P Block Elements Class 12 | Part 2 | Group 16 | Chapter 7 | Chemistry |

P Block Elements Class 12 | Part 2 | Group 16 | Chapter 7 | Chemistry | Class 12 | CBSE| 

P Block Elements Class 12

Part 2

Group 16

Group 16 of the periodic table consists of five elements viz, Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po). 

Symbol	Element	Atomic No	Electronic configuration
O	Oxygen	8	[He] 2s2 2p4
S	Sulphur	16	[Ne] 3s2 3p4
Se	Selenium	34	[Ar] 3d10 4s2 4p4
Te	Tellurium	52	[Kr] 4d10 5s2 5p4
Po	Polonium	84	[Xe] 4f14 5d10 6s2 6p4

• These elements are collectively called chalcogens.
• The general electronic configuration is ns² np⁴ . 
• The number of valence electrons are 6. 

Question: Why are the group 16 elements called chalcogens?

Answer: Chalcogen means ore forming. The elements of group 16 are called chalcogens because many metals are found as oxides and sulphides and few as selenides and tellurides. 

Atomic Properties

1)Atomic and Ionic Radii

Question: Atomic Radii of elements of  group 16 are smaller than corresponding group 15 elements. Explain ? 

Answer: The smaller radii of group 16 elements compared to group 15 elements are due to increased nuclear charge of the group 16 elements which results in greater attraction of the electrons by the nucleus. 

Question: Why does, down the group, the atomic size of elements of group 16 increases ?

Answer: The increase in the atomic radii of group 16 element down the group is primarily due to increase in number of electron shells. 

2)Ionization Enthalpy

Question: Why is the first ionization enthalpies of elements of group 16 unexpectedly lower than group 15 elements?

Answer: This is due to the relatively symmetrical and more stable electronic configuration of the elements of group 15 as compared to the elements of group 16 (oxygen family). In other words, group 15 elements have half filled stable electronic configuration. So they require high value of ionization enthalpy for removal of first electron. 

Question: Why are second ionization enthalpies of group 16 elements higher than those of corresponding elements of group15? 

Answer: This is because after removal of one electron , group 16 elementys attain stable electronic configuration. So removal of second electron is difficult. Hence, II ionization enthalpies of group 16 elements is high. 

Question: Why does the values of ionization enthalpy of group 16 elements decrease down the group?

Answer: This is due to increasing shielding effect and size as we move down the group. 

3)Electron Gain Enthalpy

Question: Why have group16 elements large negative electron gain enthalpies?

Answer: Group 16 elements require only two electrons to attain stable noble gas configuration. Therefore, they have a high tendency to accept two additional electrons and hence have large negative electron gain enthalpies. 

Question: Why is electron gain enthalpy of oxygen least negative in this group? 

Answer: This is due to its small size. Its electrons which are present in outermost shell apply repulsion for incoming electron and hence the incoming electrons are not accepted with same ease as in case of other elements of this group. 

Question: Why does, down the group, negative value of E.G.E decrease after sulphur?

Answer: After sulphur , size increases as well as shielding effect increases due to which effective nuclear charge is decreased. So electron accepting tendency decreases resulting in decreases electron gain enthaply. 

Electronegativity

Question: Why do group 16 elements have higher electronegativity than corresponding group 15 elements?

Answer: This is due to smaller atomic size of group 16 elements than group 15 elements. Moreover, group 16 elements need only two eletcron to attain the stable noble gas configuration. 

Question: Why is electronegativity of sulphur much lower than that of oxygen?

Answer: This is probably due to an unexpected increase in size of sulphur (107 pm) as compared to that of oxygen (66 pm). 

Note: Oxygen is the second most electronegative element.

Physical Properties

Non- metallic / metallic character

Because of high ionization enthalpy values, the elements of group 16 are less metallic. However, as we move down the group, ionization enthalpy decreases and hence the metallic character increases.

Question; Explain the trend of non metallic / metallic character in group 16 elements?

