ ($\sim 98\%$).

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Properties

• Colourless liquid. F.p. 231.4 K, b.p. 355.6 K. Laboratory grade is $\sim 68\%$ by mass.
• Gaseous $\textsf{HNO}_3$ is a planar molecule.
• Strong acid in aqueous solution ($\textsf{HNO}_3 + \textsf{H}_2\text{O} \to \textsf{H}_3\text{O}^{+} + \textsf{NO}_3^{-}$).
• Strong oxidising agent, attacks most metals (except noble metals like Au, Pt). Products depend on acid concentration, temperature, metal reactivity.
  • With dilute $\textsf{HNO}_3$ (metals like Cu, Zn): Produces NO ($\text{N}_2\text{O}$ with Zn).
  • With concentrated $\textsf{HNO}_3$ (metals like Cu, Zn): Produces $\textsf{NO}_2$.
  • Metals like Cr, Al become passive (form protective oxide film) in concentrated $\textsf{HNO}_3$.
• Oxidises non-metals and their compounds (I$_2$ to $\textsf{HIO}_3$, C to $\textsf{CO}_2$, S to $\textsf{H}_2\text{SO}_4$, P to $\textsf{H}_3\text{PO}_4$).

Brown Ring Test for nitrates: Based on reduction of nitrates to NO by $\textsf{Fe}^{2+}$, and NO reacting with $\textsf{Fe}^{2+}$ to form a brown complex $[\textsf{Fe(H}_2\text{O)}_5\text{(NO)}]^{2+}$.

Uses: Manufacture of fertilisers ($\textsf{NH}_4\text{NO}_3$), explosives (nitroglycerin, TNT), pigments, paints, dyes, etching metals, oxidiser in rocket fuels.

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Phosphorus — Allotropic Forms

Phosphorus exists in several allotropic forms, the most important being white, red, and black phosphorus.

• White phosphorus ($\textsf{P}_4$): Translucent, white, waxy solid. Poisonous. Insoluble in water, soluble in $\textsf{CS}_2$. Glows in dark (chemiluminescence). Reacts with boiling $\textsf{NaOH}$ solution in inert atmosphere to give phosphine ($\textsf{PH}_3$) and sodium hypophosphite ($\textsf{NaH}_2\text{PO}_2$). Less stable and more reactive due to angular strain (60$^\circ$) in the tetrahedral $\textsf{P}_4$ molecule. Readily ignites in air to form $\textsf{P}_4\text{O}_{10}$.
• Red phosphorus: Obtained by heating white P at 573 K in inert atmosphere. Iron grey lustre. Odourless, non-poisonous, insoluble in water and $\textsf{CS}_2$. Less reactive than white P. Does not glow in dark. Polymeric structure (chains of $\textsf{P}_4$ tetrahedra).
• Black phosphorus: More stable allotrope. Two forms: $\alpha$-black P (from red P at 803 K, opaque monoclinic/rhombohedral, sublimes, does not oxidise in air) and $\beta$-black P (from white P at 473 K under high pressure, does not burn in air up to 673 K).

At elevated temperatures ($\sim 1000$ K), sulphur exists as $\textsf{S}_2$ molecules which are paramagnetic like $\textsf{O}_2$ (due to unpaired electrons in $\pi^*$ antibonding orbitals).

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Phosphine

Phosphine ($\textsf{PH}_3$) is a hydride of phosphorus.

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Preparation

• By reaction of calcium phosphide ($\textsf{Ca}_3\text{P}_2$) with water or dilute $\textsf{HCl}$.
  • $\textsf{Ca}_3\text{P}_2\text{(s)} + 6\textsf{H}_2\text{O(l)} \to 3\textsf{Ca(OH)}_2\text{(aq)} + 2\textsf{PH}_3\text{(g)}$
  • $\textsf{Ca}_3\text{P}_2\text{(s)} + 6\textsf{HCl(aq)} \to 3\textsf{CaCl}_2\text{(aq)} + 2\textsf{PH}_3\text{(g)}$
• Laboratory preparation: By heating white phosphorus ($\textsf{P}_4$) with concentrated $\textsf{NaOH}$ solution in an inert atmosphere ($\textsf{CO}_2$). This is a disproportionation reaction (see Intext Q 7.8).
  • $\textsf{P}_4\text{(s)} + 3\textsf{NaOH(aq)} + 3\textsf{H}_2\text{O(l)} \to \textsf{PH}_3\text{(g)} + 3\textsf{NaH}_2\text{PO}_2\text{(aq)}$ (Sodium hypophosphite)

Pure phosphine is non-inflammable, but it becomes inflammable due to impurities like $\textsf{P}_2\text{H}_4$ or $\textsf{P}_4$ vapors. Purification involves absorbing it in HI to form phosphonium iodide ($\textsf{PH}_4\text{I}$), which on treatment with KOH liberates pure $\textsf{PH}_3$.

Uses: Spontaneous combustion used in Holme's signals (containers with $\textsf{CaC}_2$ and $\textsf{Ca}_3\text{P}_2$ thrown in sea produce $\textsf{C}_2\text{H}_2$ and $\textsf{PH}_3$, which ignite). Also used in smoke screens.

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Properties

• Colourless gas, highly poisonous, rotten fish smell.
• Explodes on contact with traces of oxidising agents ($\textsf{HNO}_3$, $\textsf{Cl}_2$, $\textsf{Br}_2$ vapours).
• Slightly soluble in water. Aqueous solution decomposes in light to red P and $\textsf{H}_2$.
• Reacts with metal salt solutions (e.g., $\textsf{CuSO}_4$, $\textsf{HgCl}_2$) to form corresponding phosphides.
• Weakly basic, forms phosphonium compounds with acids ($\textsf{PH}_3 + \textsf{HBr} \to \textsf{PH}_4\text{Br}$).

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Phosphorus Halides

Phosphorus forms two types of halides: $\textsf{PX}_3$ ($\textsf{X} = \textsf{F}, \textsf{Cl}, \textsf{Br}, \textsf{I}$) and $\textsf{PX}_5$ ($\textsf{X} = \textsf{F}, \textsf{Cl}, \textsf{Br}$).

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Phosphorus Trichloride

Phosphorus trichloride ($\textsf{PCl}_3$) is obtained by passing dry chlorine over heated white phosphorus or by action of thionyl chloride ($\textsf{SO}_2\text{Cl}_2$) on phosphorus.

• $\textsf{P}_4\text{(s)} + 6\textsf{Cl}_2\text{(g)} \to 4\textsf{PCl}_3\text{(l)}$
• $\textsf{P}_4\text{(s)} + 8\textsf{SO}_2\text{Cl}_2\text{(l)} \to 4\textsf{PCl}_3\text{(l)} + 4\textsf{SO}_2\text{(g)} + 2\textsf{S}_2\text{Cl}_2\text{(l)}$

Properties: Colourless oily liquid. Fumes in moist air due to hydrolysis to phosphorous acid and $\textsf{HCl}$ fumes.

