CBSE Class 11 Chemistry The p-Block Elements Worksheet Set 01

Oxoacids of sulphur

Sulphur forms a large no. of oxoacids like Sulphurous acid (\( \text{H}_2\text{SO}_3 \)), Dithionous acid (\( \text{H}_2\text{S}_2\text{O}_4 \)), Sulphuric acid (\( \text{H}_2\text{SO}_4 \)), Pyrosulphuric acid (Oleum, \( \text{H}_2\text{S}_2\text{O}_7 \)), Peroxomonosulphuric acid (Caro’s acid, \( \text{H}_2\text{SO}_5 \)), Peroxodisulphuric acid (Marshell’s acid, \( \text{H}_2\text{S}_2\text{O}_8 \)) etc. structure of some oxoacids are:

Sulphurous acid (\( \text{H}_2\text{SO}_3 \))      Sulphuric acid (\( \text{H}_2\text{SO}_4 \))      Peroxodisulphuric acid (\( \text{H}_2\text{S}_2\text{O}_8 \))      Pyrosulphuric acid (Oleum) (\( \text{H}_2\text{S}_2\text{O}_7 \))

Sulphuric Acid (\( \text{H}_2\text{SO}_4 \))

The most important oxoacid of sulphur is sulphuric acid which is also known as the ‘King of Chemicals’.

Manufacture

Sulphuric acid is manufactured by the Contact Process which involves three steps:

• (i) burning of sulphur or sulphide ores in air to generate \( \text{SO}_2 \).

  \( \text{S(s)} + \text{O}_2\text{(g)} \rightarrow \text{SO}_2\text{(g)} \)

  Or,

  \( 4\text{FeS}_2\text{(s)} + 11\text{O}_2\text{(g)} \rightarrow 2\text{Fe}_2\text{O}_3\text{(s)} + 8\text{SO}_2\text{(g)} \)

• (ii) conversion of \( \text{SO}_2 \) to \( \text{SO}_3 \) by the reaction with oxygen in the presence of a catalyst (\( \text{V}_2\text{O}_5 \))

  \( 2\text{SO}_2 + \text{O}_2 \rightarrow 2\text{SO}_3 \)

• (iii) absorption of \( \text{SO}_3 \) in \( \text{H}_2\text{SO}_4 \) to give Oleum (\( \text{H}_2\text{S}_2\text{O}_7 \)).

  \( \text{SO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{H}_2\text{S}_2\text{O}_7 \)

• (iv) Dilution of oleum with water gives \( \text{H}_2\text{SO}_4 \) of the desired concentration.

  \( \text{H}_2\text{S}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{H}_2\text{SO}_4 \)

Properties

Sulphuric acid is a colourless, dense, oily liquid. It dissolves in water with the evolution of a large quantity of heat. Hence, for diluting the acid, the concentrated acid must be added slowly into water with constant stirring.

Chemical properties: The chemical reactions of sulphuric acid are due to the following reasons:

• (a) its low volatility
• (b) strong acidic character
• (c) strong affinity for water and
• (d) its ability to act as an oxidising agent.

In aqueous solution, sulphuric acid ionises in two steps.

\( \text{H}_2\text{SO}_4\text{(aq)} + \text{H}_2\text{O(l)} \rightarrow \text{H}_3\text{O}^+\text{(aq)} + \text{HSO}_4^-\text{(aq)} \)
\( \text{HSO}_4^-\text{(aq)} + \text{H}_2\text{O(l)} \rightarrow \text{H}_3\text{O}^+\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \)

So it is dibasic and forms two series of salts: normal sulphates and acid sulphates.

Because of its low volatility sulphuric acid can be used for the manufacture of more volatile acids from their corresponding salts.

\( 2\text{MX} + \text{H}_2\text{SO}_4 \rightarrow 2\text{HX} + \text{M}_2\text{SO}_4 \) (where \( \text{X} = \text{F, Cl, NO}_3 \) etc. and \( \text{M} \) is a metal)

Concentrated sulphuric acid is a strong dehydrating agent and drying agent. Many wet gases can be dried by passing them through sulphuric acid. Sulphuric acid removes water from organic compounds

e.g.: \( \text{C}_{12}\text{H}_{22}\text{O}_{11} + \text{H}_2\text{SO}_4 \rightarrow 12\text{C} + 11\text{H}_2\text{O} \)

Hot concentrated sulphuric acid is a moderately strong oxidising agent. It oxidises both metals and non- metals and the acid itself reduces to \( \text{SO}_2 \).