Answer: Oxygen →Non metal

Sulphur → Non metal and is an insulator

Selenium → metalloid and hence semiconductor

Tellurium →  metalloid and hence semiconductor

Polonium → metallic but radioactive with a short half life

Melting and Boiling points

Question: Why do melting and boiling point and densities increase regularly as we move down the group up to tellurium?

Answer: As we move down the group, atomic size increases. As a result, van der Waals force of attraction among their atom also increase and hence melting and boiling point regularly increase from O to Te. 

Question: Why are the melting and boiling points of polonium lower than those of tellurium? 

Answer: Due to maximum number of intervening d and f electrons, the inert pair effect is also maximum. Consequently, the s valence electrons in polonium are less available as compared to those in tellurium. As a result, van der Waals force of attraction will be weaker in Po than in Te and therefore, m.p and b.p of Po will be lower than that of Te.  

Elemental State

Question: Why does oxygen exist as diatomic gas at room temperature? 

Answer: Due to small size and high electronegativity , oxygen atom form pπ—pπ double bond with other oxygen atom to form O = O molecule. The intermolecular force of attraction between oxygen molecules are weak van der Waals forces and hence oxygen exists as diatomic gas at room temperature. 

Question: Why do other elements (S, Se and Te) exist as octa atomic molecule?

Answer: Elements of group 16 except oxygen do not form pπ—pπ double bond due to their large size. Thus, they prefer to form single bond and have complex structures. Example S, Se and Te exists as octa-atomic molecules (S₈, Se₈ and Te₈) having puckered 8-membered crown shaped rings.

Multiple Bonding

Oxygen form pπ—pπ double bonds but other elements do not. 

Sulphur and other elements of this group possess d orbitals and hence form pπ—dπ multiple bonds. To obtain effective pπ—dπ overlap, the size of d orbital must be similar to that of p orbital.

Tendency of these elements to form pπ—dπ multiple bonds decreases from S to Se. 

Catenation

Question: Why Sulphur has stronger tendency for catenation that oxygen?

Answer: Due to its small size, the lone pair of electron on the oxygen atom repel the bond pair of the O—O bond to a greater extent than the lone pair of electrons on the sulphur atoms in S—S bond. As a result, S—S bond is much stronger than O—O bond and hence sulphur has a much stronger tendency for catenation than oxygen. 

Question: Why catenation tendency decreases from S to Po?

Answer: As the size of the atom increases down the group from S to Po, the strength of the element- element bond decreases and hence tendency for catenation decreases accordingly.

Allotropy

All the elements of this group show allotropy. 

Question: Why Se and Te are called photosensitive elements?

Answer: The grey form of selenium and tellurium consists of parallel chains held by weak metallic bonds. In the presence of light, weak metallic bonded electrons are excited and as a result, number of free electrons increases and so does the electricity. Thus, these elements conduct electricity significantly in presence of light. That is why Se and Te are called photosensitive elements. 

Chemical Properties

Oxidation States

Negative Oxidation State: 

• All metal oxides are ionic and contain O²¯ ions in which oxygen shows an oxidation state of –2. 
• Since the electronegativities decrease as we move down the group, the tendency of these elements to show -2 oxidation state decrease from sulphur to polonium. 
• Oxygen also show an oxidation state of –1 in peroxides such as H₂O₂, –1/2 in superoxides such as KO₂ , zero in O₂ and O₃, +1 in O₂F₂ and +2 in OF₂. 
• The least electronegative element polonium does not exhibit negative oxidation state at all. It shows positive oxidation state only.

Positive Oxidation State:

• Oxygen does not show positive oxidation state except in O₂F₂ and OF₂.
• Other elements of this group show positive oxidation state of +2, +4 and +6 due to promotion of electrons to vacant d orbital.
• S, Se , Te and Po with oxygen are tercovalent (+4 oxidation state). These +4 compounds show both oxidising as well as reducing properties.
• The compounds of S, Se, Te and Po with fluorine show an oxidation state of +6. These compounds in +6 oxidation state show only oxidising properties.

