3. GROUP 15 ELEMENTS: THE NITROGEN FAMILY
• The p-block elements are put to the right-hand side of the periodic table in
groups from 13 to 18.
• In the p-block elements, the separating electron enters the valence p
subshell.
• In this manner, in these elements, the n p subshell is step by step filled.
• The general valence shell electronic setup of p-block elements is ns2 np1-6.
• The electronic design of helium is 1s2.
• Disregarding the way that helium does not have p orbitals, it is a p-block
element since it takes after that of other p-block elements of the eighteenth
group concerning their physical and compound properties.
• P-block elements are generally non-metals, while the remaining are
metalloids and metals.
4. Elements of group 15 with their atomic number, electronic
configuration, group number and period number.
5. OCCURRENCE:
• Group 15 elements include nitrogen, phosphorus, arsenic,
antimony and bismuth.
• Nitrogen is the real constituent of the world's air, and
records for 78% of it by volume.
• It is the primary member of this group and happens in a
free state as a diatomic gas, N2.
• Minerals of Nitrogen: Indian saltpeter (KNO3) and Chile
saltpeter (NaNO3).
• It is additionally found as the fundamental constituent of
proteins, nucleic acids, amino acids, and catalysts.
6. • The following element in the group is phosphorus.
• It is the eleventh most copious element in the crust of the earth.
• In the consolidated state, it occurs as minerals as phosphates.
• Example: Fluoroapatite 3Ca3(PO4)2.CaF2, Chlorapatite 3Ca3(PO4)2.CaCl2,
and Hydroxyapatite 3Ca3(PO4)2.Ca(OH)2.
• Phosphate groups are constituents of nucleic acids, that is, DNA and
RNA.
• Around 60% of bones and teeth are made out of phosphates.
• Phosphoproteins are available in egg yolk, milk, and bone marrow.
• The rest of the elements of the group, that is, arsenic, antimony, and
bismuth, mostly happen as sulfides.
• Example: Stibnite, Arsenopyrite, and bismuth glance.
7.
8. ELECTRONIC CONFIGURATION OF
GROUP 15 ELEMENTS
The valence shell electronic arrangement for these elements becomes ns2
np3. There are five electrons in the valence shell of these elements.
Because of the precisely half-filled electronic arrangement of the 'n p'
subshell, the elements of this group are genuinely steady and stable.
11. (1)ATOMIC AND IONIC RADII:
(a) The atomic (covalent) and ionic radii (in a particular oxidation state)
of the elements of nitrogen family are smaller than the
corresponding elements of group 14.
Reason: The electrons in the same shell do not screen each other as a
result the effective nuclear charge increases and hence the electrons are
more strongly attracted towards the nucleus.
(b) On moving down the group, the covalent radii increase with increase
in atomic number.
Reason: The considerable increase in covalent radius from N to P is due
to addition of new energy shell but also due to strong shielding effect of
s and p-electrons.
Small increase from As to Bi is due to the poor shielding of the valence
electrons by the d- and f- electrons present in the inner shells.
12. (2) IONIZATION ENTHALPY
These elements demonstrate higher values of ionization enthalpy when
contrasted with group 14 elements.
This is because of their higher atomic charge, smaller nuclear radii, and
stable half-filled electronic setups.
The decrease in the values of ionization energies as we move down the
group is due to gradual increase in the atomic size which reduces the
force of attraction on the electrons by the nucleus.
The successive ionization enthalpies of these elements increase in the
order:
ΔiH1 < ΔiH2 < ΔiH3
13. (3) ELECTRO-NEGATIVITY
• Electro negativity is the inclination of a particle to
pull in a shared pair of electrons more towards itself.
• The electro negativity diminishes bit by bit on
moving down the group because of the increase in
atomic radius.
• But the decrease is not regular.
14. PHYSICAL PROPERTIES: GROUP-15
(1) Metallic character: The elements of group 15 are less metallic
than the corresponding elements of group 14.
On moving down the group, the metallic character increases.
Reason: Due to increased nuclear charge and higher electronegativity
the elements of group 15 are less metallic than 14.
On moving down the electronegativity decreases as a result the valence
electrons are lost more readily .
N, P As, Sb Bi
Non-metals Metalloids Metal
Metallic character increases
15. (2) MELTING AND BOILING POINTS: The melting points
of group 15 elements first increase from nitrogen to Arsenic
and then decrease to antimony and bismuth.
The boiling points however increase regularly as we move
from N to Bi.
Reason: The melting points increase down the group from N to
As due to increase in atomic size.
The unexpected decrease in the atomic size of Sb and Bi even
though the atomic size increases is due to their tendency to
form three covalent bonds instead of five covalent bonds due
to inert pair effect.
16. (3) DENSITY: The density of the elements of group 15
increases regularly from top to bottom.
