Chemistry Answer the following in one or two sentence.

• Oxygen being more electronegative, is capable of forming hydrogen bonding in the compound $H_2O$.
• The other elements $S, Se$ and $Te$ of Group $16$, being less electronegative do not form hydrogen bonds.
• Thus, hydrogen bonding is not present in the other hydrides $H_2S, H_2Se$ and $H_2Te$.

• Oxygen combines with the most electronegative element fluorine to form $OF_2$ and exhibits positive oxidation state $(+ 2)$. Since, oxygen does not have vacant J-orbitals it cannot exhibit higher oxidation states.
• Sulphur has vacant d-orbitals and hence can exhibit + $6$ oxidation state to form $SF_6$​​​​​​​.

• for sterilization of drinking water.
• bleaching wood pulp required for the manufacture of paper and rayon, cotton and textiles are also bleached using chlorine.
• in the manufacture of organic compounds like $CHCl _3, CCl _4$ DDT, dyes and drugs.
• in the extraction of metals like gold and platinum.
• in the manufacture of refrigerant like Freon (i.e., $CCl _2 F_2$ ).
• in the manufacture of several poisonous gases like mustard gas $\left( Cl ^{-} C _2 H _4- S - C _2 H _4- Cl \right)$, phosgene $\left( COCl _2\right)$ used in warfare.
• in the manufacture of tear gas $\left( CCl _3 NO _2\right)$.

$
Mg _{( s )}+ Br _{2( l )} \longrightarrow MgBr _{2( s )}\ \ (Magnesium \ bromide)
$

• in interhalogen compound $ICl$ and $I_2$ the atoms are bonded by covalent bonds.
• The covalent bond between dissimilar atoms, $I$ and $Cl$ atoms in $ICl$ is weaker than between similar atoms in $I_2.$
• Therefore bond dissociation enthalpy of ICl bond is less than that of $I_2$ bond.
  Hence $ICl$ is more reactive than $I_2.$

• $ICl$ is used as a halogenating agent and also in the estimation of iodine number of fats and oils.
• In the preparation of polyhalides, interhalogen compounds are used.
• $ClF_3 $ and $BrF_3$_ are widely used as fluorinating agent. They are used in the separation of $^{235}U$ isotope by fluorinating. $^{235}U_{(S)} + 3ClF_{3(l)} – ^{235}UF_{6(g)} + 3ClF_{(g)}$
• In propellants $ClF_3 $ and $IF_3 $ are used as oxidisers.
• They are used as non$-$aqueous solvents.
• The $XX’$ interhalogen compounds are used as a catalyst for the oxidation of As$(III)$

• in the manufacture of chlorine and ammonium chloride,
• to manufacture glucose from com, starch
• to manufacture dye
• in mediClne and galvanising
• as an important reagent in the laboratory
• to extract glue from bones and for the purification of bone black.
• for dissolving metals, $Fe +2 HCl _{( aq )} \rightarrow FeCl _2+ H _{2( g ) Z }$

• Chlorine reacts with the excess of ammonia to form ammonium chloride, $NH_4Cl$ and nitrogen.
  $3 Cl _2+\underset{\text { (excess) }}{8 NH _3} \longrightarrow 6 NH _4 Cl + N _2$
• When chlorine is in excess, then with ammonia it forms explosive nitrogen trichloride, $NCl_3.$
  $\underset{\text { (excess) }}{3 NH _2}+3 NCl + NCl _3$

• On a large scale chlorine is obtained by an electrolytic process.
• The electrolyte is brine, a concentrated solution of NaCl.
• Nelson’s two compartments diaphragm cell is used for electrolysis in which stout graphite rod is used as an anode while U shaped steel vessel is used as a cathode.
• Reactions :
  $\begin{aligned} & NaCl \rightarrow Na ^{+}+ Cl ^{-}
  \\ & H _2 O \rightleftharpoons H ^{+}+ OHAt \text { cathode : }
  \\ & 2 H _2 O +2 e ^{-} \rightarrow H _2+2 HH ^{-}
  \\ & Na ^{+}+ OH ^{-} \rightarrow NaOHAt \text { anode : }
  \\ & 2 Cl \rightarrow 2 Cl +2 e ^{-}
  \\ & 2 Cl \rightarrow Cl _{2( g )}\end{aligned}$

• In the contact process of the manufacture of sulphuric acid, $SO_2$ is oxidised to $SO_3$ by heating the mixture on the heterogeneous catalyst $V_2O_5.$
  $2 SO _{2( g )}+ O _{2( g )} \rightleftharpoons 2 SO _{3( g )} \Delta H =-196.6 kJ$
• The forward reaction is exothermic and there is a decrease in number of moles and volume.
• The optimum conditions to maximise the yield of $H_2SO_4$ are pressure of $2$ bar and temperature around $720 K.$

