P block elements (Group 15th and 16th)
Physical properties:
Properties N – Family (15th group) O – Family (16th group)
Physical state 1. N 2 – Diatomic gas 1. O2 – Diatomic gas
2. P4 – Soft vaxy solid 2. S8 – octaatomic solid
3. (As4, Sb4) - solid 3. (Se, Te, Po) – Solid
4. Bi – metallic solid
Ionization Top – Bottom Top – Bottom
energy 1. Decreases 1. Decreases
2. I.E. of N is greater than
corresponding elements of O
family due to stable 2p3
configuration.
Electronegativity Top – Bottom Top – Bottom
1. Decreases 1. Decreases
2. It increases in period 2. Oxygen is 2nd most E.N.
element in P.T.
Electron affinity Top – Bottom Top – Bottom
1. It generally decreases 1. It generally decreases
2. But E.A of Nitrogen is less 2. But E.A of sulphur is less
than that of Phosphorus than that of oxygen
because of small size of because of small size of
Nitrogen oxygen
3. P>As>Sb>Bi>N 3. S>Se>Te>Po>O
Atomic radius & Top – Bottom Top – Bottom
Ionic radius 1. Increases 1. Increases
Metallic Top – Bottom Top – Bottom
character 1. Increases (due to decrease in 1. Increases
I.E. and increase in atomic 2. (O, S) – Non metal
size. 3. (Se, Te) – semi metal
2. (N, P) – Non Metal 4. Po - Metal
3. (As, Sb) – Semi metal
4. Bi - Metal
Density Top – Bottom Top – Bottom
1. Increase ( as mass dominates 1. Increase ( as mass
over volume or size) dominates over volume or
size)
Electrical and Top – Bottom Top – Bottom
thermal 1. Increases 1. Increases
conductivity 2. (N, P) – Non conductor 2. (O, S) – Non cond.
3. (As, Sb) – Semi conductor 3. (Se, Te) – Semi cond.
4. Bi – good conductor 4. Po – good cond.
Melting Point Top – Bottom Top – Bottom
1. Generally increases 1. Increases
2. N<P<As>Sb>Bi 2. In period M.P. decreases
3. From As to Bi M.P. decreases 3. But M.P. of Te>Bi
due to decrease in the
strength of metallic bonding.
Boiling Point Top – Bottom Top – Bottom
1. Increases 1. Increases
Allotropy 1. Except Nitrogen all other
elements in the group show
allotropy
Chemical Properties:
Properties N- Family O- Family
Anamolus 1. Due to smaller size, High I.E., 1. Same
behavior High E.N. and unavailability of 2. Same
vecant d orbital : Nitrogen 3. O-O < S-S (Bond enthalpy)
differs from rest of the elements 4. O = O > S = S (due to small size
in some property &interelectronic repulsion bond
2. Nitrogen has ability to form Pπ- energy)
Pπ multiple bond with itself or 5. Oxygen exist as O2 but sulphur
with other smaller atoms like exist as S 8 due to high S-S B.E.
carbon & oxygen N = O 6. Same for sulphur (Ex- S02, S03,
3. N-N < P-P ( bond, enthalpy) H2SO4etc)
4. N ≡ N > P ≡ P (due to small size 7. Oxygen can not form dπ – pπ
&interelectronic repulsion bond bond
energy)
5. Nitrogen exist as N2 but
phosphorus exist as P4 due to
above reason
6. Phosphorus has a ability to form
dπ – pπ bond and this ability
decreases down the group ( ex –
H3PO4)
7. Nitrogen can not form dπ – pπ
bond
Oxygen state 1. Common O.S. are +3, -3, +5 1. -2, +4, +6
2. All forms +3, +5O.S. I covalent 2. All form +4 and +6 O.S. in
compound covalent compound.
3. Nitrogen form +5 O.S. by 3. +6 is more covalent than +4
formation of coordination bond. 4. Oxygen can not form +4 and +6
(Ex: HNO3 , N2O5) state due to unavailability of d
4. +5 O.S. is more covalent than +3 orbitals.
5. Stability of +5 O.S. decreases 5. All forms -2 O.S in hydrides
due to inert pair effect. (H2O, H 2S)
(The inert pair effect is the 6. Oxygen generally forms -2 in
tendency of the electrons in the ionic oxides, -1 in peroxides, -
outermost s atomic orbital to 1/2 in super oxides and +1 or +2
remain un-ionized or unshared in oxygen fluorides.
in compounds of post-transition
metals.)