(4) ALLOTROPY: Except nitrogen and bismuth, all the
elements of this group show allotropy.
Phosphorus exists in three allotropic forms, white, red and
black phosphorus.
Arsenic and antimony exist in two allotropic forms,
yellow and grey.
17. CHEMICAL PROPERTIES
(1) OXIDATION STATES:
Every element of group 15 has 5 electrons in their outermost circle.
They require just 3 electrons to finish their octet setup.
The octet can be accomplished either by picking up 3 electrons or by sharing 3 electrons
by a method for covalent bonds.
Accordingly, the basic negative oxidation state of these elements is - 3. Moving down
the group, the propensity to display - 3 oxidation state diminishes due to decrease in their
electronegativity and ionization enthalpy.
Group 15 elements additionally indicate positive oxidation states of +3 and +5 by
developing covalent bonds.
Because of the inert pair effect, the +5 oxidation state stability diminishes down the
group, while that of +3 oxidation state increments.
Nitrogen has just s-and p-orbitals, yet no d-orbitals in its valence shell. Therefore,
nitrogen can demonstrate a most extreme covalency of 4.
18. Phosphorus and the rest of the elements can display a covalency of five and a
most extreme covalency, additionally called extended covalency of six.
This is conceivable as a result of the nearness of empty d-orbitals in the valence
shell.
Every one of the compounds of group 15 elements, which display a +5 oxidation
state, are covalent.
In case of +3 oxidation state, both ionic and covalent compounds are formed.
19.
20. Nitrogen, in light of its small size, high electro negativity and solid
propensity to shape p pi – p pi numerous bonds, it shows different oxidation
states from - 3 to +5.
COMPOUND OXIDATION STATE
NH3 Ammonia -3
N2H4 Hydrazine -2
NH2OH Hydroxylamine -1
N2 Dinitrogen 0
N2O Nitrous oxide 1
NO Nitric oxide 2
N2O3 Nitrogen trioxide 3
N2O4 Nitrogen tetroxide 4
N2O5 Nitrogen pentoxide 5
22. (1) Bond Angles: All the hydrides of group 15 elements have pyramidal
structures and their bond angles decreases due to increase in size and
decrease in electronegativity.
(2) Boiling points: The boiling points increase regularly as we move
from PH3 to BiH3. However the boiling point of NH3 is higher than
those of PH3 and AsH3.
The abnormally high boiling point of NH3 is due to intermolecular H-
bonding. As we move from PH3 to BiH3 the molecular mass increases as
a result the van der Waals force of attraction increase and boiling point
increase.
(3) Melting Points: The melting point of NH3 is the highest due to
intermolecular H-bonding.
(4) Solubility: NH3 forms H-bonds with water while other hydrides do
not form H-bonds. Therefore NH3 is soluble in H2O while other
hydrides are insoluble.
23. (5) Thermal Stability: The thermal stability of the trihydrides
of group 15 elements decreases as we go down the group due
to decrease in Bond dissociation enthalpy.
(6) Reducing Character: The reducing character of the
hydrides of group 15 elements increases down the group. As
we go down the group the thermal stability decreases so the
tendency to release hydrogen increases and hence reducing
character increases.
Bismuth is the strongest reducing agent.
(7) Basic Nature: The basic character of hydrides decreases
down the group. With the increase in size, the electron density
on the central atom decreases and consequently its tendency to
donate a pair of electrons decreases.
24. 2. Reactivity towards Oxygen.
Oxidation State
OXIDES
N P As Sb Bi
+1 N2O - - - -
Nitrous oxide
+2 NO - - - -
Nitric oxide
+3 N2O3 P4O6 As4O6 Sb4O6 Bi2O3
Dinitrogen trioxide
+4 N2O4 P4O8 - - -
Dinitrogen tetroxide
+5 N2O5 P4O10 As4O10 Sb4O10 -
Dinitrogen pentoxide
25. Oxides of non-metals are acidic, those of metalloids are amphoteric
while those of metals are basic.
Greater the electronegativity of the element more acidic is the
oxide.
Among the oxide of the same element, higher the oxidation
state more is its acidic strength.
N2O < NO < N2O3 < N2O4 < N2O5
Oxides of N2O and NO are neutral
N2O3 > P4O6 > As4O6 > Sb4O6
As4O6 and Sb4O6 are amphoteric while Bi2O3 is basic in nature.
N2O5 > P4O10 > As4O10 > Sb4O10
26. 3. Reactivity towards Halogens.
• On reaction with halogens all the elements of group 15 form trihalides
and pentahalides with the general formula EX3 and EX5.
• 2E + 2X2 → 2EX3
• 2E + 5X2 → 2EX5
• Example: NF3, PF3, AsF3, SbF3 and BiF3 are trihalides.