\( \text{Cu} + 2\text{H}_2\text{SO}_4\text{(conc.)} \rightarrow \text{CuSO}_4 + \text{SO}_2 + 2\text{H}_2\text{O} \)
\( \text{S} + 2\text{H}_2\text{SO}_4\text{(conc.)} \rightarrow 3\text{SO}_2 + 2\text{H}_2\text{O} \)
\( \text{C} + 2\text{H}_2\text{SO}_4\text{(conc.)} \rightarrow \text{CO}_2 + 2\text{SO}_2 + 2\text{H}_2\text{O} \)

Uses:

The important uses of Sulphuric acid are:

1) In the manufacture of fertilizers 2) in petroleum refining 3) in the manufacture of pigments, paints and dyestuff intermediates 4) in detergent industry 5) in metallurgical applications 6) as electrolyte in storage batteries 7) in the manufacture of nitrocellulose products and 8) as a laboratory reagent.

Group 17 Elements

Fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At) are the members of Group 17. They are collectively known as the halogens (means salt producers). They are highly reactive non-metallic elements. All these elements have seven electrons in their outermost shell (\( ns^2np^5 \)) and so they do not readily lose their electron. So they have very high ionisation enthalpy.

Halogens have maximum negative electron gain enthalpy in the corresponding periods. This is due to the fact that the atoms of these elements have only one electron less than stable noble gas configurations. Electron gain enthalpy of these elements decreases down the group. However, the negative electron gain enthalpy of fluorine is less than that of chlorine. It is because, in fluorine the incoming electron goes to the 2p sub shell, but in Cl it enters in to the 3p sub shell. Due to the compactness of 2p sub shell compared to 3p sub shell, the electron – electron repulsion is greater in fluorine than in chlorine. So F does not easily gains electron.

Halogens have very high electronegativity. The electronegativity decreases down the group. Fluorine is the most electronegative element in the periodic table.

All halogens have characteristic colour. For example, \( \text{F}_2 \) has yellow, \( \text{Cl}_2 \)-greenish yellow, \( \text{Br}_2 \)-red and \( \text{I}_2 \), violet colour. This is due to absorption of radiations in visible region which results in the excitation of outer electrons to higher energy level.

The bond dissociation enthalpy of \( \text{F}_2 \) is low. This is due to the relatively large electron-electron repulsion among the lone pairs in \( \text{F}_2 \) molecule.

All the halogens are highly reactive. They react with metals and non-metals to form halides. The reactivity of the halogens decreases down the group.

Halogens are strong oxidising agents since they readily accept electron. \( \text{F}_2 \) is the strongest oxidising halogen and it oxidises other halide ions in solution or in the solid phase.

Oxidation states

All the halogens exhibit –1 oxidation state. Chlorine, bromine and iodine also show + 1, + 3, + 5 and + 7 oxidation states in their oxides, oxy acids and in inter halogen compounds. Due to the absence of vacant d orbitals and the maximum electronegativity, fluorine exhibits only –1 oxidation state.

Anomalous behavior of fluorine

Due to the small size, highest electronegativity, low F-F bond dissociation enthalpy, and non availability of d orbitals in valence shell, fluorine shows properties different from other halogens.

Some of the anomalous properties of fluorine are:

• 1. Ionisation enthalpy, electronegativity, enthalpy of bond dissociation and electrode potentials are higher for fluorine than expected.
• 2. Ionic and covalent radii, m.p. and b.p. and electron gain enthalpy are quite lower than expected.
• 3. Most of the reactions of fluorine are exothermic (due to the small and strong bond formed by it with other elements).
• 4. F forms only one oxoacid while other halogens form a number of oxoacids.
• 5. Hydrogen fluoride is a liquid due to strong hydrogen bonding. While the hydrogen halides of other elements are gases.

Hydrides of halogens

Halogens react with hydrogen to give hydrogen halides which dissolve in water to form hydrohalic acids. The acidic strength of these acids varies in the order: \( \text{HF} < \text{HCl} < \text{HBr} < \text{HI} \). The stability of these halides decreases down the group due to decrease in bond dissociation enthalpy from \( \text{HF} \) to \( \text{HI} \).

Chlorine (\( \text{Cl}_2 \))

Preparation:

It can be prepared by any one of the following methods:

• (i) By heating manganese dioxide with concentrated hydrochloric acid.

  \( \text{MnO}_2 + 4\text{HCl} \rightarrow \text{MnCl}_2 + \text{Cl}_2 + 2\text{H}_2\text{O} \)

  Conc. HCl can be replaced by a mixture of common salt and concentrated \( \text{H}_2\text{SO}_4 \):

  \( 4\text{NaCl} + \text{MnO}_2 + 4\text{H}_2\text{SO}_4 \rightarrow \text{MnCl}_2 + 4\text{NaHSO}_4 + 2\text{H}_2\text{O} + \text{Cl}_2 \)

• (ii) By the action of HCl on potassium permanganate.

  \( 2\text{KMnO}_4 + 16\text{HCl} \rightarrow 2\text{KCl} + 2\text{MnCl}_2 + 8\text{H}_2\text{O} + 5\text{Cl}_2 \)

Manufacture of chlorine

• (i) Deacon’s process: By oxidation of hydrogen chloride gas by atmospheric oxygen in the presence of \( \text{CuCl}_2 \) (catalyst) at 723 K.