• $H _2 SO _4$ is a dibasic acid. In aqueous, solution it dissociates in two steps as follows :
  $ H _2 SO _{4( aq )} \rightleftharpoons H ^{+}{ }_{( aq )}+ HSO ^{-}{ }_{4( aq )} Ka _1>10$
  $ HSO _{( aq )} \rightleftharpoons H ^{+}( aq )+ SO ^{2-}{ }_{4( aq )} Ka _2=1.2 \times 10^{-2}$
• Neutral $H _2 SO _4$ molecule has more tendency to lose proton $(H^+)$ than anionic Lowry$-$Bronsted base $HSO -4$.
• Therefore second dissociation constant $Ka_2$ is smaller than first dissociation constant $Ka_1.$

• $FeSO _4$ is oxidised to $Fe _2\left( SO _4\right)_3$ by hot and concentrated $H _2 SO _4$. $2 FeSO 4+2 H 2 SO 4 \rightarrow \Delta Fe 2( SO 4) 3+ SO 2+2 H 2 O$
• $Hl$ is oxidised to $I 2$ by hot and concentrated $H 2 S 04$.
  \92 HI + H _2 SO _4 \rightarrow 2 H _2 O + SO _2+ I _2$

• Sulphur dioxide, $SO_2$ reduces halogens $(X^\circ )$ to haloacids $(X^{1-}).$
  $I_2 + 2H_2O + SO_{2 }\rightarrow 2HI + H_2SO_4$
• $SO_2$_ gas when passed through an acidified solution of $KMnO_4$ it is decolourised due to reduction $(Mn^{7+} to Mn^{2+}).$
  $2KMnO_4 + 2H_2O + 5SO_2 \rightarrow K_2SO_4 + 2MnSO_4 + 2H_2SO_4$
• Sulphur dioxide $(SO_2)$ reduces $Fe^{3+} $ to $Fe^{2+}$
  $2FeCl_3 + SO_2 + 2H_2O \rightarrow 2FeCl_2 + H_2SO_4 + 2HCl$

• The coolents used in refrigerants are generally chlorofluorocarbons, $CFCs$ also known as Freon, for example $CF_2Cl_2.$
• $CFCs$ have very long lifetime about $20$ to $100$ years.
• $CF_2Cl_2$_ undergoes photochemical decomposition giving free radicals which react and destroy ozone.$ CF _2 Cl _2 \xrightarrow{h \nu} CF _2 Cl \cdot+ Cl \cdot$
  $ Cl \cdot+ O _3 \longrightarrow ClO \cdot+ O _2$
  $ ClO \cdot+ O _3 \longrightarrow Cl \cdot+2 O _2$

• When a slow dry stream of oxygen is passed through a silent electric discharge, oxygen is converted into ozone $($about $10\%). $ The mixture is called ozonised oxygen. $3 O _2 \xrightarrow{\text { electric } \text { discharge }} 2 O _3 \Delta H=+142 kJ$
• It is an endothermic reaction.
• Silent electric discharge prevents the decomposition of ozone.

• In the atmosphere, ozone is naturally formed through photochemical reactions.
• Oxygen present in the lower mesosphere on the absorption of solar radiations, is dissociated into two oxygen atoms which oxidise oxygen to ozone.
• One atomic oxygen combines with molecular oxygen to form $O _3$.
  $\begin{aligned} & O _2 \xrightarrow{\text { U.V. light }} O + O
  \\ & O _2+ O \longrightarrow O _3\end{aligned}$

• acidic oxides, $CO _2, SO _2$, etc.
• Basic oxides, $CaO , BaO$, etc.
• Amphoteric oxides, $Al _2 O _3, ZnO$, etc.
• Neutral oxides, $NO , N _2 O , CO$, etc.

• Dioxygen is essential for sustaining life. It is used in hospitals for artificial respiration.
• It is used in oxy-hydrogen $\left(2800^{\circ} C \right)$ and oxy-acetylene $\left(3200^{\circ} C \right)$ torches used for cutting and welding metals.
• Oxygen cylinders are used in hospitals, for high altitude flying and in mountaineering.
• It is used in the combustion of fuels. In rockets, hydrazine in oxygen is used as a fuel as it provides tremendous thrust.

