Group 16

Electronic configuration- ns2np4

Radii- decreases down the group Ionization enthalpy- decreases down the group

I.E. (gr 15) > I.E. (gr 16) (Extra stable)

Electron gain enthalpy- O < S Less negative due to e - repulsion in small 2p orbital of O

Electronegativity- decreases down the group.

Physical properties- MP and BP increases with increase in atomic number. Large difference b/w MP and BP of O and S because

O2 - diatomic S8 – polyatomic

Chemical properties- Oxidation states

O -2 +2 more common S -2 +2 +4 +6Se -2 +2 +4 +6Te -2 +2 +4 +6Po -2 +2 +4 +6

Stability decreases down the group

Stability increases down the group

-2 stability decreases down the group OF2 (+2) ,O2F2

S, Se, Te : +4 with O S, Se, Te : +6 with F +4 and +6 oxidation state have covalent bond

Oxygen anomalies – Small size, high electronegativity – H-bond, H-bond in H2O not in H2S Covalency of O is max. 4 because of non- availability of d- orbitals.

Q. SF6 is known but SH6 is not?A. It is due to high oxidation state (+6), S can combine only with highly electronegative F.

React towards H2 –H2O < H2S < H2Se < H2Te (acidic character)H2O < H2S < H2Se < H2Te (reducing behav.)H2O > H2S > H2Se >H2Te (thermal stability)

Reducing property -H2S < H2Se < H2Te H2O does not show reducing property

Reaction with O –EO2, EO3 (both are acidic in nature)Example- O3, SO2, SeO2, SO3, SeO3, TeO3

Reducing property-SO2 > SeO2 > TeO2

(R.A) (O.A)

Q. Why H2O is liquid and H2S is gas?Q. Why H2S is less acidic than H2Te?

React towards Halogens –EX6, EX4, EX2

Note- only EF6 is stable (all gaseous octahedral structure eg: SF6 exceptionally stable due to steric reason)Stability order of halides –F- > Cl- > Br- > I-

Q. Why SF6 is highly stable?A. Because it is sterically protected

EX4 : sp3d, see-saw structure SF4 (gas)

SeF4 (liquid) TeF4 (solid)

Tetrahalides act as Lewis base by having lone pair and Lewis Acid due to extension of co-od number. Following reaction supports above facts, SF4 + BF3 → SF4→BF4

SF4 +2F-→SF6-2

EX2 : sp3 All except Se forms ECl2 and EBr2

EX : monohalides S2F2, S2Cl2, SBr2, Se2Cl2, Se2Br2

2Se2Cl2 → SeCl4 + 3Se (Disproportionation)

Preparation of O2 – Lab. preparation:

i. 2KClO3 △ /MnO2→ 2KCl + 3O2 (NO3

- or MnO4- also used as catalyst)

ii. 2Ag2O(s) → 4Ag(s) + O2 (g)

2HgO(s) → 2Hg(s) + O2 (g)

less reactive metals

2Pb3O4(s) → 6PbO(s) + O2 (g)

higher oxides 2PbO2(s) → 2PbO(s) + O2

iii. 2H2O2 decomposition / finely dividedMnO2→ 2H2O + O2

iv. Electrolysis of water

Properties of O2 –

Colourless, odourless Paramagnetic, despite of having even e- (e-s in π*2px1, π*2py1)

2Ca + O2 → 2CaO

4Al + 3O2 → 2Al2O3

P4 + 5O2 →

P4O10 exothermic reaction

C + O2 →

CO2

2ZnS + 3O2 →

2ZnO + 2SO2

CH4 + 2O2 →

CO2 + 2H2O

2SO2 + O2 V 2O 5→

2SO3

4HCl + O2 →

2Cl2 + 2H2O

Uses:

Respiration Oxyacetylene welding Metal manufacturing Mountaineering Oxidizers in fuel

Simple oxides: An oxide is a binary compound of oxygen with other element.

Oxides

Simple Mixed

MgO, Al2O3 Pb3O4, Fe3O4

Metals in their high Non-metal oxides oxidation state

Acidic oxide: SO2, Cl2O7, CO2, N2O5, Mn2O7, CrO3, V2O5

An oxide which combine with water to give acid

Eg. SO2 + H2O →

H2SO3

Basic: CaO

Eg. CaO + H2O →

Ca(OH)2

Amphoteric: both acidic and basic

Al2O3 + 6HCl + 9H2O →

2[Al(H2O)6]+ + 6Cl-

Al2O3 + 6NaOH + 3H2O →

2Na3[Al(H2O)6](aq)

Neutral oxides: neither acidic nor basic

Eg. CO, NO, N20

Ozone (O3) –

O3 is thermodynamically less stable

3O2 →

2O3 (△H= +142)

△S = -ve, △H= +ve, △G= +ve, therefore O3 is unstable

△G for O3 → O2 will be negative

At high concentration O3 is explosive

O3 as oxidising agent –

O3 (powerful oxidising agent) →

O2 + O (nascent oxygen)

Eg. PbS + 403 →

PbSO4(s) + 4O2(g)

2I- + H2O + O3 →

2OH- + I2 + O2(g)