• $\textsf{PCl}_3\text{(l)} + 3\textsf{H}_2\text{O(l)} \to \textsf{H}_3\text{PO}_3\text{(aq)} + 3\textsf{HCl(aq)}$

Reacts with compounds having –OH groups (alcohols, carboxylic acids) to replace –OH with –Cl.

• $3\textsf{CH}_3\text{COOH} + \textsf{PCl}_3 \to 3\textsf{CH}_3\text{COCl} + \textsf{H}_3\text{PO}_3$ (Acetyl chloride)
• $3\textsf{C}_2\text{H}_5\text{OH} + \textsf{PCl}_3 \to 3\textsf{C}_2\text{H}_5\text{Cl} + \textsf{H}_3\text{PO}_3$ (Chloroethane)

Structure: Pyramidal shape (sp$^3$ hybridisation, 3 bond pairs, 1 lone pair).

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Phosphorus Pentachloride

Phosphorus pentachloride ($\textsf{PCl}_5$) is prepared by the reaction of white phosphorus with excess dry chlorine or by action of $\textsf{SO}_2\text{Cl}_2$ on phosphorus.

• $\textsf{P}_4\text{(s)} + 10\textsf{Cl}_2\text{(g)} \to 4\textsf{PCl}_5\text{(s)}$
• $\textsf{P}_4\text{(s)} + 10\textsf{SO}_2\text{Cl}_2\text{(l)} \to 4\textsf{PCl}_5\text{(s)} + 10\textsf{SO}_2\text{(g)}$

Properties: Yellowish white powder. Hydrolyses in moist air to $\textsf{POCl}_3$ and finally to phosphoric acid ($\textsf{H}_3\text{PO}_4$).

• $\textsf{PCl}_5\text{(s)} + \textsf{H}_2\text{O(l)} \to \textsf{POCl}_3\text{(l)} + 2\textsf{HCl(g)}$
• $\textsf{POCl}_3\text{(l)} + 3\textsf{H}_2\text{O(l)} \to \textsf{H}_3\text{PO}_4\text{(aq)} + 3\textsf{HCl(g)}$

Sublimes on heating, decomposes on stronger heating to $\textsf{PCl}_3$ and $\textsf{Cl}_2$.

• $\textsf{PCl}_5\text{(s)} \xrightarrow{\text{Heat}} \textsf{PCl}_3\text{(g)} + \textsf{Cl}_2\text{(g)}$

Reacts with compounds containing –OH group (alcohols, carboxylic acids) to convert them to chloro derivatives.

• $\textsf{C}_2\text{H}_5\text{OH} + \textsf{PCl}_5 \to \textsf{C}_2\text{H}_5\text{Cl} + \textsf{POCl}_3 + \textsf{HCl}$
• $\textsf{CH}_3\text{COOH} + \textsf{PCl}_5 \to \textsf{CH}_3\text{COCl} + \textsf{POCl}_3 + \textsf{HCl}$

Reacts with finely divided metals on heating to give chlorides.

Uses: Synthesis of organic chloro compounds.

Structure: In gaseous and liquid phases, it has a trigonal bipyramidal structure (sp$^3$d hybridisation). Three equatorial P-Cl bonds are equivalent, two axial P-Cl bonds are longer and weaker due to greater repulsion from equatorial bond pairs.

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Oxoacids Of Phosphorus

Phosphorus forms a variety of oxoacids, which contain phosphorus, oxygen, and hydrogen. The composition of these acids is often related by loss or gain of $\textsf{H}_2\text{O}$ or O atoms. Important oxoacids are listed in Table 7.5.

Name	Formula	Oxidation state of phosphorus	Characteristic bonds and their number	Preparation
Hypophosphorous acid (Phosphinic)	H$_3$PO$_2$	+1	One P–OH, Two P–H, One P=O	white P$_4$ + alkali
Orthophosphorous acid (Phosphonic)	H$_3$PO$_3$	+3	Two P–OH, One P–H, One P=O	P$_2$O$_3$ + H$_2$O
Pyrophosphorous acid	H$_4$P$_2$O$_5$	+3	Four P–OH, One P–O–P, Two P=O	PCl$_3$ + H$_3$PO$_3$
Hypophosphoric acid	H$_4$P$_2$O$_6$	+4	Four P–OH, Two P=O, One P–P	red P$_4$ + alkali
Orthophosphoric acid (Phosphoric acid)	H$_3$PO$_4$	+5	Three P–OH, One P=O	P$_4$O$_{10}$+H$_2$O
Pyrophosphoric acid	H$_4$P$_2$O$_7$	+5	Four P–OH, Two P=O, One P–O–P	heat phosphoric acid
Metaphosphoric*	(HPO$_3$)$_n$	+5	One P–OH, Two P=O, One P–O–P (in trimer (HPO$_3$)$_3$)	phosphorus acid + Br$_2$, heat in a sealed tube

* Exists in polymeric forms only. Characteristic bonds of (HPO$_3$)$_3$ have been given in the Table.

Key structural features of phosphorus oxoacids:

• Phosphorus is tetrahedrally surrounded by other atoms (O or H).
• All contain at least one $\textsf{P=O}$ bond and one $\textsf{P–OH}$ bond.
• Acids with P in lower oxidation states (less than +5) also contain $\textsf{P–P}$ (e.g., $\textsf{H}_4\text{P}_2\text{O}_6$) or $\textsf{P–H}$ (e.g., $\textsf{H}_3\text{PO}_2$) bonds, but not both.

Acids with P in +3 oxidation state disproportionate on heating to higher (+5) and lower (–3) oxidation states. Example: Orthophosphorous acid ($\textsf{H}_3\text{PO}_3$) disproportionates to orthophosphoric acid ($\textsf{H}_3\text{PO}_4$) and phosphine ($\textsf{PH}_3$).

$4\textsf{H}_3\text{PO}_3 \xrightarrow{\text{Heat}} 3\textsf{H}_3\text{PO}_4 + \textsf{PH}_3$

The acids with $\textsf{P–H}$ bonds (e.g., $\textsf{H}_3\text{PO}_2$, $\textsf{H}_3\text{PO}_3$) have strong reducing properties. $\textsf{H}_3\text{PO}_2$ is a strong reducing agent (two $\textsf{P–H}$ bonds). $\textsf{H}_3\text{PO}_3$ has one $\textsf{P–H}$ bond.