• All the trihalides of these elements, beside those of nitrogen, are
consistent.
• Phosphorus, arsenic and antimony shape pentahalides considering the
closeness of empty d orbitals in their valence shells.
• Nitrogen does not shape pentahalides as a result of the non-appearance
of a d-orbital in its valence shell. Pentahalides are more covalent than the
relating trihalides.
• Moreover, the covalent character of halides decreases from nitrogen to
bismuth.
27. 4. Reactivity towards Metals
• Each element of group 15 react with metals to frame their binary
compounds demonstrating -3 oxidation state with the general
equation,M3E2.
• Here, M remains for metals while E remains for an element of group
15.
• Example: calcium phosphide, calcium nitride, and so forth.
• 3M + 2E → M3E2
• 3Ca + N2 → Ca3N2
• 6Ca + P4 → 2Ca3P2
• 6Zn + 4Sb → 2Zn3Sb2 (Zinc antimonide)
• 6Mg + 4Bi → 2Mg3Bi2 (Magnesium bismuthide)
28. ANOMALOUS PROPERTIES OF NITROGEN
• Nitrogen, the first member of group 15 elements shows anomalous
behavior and differs from rest of the members of its family.
• The main reasons for this difference are as follows:
(i) Exceptionally small atomic size
(ii) High electronegativity
(iii)High ionization enthalpy
(iv) Absence of d-orbitals in its valence shell
29. 1. pπ ̶ pπ multiple bonds
Nitrogen because of its small size and high electronegativity forms pπ ̶
pπ multiple bonds with itself and with other elements having small size
and high electronegativity (C,O).
The other elements of this group do not form pπ ̶ pπ multiple bonds
because their atomic orbitals are so large and diffused that they cannot
have effective overlapping.
2. Catenation: The elements of group 15 also show the property of
catenation but to a much smaller extent than carbon.
Among the group 15 elements phosphorus has the maximum tendency
for catenation. Nitrogen has little tendency for catenation since N − N
single bond is very weak due to large interelectronic repulsions between
the lone pairs of electrons present on the N-atoms.
30. 3. Reactivity: Nitrogen is inert and unreactive in its elemental state due to
its high bond dissociation enthalpy.
Phosphorus is much more reactive than nitrogen.
4. Maximum Covalency: Nitrogen shows a maximum covalency of 4
due to absence of d-orbitals. Phosphorus and other members of group 15
have empty d-orbitals and hence can exhibit a covalency of either 5 or 6.
5. dπ ̶ pπ multiple bonds: Nitrogen does not form dπ ̶ pπ multiple bonds
due to absence of d-orbitals.
6. Nature of hydrides: Hydride of nitrogen is stable while the hydrides
of other elements are not stable.
7. Nature of halides: Except NF3, the trihalides of nitrogen are unstable.
The trihalides of other elements are stable.
8. Nature of oxides: Nitrogen forms five which are monomeric,
Phosphorus forms three which are dimeric. Arsenic and Antimony form
only two dimeric oxides.
31. DINITROGEN, N2
Commercial preparation
of dinitrogen from air.
Dinitrogen is prepared
commercially from air by
liquefaction and
fractional distillation.
When liquid air is allowed
to distill, dinitrogen
having lower boiling point
(77.2 K) distils over first
leaving behind liquid
dioxygen (boiling point 90
K).
32. LABORATORY METHOD OF PREPARING DINITROGEN
(1) By thermal decomposition of ammonium nitrite.
NH4Cl + NaNO2 NH4NO2 + NaCl
NH4NO2 + Heat N2 + 2H2O
(2) From Ammonium dichromate.
(NH4)2Cr2O7 + Heat N2 + 4H2O + Cr2O3
(3) Very Pure Dinitrogen by Barium azide.
Ba(N3)2 + Heat Ba + 3N2
33.
34. PHYSICAL PROPERTIES:
(1) Dinitrogen is a colorless, tasteless, non-toxic gas.
(2) It has two stable isotopes: 14N and 15N.
(3) It is very slightly soluble in water.
(4) It has low freezing point.
(5) It is adsorbed by activated charcoal.
(6) Nitrogen undergoes condensation to form a colourless liquid which
on solidification results in the formation of snow like mass.
35. CHEMICAL PROPERTIES:
1. Action of litmus. It is neutral towards litmus.
2. Active metals.
6Li + N2 + Heat 2Li3N
3Mg + N2 + Heat Mg3N2
3. Non-metals.
N2 + 3H2 2NH3
N2 + O2 2NO
4. Calcium carbide.
CaC2 + N2 CaCN2 + C
CaCN2 + 3H2O CaCO3 + 2NH3
cyanamide (CN22-) anion
1273 K
36. USES:
1. Manufacture of nitric acid, ammonia, calcium cynamide.
2. Provides inert atmosphere in iron and steel industry.
3. In filling electric bulbs to reduce the rate of volatilization of the
tungsten filament.
