SULPHUR AND ITS COMPOUNDS
Sulphur is a non-metal in the same family (Group VI) with oxygen in the Periodic Table. Sulphur is different from oxygen in that:
– it is a solid at room temperature
– it shows the valencies of -2, +4 and +6
The electronic configuration of sulphur is 1S2,2S2,2P6, 3S2,3P4 (2,8,6). It gains two electrons from a metallic
element to form an ionic bond or shares two pairs of electrons with other non-metallic element to form
covalent bond.
Allotropes of sulphur
Allotropy is a phenomenon whereby an element occurs in two or more free forms in the same physical
state. The different forms are called isotopes. The main allotropes of sulphur are:
– Rhombic sulphur (α sulphur)
– Monoclinic or prismatic sulphur (β sulphur)
– Amorphous sulphur (δ sulphur) and
– Plastic sulphur
Rhombic sulphur
It is the only stable allotrope of sulphur at temperature below 960C. They are bright yellow and octahedral crystals. They are made up of S8 molecules. Each S8 molecule consists of a ring of 8 atoms.
Rhombic sulphur is prepared by allowing a saturated solution of sulphur in carbon (IV) sulphide to evaporate slowly. The octahedral structure will gradually be deposited. The preparation is done in the fume cupboard because of the poisonous and highly inflammable nature of carbon (IV) sulphide.
Monoclinic sulphur
It is the only stable allotrope at temperature between 960C and 1190C. The crystals are long, thin and needle-shaped. They consist of S8 molecules. At room temperature, they slowly revert to rhombic
sulphur crystals. The S8 molecules in monoclinic sulphur are more tightly packed; hence, monoclinic sulphur is less dense than rhombic sulphur.
To prepare monoclinic sulphur, some powdered sulphur is heated in a crucible till it melts. Add more; continue heating and stirring at the same time. Repeat the process until the crucible is almost filled with molten sulphur. Allow the molten sulphur to cool. After some time, a hard crust will be formed at the top. Pierce two holes through the crust and pour off the remaining molten sulphur from inside before removing the crust. Needle-shaped crystals of monoclinic sulphur can be found as a deposit on the sides of the crucible. The relationship between rhombic and monoclinic sulphur is shown below.
Between 960C and 1190C
Rhombic sulphur monoclinic sulphur
Below 960C
Comparison between the two allotropes
Property Rhombic Monoclinic
Colour Bright yellow Amber
Shape Octahedral Needle-shaped
Density (g/cm3) 2.08 1.98
Melting point (0C) 113 119
Stability Below 960C Between 960C and 1190C
Note: 960C is the transition temperature between the two allotropes.
Amorphous sulphur
It has no regular crystalline shape. It is pale yellow in colour and can be prepared in two ways;
– By bubbling hydrogen sulphide through water for a long time and the saturated solution
exposed to the air; 2H2S(g) + O2(g) 2H2O(l) + 3S(s)
– By chemical reaction e.g action of dilute hydrochloric acid on a trioxothiosulphate (VI) solution;
S2O32-(aq) + 2H+(aq) H2O(l) + SO2(g) + S(S)
Plastic sulphur
It is generally not considered to be a true allotrope of sulphur because it is unstable and reverts to
rhombic sulphur on standing. It is regarded as a super-cooled form of sulphur. It is prepared by heating
yellow sulphur and pouring it in cold water. It rolls up into yellow ribbons which look like plastic
material. It is soft and elastic and does not dissolve in carbon (IV) sulphide.
Physical properties of sulphur
- It is a yellow solid which exists in two forms; amorphous and crystalline
- It is insoluble in water but soluble in carbon (IV) sulphide and methyl benzene (toluene)
- It is a bad conductor of heat and electricity
- The density depends on the allotropic form
5.The melting point is 1190C and the boiling point is 4440C
Chemical properties of sulphur
Sulphur is a Group VI element. It gains two electrons from Group I or II elements to form sulphide ion; S2-. It also shares electrons with other non-metals to obtain its octet structure. The oxidation state ranges from -2 to +6. The chemical properties include;
Direct combination with other elements.