• All halogens react with metals instantly to give metal halides.
• Down the group reactivity decreases from fluorine to iodine.
• $2 Na _{( s )}+ Cl _{2(1)} \rightarrow 2 NaCl _{( s )}, Mg _{( s )}+ Br _{2( s )} \rightarrow MgBr _{2( s )}$
• Halogens being highly electronegative, the metal-halogen bonds are ionic and ionic character decreases down the group. For example $M - F > M - Cl > M - Br > M - I$.
• The metal halides with higher oxidation state of the metal are more covalent than the halides with lower oxidation state of metal. For example $SnCl _4, PbCl _4, SbCl _5$ and $UF _6$ are more covalent than

• Group 18 elements are chemically inert.
• However, inert elements like krypton and xenon react directly with fluorine under appropriate conditions to give their fluorides. For example,
  $Xe _{( g )}+ F _{2( g )} \xrightarrow[1 atm ]{673 K } XeF _{2( s )}$
• The xenon fluorides $XeF _2, XeF _4$ and $XeF _6$ are colourless crystalline solids which sublime at 298 K.
• These fluorides are strong fluorinating agents.

• The halogens have a tendency to combine amongst themselves forming different compounds called interhalogen compounds.
• Interhalogens are of the type $X X, X X_3^{\prime}, X X_5^{\prime}$ and $X X_7^{\prime}$ where X is larger size halogen than $X^{\prime}$.
• They are covalent, diamagnetic, reactive and good oxidising agents.
• Preparation :
  $\begin{aligned} & Cl _2+ F _2 \xrightarrow{437 K } 2 ClF
  \\ & Br _2+3 F _2 \xrightarrow{\Delta} 2 BrF _3\end{aligned}$

• Halogens with oxygen form many oxides with different oxidation states of halogens.
• Fluorine forms $O _2 F _2$ and $OF _2$.
• Chlorine forms oxides like  $\stackrel{1+}{ Cl _2 O }, \stackrel{3+}{ Cl _2} O _3, \stackrel{5+}{ Cl }_2 O _5, \stackrel{6+}{ ClO _3}, \stackrel{7+}{ Cl _2} O _7$ All oxides are oxidising agents.
• Bromine forms oxides like $Br _2 O , BrO _2, BrO _3$ and iodine forms $I _2 O _4, I _2 O _5, I _2 O _7$

• The H-F bond is stronger than H-Cl bond.
• Hence, HF ionises less readily than HCl in $H _2 O$, to give $H ^{+}$ ions.
  Therefore, HF is a weaker acid than HCl.

• The aCldic strength of haloaClds vary in the following order, $HF < HCl < BHr < HI$
• The stability of hydrogen halides decreases down the group. $HF > HCl > HBr > HI$.
  This is because from F to I, atomic size increases and bond dissociation enthalpy of H-X decreases.

• Oxygen being more electronegative, the O-H bond is more polar and there arises association of $H _2 O$ molecules through intermolecular hydrogen bonding.
• The other hydrides of group 16 do not form H bonds and hence exist as discrete molecules.
• As a result, water shows unusual properties like high B.P, high thermal stability and weaker acidic character as compared to the other hydrides of group 16.

• Hydrogen bonding and molecular association are not present in $H _2 S$.
• The $H _2 O$ molecules are associated with each other through intermolecular hydrogen bonding.
• Hence, $H _2 S$ has a lower boiling point than $H _2 O$.

• Except $H _2 O$, all hydrides of group 16 elements are reducing agents.
• reducing power increases from $H _2 S$ to H2Te.
• reducing power of the hydrides is due to their less stability and tendency to dissociate, which increases from $H _2 S$ to $H _2 Te$.

• The thermal stability of hydrides decreases in the order of $H _2 O > H _2 S > H _2 Se > H _2 Te$.
• Since atomic size increases from O to Te, the tendency to form hydride bond, M-H decreases.
• Hence M-H bond in O-H is the strongest and in Te-H the weakest. Therefore thermal stability decreases from $H _2 O$ to $H _2 Te$.

• All the elements of group 16 react with hydrogen and form hydrides of the type $H _2 M$ where M = O, S, Se, Te and Po. For example $H _2 O , H _2 S , H _2 Se , H _2$Te and $H _2 Po$.
• The hydrides of these elements show regular trends in their physical and chemical properties.z

• The electronic configuration of oxygen is $Is^2 2 s^2 2 p^4$.
• It has two half filled p-orbitals and no d-orbitals for exCltation of electrons. Hence it cannot show higher oxidation states.
• Oxygen being highly electronegative, it mostly shows an oxidation state of – 2 only.
• Other members of the family like sulphur, have vacant d-orbitals, thereby giving four and six half- filled orbitals for bonding.
• Furthermore, they can combine with more electronegative elements.
• Hence, they show oxidation states of $+2,+4$ and 4- 6 also.

• It has the smallest size in the group.
• It has the highest electronegativity ($3.5$).
• It does not have vacant d-orbitals like other elements in group 16.