4. Dinitrogen gas-filled thermometers are used for measuring high
temperatures.
5. Liquid dinitrogen is used as a refrigerant to preserve biological
materials.
37. AMMONIA (NH3)
• Preparation:
(1) By heating Ammonium salts with a strong base.
(NH4)2SO4 + 2NaOH + Heat 2NH3 + 2H2O + Na2SO4
NH4Cl + KOH NH3 + H2O + KCl
2NH4Cl + Ca(OH)2 + Heat 2NH3 + 2H2O + CaCl2
(2) By the action of water on metal nitrides.
Mg3N2 + 6H2O 3Mg(OH)2 + 2NH3
AlN + 3H2O Al(OH)3 + NH3
Drying of Ammonia gas is done by quick lime (CaO) and not by using
concentrated H2SO4 or P2O5, since it reacts.
38. HABERS PROCESS FOR AMMONIA
• The preparation of ammonia by Haber's procedure includes direct mixing of nitrogen and hydrogen.
39. This reaction is, (a) exothermic, (b) reversible, and (c) proceeds with a decrease in
volume. As per the Le Chatelier's principle, the ideal conditions for the arrangement
of ammonia are:
Low Temperature: The temperature ought to stay as low as would be prudent, (in
spite of the fact that at strangely low temperatures, the rate of reaction turns out to
be moderate). It has been found that the temperature, which upgrades the yield of
ammonia for the reaction, is most extreme at about 500°C.
High Pressure: Since Haber's procedure continues with abatement in volume, it is
supported by high pressure. In real practice, a pressure of 200 - 900 atmospheres is
utilized.
Catalyst: An impetus is generally utilized to expand the speed of the reaction.
Finely isolated iron containing molybdenum or alumina is utilized as an impetus.
Molybdenum or alumina (Al2O3) goes about as a promoter and builds the
productivity of the impetus. A blend of iron oxide and potassium aluminate has
been found to work all the more adequately.
40. PHYSICAL PROPERTIES OFAMMONIA
● Ammonia is a gas that has no color.
● It has a sharp pungent odour having a soapy taste. At the point when
inhaled suddenly, it attacks the eyes bringing tears.
● It is lighter than air and is in this way gathered by the descending
displacement of air.
● It is very soluble in water
● It can be effortlessly melted at room temperature by applying a pressure of
around 8-10 atmospheres.
● Liquid ammonia bubbles at 239.6 K (- 33.5°C) under one-atmosphere
pressure. and is subsequently utilized as a part of refrigeration plants of ice
making machines.
41. STRUCTURE OFAMMONIA
The ammonia particle is shaped because of the overlap of orbitals of three hydrogen
and three sp3 hybrid orbitals of N in the structure as central atom.
The fourth sp3 hybrid orbital is involved by a lone pair.
This provides a trigonal pyramidal shape to ammonia particle.
The H-N-H bond edge is 107.3°, which is somewhat not exactly the tetrahedral edge
of 109°28.
42. Chemical Properties of Ammonia:
1. Thermal Stability
2. Combustibility
3. Basic Character
44. 6. With Carbon Dioxide (formation of urea)
7. Action with Heavy Metal Ions
45. Uses of Ammonia:
1. In the production of urea and rayon
2. In the production of composts, for example, ammonium nitrate, urea
diammonium phosphate, ammonium sulfate and so on.
3. As a refrigerant, in ice plants.
4. In the furniture industry, as a purging operator for furniture and glass
surfaces.
5. In the production of nitric acid by Ostwald's procedure.
6. In the production of sodium carbonate by Solvay's procedure.
46. TESTS OF AMMONIA:
● The ammoniacal odor of ammonia is effectively perceivable having a trademark pungent
smell.
● Ammonia turns wet red litmus blue and moist turmeric paper brown in color.
● A glass bar dunked in concentrated HCl when conveyed near ammonia, causes thick
white exhaust.
● When added to a solution of copper sulphate, ammonia turns the solution deep blue.
● 4 NH3 + CuSO4 → [Cu(NH3)4]SO4
● When added to Nessler's reagent (basic arrangement of K2[HgI4] ammonia gives
precipitate brown in color.
• NH4+ + 2[HgI4]2− + 4OH− → HgO·Hg(NH2)I ↓ + 7I− + 3H2O
54. (c) With Metals: Reaction with more electropositive than hydrogen.
Mg + 2 HNO3(very dil.) Mg(NO3)2 + H2
4 Zn + 10 HNO3 (cod & dil.) 4 Zn(NO3)2 + 3 H2O + NH4NO3
4 Zn + 10 HNO3 (hot & dil.) 4 Zn(NO3)2 + 5 H2O + N2O
Zn + 4 HNO3 (conc.) Zn(NO3)2 + 2 H2O + 2 NO2
Reaction with less electropositive metals than hydrogen.
