Lecture 15.
p-Elements of VI A group.
Oxygen and Sulfur
PhD. Halina Falfushynska
Preparation of Oxygen
• In the laboratory, oxygen is prepared by several ways.
• By heating oxygen containing salts such as chlorates,
nitrates, bichromates and permanganates.
KClO3→ 2KCl + 3O2 
OF2 + 2NaOH  O2 + 2NaF + H2O
2K2Cr2O7+ 8H2SO4  2K2SO4 +2Cr2(SO4)3 +3O2 +8H2O
• By the thermal decomposition of the oxides of metals
low in the electrochemical series and higher oxides of
some metals
2 HgO  2 Hg + O2 ; PbO2  Pb + O2 ;
2Pb3O4(s) → 6PbO(s) + O2(g)
Preparation of Oxygen
• Hydrogen peroxide is readily decomposed into water
and dioxygen by catalysts such as finely divided
metals and manganese dioxide
2Н2О2  2Н2О + О2 
• On large scale it can be prepared from water or air.
Electrolysis of water leads to the release of hydrogen
at the cathode and oxygen at the anode.
Cathode: 4Н2О + 4е  2Н2  + 4ОН–
Anode: 4ОН– – 4е  О2  + 2Н2О
• 2Na2O2 + 2CO2  2Na2CO3 + O2 
Preparation of Oxygen
Properties of Oxygen gas
• Dioxygen is a colourless and odourless gas.
• Its solubility in water is to the extent of 3.08 cm3
in 100 cm3 water at 293 K which is just sufficient
for the vital support of marine and aquatic life.
• It liquefies at 90 K and freezes at 55 K. It has three
stable isotopes: 16O, 17O and 18O.
• Molecular oxygen, O2 is unique in being
paramagnetic inspite of having even number of
electrons
Reactivity towards
non-metals
Dioxygen directly reacts with
nearly all non-metals
except some noble gases.
2H2 + O2  2H2О
N2 (g)+ O2 (g) (heat)  2NO (g)
P4 + 5O2 (heat)  2P4O10
Reactivity towards
metals
Dioxygen directly reacts with
nearly all metals except some
of them (Au, Pt).
4Li + O2  2Li2O
2Na + O2  Na2O2
K + O2  KO2
Its combination with other elements is often strongly
exothermic which helps in sustaining the reaction.
However, to initiate the reaction, some external heating is
required as bond dissociation enthalpy of oxgyen-oxygen
double bond is high (493.4 kJ mol–1).
A space-filling model of the water
molecule.
Catalytic Decomposition of
Hydrogen Peroxide with MnO2
H2O2 is decomposed by catalytic
amounts of manganese dioxide.
Singlet oxygen is prepared by
oxidation of hydrogen peroxide
with NaOCl.
Oxidation of Iron
Iron is burned in pure oxygen to
give Fe3O4.
Combustion of a Cigar (cellulose)
in pure Oxygen. A cigar soaked in
liquid oxygen burns under strong
evolution of heat.
Redox Amphoteric Behaviour of
Hydrogen Peroxide. H2O2 reacts
in dependence of the reactant as
oxidizing or as reducing agent.
Test Reaction of Hydrogen
Peroxide. H2O2 is identified by
formation of an orange
peroxocomplex of Ti(IV).
Oxides
Neutral
N2O, CO
Basic
Simple
MgO, Al2O3
Acidic
Mixed
Pb3O4, Fe3O4
Amphoteric
metallic oxides non-metal oxides (N2O5) Some metallic oxides
BaO
oxides of some metals in Al2O3
high oxidation state
Al2O3 + 6HCl + 9H2O →
CaO + H2O
(Mn2O7, CrO3, V2O5)
2 [Al(H2O)6]3+ + 6Cl→ Ca(OH)2
Al + 6NaOH + 3H2O →
SO3 + H2O → H2SO4
2Na3[Al(OH)6]
Ozone and its
preparation
• Ozone is an allotropic form of oxygen. It is too
reactive to remain for long in the atmosphere
at sea level.
