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Group 15 Nitrogen group

Group 15
Nitrogen Group
Group 15—The Nitrogen Group
Nitrogen and phosphorus are required by living
things and are used to manufacture various items.
• These
elements also are parts of the biological
materials that store genetic information and energy in
living organisms.
The nitrogen cycle
Ways nitrogen is lost to the cycle
❖ Most of the nitrogen cycle is soil based. Four ways how nitrogen is lost :
Bacteria change nitrate in the soil to atmospheric nitrogen
II. Volatilization
Turns urea fertilizers and manures on the soil surface into gases
III. Runoff
Carries the nitrogen in fertilizers and manure and the nitrogen in
the soil into our rivers and streams
IV. Leaching
Carries nitrates so deep into the soil that plants can no longer use
them, producing a dual concern — for lost fertility and for water
quality, as nitrates enter the groundwater and the wells that
provide our drinking water.
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
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 ammonium
(NH4)2Cr2O7 (Heat) ⎯⎯⎯→ N2 (g) + 4 H2O (l) + Cr2O3 (s)
• Very pure nitrogen can be obtained by the thermal decomposition of
sodium or barium azide.
→ 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
Reactivity towards
All Group 15 elements form hydrides of
the type EH3 where E = N, P, As, Sb or Bi.
• All these elements form two types
of oxides: E2O3 and E2O5
N2(g) + 3 H2(g) (773 k) ==> 2 NH3(g);
ΔH = – 46.1 kJ mol–1
• The oxide in the higher oxidation
state of the element is more
acidic than that of lower oxidation
Р4 + 6 Н2 (heat, p) ==> 4 РН3
• The stability of hydrides decreases
from NH3 to BiH3
• Reducing character of the hydrides
• 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
• 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
Reactivity towards
• These elements react to form two
series of halides: EX3 and EX5
• Nitrogen
pentahalide. Pentahalides are more
covalent than trihalides.
Reactivity towards
• All these elements react with
metals to form their binary
compounds exhibiting –3 oxidation
state. E.g.
‒ Ca3N2 (calcium nitride)
• All the trihalides of these elements
except those of nitrogen are stable
‒ Ca3P2 (calcium phosphide)
• In case of nitrogen, only NF3 is
known to be stable.
‒ Zn3Sb2 (zinc antimonide)
‒ Na3As2 (sodium arsenide)
‒ Mg3Bi2 (magnesium bismuthide)
P4 + 6 Cl2 ==> 4 PCl3
3 PCl5 + 2 P ==> 5 PCl3
3 PCl5 + P2O5 ==> 5 POCl3
3 Mg + N2 ==> Mg3N2
Mg3N2 + 6 H2O ==> 3 Mg(OH)2 + 2 NH3
• 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)2SO4 + 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
• 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 КН2РО2
PH4I + KOH → KI + H2O + PH3
(phosphonium iodide)
Nitrogen oxides
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.
Black phosphorus
• It has two forms α-black phosphorus and β-black
• 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
• 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.
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.
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.