Cu + 4 HNO3(conc.) Cu(NO3)2 + 2 H2O + 2 NO2
3 Cu + 8 HNO3(dil.) 3 Cu(NO3)2 + 4 H2O + 2 NO
55. Reaction with noble metals.
Noble metals like gold and platinum do not react with conc. HNO3.
These metals dissolve in aqua regia ( 3 parts of conc. HCl + 1 part of
conc. HNO3) forming their respective chlorides.
3 HCl + HNO3 NOCl + 2 H2O + 2 Cl
Au + 3 Cl AuCl3
Pt + 4 Cl PtCl4
59. WHITE PHOSPHORUS
It is obtained when phosphate rock is heated with coke at 1773 K.
2 Ca3(PO4)2 + 6 SiO2 6 CaSiO3 + P4O10
P4O10 + 10 C P4 + 10 CO
Properties: 1. It is a soft, translucent waxy white solid.
2. It is very poisonous. 3. Melting point 317 K and boiling point 553 K.
4. It is insoluble in water.
5. P4 + 3 NaOH + 3 H2O PH3 + 3 NaH2PO2.
60. RED PHOSPHORUS
It is obtained by heating white phosphorus at 573 K in an inert atmosphere.
P4(s) P4(s)
White Red
It is a hard crystalline solid.
It is non-poisonous in nature.
It is insoluble in water as well in organic solvents such as CS2.
It is denser than white phosphorus.
Being polymeric, it is less reactive than white phosphorus.
It does not react with caustic alkalies.
61. BLACK PHOSPHORUS
It has two forms: α - black phosphorus and ꞵ - phosphorus.
α - black phosphorus is formed when red phosphorus is heated in a
sealed tube at 803 K.
Red phosphorus α - black phosphorus
803 K sealed
tube
62. PHOSPHINE, PH3
Preparation:
1. Calcium phosphide treated with water or dil HCl.
Ca3P2 + 6 H2O 3 Ca(OH)2 + 2 PH3
Ca3P2 + 6 HCl 3 CaCl2 + 2 PH3
2. From phosphorus acid.
4 H3PO3 3 H3PO4 + PH3
3. From phosphonium iodide.
PH4I + NaOH NaI + H2O + PH3
4. Laboratory method from white phosphorus.
P4 + 3 NaOH + 3 H2O PH3 + 3 NaH2PO2
64. USES OF PHOSPHINE:
(i) As Holme’s signals in deep seas and oceans for signalling danger
points to steamers. Containers containing a mixture of of calcium
phosphide and calcium carbide are pierced and thrown into the sea.
(ii) For the production of smoke screens. Calcium phosphide reacts with
water producing phosphine which burns in air to give clouds of
P4O10 which act as smoke screens.
66. Preparation: It is prepared by the action of thionyl chloride on white
phosphorus.
P4 + 8 SOCl2 4 PCl3 + 4 SO2 + 2 S2Cl2
It is also prepared by heating white phosphorus in a current of dry
chlorine.
P4 + 3 Cl2 4 PCl3
Properties: (i) Action of water.
PCl3 + 3 H2O H3PO3 + 3 HCl
3 CH3CH2OH + PCl3 3 CH3CH2Cl + H3PO3
3 CH3COOH + PCl3 3 CH3COCl + H3PO3
67. PHOSPHORUS PENTACHLORIDE, PCl5
When white phosphorus reacts with excess of dry chlorine, phosphorus
pentachloride is produced.
P4 + 10Cl2 → 4 PCl5
It can also be prepared by the reaction of SO2Cl2 and phosphorus.
P4 + 10 SO2Cl2 → 4 PCl5 + 10 SO2
Chemical Properties:
PCl5 + H2O → POCl3 + 2HCl
POCl3 + 3H2O → H3PO4 + 3HCl
PCl5 → PCl3 + Cl2
2Ag + PCl5 → 2AgCl + PCl3
C2H5OH + PCl5 → C2H5Cl + POCl3 + HCl
74. GROUP 16 ELEMENTS: OXYGEN FAMILY
OXYGEN (O)
SULPHUR (S)
SELENIUM (Se)
TELLURIUM (Te)
POLONIUM (Po)
75. GROUP 16 ELEMENTS: THE OXYGEN FAMILY
(CHALCOGENS)
• Group 16 in the p block is the first group which has no stable metallic elements.
76. OCCURENCE
The elements Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium
(Po) comprise of the 16th vertical column or VIA group elements in the currently
used long type of periodic table.
Oxygen and Sulphur are typical non-metals, Selenium and Tellurium metalloids and
Polonium is metal.