6. All form -3 O.S. in hydride
(Except Bi) , Ex: NH3, PH3, AsH3,
SbH3
7. All form -3 O.S. in binary
compound with metals. (Ex:
Li3N, Mg3N2, AlN, Ca3P 2, Ca3As2,
Zn3Bi2.
Reactivity 1. All form binary hydrides of the 1. All form binary hydrides of the
with formula EH3 formula EH2
Hydrogen 2. Ex: NH3(Ammonia), 2. H2O(water), H2S(hydrogen
PH3(Phosphine), AsH3(Arsine), sulphide), H2Se(hydrogen
SbH3(Stibine), BiH3( Bismuthene) selenide), H2Te(Hydrogen
3. EH3 : SP3hybridised ( 3 BP + 1 LP) teuride)
, Pyramidal shape 3. H2E: SP3hybridised (2 BP + 2 LP) ,
bent (V) shape.
4. DTG from top to bottom Bond 4. DTG from top to bottom Bond
angle, Bond enthalpy, basic angle, Bond enthalpy, basic
nature, stability, dipole moment nature, stability, dipole moment
decreases. decreases.
(NH3>PH3>AsH 3>SbH3>BiH3) (H2O<H2S<H2Se<H2Te)
5. Acidic nature, bond length, 5. Acidic nature, bond length,
reducing nature, volatile reducing nature, volatile
character increases. character increases.
(NH3<PH3<AsH 3<SbH3<BiH3) (H2O>H2S>H2Se>H2Te)
6. Boiling point 6. Boiling point and Melting point
(SbH3>NH3>AsH3>PH3) (H2O>H2Te>H2Se>H2S)
7. Melting point
(NH3> SbH3>AsH3>PH3)
Reactivity 1. All form oxide of the formula 1. All form EO2, EO3
with Oxygen E2O3, E2O 5 2. Ex: SO2, SeO2, TeO2, SO3, SeO3,
(Oxide) 2. (N2O3, N2O5), (P4O6, P4O10), TeO3
(As2O 3, As2O5), (Sb 2O3, Sb2O5), 3. EO3 more acidic EO2.
(Bi2O3, …x….) 4. T-B: Acidic decreases.
3. E2O5 is more acidic than E2O3. 5. Reducing nature of EO2 from
4. Acidic nature of oxides Top to bottom decreases.
decreases.(As E.N. decreases)
5. Stability of E2O5 decreases from
T-B.
Halides 1. Formula: EX5, EX3 (Except NX5) 1. Formula: EX2, EX4, EX6
2. EX3: SP3 hybridized (3 BP + 1LP), 2. EX6 is found in fluorides only
Pyramidal structure and its stability decreases from
3. Bond angle: NF3<NCl3<NBr3 Top to Bottom. (Ex: SF6, SeF 6etc)
4. Bond angle: NX3>PX3>AsX3>SbX3 3. EX6: SP3d2, Octahedral shape
5. EX5: SP3d hybridized (5 BP), 4. EX4 is more common in S, Se,
Trigonal bi pyramidal (TBP) Te: SF4 (Gas), SeF4 (Liquid), TeF4
6. Stability of halide decreases as (Solid) [ See saw structure ].
F>Cl>Br>I 5. Oxygen does not form OX 4, OX6
7. T-B, For same element Covalent due to absence of d orbital.
nature decreases (F<Cl<Br<I) 6. Oxygen form dihalide OF2, OCl2,
8. EX5 is more covalent than EX3 OBr2 etc. (SP3 hybridized – (2BP
9. Only BiF3 – Ionic others all are + 2 LP) Bent / V / Angular shape)
covalent. 7. Mono halide are also form with
10. These halides are soluble in S, Se etc in dimeric form. (Ex:
water due to hydrolysis O2F 2, S2Cl2, Se2Cl2.
8. These halides are soluble in
water due to hydrolysis
Group 17th and 18th
Property Halogen family Noble Gases
Electronic configuration 2
ns np 5 2
ns np 6
Atomic Radius T B TB
Ionisation Enthalpy T B T B
Electron Gain Enthalpy T B T B
F cl Br I
cl highest EA
Electronegativity T B T B
Enthalpy of Dissociation cl2 Br2 F2 I 2
MP/BP T B T B
State F2 gas All are mono atomic gases
cl2
Br2 liq .