  \( 4\text{HCl} + \text{O}_2 + \text{CuCl}_2 \rightarrow 2\text{Cl}_2 + 2\text{H}_2\text{O} \)

• (ii) Electrolytic process: Chlorine is obtained by the electrolysis of brine solution (concentrated NaCl solution). During electrolysis chlorine is liberated at the anode.

Properties

It is a greenish yellow gas with pungent and suffocating odour. It is soluble in water. It reacts with a number of metals and non-metals to form chlorides.

\( 2\text{Al} + 3\text{Cl}_2 \rightarrow 2\text{AlCl}_3 \);      \( \text{P}_4 + 6\text{Cl}_2 \rightarrow 4\text{PCl}_3 \)
\( 2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl} \);      \( \text{S}_8 + 4\text{Cl}_2 \rightarrow 4\text{S}_2\text{Cl}_2 \)
\( 2\text{Fe} + 3\text{Cl}_2 \rightarrow 2\text{FeCl}_3 \);

With excess ammonia, chlorine gives nitrogen and ammonium chloride whereas with excess chlorine, nitrogen trichloride (explosive) is formed.

\( 8\text{NH}_3\text{ (excess)} + 3\text{Cl}_2 \rightarrow 6\text{NH}_4\text{Cl} + \text{N}_2 \);
\( \text{NH}_3 + 3\text{Cl}_2\text{ (excess)} \rightarrow \text{NCl}_3 + 3\text{HCl} \)

With cold and dilute alkalies chlorine produces a mixture of chloride and hypochlorite but with hot and concentrated alkalies it gives chloride and chlorate.

\( 2\text{NaOH} + \text{Cl}_2 \rightarrow \text{NaCl} + \text{NaOCl} + \text{H}_2\text{O} \) (cold and dilute)
\( 6\text{NaOH} + 3\text{Cl}_2 \rightarrow 5\text{NaCl} + \text{NaClO}_3 + \text{H}_2\text{O} \) (hot and conc.)

With dry slaked lime it gives bleaching powder.

\( 2\text{Ca(OH)}_2 + 2\text{Cl}_2 \rightarrow \text{Ca(OCl)}_2 + \text{CaCl}_2 + 2\text{H}_2\text{O} \)

Chlorine reacts with hydrocarbons and gives substitution products with saturated hydrocarbons and addition products with unsaturated hydrocarbons.

\( \text{CH}_4 + \text{Cl}_2 \xrightarrow{\text{UV}} \text{CH}_3\text{Cl} + \text{HCl} \)
Methane      Methyl chloride
\( \text{C}_2\text{H}_4 + \text{Cl}_2 \rightarrow \text{C}_2\text{H}_4\text{Cl}_2 \)
Ethene      1,2-Dichloroethane

Chlorine water on standing loses its yellow colour due to the formation of HCl and HOCl. Hypochlorous acid (HOCl) so formed is unstable and dissociates to give nascent oxygen which is responsible for oxidising and bleaching properties of chlorine.

• (i) It oxidises ferrous to ferric, sulphite to sulphate, sulphur dioxide to sulphuric acid and iodine to iodic acid.

  \( 2\text{FeSO}_4 + \text{H}_2\text{SO}_4 + \text{Cl}_2 \rightarrow \text{Fe}_2(\text{SO}_4)_3 + 2\text{HCl} \)

  \( \text{Na}_2\text{SO}_3 + \text{Cl}_2 + \text{H}_2\text{O} \rightarrow \text{Na}_2\text{SO}_4 + 2\text{HCl} \)

  \( \text{SO}_2 + 2\text{H}_2\text{O} + \text{Cl}_2 \rightarrow \text{H}_2\text{SO}_4 + 2\text{HCl} \)

  \( \text{I}_2 + 6\text{H}_2\text{O} + 5\text{Cl}_2 \rightarrow 2\text{HIO}_3 + 10\text{HCl} \)

• (ii) It is a powerful bleaching agent; bleaching action is due to oxidation.

  \( \text{Cl}_2 + \text{H}_2\text{O} \rightarrow 2\text{HCl} + [\text{O}] \)

  Coloured substance + \( [\text{O}] \rightarrow \) Colourless substance

It bleaches vegetable or organic matter in the presence of moisture. Its bleaching action is permanent.

Uses:

It is used

• (i) for bleaching wood pulp, bleaching cotton and textiles,
• (ii) in the extraction of gold and platinum
• (iii) in the manufacture of dyes, drugs and organic compounds such as \( \text{CCl}_4 \), \( \text{CHCl}_3 \), DDT, refrigerants, etc.
• (iv) in sterilising drinking water and
• (v) preparation of poisonous gases such as phosgene (\( \text{COCl}_2 \)), tear gas (\( \text{CCl}_3\text{NO}_2 \)), mustard gas (\( \text{ClCH}_2\text{CH}_2\text{SCH}_2\text{CH}_2\text{Cl} \)).