• At a height of about 20 kilometres, it is
formed from atmospheric oxygen in the
presence of sunlight.
• Ozone layer protects the earth’s surface from
an excessive concentration of ultraviolet (UV)
radiations.
Properties of Ozone
• Ozone is thermodynamically unstable with
respect to oxygen since its decomposition into
oxygen results in the liberation of heat and an
increase in entropy.
• Ozone acts as a powerful oxidising agent
Sulfur
Rhombic sulfur
(α-sulphur)
369 K
The stable form at room
temperature. It is
insoluble in water but
dissolves to some extent
in benzene, alcohol and
ether, CS2
Monoclinic sulfur
(β-sulphur)
It is soluble in CS2.
At 369 K colourless
needle shaped
crystals of β-sulphur
are formed.
Both rhombic and monoclinic sulfur have S8
molecules. At elevated temperatures (~1000 K),
S2 is the dominant species and is paramagnetic
like O2.
Extraction of sulfur by Frasch Method
• As a mineral, native sulfur
under salt domes is thought to
be a fossil mineral resource,
produced by the action of
ancient bacteria on sulfate
deposits. In this method,
superheated water was
pumped into a native sulfur
deposit to melt the sulfur, and
then compressed air returned
the 99.5% pure melted product
to the surface.
Preparation of Sulphur
• Sodium thiosulfate reacts with sulfuric acid
Na2S2O3 + 2H2SO4  [H2S2O3] + 2NaHSO4
[H2S2O3]  S + SO2 + H2O
• SO2 + 2 H2S → 3 S + 2 H2O
• 2H2S + O2 → 2H2O +2S
• 16(NH4)HSO3  6(NH4)2SO4 + 4NH3 + 7SO2 + 3S
+ 10H2O (t=1500)
• 2Na2S2O4 (Sodium dithionite) + H2O + HCl 
3NaHSO3 + S + NaCl
Reactivity
towards metals
Sulfur reacts with metals to
form its binary compounds
exhibiting –2 oxidation state,
such as, CaS (calcium sulfide)
Na2S (sodium sulfide).
Fe + S → FeS + Q
Cu + S → CuS + Q
16 Na + S8 → 8 Na2S
Reactivity towards
non-metals
Sulfur reacts with non-metals
to form its binary compounds
exhibiting – 2, +4 and +6
oxidation states such as H2S
(Hydrogen sulfide), SO2
(Sulphur dioxide), S2N2 (cyclic
sulfur(II,IV) nitride).
S + O2  SO2 + Q
Н2 + S  H2S
S + 2Cl2 + 4NaF  SF4 + 4NaCl
Reactivity
towards acids
Sulfur reacts with acids
S + 2H2SO4 
3SO2 + 2H2O
S + 6HNO3  H2SO4 +
6NO2 + 2H2O
Reactivity towards
alkalis
Sulfur reacts with alkali
4S + 6NaOH  Na2SO3S
+ 2Na2S + 3H2O (heat)
Sulfur Dioxide
• In the laboratory it is readily generated by
treating a sulfite with dilute sulfuric acid.
Na2SO3 + H2SO4 → Na2SO4 + H2O + SO2
SO32- (aq) + H+ (aq) → H2O(l) + SO2 (g)
• Industrially, it is produced as a by-product of the
roasting of sulfide ores.
4FeS2 + 11O2 → 2Fe2O3 + 8SO2
• C + 2H2SO4 → CO2 + 2SO2 + 2H2O
• 4CaSO4 + 2C → 4CaO + 4SO2 + 2CO2
• 2H2S + 3O2 → 2H2O + 2SO2
Sulfur Dioxide properties
• Sulphur dioxide, when passed through water, forms
a solution of sulphurous acid. SO2 + H2O → H2SO3
• It reacts readily with sodium hydroxide solution,
forming sodium sulphite .
SO2 + 2NaOH → Na2SO3 + H2O
• SO2 reacts with chlorine in the presence of charcoal
(which acts as a catalyst) to give sulfuryl chloride.