Only the hydrogen atoms attached to oxygen in $\textsf{P–OH}$ bonds are ionisable and contribute to basicity. $\textsf{H}$ atoms directly bonded to P are not acidic. $\textsf{H}_3\text{PO}_2$ is monobasic (one $\textsf{P–OH}$). $\textsf{H}_3\text{PO}_3$ is dibasic (two $\textsf{P–OH}$). $\textsf{H}_3\text{PO}_4$ is tribasic (three $\textsf{P–OH}$).

Answer:

Answer:

Group 16 Elements

Group 16 includes Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te), Polonium (Po), and Livermorium (Lv). These are collectively called chalcogens (meaning 'ore producers').

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Occurrence

• Oxygen: Most abundant element in Earth's crust ($\sim 46.6\%$ by mass). $\sim 20.946\%$ by volume in dry air.
• Sulphur: $\sim 0.03-0.1\%$ in Earth's crust. Exists as sulphates ($\textsf{CaSO}_4.\textsf{2H}_2\text{O}$ gypsum, $\textsf{MgSO}_4.\textsf{7H}_2\text{O}$ epsom salt) and sulfides ($\textsf{PbS}$ galena, $\textsf{ZnS}$ zinc blende). Traces as $\textsf{H}_2\text{S}$. Present in organic materials (eggs, proteins).
• Selenium, Tellurium: Found as metal selenides and tellurides in sulfide ores.
• Polonium: Radioactive, decay product of thorium and uranium.
• Livermorium (Lv): Synthetic, radioactive. Atomic number 116. Very short half-life, chemistry not well studied.

Table 7.6 summarises properties of O, S, Se, Te, Po.

Property	O	S	Se	Te	Po
Atomic number	8	16	34	52	84
Atomic mass/g mol$^{-1}$	16.00	32.06	78.96	127.60	210.00
Electronic configuration	[He]2s$^2$2p$^4$	[Ne]3s$^2$3p$^4$	[Ar]3d$^{10}$4s$^2$4p$^4$	[Kr]4d$^{10}$5s$^2$5p$^4$	[Xe]4f$^{14}$5d$^{10}$6s$^2$6p$^4$
Covalent radius/(pm)a	66	104	117	137	146
Ionic radius, E$^{2-}$/pm	140	184	198	221	230b
Electron gain enthalpy/ $\Delta_{eg}H$ kJ mol$^{-1}$	–141	–200	–195	–190	–174
Ionisation enthalpy ($\Delta_i H_1$)/kJ mol$^{-1}$	1314	1000	941	869	813
Electronegativity	3.50	2.58	2.55	2.01	1.76
Density /g cm$^{-3}$ (298 K)	1.32c	2.06d	4.19e	6.25	–
Melting point/K	55	393f	490	725	520
Boiling point/K	90	718	958	1260	1235
Oxidation states*	–2,–1,1,2	–2,2,4,6	–2,2,4,6	–2,2,4,6	2,4

(aSingle bond; bApproximate value; cAt the melting point; d Rhombic sulphur; eHexagonal grey; fMonoclinic form, 673 K. * Oxygen shows oxidation states of +2 and +1 in oxygen fluorides OF$_2$ and O$_2$F$_2$ respectively).

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Electronic Configuration

General configuration: ns$^2$np$^4$. Six valence electrons.

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Atomic And Ionic Radii

Increase down the group due to more shells. Oxygen is exceptionally small. Ionic radii (E$^{2-}$) also increase down the group.

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Ionisation Enthalpy

Decreases down the group (increasing size). Lower than corresponding Group 15 elements due to less stable np$^4$ configuration vs Group 15's half-filled np$^3$.

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Electron Gain Enthalpy

Negative electron gain enthalpy (tendency to accept electron). Oxygen's value is less negative than sulfur's (due to small size and interelectronic repulsions). Becomes less negative from S to Po.

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Electronegativity

Decreases down the group (increasing size). Oxygen is highly electronegative (next to Fluorine). Metallic character increases down the group (O, S non-metals; Se, Te metalloids; Po metal).

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Physical Properties

O, S are non-metals. Se, Te are metalloids. Po is a metal (radioactive, short-lived). All show allotropy. Melting and boiling points increase down the group. Large difference between O and S is due to atomicity (O$_2$ diatomic, S$_8$ polyatomic).

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Chemical Properties

Exhibits various oxidation states (Table 7.6). Stability of –2 state decreases down the group (Po rarely shows –2). Oxygen shows –2 (except in OF$_2$ (+2) and $\textsf{O}_2\textsf{F}_2$ (+1)). Other elements show +2, +4, +6 (common). Stability of +6 decreases, +4 increases down the group (inert pair effect). Bonding in +4, +6 states is primarily covalent.

Anomalous behaviour of oxygen: Small size, high electronegativity, absence of d orbitals. Forms strong hydrogen bonding in $\textsf{H}_2\textsf{O}$ (unlike $\textsf{H}_2\textsf{S}$). Covalency limited to 4 (rarely exceeds 2 in practice).

(i) Reactivity with hydrogen: Form hydrides $\textsf{H}_2\textsf{E}$ ($\textsf{E} = \textsf{O}, \textsf{S}, \textsf{Se}, \textsf{Te}, \textsf{Po}$). Properties in Table 7.7.

Property	H$_2$O	H$_2$S	H$_2$Se	H$_2$Te
m.p/K	273	188	208	222
b.p/K	373	213	232	269
H–E distance/pm	96	134	146	169
HEH angle ($^\circ$)	104	92	91	90
$\Delta_f H$/kJ mol$^{-1}$	–286	–20	73	100
$\Delta_{diss}H$(H–E)/kJ mol$^{-1}$	463	347	276	238
Dissociation constanta	1.8$\times$10$^{-16}$	1.3$\times$10$^{-7}$	1.3$\times$10$^{-4}$	2.3$\times$10$^{-3}$

• Acidic character increases $\textsf{H}_2\textsf{O}$ to $\textsf{H}_2\textsf{Te}$ (due to decreasing H-E bond enthalpy).
• Thermal stability decreases $\textsf{H}_2\textsf{O}$ to $\textsf{H}_2\textsf{Po}$ (decreasing bond enthalpy).
• Reducing property increases $\textsf{H}_2\textsf{S}$ to $\textsf{H}_2\textsf{Te}$ ($\textsf{H}_2\textsf{O}$ is not reducing).

(ii) Reactivity with oxygen: Form $\textsf{EO}_2$ and $\textsf{EO}_3$ (E = S, Se, Te, Po). Acidic character. $\textsf{SO}_2$ reducing, $\textsf{TeO}_2$ oxidising. Reducing property of dioxide decreases $\textsf{SO}_2$ to $\textsf{TeO}_2$.