– With metals: sulphur combines directly with many metals to form sulphides when heated in the
absence of air. Reactive metals may not require heat if the metals and the sulphur are finely
divided; Fe(s) + S(s) FeS(s)
– With oxygen: when heated in a plentiful supply of air, sulphur burns with a bright blue flame to form SO2 and a small amount of SO3; O2(g) + S(s) SO2(g)
– With carbon: sulphur combines with coke in an electric furnace to form a colourless liquid known as carbon (IV) sulphide ( CS2) which vaporizes to form poisonous and highly inflammable fumes;
C(s) + 2S(s) CS2(l)
– With other non-metals: sulphur combines with other non-metals to form various sulphides e.g. disulphide dichloride; S2Cl2, sulphur hexafluoride SF6.
Reaction with oxidizing acids
When warmed with conc. H2SO4, sulphur is oxidized to form sulphur (IV) oxide; SO2.
2H2SO4(aq) + S(s) 2H2O(l) + 3SO2(g)
With conc. HNO3 when warmed and using bromine as catalyst, sulphur is oxidized to H2SO4.
6HNO3(aq) + S(s) H2SO4(aq) + 6NO2(g) + 2H2O(l)
Action of hot conc. Alkalis:
When sulphur is reacted with hot. Conc. Alkaline solutions, a mixture of sulphides and trioxosulphate
(IV) Are formed. In the presence of excess sulphur, a polysulphide and a trioxothiosulphate (VI) are
formed respectively.
3S(s) + 6OH–(aq) 2S2- + SO32- + 3H2O
S2- + nS Sn+1
2- (where n = 1 – 8)
SO32- + S S2O32-
Uses of sulphur
– It is used to produced SO2 for the manufacturing of H2SO4
– It is used in the vulcanization of rubber
– Sulphur and some of its products are used as fungicides and insecticides for spraying crops.
– It is used to manufacture the bleaching agent used in the pulp and paper industry
– It is used for the production of CS2, skin ointments, dyes and as sulphides in the manufacture of matches, fireworks and gunpowder.
HYDROGEN SULPHIDE
It occurs naturally in volcanic gases, sulphur springs, coal gas and gases formed during the decay of
organic matter containing sulphur.
Preparation
Both in the laboratory and commercially, H2S is produced by the action of a dilute acid on a metallic sulphide e.g.
2HCl(aq) + FeS(s) FeCl2(aq) + H2S(g)
Ionically; 2H+(aq) + FeS(s) Fe2+(aq) + H2S(g)
The gas is easily liquefied and stored in steel cylinder for sale.
Physical properties of H2S
– It is a colourless gas with a repulsive smell like that of a rotten egg.
– It is very poisonous
– It is about 1.18 times denser than air
– It is moderately soluble in water to form a very weak acidic solution
– It burns with a pale blue flame.
Chemical properties
As an acid:
When H2S dissolves in water, it ionizes slightly to form a weak, dibasic acid;
H2S(g) + H2O(l) H3O+(aq) + HS–(aq)
HS–(aq) + H2O(l) H3O+(aq) + S2-(aq)
It reacts with alkali to produce normal salt and water but when H2S is in excess, an acidic salt will be produced.
2NaOH(aq) + H2S(g) Na2S(aq) + 2H2O(l)
NaOH(aq) + H2S(g) NaHS (aq)+ H2O(l)
Excess acid salt
As a precipitating agent
When H2S is bubbled through the solution of the salts of some metals, coloured sulphides are precipitated e.g.
ZnSO4(aq) + H2S(g) → ZnS(s) + H2SO4(aq)
White
Pb(CH3COO)2(aq) + H2S(g) → PbS(s) + 2CH3COOH(aq)
Black
Reaction with oxygen
Although it doesn’t support combustion, H2S burns in a plentiful supply of air with a bright blue flame to produce SO2. With limited supply of oxygen, a deposit of sulphur may be formed.
2H2S(g) + 3O2(g) → 2H2O(l) + 2SO2(g)
Excess supply
2H2S(g) + O2(g) → 2H2O(l) + 2S(s)
Limited supply
As a reducing agent
H2S releases 2 electrons to oxidizing agents in a redox reaction. The H2S is itself oxidized to sulphur.
H2S → 2H+ + S2-
S2- → S + 2e-
– With acidified KMnO4
2KMnO4(aq) + 3H2SO4(aq) + 5H2S(g) → K2SO4(aq) + 2MnSO4(aq) + 8H2O(l) + 5S(s)
Ionically; 2MnO4–(aq) + 6H+(aq) + 5H2S(g) → 2Mn2+(aq) + 8H2O(l) + 5S(s)
Purple pale pink
– With acidified K2Cr2O7
K2Cr2O7(aq) + 4H2SO4(aq) + 3H2S(g) → K2SO4(aq) + Cr2(SO4)3 (aq) + 7H2O(l) + 3S(s)
Ionically: Cr2O72-(aq) + 8H+(aq) + 3H2S(g) → 2Cr3+(aq) + 7H2O(l) + 3S(s)
Orange green
– With halogens:
In the presence of moisture, H2S is oxidized by halogens to sulphur.