I 2 solid
Colour F2 Yellow Colourless test less
cl2 Greenish Yellow
Br2 Red
I 2 Voilet Colour
Standard Reduction Potential I 2 Br2 cl2 F2
Other cl2 Suffocating
Chemical Properties
1. Oxidation State: + 1, + 3, + 5, + 7.
Generally = – 1.
2. Oxidizing nature: F2 cl2 Br2 I 2 .
3. Hydration Energy (Smaller the ion higher HE)
F cl Br I .
4. Anomalous behavior of Fluorine
Due to smaller size.
Unavailability of d orbital.
Highest EN.
5. Reactivity towards hydrogen.
All form acid
HF, HCl, HBr, HI
Acitic strength strength HI HBr HCl HF .
H X dissociation enthalpy: H F H Cl H Br H I .
Stability: H F H Cl H Br H I .
6. Reactivily towards oxygen:
Fluorine makes only O2 F2 & OF2 in which OF2 thermally stable at 298 K.
Cl forms Cl2O, ClO2 , Cl2O6 , Cl2O7 ClO2 used as bleaching agent.
Br form Br2O, BrO2 , BrO3 .
I form I 2O4 , I 2O5 , I 2O7
Very good O. A
7. Reactivity towards metals:
Form metla halides
Ex. Mg Br2 , MgCl2 .
8. Reactivity towards halogens:
AB, AB3 , AB5 , AB 7 .
A Larger Size.
B Smaller Size.
F2 Preparation
On electrolysis:
2 H 2e H2 g
2F H 2 F2 g
2 H 2 F
H 2 F2 g
F2 g must be free from HF (due to corrosiveness).
To make it free from HF we pass NaF which absorb HF.
There should not be no moisture present in vessel otherwise F2 with react with H 2 O .
4 HF O2 .
2 F2 2 H 2O
If aqueous solution of NaF or KF is taken then no. F2 is farmed.
1
H 2O O2 2 H 2e E = –1.23 sop.
2
1
F F2 sop = – 2.87.
2
1
K 2 MnFb SbF5
2 K SbFb MnF3 F2 .
2
Reaction: (Oxidizing agent)
O2 F2 SF6 HF
NaF X 2
aX
N
O2 NaF F2 OF2 NaF H 2O
NaOH NaOH
AgF
N 2 HF
Chlorine Cl2 :
1. Heating chlorides
H MnO2 2 X
X 2 Mn 2 .
2. Deacon’s process
CuCl2
O2 HCl
723 K
Cl2 .
3. Electrolytic process
NaX aq .
Na X .
Na e
Na cathode.
2 X
X 2 2e anode.
Note: Chlorine water turns blue litmus to red but after some time it becomes colourless.
HOCl HCl .
Cl2 H 2O
O (Nasent Oxygen)
HOCl
Oxidizing property
HCl
NaCl + NaOCl
Ca OCl 2
Bleaching Powder
Ca OH 2 cl2
Powerful bleaching NaCl NaClO3
agent
Fe2 SO4 3 HCl .
(i) FeSO4 H 2 SO4 Cl2
x2 Cl x Br , I .
(ii) x Cl2
Cl2 H 2O 2 HCl O .
coloured O colourless
subtance substnace
SO2 bleaching action property is temprorary because reduction occue here.
SO2 H 2O
H 2 SO4 H
.
Reduce
Cl2 bleach by oxidation.
SO2 bleach by reduction.
Bromine Br2 .
Preporation:
MgCl2 Br2 .
MgBr2 Cl2
x Cl2
x2 Cl x
Br , I .
Properties:
Fairly soluble in water.
Forms hydrate like Cl2 .
Br2 .8 H 2O, Cl2 .8 H 2O .
Iodine I 2 .
Preparation:
x Cl2 or F2 or Br2
x2 Cl / F / Br .
Properties:
Sparingly soluble in water.
Readily soluble in Na S K iodide.
KI I 2 KI 3 .
NaI 3 .
NaI I 2
Reactions
kIO3 Cl2 S 4O62 2 I
O 2
3
o
2
S
yp
H
I2 NH 3 q
NH 3 q
KIO3 Br2 NI 3 NH 3 HI
C2 H 5OH KoH
NCOOK CHI 3 KI
Halogen acids (HCl, HBr, HI)
Preparation
1. 2 HCl .
H 2 Cl2
Pt
H 2 Br2 2 HBr
Pt
H 2 I 2
450C
2HI .