(iii) Reactivity towards halogens: Form $\textsf{EX}_6$, $\textsf{EX}_4$, $\textsf{EX}_2$. Stability decreases $\textsf{F}^{-}> \textsf{Cl}^{-} > \textsf{Br}^{-} > \textsf{I}^{-}$. Hexafluorides are stable (octahedral, e.g., $\textsf{SF}_6$ exceptionally stable). Tetrafluorides (e.g., $\textsf{SF}_4$) have see-saw geometry (sp$^3$d, one lone pair). Dihalides (except O) are tetrahedral (sp$^3$). Monohalides are dimeric (e.g., $\textsf{S}_2\text{Cl}_2$).

Answer:

Answer:

Answer:

Dioxygen

Dioxygen ($\textsf{O}_2$) is the elemental form of oxygen, a diatomic molecule.

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Preparation

• Laboratory:
  • By heating oxygen-containing salts (chlorates, nitrates, permanganates) with catalysts (e.g., $\textsf{MnO}_2$).
    • $2\textsf{KClO}_3\text{(s)} \xrightarrow{\text{Heat, } \textsf{MnO}_2} 2\textsf{KCl(s)} + 3\textsf{O}_2\text{(g)}$
  • By thermal decomposition of oxides of less reactive metals or higher oxides of some metals (Ag$_2$O, HgO, Pb$_3$O$_4$, PbO$_2$).
  • By decomposition of hydrogen peroxide ($\textsf{H}_2\textsf{O}_2$) catalyzed by $\textsf{MnO}_2$ or finely divided metals.
    • $2\textsf{H}_2\textsf{O}_2\text{(aq)} \to 2\textsf{H}_2\text{O(l)} + \textsf{O}_2\text{(g)}$
• Large scale: Electrolysis of water (produces $\textsf{H}_2$ at cathode, $\textsf{O}_2$ at anode). Industrially, by liquefaction and fractional distillation of air.

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Properties

• Colourless, odourless gas. Slightly soluble in water (supports aquatic life). Liquefies at 90 K, freezes at 55 K.
• Molecular oxygen ($\textsf{O}_2$) is paramagnetic despite having an even number of electrons (explained by Molecular Orbital Theory).
• Directly reacts with nearly all metals and non-metals (except Au, Pt, noble gases). Reactions are often exothermic. Requires initiation (due to high bond dissociation enthalpy of $\textsf{O}=\textsf{O}$ double bond).
  • With metals: Forms oxides (e.g., $2\textsf{Ca} + \textsf{O}_2 \to 2\textsf{CaO}$, $4\textsf{Al} + 3\textsf{O}_2 \to 2\textsf{Al}_2\text{O}_3$).
  • With non-metals: Forms oxides (e.g., $\textsf{P}_4 + 5\textsf{O}_2 \to \textsf{P}_4\text{O}_{10}$, $\textsf{C} + \textsf{O}_2 \to \textsf{CO}_2$).
  • With compounds: Oxidises sulfides, hydrocarbons, etc. (e.g., $2\textsf{ZnS} + 3\textsf{O}_2 \to 2\textsf{ZnO} + 2\textsf{SO}_2$, $\textsf{CH}_4 + 2\textsf{O}_2 \to \textsf{CO}_2 + 2\textsf{H}_2\textsf{O}$).
• Catalytic oxidation of some compounds (e.g., $\textsf{SO}_2$ to $\textsf{SO}_3$ with $\textsf{V}_2\text{O}_5$).

Uses: Respiration, combustion, oxyacetylene welding, manufacture of steel, in hospitals, high altitude flying, mountaineering. In rocket fuels.

Answer:

Simple Oxides

A binary compound of oxygen with another element is called an oxide. Elements form various oxides.

Simple oxides (containing one element besides oxygen) are classified by their acid-base character:

• Acidic oxides: Combine with water to form acids (e.g., $\textsf{SO}_2$, $\textsf{CO}_2$, $\textsf{N}_2\text{O}_5$, $\textsf{Cl}_2\text{O}_7$). Non-metal oxides are typically acidic. Some metal oxides in high oxidation states are also acidic ($\textsf{Mn}_2\text{O}_7$, $\textsf{CrO}_3$).
• Basic oxides: Combine with water to form bases (e.g., $\textsf{Na}_2\text{O}$, $\textsf{CaO}$, $\textsf{BaO}$). Metallic oxides are generally basic.
• Amphoteric oxides: Exhibit both acidic and basic properties (react with both acids and alkalies). Example: $\textsf{Al}_2\text{O}_3$, $\textsf{ZnO}$, $\textsf{PbO}$.
• $\textsf{Al}_2\text{O}_3\text{(s)} + 6\textsf{HCl(aq)} \to 2\textsf{AlCl}_3\text{(aq)} + 3\textsf{H}_2\text{O(l)}$ (Basic nature)
• $\textsf{Al}_2\text{O}_3\text{(s)} + 6\textsf{NaOH(aq)} + 3\textsf{H}_2\text{O(l)} \to 2\textsf{Na}_3[\textsf{Al(OH)}_6]\text{(aq)}$ (Acidic nature, or $2\textsf{Na[Al(OH)}_4]$)
• Neutral oxides: Neither acidic nor basic. Examples: $\textsf{CO}$, $\textsf{NO}$, $\textsf{N}_2\text{O}$.

Mixed oxides (e.g., $\textsf{Pb}_3\text{O}_4$, $\textsf{Fe}_3\text{O}_4$) contain more than one oxidation state of the metal.

Ozone

Ozone ($\textsf{O}_3$) is an allotropic form of oxygen. It is formed in the stratosphere by UV radiation acting on atmospheric oxygen. The ozone layer protects Earth from harmful UV radiation.

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Preparation

By passing a slow dry stream of oxygen through a silent electrical discharge. This converts about 10% of $\textsf{O}_2$ to $\textsf{O}_3$ (ozonised oxygen). Silent discharge is used to prevent decomposition of $\textsf{O}_3$ (formation is endothermic).

$3\textsf{O}_2 \to 2\textsf{O}_3; \Delta H = +142$ kJ mol$^{-1}$.

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Properties

• Pure ozone: pale blue gas, dark blue liquid, violet-black solid. Characteristic smell. Harmless in small concentrations, but toxic at higher concentrations.
• Thermodynamically unstable with respect to oxygen ($\textsf{O}_3 \to \textsf{O}_2 + \text{O}$), due to negative $\Delta H$ and positive $\Delta S$, resulting in a large negative $\Delta G$. High concentrations are explosive.
• Powerful oxidising agent due to liberation of nascent oxygen (O).
  • Oxidises $\textsf{PbS}$ to $\textsf{PbSO}_4$: $\textsf{PbS(s)} + 4\textsf{O}_3\text{(g)} \to \textsf{PbSO}_4\text{(s)} + 4\textsf{O}_2\text{(g)}$
  • Oxidises iodide ions to iodine: $2\textsf{I}^{-}\text{(aq)} + \textsf{H}_2\text{O(l)} + \textsf{O}_3\text{(g)} \to 2\textsf{OH}^{-}\text{(aq)} + \textsf{I}_2\text{(s)} + \textsf{O}_2\text{(g)}$
• Quantitative estimation of $\textsf{O}_3$ using $\textsf{KI}$ solution and titration of liberated $\textsf{I}_2$.
• Molecule is angular (bond angle $\sim 117^\circ$). Resonance hybrid. Two O-O bond lengths are identical (128 pm).