Trends in Chemical Reactivity

Reactivity toward Hydrogen (Formation of Hydrides)

General formula is H₂E where E= O, S, Se, Te, Po

All the hydrides have angular shape involving sp3 hybridization of the central atom.

Question: Why the bond angles decreases from H₂O to H₂Te?

Answer: As we move down the group form O to Te, the size of the central atom goes on increasing and its electronegativity goes on decreasing. As a result, the position of two bond pairs shifts away and away from the central atom as we move from H₂O to H₂Te. Consequently, the repulsion between bond decreases as we move from H₂O to H₂Te . Bond angle decreases in order: H₂O > H₂S > H₂Se > H₂Te 

Physical State

The hydride of O, i.e., H2O is a colourless, odourless liquid while the hydrides of all the other elements are unpleasant, foul smelling, poisonous gases. 

Acidic Nature

Question: Why acidic strength increases from H₂O to H₂Te ?

Answer: As the atomic size increases down the group, the bond length increases and hence the bond strength decreases. Consequently, the cleavage of E—H bond becomes easier. As a result, the tendency to release hydrogen as proton increases, i.e., acid strength increases down the group. 

Note: H₂O is least acidic while H₂Po is most acidic. 

Question: Why H₂O is least acidic?

Answer: It is due to intermolecular hydrogen bonding in the molecule. Hydrogen atom gets trapped in hydrogen bond. Thus release tendency of H⁺ ion is very less. 

Thermal Stability

The thermal stability of hydrides decreases from H₂O to H₂Te because as size of the atom E in H₂E increases, the bond H—E becomes weaker and thus breaks on heating. 

Reducing Character

Hydrides of all elements of this group except that of oxygen, i.e., water is reducing agent. 

Their reducing character increases from H₂S to H₂Te due to decrease in their thermal stability. 

 Melting and Boiling Points

Question: Why the m.pt and b.pt increases from H₂S to H₂Te?

Answer: This is due to increase in van der waals force of attraction with increasing molecular mass.

Question: Why H₂O has high m.pt and b.pt than other hydrides?

Answer : This is due to presence of intermolecular hydrogen bonding in H₂O.

Reactivity towards Oxygen (Formation of Oxides)

Dioxides

Sulphur, Selenium and tellurium when burnt in air form dioxides of formula EO₂. 

S₈ + 8 O₂ → 8SO₂

Question: Why SO₂ exists as gas at room temperature? 

Answer: It exists as descrete molecules even in the solid state. These molecules are held together by weak van der Waals force of attraction. Therefore, SO₂ is gas at room temperature. 

Question: SeO₂ is reducing while TeO₂ is an oxidising agent. Give reasons. 

Answer: Since +6 oxidation state of S is more stable than +4, therefore, it can readily donate electrons and hence acts as reducing agent. For example, it reduces ferric to ferrous salts. 

Fe₂(SO₄)₃ (ferric sulphate) + SO₂ + 2H₂O → 2FeSO₄ (ferrous sulphate) + 2H₂SO₄

Since due to inert pair effect, the stability of +6 oxidation state decreases while the stability of +4 and +2 oxidation state increases down the group. Thus, TeO₂ accepts electron and thus acts as an oxidising agent. 

Question: Discuss the structure and hybridization of SO₂ molecule? 

Answer: In SO₂, S is sp² hybridized. Two of three sp²– orbitals form two σ bondswith oxygen atom while the third contains the lone pair of electrons. SO₂ has bent (angular) structure.

Trioxides

All the elements of this group form trioxides of formula EO₃. Sulphur in SO₃ is sp² hybridised. The three sp2 orbitals of sulphur overlap with p orbitals of oxygen to form three S—O σ bonds. 

The molecules of SO₃ are held by weak van der waal forces of attraction. Therefore, SO₃ exists as triangular planar gaseous molecule at room temperature. 

Note: All trioxides are acidic in nature.

Reactivity toward Halogen (Formation of Halides)

The stability of the halides in any particular oxidation state decreases in the order: F > Cl > Br > I

Hexahalides

Question: Why SF₆ is chemically inert?

