The p -Block Elements

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GROUP 15 ELEMENTS
-Group 15 includes nitrogen, phosphorus, arsenic, antimony, bismuth and moscovium.
-Molecular nitrogen comprises 78% by volume of the atmosphere.

Electronic Configuration
-The valence shell electronic configuration of these elements is ns2 np3 .
-The s orbital in these elements is completely filled and p orbitals are half-filled, making their
electronic configuration extra stable.

Atomic and Ionic Radii

-Covalent and ionic (in a particular state) radii increase in size down the group

Ionisation Enthalpy
-Ionisation enthalpy decreases down the group due to gradual increase in atomic size.
-Because of the extra stable half-filled p orbitals electronic configuration and smaller size, the
ionisation enthalpy of the group 15 elements is much greater than that of group 14 elements in the
corresponding periods

Electronegativity
-The electronegativity value, in general, decreases down the group with increasing atomic size.

Physical Properties
-All the elements of this group are polyatomic
-Metallic character increases down the group.
-Nitrogen and phosphorus are non-metals, arsenic and antimony metalloids and bismuth is a metal.
-The boiling points, in general, increase from top to bottom in the group but the melting point
increases upto arsenic and then decreases upto bismuth.
-Except nitrogen, all the elements show allotropy.

Chemical Properties
-The common oxidation states of these elements are –3, +3 and +5.
-The stability of +5 oxidation state decreases down the group
-Nitrogen is restricted to a maximum covalency of 4 since only four (one s and three p) orbitals are
available for bonding
-The heavier elements have vacant d orbitals in the outermost shell which can be used for bonding.
Anomalous properties of nitrogen
-Nitrogen differs from the rest of the members of this group due to its small size, high
electronegativity, high ionisation enthalpy and non-availability of d orbitals.
-Nitrogen has unique ability to form pπ-pπ multiple bonds with itself and with other elements
having small size and high electronegativity.
-The single N–N bond is weaker than the single P–P bond because of high interelectronic repulsion
of the non-bonding electrons, owing to the small bond length.
-As a result the catenation tendency is weaker in nitrogen.
-Another factor which affects the chemistry of nitrogen is the absence of d orbitals in its valence
shell.

(i)Reactivity towards hydrogen:

-All the elements of Group 15 form hydrides of the type EH3
-The stability of hydrides decreases from NH3 to BiH3 which can be observed from their bond
dissociation enthalpy. the reducing character of the hydrides increases.
-Ammonia is only a mild reducing agent while BiH3 is the strongest reducing agent amongst all the
hydrides.
-Basicity also decreases in the order NH3 > PH3 > AsH3 > SbH3 > BiH3.
-Due to high electronegativity and small size of nitrogen, NH3 exhibits hydrogen bonding in solid as
well as liquid state. Because of this, it has higher melting and boiling points than that of PH 3

(ii) Reactivity towards oxygen:

-All these elements form two types of oxides: E2O3 and E2O5.
-The oxide in the higher oxidation state of the element is more acidic than that of lower oxidation
state.
-Their acidic character decreases down the group.

(iii) Reactivity towards halogens:

-Nitrogen does not form pentahalide due to non-availability of the d orbitals in its valence shell.
-Pentahalides are more covalent than trihalides.
-All the trihalides of these elements except those of nitrogen are stable.
-In case of nitrogen, only NF3 is known to be stable.
- Trihalides except BiF3 are predominantly covalent in nature.

(iv) Reactivity towards metals:

-All these elements react with metals to form their binary compounds.

Dinitrogen
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-Dinitrogen is a colourless, odourless, tasteless and non-toxic gas.
-It has a very low solubility in water
-low freezing and boiling points
-Dinitrogen is rather inert at room temperature because of the high bond enthalpy of N≡N bond.
-It combines with hydrogen at about 773 K in the presence of a catalyst (Haber’s Process) to form
ammonia:

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Ammonia
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-Ammonia is a colourless gas with a pungent odour.
-In the solid and liquid states, it is associated through hydrogen bonds as in the case of water and
that accounts for its higher melting and boiling points than expected on the basis of its molecular
mass.
-The ammonia molecule is trigonal pyramidal with the nitrogen atom at the apex
-Ammonia gas is highly soluble in water. Its aqueous solution is weakly basic due to the formation of
OH– ions.
NH3 (g) + H2O(l) NH+4(aq) + OH- (aq)
-The presence of a lone pair of electrons on the nitrogen atom of the ammonia molecule makes it a
Lewis base.
-It donates the electron pair and forms linkage with metal ions and the formation of such complex
compounds finds applications in detection of metal ions such as Cu2 +, Ag+ :