Nitrogen oxides and freons (CFCs) contribute to ozone depletion in the stratosphere.

Uses: Germicide, disinfectant, sterilizing water, bleaching (oils, ivory, flour), oxidising agent (in $\textsf{KMnO}_4$ manufacture).

Answer:

Sulphur — Allotropic Forms

Sulphur exhibits numerous allotropic forms, with rhombic ($\alpha$-sulphur) and monoclinic ($\beta$-sulphur) being the most significant.

• Rhombic sulphur ($\alpha$-sulphur): Stable at room temperature (below 369 K). Yellow crystals. M.p. 385.8 K, specific gravity 2.06. Insoluble in water, soluble in $\textsf{CS}_2$.
• Monoclinic sulphur ($\beta$-sulphur): Stable above 369 K. M.p. 393 K, specific gravity 1.98. Soluble in $\textsf{CS}_2$. Forms needle-shaped crystals. Transforms back to $\alpha$-sulphur below 369 K.

The temperature 369 K is the transition temperature where both forms are stable in equilibrium.

Both forms consist of $\textsf{S}_8$ molecules, which are puckered rings with a crown shape. The packing of these $\textsf{S}_8$ rings differs in the two allotropes.

Other forms with different ring sizes ($\textsf{S}_6$ to $\textsf{S}_{20}$) exist. At elevated temperatures ($\sim 1000$ K), $\textsf{S}_2$ is the dominant species, which is paramagnetic like $\textsf{O}_2$.

Answer:

Sulphur Dioxide

Sulphur dioxide ($\textsf{SO}_2$) is a significant oxide of sulphur.

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Preparation

• By burning sulphur in air or oxygen (forms some $\textsf{SO}_3$ as well).
  • $\textsf{S(s)} + \textsf{O}_2\text{(g)} \to \textsf{SO}_2\text{(g)}$
• Laboratory: By treating a sulphite ($\textsf{SO}_3^{2-}$) with dilute $\textsf{H}_2\text{SO}_4$.
  • $\textsf{SO}_3^{2-}\text{(aq)} + 2\textsf{H}^{+}\text{(aq)} \to \textsf{H}_2\text{O(l)} + \textsf{SO}_2\text{(g)}$
• Industrially: As a by-product of roasting sulfide ores.
  • $4\textsf{FeS}_2\text{(s)} + 11\textsf{O}_2\text{(g)} \to 2\textsf{Fe}_2\text{O}_3\text{(s)} + 8\textsf{SO}_2\text{(g)}$

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Properties

• Colourless gas, pungent smell, highly soluble in water. Liquefies easily.
• Forms sulphurous acid with water ($\textsf{SO}_2 + \textsf{H}_2\text{O} \rightleftharpoons \textsf{H}_2\text{SO}_3\text{(aq)}$).
• Reacts with alkalies similar to $\textsf{CO}_2$ (forms sulfites and hydrogen sulfites).
• Acts as a reducing agent when moist. Converts $\textsf{Fe}^{3+}$ to $\textsf{Fe}^{2+}$, decolourises acidified $\textsf{KMnO}_4$ solution (test for $\textsf{SO}_2$).
• Oxidised to $\textsf{SO}_3$ by $\textsf{O}_2$ (with $\textsf{V}_2\text{O}_5$ catalyst), or by $\textsf{Cl}_2$ (with charcoal catalyst).
• Molecule is angular (bent shape). Resonance hybrid of two canonical forms. Both $\textsf{S–O}$ bonds are equal due to resonance.

Uses: Refining petroleum and sugar, bleaching wool and silk, anti-chlor, disinfectant, preservative. Manufacture of $\textsf{H}_2\text{SO}_4$, sodium hydrogen sulphite. Liquid $\textsf{SO}_2$ as solvent.

Answer:

Oxoacids Of Sulphur

Sulphur forms several oxoacids, many of which are unstable and known only in solution or as salts. Their structures involve sulphur atoms bonded to oxygen atoms (P=O), hydroxyl groups (P–OH), and sometimes other sulphur atoms (P–P) or peroxide linkages (O–O).

Structures of some important oxoacids are shown in Fig. 7.6.

All sulphur oxoacids contain at least one $\textsf{S=O}$ bond and one $\textsf{S–OH}$ bond. Acidity is due to ionisation of $\textsf{H}$ from $\textsf{S–OH}$ bonds. Number of $\textsf{S–OH}$ groups determines basicity.

Sulphuric Acid

Sulphuric acid ($\textsf{H}_2\text{SO}_4$) is one of the most important industrial chemicals.

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Manufacture

Manufactured by the Contact Process, involving three main steps:

1. Burning sulphur or sulfide ores to produce $\textsf{SO}_2$.
2. Catalytic oxidation of $\textsf{SO}_2$ to $\textsf{SO}_3$ using $\textsf{O}_2$ in the presence of $\textsf{V}_2\text{O}_5$ catalyst.
   • $2\textsf{SO}_2\text{(g)} + \textsf{O}_2\text{(g)} \xrightarrow{\text{V}_2\text{O}_5} 2\textsf{SO}_3\text{(g)}; \Delta_r H^0 = -196.6$ kJ mol$^{-1}$
   This step is exothermic and reversible. Favourable conditions (Le Chatelier's principle): low temperature and high pressure. Plant operates at $\sim 720$ K and 2 bar pressure (compromise for rate).
3. Absorption of $\textsf{SO}_3$ in concentrated $\textsf{H}_2\text{SO}_4$ to form Oleum ($\textsf{H}_2\text{S}_2\text{O}_7$).
   • $\textsf{SO}_3\text{(g)} + \textsf{H}_2\text{SO}_4\text{(aq)} \to \textsf{H}_2\text{S}_2\text{O}_7\text{(l)}$
   Oleum is then diluted with water to get $\textsf{H}_2\text{SO}_4$ of desired concentration.
   • $\textsf{H}_2\text{S}_2\text{O}_7\text{(l)} + \textsf{H}_2\text{O(l)} \to 2\textsf{H}_2\text{SO}_4\text{(aq)}$

$\textsf{SO}_2$ is purified before catalytic oxidation. Process is continuous to reduce cost.

