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Liquid nitrogen

VA group. Nitrogen and Phosphorous and their compounds. 

Lecture 14

Group 15—The Nitrogen Group

• Phosphorous compounds

are essential ingredients for

healthy teeth and bones.

• Plants also need phosphorus,

so it is one of the nutrients in

most fertilizers.

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Atomic and Physical Properties of Group 15 Elements

Electronic Configuration.

The valence shell electronic configuration of these elements is ns2np3

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.

There is a considerable increase in covalent radius from N to P

However, from As to Bi only a small increase in covalent radius is

observed. This is due to the presence of completely filled d‐ and/or f‐

orbitals in heavier members.

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Ionisation Enthalpy

• Ionisation enthalpy decreases down the group due to gradual

increase in atomic size.

• Ionisation enthalpy of group 15 is much greater than that of

group 14 elements in the corresponding periods.

extra stable half‐filled p‐orbitals electronic configuration

smaller atomic size

The order is ΔH1 < ΔH2 < ΔH3

Physical Properties

• All the elements of this group are polyatomic

• Dinitrogen is a diatomic gas while all others are solids

• Metallic character increases down the group.

• Nitrogen and phosphorus are non‐metals, arsenic and antimony

are metalloids and bismuth is a metal. Due to:‐

(i) decrease in ionisation enthalpy

(ii) increase in atomic size.

• The boiling points, in general, increase from top to bottom in the

group but the melting point increases up to arsenic and then

decreases up to bismuth

• Except nitrogen, all the elements show allotropy

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Preparation of dinitrogen (N2)• In the laboratory, dinitrogen is prepared by treating an aqueous

solution of ammonium chloride with sodium nitrite.

NH4CI(aq) + NaNO2 (aq)→ N2 (g) + 2 H2O (l) + NaCl (aq)

• It can be obtained by the thermal decomposition of ammoniumdichromate.

(NH4)2Cr2O7 (Heat) → N2 (g) + 4 H2O (l) + Cr2O3 (s)

• Very pure nitrogen can be obtained by the thermal decomposition ofsodium or barium azide.

Ba(N3)2 → Ba + 3 N2

Air (4 N2 + O2) + C→ 4 N2 + CO2

NH3 + 3 O2 → 2 N2 + 6 H2O

2 NH3 + 3 Cl2 → N2 + 6 HCl

Reactivity towards hydrogen

All Group 15 elements form hydrides of

the type EH3 where E = N, P, As, Sb or Bi.

N2(g) + 3 H2(g) (773 k) ==> 2 NH3(g);

ΔH = – 46.1 kJ mol–1

Р4 + 6 Н2 (heat, p) ==> 4 РН3

• The stability of hydrides decreases

from NH3 to BiH3

• Reducing character of the hydrides

increases.

• NH3 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

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. The oxides of

the type E2O3 of nitrogen and

phosphorus are purely acidic

N2 (g) + O2 (g) (heat) ==> 2 NO (g)

P4 + 5 O2 (heat) ==> 2 P4O10

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Reactivity towards halogens

• These elements react to form two

series of halides: EX3 and EX5

• Nitrogen does not form

pentahalide. 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.

P4 + 6 Cl2 ==> 4 PCl3

3 PCl5 + 2 P ==> 5 PCl3

3 PCl5 + P2O5 ==> 5 POCl3

Reactivity towards metals

• All these elements react with

metals to form their binary

compounds exhibiting –3 oxidation

state. E.g.

‒ Ca3N2 (calcium nitride)

‒ Ca3P2 (calcium phosphide)

‒ Na3As2 (sodium arsenide)

‒ Zn3Sb2 (zinc antimonide)

‒ Mg3Bi2 (magnesium bismuthide)

3Mg + N2 ==> Mg3N2

Mg3N2 + 6H2O ==> 3Mg(OH)2 + 2NH3

Ammonia• Ammonia is present in small quantities in air and soil where it is

formed by the decay of nitrogenous organic matter e.g., urea.

NH2CONH2 + 2 H2O→ (NH4)2CO3 → 2 NH3 + H2O + CO2

• On a small scale ammonia is obtained from ammonium salts

which decompose when treated with caustic soda or lime.

2 NH4Cl + Ca(OH)2 → 2 NH3 + 2 H2O + CaCl2

(NH4)2 SO4 + 2 NaOH→ 2 NH3 + 2 H2O + Na2SO4

• On a large scale, ammonia is manufactured by Haber’s process.

N2(g) + 3H2(g)→ 2NH3(g); ΔH = – 46.1 kJ mol‐1

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Flow chart for the manufacture of ammonia

Phosphine

• Phosphine is prepared by the reaction of calcium phosphide

with water or dilute HCl

Ca3P2 + 6 H2O→ 3 Ca(OH)2 + 2 PH3

Ca3P2 + 6 HCl→ 3 CaCl2 + 2 PH3

• In the laboratory, it is prepared by heating white phosphorus

with conc. NaOH solution in an inert atmosphere of CO2.

Р4 + 3 КОН + 3 Н2О→ РН3 + 3 КН2РО4

PH4I + KOH → KI + H2O + PH3

(phosphonium iodide)

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Properties of Ammonia

• Ammonia gas is highly soluble in water. Its aqueous solution is weakly basic due to the formation of OH– ions.