77. The initial four elements of the group are together termed as chalcogens or ore-
forming elements, on the grounds that an extensive number of metal ores are found
in the earth's crust as sulphiides or oxides.
• Sulphate Ores:
• Includes gypsum, Epsom salt MgSO4.7H2O, CaSO4.2H2O, and barytes, BaSO4.
• Sulphide Ores:
• Includes galena (PbS), zinc blende (ZnS), and copper pyrites (CuFeS2).
• Sulphur can also be seen in many organic substances like mustard, eggs, seeds,
onion, wool, garlic, and hair.
• Selenium and tellurium are found in sulphides ores as metal selenides and
tellurides.
• Polonium is a radioactive element.
79. ATOMIC PROPERTIES
1. Atomic and Ionic Radii:
The atomic and ionic radius increases as we move from oxygen to
polonium.
The atomic and ionic radii of the elements of group 16 are smaller than
those of the corresponding elements of group 15.
Reason: Increased Nuclear Charge
2. Ionization enthalpy:
The first ionization enthalpies of the elements of group 16 are
unexpectedly lower while their second ionization enthalpies are higher
than those of the corresponding elements of group 15.
Reason: Stable electronic configuration .
80. 3. Electron gain enthalpy:
Group 16 have large negative electron gain enthalpies next to
halogens.
Oxygen has the least negative value in this group due to its small size.
4. Electronegativity:
Group 16 have higher values of electronegativity than the corresponding
elements of group 15.
Electronegativity decreases down the group.
81. PHYSICAL PROPERTIES: GROUP 16
1. Non-metallic/metallic character.
Because of high ionization enthalpy values the elements of group 16 are
less metallic, moving down the group the metallic character increases.
Oxygen is the most non-metallic, Sulphur is a typical non-metal.
Se and Te are metalloids and are semiconductors.
Polonium is metallic but is radioactive.
2. Melting and Boiling Points.
The melting points, boiling points and densities increase regularly as we
go down the group upto tellurium.
Melting and boiling point of Polonium is lower than those of tellurium.
82. 3. Elemental state:
Oxygen exists as diatomic gas at room temperature while other elements exist
as octaatomic solids.
Reason: Small size and high electronegativity and forming pπ ̶ pπ double
bonds.
4. Catenation:
Because of stronger S ─ S bonds as compared to O ─ O bonds, sulphur has
a stronger tendency for catenation than oxygen.
5. Allotropy:
• Each one of the elements of group 16 display allotropy. Oxygen has two
allotropes: Oxygen & Ozone.
• Sulphur exist as many allotropic forms but only two of them are stable:
Rhombic Sulphur & Monoclinic Sulphur
• Selenium and Tellurium is found in both amorphous and crystalline forms.
84. TRENDS IN CHEMICAL REACTIVITY
1. FORMATION OF HYDRIDES:
All the elements of group 16 form hydrides of the general formula H2E.
H2O, H2S, H2Se, H2Te, H2Po.
(a) Bond angle decreases from H2O to H2Te.
(b) Physical state: H2O (liquid) while other are gases.
(c) Acidic character: Increases from H2O to H2Te.
(d) Thermal stability: Decreases from H2O to H2Te.
(e) Reducing character: Increases from H2O to H2Te.
85. 2. FORMATION OF OXIDES:
All the elements of this group form two types of oxides.
EO2 and EO3.
Elements Dioxide (EO2) Trioxide (EO3)
O ̶ O3
S SO2 SO3
Se SeO2 SeO3
Te TeO2 TeO3
Po PoO2 ̶
86. Simple Oxides
• Simple oxides are oxides that carry only that number of oxygen atoms
as is allowed by the normal valency of its metal.
• Example: H2O, MgO & Al2O3.
Mixed Oxides
• Two simple oxides combine to form mixed oxides.
• Example: Lead dioxide (PbO2) and lead monoxide (PbO) together
form the mixed oxide Red lead (Pb3O4).
• Ferric oxide (Fe2O3) and ferrous oxide (FeO) together form the mixed
oxide Ferro-ferric oxide (Fe3O4).
87. ACIDIC OXIDE : An oxide that reacts with water to give an acid is called as
Acidic Oxide. Example: Oxides of non-metals, such as SO2, CO2, SO3, Cl2O7, P2O5,
& N2O5, or oxides of metals with high oxidation states, such as CrO3, Mn2O7,
&V2O5 are acidic in nature.
BASIC OXIDE: An oxide that reacts with water to give a base is called as a basic
oxide.
• Example: Oxides of most metals, such asNa2O, CaO, BaO, are basic in nature.
AMPHOTERIC OXIDES: Some metallic oxides display dual behavior that is
they show both the characteristics of acid as well as base. These metal oxides are
known as Amphoteric Oxides. They can react with both alkalis as well as acids.