Question; Why SF₆ is not hydrolysed but TeF₆ is easily hydrolysed?

Answer: Its chemical inertness is due to the reason that six F atoms protect the sulphur atom from attack by the reagents to such an extent that even thermodynamically most favourable reactions like hydrolysis do not occur. However, as the size of central atom increases, steric hindrance decreases and hence the hydrolysis become easy. 

The order of hydrolysis of hexafluorides is: SF₆ > SeF₆ > TeF₆ 

All hexafluorides have octahedral structure. 

Tetrahalides

Amongst tetrahalides, tetrafluorides are most stable. SF₄ is a gas, SeF₄ is a liquid while TeF₄ is a solid. These fluorides have sp³d hybridization and trigonal bipyramidal structure. 

Question: Why SF₄ is easily hydrolysed while SF₆ does not?

Answer: SF₆ does not undergo hydrolysis because six fluorine atoms protect the sulphur atom from attack of water due to steric hindrance, SF₄ readily undergo hydrolysis because four F atoms cannot protect the S atom from attack by water. 

Question: How tetrahalides of group 16 elements can act both as lewis acid as well as lewis base?

Ans: The tetrahalides can act both as Lewis bases due to the presence of a lone pair of electrons and lewis acids because the central atom can extend its co-ordination number to six. 

Dihalides

All elements except selenium form dihalides. These dihalides are formed by sp³ hybridisation. However, due to presence of two lone pair of electrons they have bent structure.

Question: Why SCl₂ and H₂O show different bond angles while both have bent (angular) shape?

Answer: In SCl₂, the bond angle is little smaller (103º) as compared to that in H₂O (104.5º) . This is due to reason that S is less electronegative than O. As a result, bond pair of two S—Cl bonds lie away from S atom in SCl₂ as compared to those of O—H bonds in H₂O. Consequently, bond pair – bond pair repulsion decreases and hence the bond angle decreases to 103º in SCl₂ from 104.5º in  H₂O. 

Dioxygen

Preparation

By heating oxygen containing salts: 

2KClO₃ (Pot. Chlorate) → 2KCl (Pot. Chloride)  +3O₂

2KMnO₄ (pot. Permanganate) → K₂MnO₄ (Pot. Manganate) + MnO₂ +O₂

By thermal decomposition of oxides of metal: 

2HgO (Mercuric oxide) → 2Hg (Mercury) + O₂

2Ag₂O (Silver oxide) → 4Ag (Silver) + O₂

2Pb₃O₄ (Red lead) → 6PbO (Litharge) + O₂

3MnO₂ (Manganese dioxide) → Mn₃O₄ (Trimanganic tetroxide) + O₂

Physical Properties

• Dioxygen is a colourless, odourless and tasteless gas.
• It is paramagnetic inspite of having even number of electrons.
• It has three isotopes: ¹⁶₈O , ¹⁷₈O and ¹⁸₈O
  

Chemical Properties

Reaction with Metals: 

4Na + O₂ → 2Na₂O 

2Ca +O₂ → 2CaO

2Mg + O₂ → 2MgO

4Al +3O₂ → 2Al₂O₃

2Fe +3O₂ → 2Fe₂O₃ 

Reaction with non-metals: 

P₄ + 5O₂ → P₄O₁₀

S + O₂ → SO₂

N₂ +O₂ → 2NO

Reaction with sulphur dioxide:                        2SO₂ + O₂ → 2SO₃

Reaction with Hydrogen chloride:                4HCl +O₂ → 2H₂O + Cl₂

Question: Why dioxygen requires initiation by external heating?

Answer:  The bond dissociation energy of dioxygen is high and hence the reaction requires initiation by external heating. However, when the reaction starts, it continues on its own. This is due to the reason that the chemical reactions of dioxygen are exothermic and the heat liberated during the reaction is sufficient ot carry out the reactions. 

Ozone, O₃

It is an allotropic form of oxygen and also called trioxygen.