Cl2(g) + H2S(g) → 2HCl(g) + S(s)
– With FeCl3:
2FeCl3(aq) + H2S(g) → 2FeCl2(aq) + 2HCl(aq) + S(s)
Brownish yellow green
– With SO2;
H2S is a stronger reducing agent than SO2. It reduces SO2 to sulphur in the presence of moisture.
SO2(g) + 2H2S(g) → 2H2O(l) + 3S(s)
– With oxidizing agents
Oxidizing acids convert hydrogen sulphide to sulphur. Meanwhile, conc. HNO3 being a very strong oxidizing acid converts H2S to H2SO4
H2SO4(aq) + H2S(g) → 2H2O(l) + SO2(g) + S(s)
2HNO3(aq) + H2S(g) → 2H2O(l) + 2NO2(g) + S(s)
8HNO3(aq) + H2S(g) → 4H2O(l) + 8NO2(g) + H2SO4(aq)
Test for H2S
The gas can be suspected by its odour. It smells like rotten egg. To confirm the gas, moisten a piece of filter paper with Pb(NO3)2 solution and drop it in a gas jar containing the unknown gas. If the gas is H2S, the paper will turn black due to the formation of PbS.
Pb(NO3)2(aq) + H2S(g) → PbS(s) + 2HNO3(aq)
Note: lead (II) ethanoate also gives the same reaction with H2S.
Uses of H2S
It is used in the analysis of ores and metals. The presence of certain metals is often indicated by its colour. Also the sulphides of metals have different solubility in acids and alkalis.
SULPHIDES
Almost all sulphides except Na, K and NH4 – sulphides are insoluble in water. Some of the insoluble sulphides are however soluble in dilute HCl while others are not.
Preparation of sulphides
– Neutralization: the three soluble sulphides are prepared by bubbling H2S through the corresponding alkalis e.g.
2NaOH(aq) + H2S(g) → Na2S(aq) + 2H2O(l)
– By direct combination: FeS and ZnS can be prepared by heating a mixture of the metal and
sulphur e.g. Fe(s) + S(s) → FeS(s)
– By precipitation: insoluble sulphides are prepared by bubbling H2S through a solution of the
corresponding salts e.g. SnCl2(aq) + H2S(g) → SnS(s) + 2HCl(aq)
Dark brown
Properties of sulphides
Reaction with oxygen
Except mercury (II) sulphide which yields the metallic mercury, most metallic sulphides when heated in air or oxygen produce the corresponding metallic oxides and sulphur (IV) oxide. This reaction is used for
the extraction of metals from their naturally occurring sulphides.
2PbS(s) + 3O2(g) → 2PbO(s) + 2SO2(g)
HgS(s) + O2(g) → Hg(l) + SO2(g)
Reaction with dilute hydrochloric acid
When heated with HCl(aq), most sulphides yield H2S. This reaction is used to test for sulphide. However, sulphides which are insoluble in dilute hydrochloric acid must first be fused with sodium
trioxocarbonate (IV) before the test is carried out.
PbS(S) + 2HCl(aq) → PbCl2(s) + H2S(g)
Uses of sulphides
– FeS is reacted with dilute hydrochloric acid in the kipp’s apparatus to obtain an immediate supply of H2S in the lab.
– ZnS is used to coat the inside of television screen as it fluoresces when exposed to light
– Tin (IV) sulphide is used in paints
Tutorial questions
– What is allotropy? List four allotropes of sulphur
– State four physical properties of hydrogen sulphide
– With two balanced chemical equations, show the reducing action of hydrogen sulphide
– State five physical properties of sulphur
– Describe the laboratory test for S2-.
TOPIC: SULPHUR AND ITS COMPOUNDS
TRIOXOSULPHATE (IV) ACID – H2SO3
It is an unstable acid formed when SO2 dissolves in water. It is dibasic and decomposes readily to yield SO2 and water.
SO2(g) + H2O(l) → H2SO3(aq)
Preparation of H2SO3
It is prepared by adding conc. HCl to Na2SO4 and heating the mixture gently. The SO2 liberated is then dissolved in water to form H2SO3.