• The bond dissociation enthalpies of halogen molecules decrease down the group with an increase in atomic size.
• However, the bond dissociation enthalpy of fluorine is lower than that of chlorine and bromine because the F – F bond is weak.
• The bond dissociation enthalpies of the halogen molecules decreases in the order,
  $Cl - Cl > Br - Br > F - F > I -1$..

• Halogens are soluble in various organic solvents such as chloroform, carbon disulphide and carbon tetrachloride
• All halogens react with hydrocarbons and form mono or multisubstituted compounds.
• The reactivity decreases down the group from fluorine to iodine.
  $\begin{aligned} & CH _4+2 F _2 \rightarrow C +4 HF
  \\ & C +2 F _2 \rightarrow CF _4
  \\ & CH _4+ Cl _2 \rightarrow CCl _4+4 HCl \end{aligned}$

• Halogens react with water forming halo acids and haloxyaClds.
• The reactivity decreases from fluorine to iodine.
• $\begin{aligned} & 2 F _2+2 H _2 O \rightarrow 4 HF + O _2
  \\ & 3 F _2+3 H _2 O \rightarrow 6 HF + O _3
  \\ & X _{2( g )}+ H _2 O _{(1)} \rightarrow HX _{( aq )}+ HOX ( X = Cl , Br ) \\ & \end{aligned}$

• Group 18 elements have very low melting points and boiling points.
• The melting points and boiling points increase down the group from He to Rn.
• As the size of the atoms increases on moving down the group, the magnitude of the van der Waals forces increase from He to Rn.
• Thus, the melting and boiling points increase from He to Rn. Helium has the lowest boiling point $(4.2 K )$ of any known substance.
• The boiling points increase in the order He $He < Ne < Ar < Kr < Xe < Rn$

• Since halogens have eletronic configuration, $n s^2$ ,$n p^5$, they have tendency to accept one electron to complete an octet.
• Group 17 elements have the lowest atomic radii and high effective nuclear charge.
• Therefore halogens have high values of electronegativity.
• The decreasing order of electronegativity is, $F > Cl > Br > I$

• The atomic or ionic radii decreases across a period with increase in atomic number and increase in the effective nuclear charge $\left(Z_{\text {eff }}\right)$.
• The ionisation enthalpy also increases across a period with increase in atomic number. This is due to addition of electrons to the same shell, while moving across a period.
• Inert element at the end of period has the highest value of ionisation enthalpy.

• In general, group 18 elements have high values of ionisation enthalpy.
• In a period, each noble gas has the highest ionisation enthalpy.
• The noble gases have electronic configuration, $n s^2 n p^6$, they have complete octet with paired electrons and very stable closed-shell electronic configuration. Therefore high energy is required to remove the electron from valence shell.
• Down the group ionisation enthalpy decreases.

• Group 18 elements have electronic configuration $n s^2 n p^6$ and complete octet of electrons, hence they have no tendency to accept electrons.
• Therefore, they have a large positive electron gain enthalpy.

• The atomic radii of group 18 elements is larger than the atomic radii of group 17 elements.
• Down the group from He to Xe, atomic radii increases due to increase in the number of quantum shells.
• The atomic radii increase in the order $He < Ne < Ar < Kr < XeZ$

• As compared to other elements (K to Kr) of 4th period, bromine (Br) has the lowest atomic size.
• It has electronic configuration, $35 Br [ Ar ]^{18} 3 d^{10} 4 s^2 4 p^5$.
• It has seven valence electrons and it needs one electron to complete octet.
• Hence Br has the strongest tendency to gain an electron and energy released is maximum in the period. $Br _{( g )}+ e ^{-} \rightarrow Br _{( g )} \Delta_{ eg } H =-325 kJ mol ^{-1}$
  

• Oxygen has low atomic size.
• It has high electronegativity $(3.5)$.
• It has high electron density.
• The incoming electron is repelled due to high electron density hence oxygen has less negative value of electron gain enthalpy as compared to other group 16 elements. $O _{( g )}+ e ^{-} \rightarrow O ^{-}{ }_{( g )} \Delta_{ eg } H =-141 kJ mol ^{-1}$

• The most important source of noble gases is the atmosphere where they make up about $1 \%$ by volume of air. Argon is the major constituent.
• All noble gases occur in nature except Radon. Radon is a decay product of radioactive ${ }^{226} Ra$..
• Helium on a large scale is obtained from natural gas. Helium and neon are found in the minerals pitchblende, monazite and cleveite.
• Xenon and Radon are the rarest noble gases.

• Sulphides : Galena (PbS), Zinc blende (ZnS), Copper pyrites $\left( CuFeS _2\right)$
• Sulphates : Gypsum $\left( CaSO _4 \cdot 2 H _2 O \right)$, Epsom salt $\left( MgSO _4 \cdot 7 H _2 O \right)$, Baryte $\left( BaSO _4\right)$