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Uses:NCERT

Oxides of Nitrogen
-7.3 and 7.4 table form ncert

Nitric Acid
-Nitrogen forms oxoacids such as
H2N2O2 = hyponitrous acid
HNO2 =nitrous acid
HNO3 =nitric acid
Preparation-NCERT
-It is a colourless liquid.
-In the gaseous state, HNO3 exists as a planar molecule with the structure as shown.
-In aqueous solution, nitric acid behaves as a strong acid giving hydronium and nitrate ions.
-Concentrated nitric acid is a strong oxidising agent and attacks most metals except noble metals
such as gold and platinum.
-Zinc reacts with dilute nitric acid to give N2O and with concentrated acid to give NO2.
-Some metals (e.g., Cr, Al) do not dissolve in concentrated nitric acid because of the formation of a
passive film of oxide on the surface.
-Concentrated nitric acid also oxidises non–metals and their compounds. Iodine is oxidised to iodic
acid, carbon to carbon dioxide, sulphur to H2SO4 , and phosphorus to phosphoric acid.

Brown Ring Test:

Phosphorus — Allotropic AllotropicAllotropic Forms

-Phosphorus is found in many allotropic forms, the important ones being white, red and black.
White phosphorus
-is a translucent white waxy solid.
-It is poisonous, insoluble in water but soluble in carbon disulphide
-It dissolves in boiling NaOH solution in an inert atmosphere giving PH3
-White phosphorus is less stable and therefore, more reactive than the other solid phases under
normal conditions because of angular strain in the P4 molecule where the angles are only 60°.
-It readily catches fire in air to give dense white fumes of P 4O10
P4 + 5O2 P4O10
Red phosphorus
- obtained by heating white phosphorus at 573K in an inert atmosphere for several days.
- When red phosphorus is heated under high pressure, a series of phases of black phosphorus is
formed.
- It is odourless, nonpoisonous and insoluble in water as well as in carbon disulphide.
- red phosphorus is much less reactive than white phosphorus.
- It is polymeric, consisting of chains of P4 tetrahedra linked together.
Black phosphorus
- two forms α-black phosphorus and β-black phosphorus.
- α-Black phosphorus is formed when red phosphorus is heated in a sealed tube at 803K.
- β-Black phosphorus is prepared by heating white phosphorus at 473 K under high pressure. It
does not burn in air upto 673 K.

Phosphine
Preparation-NCERT
- It is a colourless gas with rotten fish smell and is highly poisonous.
- It is slightly soluble in water.
- Phosphine is weakly basic and like ammonia

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Group 16 Elements
-Known as group of chalcogens.
- Most copper minerals contain either oxygen or sulphur and frequently the other members of the
group.
Electronic Configuration- ns2 np4

Atomic and Ionic Radii

- Due to increase in the number of shells, atomic and ionic radii increase from top to bottom in the
group

Ionisation Enthalpy
- Ionisation enthalpy decreases down the group.

Electron Gain Enthalpy

- Because of the compact nature of oxygen atom, it has less negative electron gain enthalpy than
sulphur. However, from sulphur onwards the value again becomes less negative upto polonium.

Electronegativity
- Next to fluorine, oxygen has the highest electronegativity value amongst the elements.
- Within the group, electronegativity decreases with an increase in atomic number.
-This implies that the metallic character increases from oxygen to polonium.

Physical Properties
- Oxygen and sulphur are non-metals, selenium and tellurium metalloids, whereas polonium is a metal.
- The melting and boiling points increase with an increase in atomic number down the group.
- The large difference between the melting and boiling points of oxygen and sulphur may be
explained on the basis of their atomicity; oxygen exists as diatomic molecule (O2 ) whereas sulphur
exists as polyatomic molecule (S8 ).

Chemical Properties
-The stability of -2 oxidation state decreases down the group.
- The stability of + 6 oxidation state decreases down the group and stability of + 4 oxidation state
increases.