NEUTRAL OXIDES: Neutral Oxides, as name suggests do not exhibit any
tendency to form salts either with acids or bases.
• Example: Nitrous oxide and Carbon monoxide are neutral oxides.
88. ANOMALOUS BEHAVIOUR OF OXYGEN:
Reason: Small size, Higher electronegativity, Non-availability of d-orbitals.
Points of differences:
(i) Physical state: Oxygen gas others solid.
(ii) Atomicity: Oxygen diatomic, others octaatomic.
(iii) Oxidation states: Oxygen shows: − 2, + 1 and + 2.
Other elements show +2, +4 and +6.
(iv) Nature of compounds: Oxygen because of high electronegativity is more
ionic in its compound and also form covalent.
Others mainly form covalent.
(v) Multiple bonds: Oxygen forms multiple bonds but other elements not.
(vi) Magnetic nature: Dioxygen is paramagnetic while others are diamagnetic.
89. DIOXYGEN (O2)
It is prepared by the following methods:
1. Decomposition of oxygen-rich compounds.
2 KClO3 2 KCl + 3 O2
2 KMnO4 K2MnO4 + MnO2 + O2
2HgO → 2Hg + O2
2Ag2O → 4Ag + O2
2PbO2 → 2PbO + O2
2BaO2 → 2BaO + O2
93. PHYSICAL PROPERTIES OF DIOXYGEN
● Dioxygen is a tasteless, colorless and scentless gas.
● It is marginally heavier than air.
● It is marginally soluble in water. This little amount of dioxygen dissolved is
quite adequate to support the aquatic and marine life.
● Under pressure, it can be condensed to a light blue fluid by compacting the gas
at 90K. It can likewise be solidified into a bluish white solid at 55K.
94. CHEMICAL PROPERTIES:
1. Action on litmus: Neutral
2. Metals: Reacts with all metals (except noble metals such as gold and
platinum)
3. Non-Metals: H2, N2, S, C, P4.
4. Compounds:
SO2 + O2 SO3
NH3 + O2 NO + H2O
HCl + O2 H2O + Cl2
CS2 + O2 CO2 + SO2
ZnS + O2 ZnO + SO2
CH4 + O2 CO2 + H2O
95. USES OF DIOXYGEN
● The fundamental significance of dioxygen lies in its support to
key procedures. For Example: respiration and combustion.
● Dioxygen blended with carbon dioxide or helium is utilized for
artificial respiration.
● It is utilized as a part in manufacturing many metals.
● It is utilized as a part of oxy-acetylene welding and metal
cutting.
● It is utilized to oxidize ammonia in the nitric acid preparation.
● Oxygen barrels are broadly utilized as a part of healing facilities,
high-height flying and in mountaineering.
97. PREPARATION OF OZONE
Ozone is set up by passing a silent electric discharge through dry, unadulterated, and
cold oxygen in an extraordinary device called the ozoniser.
This way concentration of approximately 10% of ozone can be obtained.
The ozone formation is an endothermic process; it must be completed at high
temperature. This is the reason; it is set up by method of silent electric discharge.
99. USES OF OZONE:
● Ozone is used as an antiseptic.
● Ozone is moreover used as a disinfectant. For Example, it is used
in filtration of drinking water.
● It is used as a delicate dying agent for fading oils, starch, ivory,
wax, flour, and delicate textures
● It is used as a part of the produce of potassium permanganate
from potassium manganate.
● It is moreover used as a part of the creation of manufactured
camphor and artificial silk.
101. SOURCES OF SULPHUR
Sulphur can be found from the accompanying sources:
(1) Extraction from underneath the earth outside layer - this is the most imperative
source.
(2) From natural gas - this is the second most imperative source. The characteristic
gas is found in a place called Lacq, in southern France.
(3) From different procedures - example, as a by-result of the refining of petroleum
and decontamination of crude coal gas.
PHYSICAL PROPERTIES : Sulphur is a yellow solid and is insoluble in water
yet soluble in toluene (methyl benzene) and carbon disulphide. Sulphur is a non-
metal and therefore a poor channel of electricity and heat.
The boiling point of sulphur is 444oC. At the point when sulphur vapor is
consolidated, a fine powder, which shapes a pattern resembling flower is gotten -
this is called 'Flower of Sulphur’.
102.
103. ALLOTROPES OF SULPHUR
• Rhombic sulphur takes shape at a temperature beneath 96oC, while
monoclinic sulphur solidifies at a temperature over 96oC.
• The temperature, 96oC is known as the transitional temperature between
the two structures.
104. RHOMBIC SULPHUR: (i) Yellow, translucent crystals. (ii) Melting point of 114oC
(iii) Density of 2.08 gcm3 (iv) Stable at temperatures below 96oC
105. MONOCLINIC SULPHUR: (i) Transparent, amber crystals (ii) Melting point of
119oC (iii) Density of 1.98gcm3 (iv) Unstable at temperatures below 96oC, reverting
to rhombic form.