Estimation of O3 volumetrically –

O3 + KI borate buffer / pH 9.7→

I2 (titrated against Na2S2O3)

Depletion of ozone layer –

NO(g) + O3(g) depletion→

NO2(g) + O2(g)

Note – NO is released from jet engines which combines rapidly with ozone

Freons + O3 →

depletion, freons are the substances released from sprays and

refrigerants

Structure of O3 –

Bond length- 128 pm Bond angle- 117°

Uses of O3 –i. Germicideii. Disinfectant-iii. Sterilizing wateriv. Bleaching oils, ivory, flour, starch etc. v. Oxidising agentvi. Manufacturing of KMnO4

Sulphur – Allotropes of S :

α Sulphur (stable under 369k)369k−equilibrium⇔

β Sulphur (stable above 369k)

Also known as rhombic sulphur also known as monoclinic sulphur)

i. Yellow in colour Yellow in colourii. MP 385.8k MP 393k

iii. Density- 2.08 g/cm 1.98 g/cmiv. Insoluble in water Insoluble in water v. Soluble in certain extent in Soluble in certain

extent inBenzene, alcohol and ether Benzene, alcohol and ether

vi. Readily soluble in CS2 Soluble in CS2

vii. S8 S8

viii. Crowned puckered structure Crowned puckered structure

Other allotropes of S has 6-20 S atoms

At high temp. above 1000K S8(s) →

S2(g) (paramagnatic)

Q. Why S in vapour phase is paramagnetic?

Sulphur dioxide – SO2 Preparation:

S(s) + O2 △

6−8% SO3→

SO2(g)

In lab. –

SO3-2(aq) + 2H+(aq)

→ H2O + SO2(g)

{dil.H2SO4}

Industrially –

4FeS2 + 1102(g) △ /roasting→

2Fe2O3 + 8SO2(g)

Sulphide by product Ore

Properties: Colourless gas Pungent smell Highly soluble in water Liquefies at room temp. at 2atm BP – 263k

Reaction with water – SO2(g) + H2O(aq) →

H2SO3(aq)

Acid

Reaction with NaOH –

SO2 + NaOH →

Na2SO3 + H2O

Na2SO3 + H2O + SO2 →

2NaHSO3

Excess

Note- behaviour of SO2 similar to that of CO2

Reaction with Cl2 and O2 –

SO2 + Cl2 charcoal→

SO2Cl2 (sulphuryl chloride)

2SO2 + O2 N 2O 5→

2S03

SO2 as a reducing agent – Its reducing behaviour is due to liberation of nascent hydrogen hence a temporary bleaching agent.

Fe+3 + SO2 + 2H2O →

2Fe+2 + SO4-2 + 4H+ where SO2 is reducing agent

2MnO4- (+7) + 5SO2 + 2H2O

→ 5S04

-2 + 4H+ + 2Mn+2 (+2)

(Violet) (No colour)

Note: Above reaction is used for detection of SO2

Structure of SO2 –

Angular

Both S-O bond are same (resonance)

Uses –

Refining of petroleum and sugar

Bleaching wool and silk

Antichlor, disinfectant and preservative

Manufacture of H2SO4, NaHSO3, CaHSO3 etc.

Liquid SO2 as a solvent to dissolve a number of chemicals.

Q. how is presence of SO2 detected?

Oxoacids of sulphur –

Sulphuric acid –

Contact process:

Note: Formation of SO3 is key step during contact process. Optimum condition required is

Temperature of 720K

Pressure 2 bar

Catalyst- V2O5

Q. Why water is not directly added to SO3 during prep. Of H2SO4?

Properties –

Colourless Dense Oily liquid 1.84 g/cm density FP- 283k, BP- 611k

Reaction with water:

H2SO4 + H2O →

large amount of heat

Note: add H2SO4 to water instead of adding water to H2SO4 to dilute it (with constant stirring).

Chemical characteristics of H2SO4 – Low volatility Strong acidic character Strong affinity for water Oxidising agent

H2SO4 ionizes in water as:

H2SO4 + H2O →

H30+ + HSO4- Ka1 > 10 (very large)

(From NaHSO4)

H2SO4 + H2O → H30+ + SO4

2- Ka2 = 1.2 x 10-2

(From Na2SO4)

Note – H2SO4 has two Ka value.Q. Why is Ka2 << Ka1 for H2SO4 in water?

Preparation of other acids from H2SO4 –

2MX + H2SO4 → 2HX + M2SO4

(X= F-, Cl-, NO3-) (M= metal)

H2SO4 as dehydrating agent –

Many gases can be dried by passing through H2SO4

C12H22O11 H 2SO 4 /Dehydration→

12C + H20

H2SO4 as oxidising agent – H3PO4, H2SO4, HNO3

Ex.

i. Cu + 2H2SO4 (conc.) → CuSO4 + SO2 + 2H2O

ii. 3S + 2H2SO4 (conc.) → 3SO2 + 2H2O

iii. C + 2H2SO4 (conc.) → CO2 + 2SO2 + 2H2O

Uses – ( King of Chemicals) Manufacturing of ammonium sulphate, superphosphate Petroleum refining Paints and dyes stuff Detergent Metallurgy Storage battery Nitrocellulose products

Lab reagent