NH3 (g) + H2O (l) → NH4+ (aq) + OH– (aq)

3 CuO + 2 NH3 → 3 Cu + 3 H2O + N2

• It forms ammonium salts with acids, e.g., NH4Cl, (NH4)2SO4, etc. As a weak base, it precipitates the hydroxides of many metals from their salt solutions. 

2 FeCl3 + 3 NH4OH  → Fe2O3xH2O (brown ppt) + 3 NH4Cl

• The ammonia molecule can act is a Lewis base

Cu2+ (aq, blue) + 4 NH3(aq, deep blue) → [Cu(NH3)4]2+(aq)

Ag+ (aq, colourless) Cl‐ → AgCl (s, white ppt)

AgCl (s, white ppt) + 2 NH3 (aq) → [Ag(NH3)2]Cl)(aq,colourless)

Nitrogen oxides

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Nitric acid• In the laboratory, nitric acid is prepared by heating KNO3 or

NaNO3 and concentrated H2SO4 in a glass retort:

NaNO3 + H2SO4 → NaHSO4+ HNO3

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Preparation of Nitric acid

Brown Ring Test

• The familiar brown ring test for nitrates depends on the ability of Fe2+  to reduce nitrates to nitric oxide, which reacts with Fe2+  to form a brown coloured complex.

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Test for Ammonia

Test for Ammonia using Nessler's Agent. Ammonia is tested in a 1:10 dilution row using K2[HgI4].

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Reactivity Trends Organometallic Compounds

• All act as Lewis bases with basicity

• All three elements form a wide range decreasing As > Sb > Bi

of organometallic compounds

• Hydrides of As, Sb, Bi are exceedingly the +3 and +5 state.

• The are poisonous and thermally unstable

• Can form both M‐C single bonds and M=C double bonds

• As, Sb, Bi react with halides to form MX3, MX5,

• halide complexes of MIII Stability of M‐C bond strength and

MVdecreases in the order of As > Sb > Bi

• Aryl more stable than alkyl

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Phosphorus. Allotropic Forms

• White phosphorus is a translucent white waxy solid. It is poisonous, insoluble in water but soluble in carbon disulphide and has chemiluminescence. 

Р4 + 3 NaOH + 3 H2O→ PH3 + NaH2PO2

Р4 + 3О2 → 2Р2О3

Р2О3 + О2 → Р2О5

3 Mg + 1/2P4 →Mg3P2

2 P + 3 Cl2 → 2 PCl3

15P

1s2

2s2 2p6

3s2 3p3 3d0

White phosphorus

White phosphorus exposed to air glows in the darkness

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Red phosphorus

• It is 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 are formed.

• Red phosphorus possesses iron grey lustre. It is

odourless, non‐poisonous and insoluble in water as

well as in carbon disulphide.

• Chemically, red phosphorus is much less reactive than

white phosphorus. It does not glow in the dark.

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Black phosphorus

• It has two forms  α‐black phosphorus and  β‐blackphosphorus.  

• α‐Black phosphorus is formed when red phosphorus is heated in a sealed tube at 803K. It can be sublimed in air and has opaque monoclinic or rhombohedral crystals.  It does not oxidise in air. 

• β‐Black phosphorus is prepared by heating white phosphorus at 473 K under high pressure. It does not burn in air up to 673 K.

Phosphorus Tri‐ and Pentachlorides

They are obtained by the action of thionyl chloride with white

phosphorus.

P4 + 8 SOCl2 → 4 PCl3 + 4 SO2 + 2S2Cl2

P4 + 10 SOCl2 → 4 PCl5 + 10 SO2

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Applications of nitrogen compounds

• As a modified atmosphere, pure or mixed with carbon

dioxide, to preserve the freshness of packaged or bulk foods

• Nitrogen can be used instead of CO2 to pressurize kegs of

some beers, in particular, stouts and British ales, due to the

smaller bubbles it produces, which make the dispensed beer

smoother and headier

• Liquid nitrogen is used in the cryopreservation of blood,

reproductive cells (sperm and egg), and other biological

samples. It is used in the clinical setting in cryotherapy to

remove cysts and warts on the skin.

Applications of nitrogen compounds

• Nitrous oxide (N2O), "laughing gas“, was discovered early in

the 19th century to be a partial anesthetic, though it was not

used as a surgical anesthetic until later.

• Nitrogen‐containing drugs are drugs derived from plant

alkaloids, such as morphine (there exist many alkaloids known

to have pharmacological effects; in some cases, they appear

natural chemical defenses of plants against predation). Drugs

that contain nitrogen include all major classes of antibiotics

and organic nitrate drugs like nitroglycerin and nitroprusside

that regulate blood pressure and heart action by mimicking

the action of nitric oxide.

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Applications of phosphorous compounds

• White phosphorus, called "WP" (slang term "Willie Peter") is

used in military applications as incendiary bombs, for

smoke‐screening as smoke pots and smoke bombs, and in

tracer ammunition.

• The spontaneous combustion of phosphine is technically

used in Holme’s signals. Containers containing calcium

carbide and calcium phosphide are pierced and thrown in

the sea when the gases evolved burn and serve as a signal.

• Phosphine is also used in smoke screens.