106. PLASTIC SULPHUR: It is obtained by pouring molten sulphur into
cold water a soft rubber like mass called plastic sulphur is obtained.
111. USES:
● Fertilizers like ammonium sulfate and superphosphate
● Dyes, shades, and paints
● Explosives, for example, TNT
● Other imperative chemicals like hydrochloric acid, phosphoric acid,
nitric acid, and sodium carbonate
● It is utilized as a part of the refining of petroleum
● As a pickling agent
● As a laboratory agent, and an oxidizing and dehydrating agent
112. GROUP 17 ELEMENTS: HALOGEN FAMILY
FLUORINE (F)
CHLORINE (Cl)
BROMINE (Br)
IODINE (I)
ASTATINE (At)
114. ATOMIC PROPERTIES:
(1)Atomic and Ionic radii: Smallest radii in their respective
periods.
(2)Ionization enthalpy: Halogens having high ionization
enthalpy next to 18th group elements in each period.
(3)Electron gain enthalpy: Halogens have highest negative
electron gain enthalpy in their respective periods.
Chlorine has the highest negative electron enthalpy than
fluorine.
(4) Electronegativity: Halogens are highly electronegative
elements.
115. PHYSICAL PROPERTIES
1. Atomicity: All halogens exist as diatomic molecules.
2. Physical state: F2 and Cl2 are gases at room temperature, Br2 is a
liquid whereas I2 is a solid.
Melting and boiling points increases with increase in atomic number.
3. Non-metallic character: Due to their high ionization enthalpies and
high electronegativities halogens are non-metallic in nature.
Non-metallic character decreases as we move down the group.
4. Bond dissociation enthalpy: Fluorine is very small and hence the
electron-electron repulsion is very high.
Cl2 > Br2 > F2 > I2
116. 5. Colour:
Fluorine: Pale yellow , Chlorine: Greenish yellow
Bromine: Reddish brown , Iodine: Deep violet
6. Nature of bonds:
As the electronegativity decreases from fluorine to iodine the
to form ionic bonds decreases while the tendency to form covalent bond
increases.
117. CHEMICAL PROPERTIES:
1. Oxidation states:
Fluorine shows a negative oxidation state of − 1 except in HOF where
it shows + 1.
Other halogens in addition to − 1 and + 1 also shows +3, +5 and +7.
Cl also shows +4 (ClO2) and +6 (Cl2O6) while Br shows +4 (BrO2).
2. Oxidising power:
Halogens have a strong tendency to gain 1 electron hence they act as
strong oxidizing agents.
Oxidising power decreases from F2 to I2.
F2 can oxidise Cl−, Br− and I− from their solutions.
Cl2 ???
Br2 ???
118. 3. Reactivity towards Hydrogen:
HF, HCl, HBr, HI
a. Physical state: HF is liquid others are gases.
b. Melting & boiling points: HF has the highest boiling point due to
extensive intermolecular hydrogen bonding.
c. Bond length: HF < HCl < HBr < HI
d. Bond strength: HF > HCl > HBr > HI
e. Thermal stability: HF > HCl > HBr > HI
f. Acid strength: HF < HCl < HBr < HI
g. Reducing power: HF < HCl < HBr < HI
Greater the bond dissociation energy more stable is the halogen and
hence weaker is the reducing agent.
125. INTERHALOGEN COMPOUNDS (X needs to be less electronegative than X')
Type Formula Structure
XX' ClF Linear
BrF Linear
IF Linear
BrCl Linear
ICl Linear
IBr Linear
XX'3 ClF3 Bent T-shaped
BrF3 Bent T-shaped
IF3 Bent T-shaped
ICl3 Bent T-shaped
XX’5 IF5 Square pyramidal
BrF5 Square pyramidal
ClF5 Square pyramidal
XX’7 IF7 Pentagonal bipyramidal
129. 1. Monoatomic nature.
2. The atomic radii of noble gases are by far the largest in
their respective periods. Because noble gases have only
van der waals radii which are larger than covalent radii.
3. Ionisation enthalpy: Highest in their respective periods.
4. Electron gain enthalpy: Noble gases have completely
filled subshells. As a result they have positive electron
gain enthalpy.
5. Melting and boiling points are very low.
130. CHEMICAL PROPERTIES:
The noble gases have high ionization enthalpies and positive electron
gain enthalpies so they don’t have tendency to lose nor to gain electrons.
In 1962 Neil Bartlett observed that platinum hexafluoride (PtF6) reacts
with dioxygen and the first ionization enthalpy of xenon is fairly close to
that of O2 molecule.
So compounds of Xenon with fluorine and oxygen have been prepared.