2. By RXN with P4.
P4 Br2 I 2
PBr3 PI3 .
PBr3 PI3 H 2O
H 3 PO3 HBr .
3. Passing H 2 S / SO2 into solution of halogen.
HX S .
H 2 S X 2
HX H 2 SO4 .
nSO2 X 2 H 2O
Cl H 3O
HCl H 2O l
GROUP 17
Hydrogen Halides
Action of SiO2/Glass
SiO2 HF H 2 SiO3 H 2 SiF6 .
SiF4 2 H 2O
Na2 SiF6 H 2O .
Na2 SiO3 HF
Action of AgNO3 Solution
Agf [Soluble]
AgCl [White ppt. insolube in HNO3, but soluble in NH4OH]
AgBr [Pale yellow solube in HNO3, but less soluble in NH4OH]
Agl [yellow soluble in both HNO3 and NH4OH]
Pb 2 salt
HF No Precipitate.
PbCl2 , PbBR2 white , Pbl2 Yellow .
All soluble in hot water
Hg 2 and Hg .
HF
No precipitate
HX Precipitate
CuSO4 Solution
Only HI reacts
HI CuSO4 Cul I 2 .
Cul2
BaCl2 , SrCl2 and CaCl2 solution.
Only HF reacts to produce HF2 (white ppt).
Others – No precipitate
Reaction with MnO2 and H 2 SO4 .
MnSO4 H 2O O .
MnO2 H 2 SO4
HX + O H2O + X2.
HF does not give this reaction.
BLEACHING POWDER
Mixed salt of calcium hypo – chlorite Ca OCl 2 3H 2O and basic calcium carbonate
CaCl2 Ca OH 2 .H 2O .
Properties
Aqueous solution gives positive test for Ca 2 , Cl and OCl .
On standing, decomposes into CaCl2 and Ca ClO3 2 .
Ca ClO3 2 CaCl2 .
CaOCl2
Reaction with Acid
CaOCl2 H
Ca 2 CaCl2 HOCl .
HCl O (With limitd quantity of Acid)
HOCl
CaOCl2 H
Ca 2 H 2O Cl2 (with excess of Acid)
CaCO3 Cl2 .
CaOCl2 CO2
Oxidizing nature and Bleaching nature
CaCl2 O .
CaOCl2
1. Proparation of HCl, HBr, HI
A. H 2 Cl2 2 HCl .
pt
B. H 2 Br2 2 HBr .
Pt
C. H 2 I 2
450 c
2 HI .
2. By reaction of P4.
P4 6 Br2 6 I 2
4 PBr3 4 PI 3 .
PBr3 PI 3 3H 2O
3HBr HI H 3 PO3 .
3. By passing H 2 S & SO2 into solution of halogen.
2 HX S .
H 2 S X 2
2 HX H 2 SO4 .
SO2 2 H 2O X 2
Reducing nature: HI > HBr > HCl.
Stability: HF > HCl > HBr > HI.
Acidic nature: HI> HBr > HCl > HF.
Detection of Cation:
Fe
HF preparation
1. Laboratory method
kHF2 kF HF .
2. Industrial Method.
CaF2 H 2 SO4 CaSO4 2 HF .
Note: Aqous
HF
corrosive to glass
stored in wax lined bottles or vessel made of copper or monel.
high B.P. (Due to H – Bonding)
NaCl
traces
ClO2 NaHSO4
light
ClO2 ClO O
H
ClO2 O3 Cl2O6 O2
H
ClO2 HI I 2 Cl
ClO2 O2
ClO2 H 2O2
Oxides of chlorine:
CaOCl2 (Bleaching Powder)
Preparation:
40C
Ca OH 2 Cl2 Ca OCl Cl H 2O .
Properties:
a) Pale yellow powder.
b) Soluble in water.
c) Long standing undergoes auto oxidation.
CaOCl2 Ca ClO3 2 SCaCl2 .
d) Losses O in H 2 SO4 (dilute)
CaCl2 CaSO4 HClO .
CaOCl2 H 2 SO4
HClO
HCl O
.
H2S
S + CaCl2
KI I2 + CaCl2
Oxidizing NH3
Property N2 + CaCl2
NO 2
NO3
Fe+2 Fe+3
H+
Note: When CaOCl2 react with dilute acid or CO2 then it releases available chlorine:
Cl2 CaCl2 .