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Properties

• Colourless, dense, oily liquid. F.p. 283 K, b.p. 611 K. Specific gravity 1.84 at 298 K.
• Dissolves in water with large heat evolution (exothermic). Add concentrated acid slowly to water with stirring.
• Properties: low volatility, strong acidic character, strong affinity for water (dehydrating agent), oxidising agent.
• Ionises in two steps in water ($\textsf{H}_2\text{SO}_4 + \textsf{H}_2\text{O} \to \textsf{H}_3\text{O}^{+} + \textsf{HSO}_4^{-}$ (very strong first ionisation), $\textsf{HSO}_4^{-} + \textsf{H}_2\text{O} \rightleftharpoons \textsf{H}_3\text{O}^{+} + \textsf{SO}_4^{2-}$ (weaker second ionisation)). Forms normal and acid sulphates.
• Low volatility allows making more volatile acids from their salts ($\textsf{MX} + \textsf{H}_2\text{SO}_4 \to \textsf{HX} + \textsf{MHSO}_4$ or $\textsf{M}_2\text{SO}_4$).
• Strong dehydrating agent: removes water from compounds (charring of carbohydrates).
• Moderately strong oxidising agent (intermediate between $\textsf{H}_3\text{PO}_4$ and $\textsf{HNO}_3$). Oxidises metals (Cu), non-metals (S, C) to their oxides, getting reduced to $\textsf{SO}_2$.

Uses: Manufacture of fertilisers, pigments, paints, dyes, detergents, petroleum refining, pickling metals, storage batteries, manufacture of nitrocellulose products, laboratory reagent.

Answer:

Answer:

Group 17 Elements

Group 17 includes Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At), and Tennessine (Ts). These are called halogens (salt producers). They are highly reactive non-metallic elements. Astatine and Tennessine are radioactive.

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Occurrence

• Fluorine: Mainly as insoluble fluorides ($\textsf{CaF}_2$ fluorspar, $\textsf{Na}_3\text{AlF}_6$ cryolite). Small amounts in soil, water, plants, bones, teeth.
• Chlorine, Bromine, Iodine: In seawater as chlorides, bromides, iodides of Na, K, Mg, Ca. $\textsf{NaCl}$ is main component ($\sim 2.5\%$). Dried up seas contain these salts (e.g., $\textsf{KCl.MgCl}_2.\textsf{6H}_2\text{O}$ carnallite). Seaweeds contain iodine. Chile saltpetre contains sodium iodate.
• Astatine (At) and Tennessine (Ts): Radioactive.

Table 7.8 provides properties of F, Cl, Br, I, At.

Property	F	Cl	Br	I	At$^\text{a}$
Atomic number	9	17	35	53	85
Atomic mass/g mol$^{-1}$	19.00	35.45	79.90	126.90	210
Electronic configuration	[He]2s$^2$2p$^5$	[Ne]3s$^2$3p$^5$	[Ar]3d$^{10}$4s$^2$4p$^5$	[Kr]4d$^{10}$5s$^2$5p$^5$	[Xe]4f$^{14}$5d$^{10}$6s$^2$6p$^5$
Covalent radius/pm	64	99	114	133	–
Ionic radius X$^{-}$/pm	133	184	196	220	–
Ionisation enthalpy/kJ mol$^{-1}$	1680	1256	1142	1008	–
Electron gain enthalpy/kJ mol$^{-1}$	–333	–349	–325	–296	–
Electronegativityb	4	3.2	3.0	2.7	2.2
$\Delta_{hyd}H$(X$^{-}$)/kJ mol$^{-1}$	515	381	347	305	–
	F$_2$	Cl$_2$	Br$_2$	I$_2$	–
Melting point/K	54.4	172.0	265.8	386.6	–
Boiling point/K	84.9	239.0	332.5	458.2	–
Density/g cm$^{-3}$	1.5 (85K)c	1.66 (203K)c	3.19(273K)c	4.94(293K)d	–
Distance X – X/pm	143	199	228	266	–
Bond dissociation enthalpy/(kJ mol$^{-1}$)	158.8	242.6	192.8	151.1	–
E$^\circ$/Ve	2.87	1.36	1.09	0.54	–

(a Radioactive; b Pauling scale; c For the liquid at temperatures (K) given in the parentheses; d solid; e The half-cell reaction is X$_2$(g) + 2e$^-$ $\to$ 2X$^-$(aq)).

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Electronic Configuration

General configuration: ns$^2$np$^5$. One electron short of noble gas configuration.

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Atomic And Ionic Radii

Smallest atomic radii in their periods (max effective nuclear charge). Radii increase down the group (more shells). Fluorine has exceptionally small atomic radius.

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Ionisation Enthalpy

Very high ionisation enthalpy (tendency to lose electron). Decreases down the group (increasing size).

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Electron Gain Enthalpy

Maximum negative electron gain enthalpy in their periods (readily accept electron to achieve noble gas configuration). Becomes less negative down the group. Fluorine's value is less negative than chlorine's (small size of F atom, interelectronic repulsions in 2p orbitals).

Answer:

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Answer:

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Electronegativity

Very high electronegativity, decreases down the group. Fluorine is the most electronegative element.

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Physical Properties

Gradation in properties: F$_2$ (yellow gas), Cl$_2$ (greenish yellow gas), Br$_2$ (red liquid), I$_2$ (violet solid). M.p. and b.p. increase with atomic number. All are coloured due to light absorption. Soluble in organic solvents. Anomalously low bond dissociation enthalpy of F$_2$ compared to Cl$_2$ is due to electron repulsion in small F atom.

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Chemical Properties

• Highly reactive, reactivity decreases down the group. React with metals and non-metals to form halides.
• Strong oxidising agents (readily accept electron). F$_2$ is the strongest. A halogen oxidises halide ions of higher atomic number.
• Reactions with water: F$_2$ oxidises water to $\textsf{O}_2$. Cl$_2$ and $\textsf{Br}_2$ form hydrohalic and hypohalous acids. $\textsf{I}_2$ reaction with water is non-spontaneous (reverse reaction can occur).

Anomalous behavior of fluorine: Small size, highest electronegativity, low F-F bond dissociation enthalpy, no d orbitals. Leads to higher ionization enthalpy, electronegativity, electrode potential, and lower ionic/covalent radii, m.p./b.p., bond dissociation enthalpy than expected. Forms only one oxoacid (HOF). HF is liquid due to strong hydrogen bonding (others are gases).

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Chlorine

Chlorine ($\textsf{Cl}_2$) is a greenish yellow gas.

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Preparation

• Laboratory: By heating $\textsf{MnO}_2$ with concentrated $\textsf{HCl}$ or a mixture of $\textsf{NaCl}$ and concentrated $\textsf{H}_2\text{SO}_4$ with $\textsf{MnO}_2$. Also by action of $\textsf{HCl}$ on $\textsf{KMnO}_4$.