Anomalous behaviour of oxygen

- The anomalous behaviour of oxygen, like other members of p-block present in second period is due
to its small size and high electronegativity.
- One typical example of effects of small size and high electronegativity is the presence of strong
hydrogen bonding in H2O which is not found in H2S.

(i)Reactivity with hydrogen

- All the elements of Group 16 form hydrides of the type H2E.
- Acidic character increases from H2O to H2Te
- The increase in acidic character can be explained in terms of decrease in bond enthalpy for the
dissociation of H–E bond down the group.
- Owing to the decrease in enthalpy for the dissociation of H–E bond down the group, the thermal
stability of hydrides also decreases from H2O to H2Po.

(ii) Reactivity with oxygen:

- Ozone (O3 ) and sulphur dioxide (SO2 ) are gases while selenium dioxide (SeO 2 ) is solid.

(iii) Reactivity towards the halogens:

- The stability of the halides decreases in the order F- > Cl- > Br- > I-
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All hexafluorides are gaseous in nature.
- Sulphur hexafluoride, SF6 is exceptionally stable for steric reasons.
- Amongst tetrafluorides, SF4 is a gas, SeF4 a liquid and TeF4 a solid.

Dioxygen
Preparation-NCERT
- Dioxygen is a colourless and odourless gas.
- Dioxygen directly reacts with nearly all metals and non-metals except some metals ( e.g., Au, Pt)
and some noble gases.
- Its combination with other elements is often strongly exothermic which helps in sustaining the
reaction.

Ozone
-Ozone is an allotropic form of oxygen.
- It is too reactive to remain for long in the atmosphere at sea level.
- Preparation=NCERT
- Pure ozone is a pale blue gas, dark blue liquid and violet-black solid.
- Ozone has a characteristic smell and in small concentrations it is harmless.
- Ozone is thermodynamically unstable with respect to oxygen since its decomposition into oxygen
results in the liberation of heat (∆H is negative) and an increase in entropy (∆S is positive).

Sulphur — Allotropic Forms

-Sulphur forms numerous allotropes of which the yellow rhombic (α-sulphur) and monoclinic (β -
sulphur) forms are the most important.
- The stable form at room temperature is rhombic sulphur.
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Oxoacids of Sulphur
-Sulphur forms a number of oxoacids such as H2SO3 , H2S2O3 , H2S2O4 , H2S2O5 , H2SxO6 (x =
2 to 5), H2SO4 , H2S2O7 , H2SO5 , H2S2O8 .
- Some of these acids are unstable and cannot be isolated.

Sulphuric Acid
-Sulphuric acid is one of the most important industrial chemicals worldwide.
- Sulphuric acid is manufactured by the Contact Process which involves three steps:
(i) Burning of sulphur or sulphide ores in air to generate SO 2 .
(ii) Conversion of SO2 to SO3 by the reaction with oxygen in the presence of a catalyst (V2O5 ),
(iii) Absorption of SO3 in H2SO4 to give Oleum (H2S2O7).
- The SO2 produced is purified by removing dust and other impurities such as arsenic compounds.
- The key step in the manufacture of H2SO4 is the catalytic oxidation of SO2 with O2 to give SO3
in the presence of V2O5 (catalyst).
- The reaction is exothermic, reversible and the forward reaction leads to a decrease in volume.
- low temperature and high pressure are the favourable conditions for maximum yield.
- Sulphuric acid is a colourless, dense, oily liquid with a specific gravity of 1.84 at 298 K.
- The chemical reactions of sulphuric acid are as a result of the following characteristics: (a) low
volatility (b) strong acidic character (c) strong affinity for water and (d) ability to act as an
oxidising agent. In aqueous solution, sulphuric acid ionises in two steps.
Group 17 Elements
-Fluorine and chlorine are fairly abundant while bromine and iodine less so.

Electronic Configuration
- ns2 np5
Atomic and Ionic Radii
- The halogens have the smallest atomic radii in their respective periods due to maximum effective
nuclear charge..
- The atomic radius of fluorine like the other elements of second period is extremely small.
- Atomic and ionic radii increase from fluorine to iodine due to increasing number of quantum shells.

Ionisation Enthalpy
- They have little tendency to lose electron.
- very high ionisation enthalpy. Due to increase in atomic size, ionisation enthalpy decreases down
the group.

Electron Gain 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.