CaOCl2 HCl
Cl2 CaSO4 .
CaOCl2 H 2 SO4
CaOCl2 CO2
Cl2 CaCO3
Available
Chlorine
OXY ACIDS OF HALOGENS:
FLUORINE:
HOF Hypo florous acid
CHLORINE: BROMINE:
IODINE:
ACIDIC STRENGTH:
HClO HBrO HIO .
HClO3 HBrO3 HIO3 .
HClO4 HBrO4 HIO4 .
Q. Which one is stronger acid.
H 2 SO4 , HClO4 .
Q. Oxidizing Power
BrO4 IO4 ClO4 .
INTERHALOGEN COMPOUNDS
AB AB3 AB5 AB7
ClF . ClF3 . ClF5 . IF7 .
BrF . BrF3 . BrF5 .
ICl . ICl3 . IF5 .
IF IF3 .
PROPERTIES:
1. Can be gas, liquid, solid.
Gas ClF , IF7 .
BrF3 .
Liquid
ICl , IBr , IF3 , ICl3 .
Solid
2. ftlesa F gS
os jxafgu gSaA ftlesa Br, I, Cl gS os jafxu gSaA
Intensity (frozrk) (darkness) as mw .
3. All are covalent (lgl;ksth) Molecule (v.kq).
4. izfrpqEcfd;A
5. rkfi; LFkkf;Rork as EN es vraj .
IF BrF ClF ICl IBr BrCl .
6. ftruk T;knk AB /kzqfo; (Polar) mruk fg LFkk;h (Stable).
7. fØ;kf’kyrk dk Øe.
ClF3 BrF3 IF7 BrF5 BrF .
8. Hydrolysis (ty;kstu)
HB HOA .
AB H 2O
HCl HOBr .
BrCl H 2O
HCl HIO2 .
ICl3 H 2O
HF HIO3 .
IF5 H 2O
SODIUM THIO SULPHATE Ha2 S 2O3 5 H 2O HYPO
PREPARATION:
Na2 S3 O2
Na2 S 2O3 Na2 SO4
S
Na2 SO3
boiling
Na2 S 2O3
SO2 S I2
NaOH Na2 SO3 Na2 S 2O3
Na2 S2O3
PPT RXN:
H 2O
PbS
H 2O
Ag 2 S
REDUCING BEHAVIOR:
I2
Na2S4O6 + NaI
Cl2
Na2S2O3 NaHSO4 + HCl
kMnO4
k2SO4 + Na2SO4 + Mn2O3
COMPLEXATION RXN:
Excess
Ag 2 S 2O3 Na3 Ag S2O3 2
Na3 Ag S2O3 2
x Cl , Br , I
Excess
AuCl
Na2 S 2O3
Au2 S 2O3 Na3 Au S 2O3 2
Na4 Cu6 S 2O3 5
Na3 Bi S 2O3 3 Bi2 S3
Fe 3
Fe S2O3 2
Group – 18
FLUORIDES OF XENON
Reaction with HF– This reaction takes place as follows
XeF2 and XeF4 dissolves in HF without any XeF6 H 2O XeO3 HF .
rection.
XeO3 OH HXeO4 .
XeF6 dissolves in HF to form an equilibrium.
XeF6 HF XeF5 HF2 . HXeO4 OH XeO64 Xe O2 H 2O .
Reaction with water XeF6 reacts with silica, so can’t be stored in glass
XeF4 2 H 2O Xe HF O2 . XeF6 SiO2 XeOF4 SiF4 .
XeF2 and XeF4 acts as fluorinating agent, but XeF6
XeF2 OH Xe F O2 [Base catalysed]
accepts F– from alkali metal fluorides [except LiF]
XeO3 HF Xe O2 .
XeF4 H 2O All these can act as oxidizing agent.
XeOF2 HF [Partial hydrolysis]
XeF4 H 2O XeFx Xe .
XeOF4 / XeO2 F2 / XeO3 .
XeF6 H 2O
XeF6 OH
XeO64 Xe F O2 .
OXIDES OF XENON
XeF6 + H2O XeF4 + H2O
HXeO4 OH– XeO3 Powerful oxidizing agent
OH– Pu2++XeO3+H+ Pu4+ + Xe +H2O
XeO64 Xe O2 Explosive solid
H2SO4
Powerful oxidizing agent Room temp.
Xe + O2
Cl– Cl2 Unstable gas
Mn2+ MnO4
H2O O2