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Properties

• Greenish yellow gas, pungent, suffocating odour. $\sim 2.5$ times heavier than air. Soluble in water.
• Reacts with metals and non-metals to form chlorides (e.g., Al, Na, Fe, P$_4$, S$_8$).
• High affinity for hydrogen, reacts with hydrogen-containing compounds to form $\textsf{HCl}$ (e.g., $\textsf{H}_2\textsf{S}$, hydrocarbons).
• With cold, dilute alkalies: forms chloride and hypochlorite ($\textsf{NaCl} + \textsf{NaOCl}$).
• With hot, concentrated alkalies: forms chloride and chlorate ($\textsf{NaCl} + \textsf{NaClO}_3$). This is a disproportionation reaction for chlorine (0 to –1 and +5).
• With dry slaked lime $\textsf{Ca(OH)}_2$: forms bleaching powder ($\textsf{Ca(OCl)}_2.\textsf{CaCl}_2.\textsf{Ca(OH)}_2.\textsf{2H}_2\text{O}$).
• With hydrocarbons: substitution in saturated, addition in unsaturated hydrocarbons.
• Chlorine water loses color on standing (forms $\textsf{HCl}$ and $\textsf{HOCl}$). $\textsf{HOCl}$ gives nascent oxygen, responsible for oxidising and bleaching properties.
• Powerful oxidising agent (due to nascent oxygen). Oxidises $\textsf{Fe}^{2+}$ to $\textsf{Fe}^{3+}$, sulphite to sulphate, $\textsf{SO}_2$ to $\textsf{H}_2\text{SO}_4$, $\textsf{I}_2$ to $\textsf{HIO}_3$.
• Powerful bleaching agent for vegetable/organic matter in presence of moisture (permanent bleaching).

Uses: Bleaching (woodpulp, cotton), extraction of gold/platinum, manufacture of dyes, drugs, organic compounds ($\textsf{CCl}_4$, $\textsf{CHCl}_3$, DDT, refrigerants), sterilizing water, preparing poisonous gases (phosgene, tear gas, mustard gas).

Answer:

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Manufacture Of Chlorine

• Deacon’s process: By oxidation of $\textsf{HCl(g)}$ with atmospheric $\textsf{O}_2$ over $\textsf{CuCl}_2$ catalyst at 723 K.
  • $4\textsf{HCl(g)} + \textsf{O}_2\text{(g)} \xrightarrow{\text{CuCl}_2} 2\textsf{Cl}_2\text{(g)} + 2\textsf{H}_2\text{O(l)}$
• Electrolytic process: By electrolysis of concentrated $\textsf{NaCl}$ solution (brine). $\textsf{Cl}_2$ is liberated at the anode. This is also a source of $\textsf{NaOH}$ and $\textsf{H}_2$.

Answer:

Hydrogen Chloride

Hydrogen chloride ($\textsf{HCl}$) is a pungent smelling gas, highly soluble in water forming hydrochloric acid.

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Preparation

• Laboratory: By heating $\textsf{NaCl}$ with concentrated $\textsf{H}_2\text{SO}_4$.
  • $\textsf{NaCl(s)} + \textsf{H}_2\text{SO}_4\text{(conc)} \xrightarrow{\text{420 K}} \textsf{NaHSO}_4\text{(s)} + \textsf{HCl(g)}$
  • $\textsf{NaHSO}_4\text{(s)} + \textsf{NaCl(s)} \xrightarrow{\text{823 K}} \textsf{Na}_2\text{SO}_4\text{(s)} + \textsf{HCl(g)}$ (Higher temp step)

$\textsf{HCl}$ gas can be dried by passing through concentrated $\textsf{H}_2\text{SO}_4$.

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Properties

• Colourless, pungent gas. Easily liquefied, freezes to white solid. Extremely soluble in water.
• Aqueous solution (hydrochloric acid) is a strong acid ($\textsf{HCl} + \textsf{H}_2\text{O} \to \textsf{H}_3\text{O}^{+} + \textsf{Cl}^{-}$).
• Reacts with $\textsf{NH}_3$ to give white fumes of $\textsf{NH}_4\text{Cl}$ (ammonium chloride).
• Forms aqua regia with concentrated $\textsf{HNO}_3$ (3 parts conc. $\textsf{HCl}$ : 1 part conc. $\textsf{HNO}_3$), used for dissolving noble metals (Au, Pt). Aqua regia contains $\textsf{HNO}_3$, $\textsf{HCl}$, and also NOCl and $\textsf{Cl}_2$. The noble metals react with the $\textsf{Cl}$ to form soluble complexes ($\textsf{AuCl}_4^{-}$, $\textsf{PtCl}_6^{2-}$).
• Decomposes salts of weaker acids (carbonates, sulfites).

Uses: Manufacture of $\textsf{Cl}_2$, $\textsf{NH}_4\text{Cl}$, glucose. Extracting glue, purifying bone black, medicine, laboratory reagent.

Answer:

Oxoacids Of Halogens

Halogens form several oxoacids, where the halogen is the central atom bonded to oxygen and hydrogen (as –OH groups). Most are unstable and exist only in solution or as salts. Fluorine forms only one oxoacid (HOF or HFO) due to its high electronegativity and small size.

Structures of oxoacids of halogens are based on VSEPR theory (Fig. 7.8).

Oxidation states of halogens (except F) in oxoacids range from +1 to +7.

Oxidation State	Halic (I) acid (Hypohalous acid)	Halic (III) acid (Halous acid)	Halic (V) acid (Halic acid)	Halic (VII) acid (Perhalic acid)
+1	HOF (Hypofluorous acid) HOCl (Hypochlorous acid) HOBr (Hypobromous acid) HOI (Hypoiodous acid)	–	–	–
+3	–	HOClO (Chlorous acid)	–	–
+5	–	–	HOClO$_2$ (Chloric acid) HOBrO$_2$ (Bromic acid) HOIO$_2$ (Iodic acid)	–
+7	–	–	–	HOClO$_3$ (Perchloric acid) HOBrO$_3$ (Perbromic acid) HOIO$_3$ (Periodic acid)

Acidic strength of oxoacids increases with increasing oxidation state of the halogen. For oxoacids of the same halogen, acidic strength increases with number of non-hydroxylated oxygen atoms.

Answer:

Interhalogen Compounds

Interhalogen compounds are formed when two different halogens react with each other. General formulas are $\textsf{XX}'$, $\textsf{XX}'_3$, $\textsf{XX}'_5$, and $\textsf{XX}'_7$, where X is the larger, more electropositive halogen, and X' is the smaller, more electronegative halogen. The number of X' atoms increases with the ratio of atomic radii between X and X'. Example: $\textsf{IF}_7$.

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Preparation

By direct combination of halogens or by reaction of a halogen with a lower interhalogen compound under specific conditions.