Preparation

Ozone is obtained by passing silent electric discharge through pure, cold and dry oxygen in a specially designed appartus called ozoniser. During this reaction, conversion of dioxygen to ozone is 10% and the product is called ozonised oxygen.

Since the formation of ozone from dioxygen is an exothermic process, therefore, it is necessary to use silent electric discharge in its preparation. A silent electric discharge produces less heat and thus prevent the decomposition of ozone back to oxygen.

Structure of Ozone

The central oxygen atom is sp2 hybridized containing a lone pair of electrons. As a result, ozone has angular structure with bond angle of 117º. 

Question: Account for following : The two O—O bond lengths in the ozone molecule are equal.

Answer: Ozone may be regarded as resonance hybrid of the two resonating structures. Because of resonance , both the O—O bond lengths have partial double bond character. In other words, both the O—O bond lengths are equal.

Physical Properties

• Ozone is a pale blue gas having a strong characteristic smell. 
• In small concentration, it is harmless. However, if concentration rises above 100 ppm breathing becomes uncomfortable.
• Ozone is diamagnetic. 
• It is about 1.67 times heavier than air. 

Chemical Properties

Decomposition: Ozone is thermodynamically unstable as compared to oxygen since its decomposition into oxygen results in liberation of heat (ΔH is -ve) and an increase in entropy (ΔS is +ve). These two effect reinforce each other esulting in large negative gibb’s free energy for decomposition of ozone to oxygen.

2O₃ (Ozone) → 3O₂  

Question: How is O₃ estimated quantitatively? 

Answer: When O₃ is treated with excess of KI buffered with borate buffer, I₂ is liberated quantitatively. 

2I¯ +H₂O +O₃ → 2OH‾ + I₂ + O₂

The I₂ thus liberated is treated against a standard solution of sodium thiosulphate using starch as an indicator and the amount of ozone can thus be calculated. 

2Na₂S₂O₃ (Sod. Thiosulphate) + I₂ → Na₂S₄O₆ (Sod. Tetrathionate) +2NaI

Oxidising Agent: Ozone is thermodynamically unstable thus it decomposes to give O₂ and nascent oxygen. This [O] is responsible for its oxidising power. 

Ozone oxidises black lead sulphite to white lead sulphate

PbS +4 O₃ → PbSO₄ +4O₂

Ozone oxidises moist iodine to iodic acid

I₂ +5O₃ +H₂O → 2HIO₃ +5O₂

Bleaching Action: Because of its oxidising action, it acts as mild bleaching agent and bleaches vegetable colouring matter.

Vegetable coloured matter + O₃ → Colourless oxidised matter + O₂

Question: Which Aerosols deplete ozone?

Answer: The aerosols responsible largely for depletion of ozone layer are

• Nitric oxide emitted from the exhaust system of supersonic jet planes 
• Freons (chlorofluorocarbons) which are used in aerosol spray and refrigerants.

Sulphur – Allotropic Forms

Rhombic or Orthorhombic sulphur or α – Sulphur	Monoclinic sulphur or β sulphur
The S8 molecules are arranged in such a way that ring fits into each other.	The S8 molecules are arranged in such a way that rings are stacked on the top of each other.
Its density is high due to close packing.	Its density is less than α sulphur.
It is bright pale yellow solid.	It is dull yellow.
It is opaque due to close packing.	It is transparent.
It is stable below 369K and transforms into β- sulphur below it.	It is stable above 369 K and transform into α sulphur below it.

Its specific gravity is 2.06 g cm-3 and melts at 385.8K.

Its specific gravity is 1.98 g cm-3 and melts at 393 K.

Sulphur Dioxide

Preparation 

It is formed together with little sulphur trioxide when sulphur is burnt in air or oxygen.

S₈ + 8 O₂ → 8SO₂

In industry, it is obtained as by product of roasting of sulphide ores

4FeS₂ +11O₂ → 2Fe₂O₃ + 8SO₂

Properties

Colour, Smell and Solubility: It is colourless gas with pungent smell and is highly soluble in water. It liquifies at romm temperature under 2 atm pressure and boils at 263 K. 