SO2(g) + Cl2 (g) → SO2Cl2(l)
• When moist, SO2 behaves as a reducing agent.
J2 + SO2 + H2O ® 2HJ + H2SO4
Sulfur trioxide preparation
2SO2 + O2 → 2SO3
• Sodium pyrosulfate decomposes, when sulfuric
acid is present:
Na2S2O7 (H+) → Na2SO4 + SO3
• Dehydration at 315°C and Cracking at 460°C:
2 NaHSO4 → Na2S2O7 + H2O;
Na2S2O7 → Na2SO4 + SO3
Sulfur trioxide preparation
Sulfur trioxide fuming in air
Sulfur trioxide properties
• SO3 is the anhydride of H2SO4:
SO3 (l) + H2O (l) → H2SO4 (l) (−88 kJ mol−1)
• At or above 340 °C, sulfuric acid, sulfur trioxide, and
water coexist in significant equilibrium concentrations.
• Sulfur trioxide also reacts with sulfur dichloride to
yield the useful reagent, thionyl chloride.
SO3 + SCl2 → SOCl2 + SO2
• SO3 is a strong Lewis acid readily forming crystalline
complexes with pyridine, dioxane and trimethylamine
which can be used as sulfonating agents.
Oxoacids of Sulphur
Oxoacids of Sulphur
Manufacture of sulfuric acid by the
Contact Process
• (i) burning of sulphur or sulphide ores in air to
generate SO2.
• (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 key step in the manufacture of H2SO4 is the
catalytic oxidation of SO2 with O2 to give SO3 in
the presence of V2O5 (catalyst).
Flow diagram for the manufacture of
sulfuric acid by the Contact Process
Properties of Sulfuric acid
• In aqueous solution, sulphuric acid ionises in two
steps.
• The acid forms two series of salts: normal sulphates
(CuSO4) and acid sulphates (NaHSO4).
• Sulphuric acid, because of its low volatility can be used
to manufacture more volatile acids from their
corresponding salts.
2 MX + H2SO4 → 2 HX + M2SO4 (X = F, Cl, NO3)
Sulfuric acids
• Both metals and non-metals are oxidised by
concentrated sulphuric acid, which is reduced to SO2.
• 2 H2SO4 + 2 e- → SO2 + 2 H2O + SO42• 2Ag + 2H2SO4(conc)  Ag2SO4 + SO2 + 2H2O
• H2SO4(conc) + H2S  S¯ + SO2 + 2H2O
• 3H2SO4(conc) + 2KBr  SO2 + Br2 + 2H2O + 2KHSO4
• 2H2SO4(conc) + C  2SO2 + CO2 + 2H2O
• Cu + 2 H2SO4 → SO2 + 2 H2O + SO42- + Cu2+
Sulfuric acids
• Dilute sulfuric acid reacts with metals via a single
displacement reaction as with other typical acids,
producing hydrogen gas and salts (the metal sulfate). It
attacks reactive metals (metals at positions above copper in
the reactivity series) such as Fe, Al, Zn,Mn, Mg and Ni.
Fe (s) + H2SO4 (aq) → H2 (g) + FeSO4 (aq)
• As shown above, the concentrated sulfuric acid is a strong
oxidizing agent, especially when hot in which the acid itself
carries out reduction whereas the dilute one acts a typical
acid where the H+ is the oxidizing agent so the oxidizing
property of it is not as powerful as the concentrated one.
• Lead and tungsten, however, are resistant to sulfuric acid.
Preparation of Metal Sulphides
from the Elements. HgS, Cu2S,
ZnS and Al2S3 are prepared by
conversion of the metals with
sulphur.
Redox Reactions of the Sulphite
Ion. Demonstration of Redoxamphotheric Behaviour of
Sulphite
Bleaching of a red rose with SO2.
The reaction of the plant
pigment with sulphur dioxide
leads to a stable adduct.
Decomposition of
Thiosulphuric Acid. Sodium
thiosulphate is decomposed in
acidic solution.