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

Physical Properties
-Fluorine and chlorine are gases, bromine is a liquid and iodine is a solid.
- Their melting and boiling points steadily increase with atomic number.
- All halogens are coloured.
- By absorbing different quanta of radiation, they display different colours. For example, F2 , has
yellow, Cl2 , greenish yellow, Br2 , red and I2 , violet colour.
- Fluorine and chlorine react with water. Bromine and iodine are only sparingly soluble in water but
are soluble in various organic solvents such as chloroform, carbon tetrachloride, carbon disulphide
and hydrocarbons to give coloured solutions.
- One curious anomaly we notice ,the smaller enthalpy of dissociation of F2 compared to that of Cl2
whereas X-X bond dissociation enthalpies from chlorine onwards show the expected trend: Cl – Cl >
Br – Br > I – I.
- A reason for this anomaly is the relatively large electron-electron repulsion among the lone pairs
in F2 molecule where they are much closer to each other than in case of Cl 2 .

Chemical Properties
-All the halogens exhibit –1 oxidation state.
- The higher oxidation states of chlorine, bromine and iodine are realised mainly when the halogens
are in combination with the small and highly electronegative fluorine and oxygen atoms.
- 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.
- F2 is the strongest oxidising halogen and it oxidises other halide ions in solution or even in the
solid phase.

Anomalous behaviour of fluorine-NCERT

(i)Reactivity towards hydrogen:

- They all react with hydrogen to give hydrogen halides but affinity for hydrogen decreases from
fluorine to iodine.
- Hydrogen halides dissolve in water to form hydrohalic acids.
- The acidic strength of these acids varies in the order: HF < HCl < HBr < HI.
- The stability of these halides decreases down the group due to decrease in bond (H–X)
dissociation enthalpy in the order: H–F > H–Cl > H–Br > H–I.

(ii) Reactivity towards oxygen:

- Halogens form many oxides with oxygen but most of them are unstable.
- Fluorine forms two oxides OF2 and O2F2 .
- only OF2 is thermally stable at 298 K.
- O2F2 oxidises plutonium to PuF6 and the reaction is used in removing plutonium as PuF6 from spent
nuclear fuel.
- Chlorine, bromine and iodine form oxides in which the oxidation states of these halogens range
from +1 to +7

(iii) Reactivity towards metals:

- Halogens react with metals to form metal halides. For example, bromine reacts with magnesium to
give magnesium bromide.
- The ionic character of the halides decreases in the order MF > MCl > MBr > MI where M is a
monovalent metal.
- If a metal exhibits more than one oxidation state, the halides in higher oxidation state will be
more covalent than the one in lower oxidation state.
(iv) Reactivity of halogens towards other halogens:
- Halogens combine amongst themselves to form a number of compounds known as interhalogens of
the types XX ′ , XX3 ′ , XX5 ′ and XX7 ′ where X is a larger size halogen and X′ is smaller size
halogen.

Chlorine –NCERT

Hydrogen Chloride
Preparation-NCERT
- It is a colourless and pungent smelling gas.
- It is easily liquefied to a colourless liquid (b.p.189 K) and freezes to a white crystalline solid (f.p.
159 K).
- It is extremely soluble in water and ionizes
- Its aqueous solution is called hydrochloric acid.
- High value of dissociation constant (Ka ) indicates that it is a strong acid in water.
- It reacts with NH3 and gives white fumes of NH4Cl.
- When three parts of concentrated HCl and one part of concentrated HNO 3 are mixed, aqua regia
is formed which is used for dissolving noble metals, e.g., gold, platinum.

Interhalogen Compounds
-When two different halogens react with each other, interhalogen compounds are formed.
- They can be assigned general compositions as XX′ , XX3 ′ , XX5 ′ and XX7 ′ where X is halogen of
larger size and X′ of smaller size and X is more electropositive than X′ .
- As the ratio between radii of X and X ′ increases, the number of atoms per molecule also
increases.

Preparation-NCERT
Table 7.11.-NCERT
- These are all covalent molecules and are diamagnetic in nature.
- They are volatile solids or liquids at 298 K except ClF which is a gas.
- Their physical properties are intermediate between those of constituent halogens except that
their m.p. and b.p. are a little higher than expected.
- In general, interhalogen compounds are more reactive than halogens (except fluorine).
- This is because X–X′ bond in interhalogens is weaker than X–X bond in halogens except F–F bond.