Acidic Nature: It dissolves in water to form sulphurous acid.

SO₂ +H₂O → H₂SO₃

It readily react with sodium hydroxide solution forming sodium sulphite which then reacts with more sulphur dioxide to form sodium bisulphite.

2NaOH + SO₂ → Na₂SO₃ (Sod.sulphite) + H₂O

Na₂SO₃ + SO₂ + H₂O → 2NaHSO₃ (Sod. Bisulphate)

Reducing Properties: In presence of moisture it acts as good reducing agent. Its reducing character is due to evolution of nascent hydrogen.

It decolourises pink violet colour of acidified KMnO₄ solution  

2KMnO₄ (Pink violet) + 5 SO₂ + 2 H₂O → K₂SO₄ + 2MnSO₄ (Colourless) +2H₂SO₄

Uses

• in the refining of petroleum and bleaching of sugarcane juice.
• as disinfectant for killing germs, fungi and moulds.

Oxoacids of Sulphur 

Sulphuric Acid (H₂SO₄)

Preparation

Sulphuric acid is mostly prepared by Contact process:

1)Production of SO₂ by burning sulphur

S₈ + 8 O₂ → 8SO₂

2)Catalytic oxidation of sulphur dioxide by air to give sulphur trioxide

2 SO₂ + O₂ ⇔ 2 SO₃ 

According to Le Chatelier principle, the favourable condition for maximum yield of SO₃ are:

• High pressure: Since the forward reaction proceed with decrease in number of moles, therfore,  High pressure would favour the reaction. In actual practice, a pressure of 2 bar is used.
• Low temperature: Since the forward reaction is exothermic, therefore. low temperature will favour the rate of reaction. The reaction is carried out at optimum temperature of 720 K. 
• Purity of gases: To prevent poisoning of catalyst, the gases must be free from impurities of As₂O₃, dust particles and moisture. 
• Use of Catalyst: divanadium pentoxide is employed because it is not only cheaper but is also not easily poisoned.

3) Absorption of sulphur trioxide in 98% sulphuric acid to form oleum.

SO₃ + H₂SO₄ → H₂S₂O₇ (oleum)

4) Dilution of oleum to get sulphuric acid 

H₂S₂O₇ + H₂O → 2H₂SO₄

Physical Properties 

• It is colourless syrupy liquid. It is also known as oil of vitriol. 
• It is highly corrosive and produces burns on skin.
• It has a high boiling point (590 K) and high viscosity because its molecules are associated due to intermolecular hydrogen bonding.
• It is always diluted by adding the acid slowly ionto water with constant stirring and not by adding water to acid.

Chemical Properties

Acidic Nature: It ionizes in two steps

H₂SO₄ → H⁺ + HSO₄¯

HSO₄¯→ H⁺ + SO₄^2-

Thus, it acts as strong dibasic acid.

Dehydrating Agent: It has a strong affinity for water and thus it acts as strong dehydrating agent. 

Charring of sugar :                       C₁₂H₂₂O₁₁ → 12C (Sugar charcoal) + H₂O

Dehydration of formic acid:                            HCOOH → H₂O + CO

Oxidising Agent: Hot concentrated sulphuric acid is a moderately strong oxidising agent since it on decomposition gives nascent oxygen.

C + 2H₂SO₄ → CO₂ + 2SO₂ + 2H₂O

Cu + 2H₂SO₄ → CuSO₄ + SO₂ + 2H₂O

3S + 2H₂SO₄ → 3 SO₂ + 2H₂O

Uses

• It is used to manufacture fertilizers like ammonium sulphate and super phosphate of lime.
• Crude petroleum is treated with sulphuric acid to remove unwanted sulphur and other tarry compounds.
• In lead storage battery as electrolyte.
• A number of metals like copper and silver are extracted from their ores using sulphuric acid. 

 

# P Block Elements Class 12

# P Block Elements Class 12 Notes

# P Block Elements Class 12 NCERT Solutions

# P Block Elements Class 12 Notes in Detail

 

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