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Properties

• Covalent molecules, diamagnetic. Volatile solids or liquids at 298 K (except ClF is a gas).
• Physical properties are generally intermediate between the constituent halogens, but m.p./b.p. can be slightly higher than expected.
• More reactive than halogens (except $\textsf{F}_2$) because X–X' bond is weaker than X–X bond (except F–F bond).
• Hydrolyse readily to give a halide ion from the smaller halogen (X'-) and an oxoacid/polyatomic anion derived from the larger halogen (e.g., $\textsf{XX}' + \textsf{H}_2\text{O} \to \textsf{HX}' + \textsf{HOX}$).
• Structures are explained by VSEPR theory (Table 7.11): $\textsf{XX}'_3$ (bent T-shaped), $\textsf{XX}'_5$ (square pyramidal), $\textsf{IF}_7$ (pentagonal bipyramidal).

Type	Formula	Physical state and colour	Structure
XX′	ClF BrF IFa BrClb ICl IBr	colourless gas pale brown gas detected spectroscopically gas ruby red solid (α-form) brown red solid (β-form) black solid	Linear
XX′$_3$	ClF$_3$ BrF$_3$ IF$_3$ ICl$_3$c	colourless gas yellow green liquid yellow powder orange solid	Bent T-shaped
XX′$_5$	IF$_5$ BrF$_5$ ClF$_5$	colourless gas but solid below 77 K colourless liquid colourless liquid	Square pyramidal
XX′$_7$	IF$_7$	colourless gas	Pentagonal bipyramidal

(aVery unstable; bThe pure solid is known at room temperature; cDimerises as Cl–bridged dimer (I$_2$Cl$_6$)).

Uses: Non-aqueous solvents, fluorinating agents (e.g., $\textsf{ClF}_3$, $\textsf{BrF}_3$ for $\textsf{UF}_6$ production in uranium enrichment).

Answer:

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Answer:

Group 18 Elements

Group 18 includes Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn), and Oganesson (Og). They are called noble gases due to their low reactivity.

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Occurrence

• Most occur in the atmosphere ($\sim 1\%$ by volume, mainly Argon).
• He and Ne found in radioactive minerals (pitchblende, monazite) and natural gas (main He source).
• Xe and Rn are rare. Rn is radioactive, decay product of $^{226}\text{Ra}$.
• Oganesson (Og) is synthetic, very short-lived. Atomic number 118.

Table 7.12 provides properties of noble gases (except Og).

Propery	He	Ne	Ar	Kr	Xe	Rn*
Atomic number	2	10	18	36	54	86
Atomic mass/ g mol$^{-1}$	4.00	20.18	39.95	83.80	131.30	222.00
Electronic configuration	1s$^2$	[He]2s$^2$2p$^6$	[Ne] 3s$^2$3p$^6$	[Ar]3d$^{10}$4s$^2$4p$^6$	[Kr]4d$^{10}$5s$^2$5p$^6$	[Xe]4f$^{14}$5d$^{10}$6s$^2$6p$^6$
Atomic radius/pm	120	160	190	200	220	–
Ionisation enthalpy/kJmol$^{-1}$	2372	2080	1520	1351	1170	1037
Electron gain enthalpy/kJmol$^{-1}$	48	116	96	96	77	68
Density (at STP)/gcm$^{-3}$	1.8$\times$10$^{-4}$	9.0$\times$10$^{-4}$	1.8$\times$10$^{-3}$	3.7$\times$10$^{-3}$	5.9$\times$10$^{-3}$	9.7$\times$10$^{-3}$
Melting point/K	–	24.6	83.8	115.9	161.3	202
Boiling point/K	4.2	27.1	87.2	119.7	165.0	211
Atmospheric content (% by volume)	5.24$\times$10$^{-4}$	1.82$\times$10$^{-3}$	0.934	1.14$\times$10$^{-4}$	8.7$\times$10$^{-6}$	–

* radioactive

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Electronic Configuration

General valence configuration: ns$^2$np$^6$ (He is 1s$^2$). Full outermost shells contribute to stability.

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Ionisation Enthalpy

Very high due to stable electron configurations. Decreases down the group (increasing size).

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Atomic Radii

Increase down the group (increasing atomic number and number of shells).

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Electron Gain Enthalpy

Large positive values (or close to zero). No tendency to accept electron due to stable configuration.

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Physical Properties

• Monoatomic gases, colourless, odourless, tasteless.
• Sparingly soluble in water.
• Very low m.p. and b.p. due to weak dispersion forces. Helium has the lowest b.p.
• Helium diffuses through many materials.

Answer:

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Chemical Properties

Generally least reactive (inertness) due to:

• Completely filled valence shells (ns$^2$np$^6$).
• High ionisation enthalpies.
• Positive electron gain enthalpies.

Reactivity increases down the group due to decreasing ionisation enthalpy and increasing size (electron cloud is more easily distorted). Xenon is the most reactive among stable noble gases.

Compounds of noble gases: Reactivity was discovered by Neil Bartlett (1962) by reacting Xe with $\textsf{PtF}_6$. Compounds are formed mainly with highly electronegative elements like fluorine and oxygen. Compounds of Kr are few (only $\textsf{KrF}_2$ studied). No true compounds of Ar, Ne, He known.

• Xenon-fluorine compounds: $\textsf{XeF}_2$, $\textsf{XeF}_4$, $\textsf{XeF}_6$ formed by direct reaction of elements under specific conditions. Colorless crystalline solids, powerful fluorinating agents, sublime readily. Readily hydrolysed even by traces of water (e.g., $\textsf{XeF}_2 + 2\textsf{H}_2\text{O} \to \textsf{Xe} + 4\textsf{HF} + \textsf{O}_2$). React with fluoride ion acceptors/donors to form cationic/anionic species.
• Xenon-oxygen compounds: $\textsf{XeO}_3$ (colourless explosive solid) obtained by complete hydrolysis of $\textsf{XeF}_4$ or $\textsf{XeF}_6$. Partial hydrolysis of $\textsf{XeF}_6$ gives oxyfluorides $\textsf{XeOF}_4$ and $\textsf{XeO}_2\textsf{F}_2$.

Structures (VSEPR theory): $\textsf{XeF}_2$ (linear), $\textsf{XeF}_4$ (square planar), $\textsf{XeF}_6$ (distorted octahedral), $\textsf{XeO}_3$ (pyramidal), $\textsf{XeOF}_4$ (square pyramidal).

Answer:

Uses: Helium (non-inflammable, light) in balloons (meteorological), gas-cooled nuclear reactors, cryogenic agent (liquid He), diluent for $\textsf{O}_2$ in diving apparatus (low solubility in blood). Neon in discharge tubes, fluorescent bulbs (advertisements). Argon in inert atmosphere (welding, electric bulbs), handling air-sensitive substances. Xe and Kr in special light bulbs. Radon (radioactive) in cancer therapy.

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