Physical properties
– It is a colourless and unstable acid
– It smells strongly of SO2, which is a product of its decomposition.
– It turn blue litmus paper red.
Chemical properties
- Action on air: when exposed to air for sometimes, H2SO3 becomes oxidized by the atmospheric oxygen to H2SO4:
2H2SO3(aq) + O2(g) → 2H2SO4(aq)
- As a reducing agent and bleaching agent: H2SO3 is responsible for all the bleaching and reducing action of SO2.
- As an acid: H2SO3 as a weak dibasic acid forms both acid and normal salts with alkalis.
NaOH(aq) + H2SO3(aq) → NaHSO3(aq) + H2O(l)
2NaOH(aq) + H2SO3(aq) → Na2SO3(aq) + 2H2O(l)
Uses of H2SO3
– It is used for bleaching straw and other fabrics.
– It is used as a germicide
TRIOXOSULPHATES (IV)
These are normal salts formed when SO2 or H2SO3 reacts with excess alkalis e.g.
H2SO3(aq) + 2KOH(aq) → K2SO3(aq) + 2H2O(l)
SO2(g) + 2NaOH(aq) → Na2SO3(aq) + H2O(l)
Preparation of trioxosulphates (IV)
- By neutralization: SO2 dissolves in excess KOH and NaOH to produce K2SO3 and Na2SO3 respectively.
H2SO3 also reacts with excess of these alkalis to produce the corresponding salt.
- By precipitation: when SO2 is bubbled through a solution of metallic salt, the corresponding insoluble
trioxosulphate (IV) are precipitated out e.g.
SO2(g) + H2O(l) + Zn(NO3)2(aq) → ZnSO3(aq) + 2HNO3(aq)
Properties of trioxosulphates (IV)
- Solubility: only the trioxosulphates (IV) of NH4, K, Na and Ca are soluble in water. All other trioxosulphates (IV) are insoluble.
- Reaction with acids: they liberate SO2 when heated with conc. Or dilute acids e.g.
CaSO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + SO2(g)
- Reaction with air: when exposed to the air, trioxosulphates (IV) are slowly oxidized to tetraoxosulphates (VI). The reaction is more rapid when some other oxidizing agents are used.
- As a reducing agent: they are good reducing agents when in dilute acidified solution.
Test for trioxosulphates(IV)
– Warm the unknown substance with a dilute acid. If it is trioxosulphate(IV), SO2 will be evolved.
2H+(aq) + SO32-(aq) → SO2(g) + H2O(l)
– Add some Ba(NO3)2 to a solution of the unknown substance. A white precipitate of BaSO3 will be formed. If the unknown solution is trioxosulphate (IV), the precipitate will dissolve in dilute HNO3 with the evolution of SO2.
Ba2+(aq) + SO32- → BaSO3(s)
BaSO3(s) + 2HNO3(aq) → SO2(g) + Ba(NO3)2(aq)
TETRAOXOSULPHATE (VI) ACID – H2SO4
It is a heavy chemical. Almost all manufacturing processes make use of it directly or indirectly.
Industrial preparation of H2SO4 (contact process)
The steps involved can be summarised as follows;
– SO2 which is obtained by burning sulphur in dry air or by roasting sulphide ore is catalytically combined with oxygen to form SO3.
S(s) + O2(g) → SO2(g)
2SO2(g) + O2(g) → 2SO3(g) + heat
Note: SO2 is usually mixed with excess air and passed through an electric chamber to remove the impurities and dusts that could kill the catalyst. The gaseous mixture is then passed through conc. H2SO4 to remove moisture before the gases are delivered into the reaction chamber or contact chamber. The SO2 and O2 combine in the contact chamber in the presence of vanadium (V) oxide catalyst at a temperature of 4500C – 5000C and atmospheric pressure to yield SO3.
– The SO3 is cooled and passed into an absorption chamber where it dissolves in conc. H2SO4 to produce a very thick liquid called oleum;
SO3(g) + H2SO4(aq) → H2S2O7
Oleum
– The oleum is diluted with calculated amount of water to produce the 98% H2SO4 used in the labs. Other desired concentration can also be obtained depending on the amount of water added.
Note: SO3 is not dissolved directly in water because the process would evolve enough heat to cause the acid solution to boil, thereby producing a mist of acid droplets which would spread throughout the factory.
Contact process at a glance
O2 conc. H2SO4 H2O
S + O2 SO2 SO3 H2S2O7 H2SO4
Physical properties of H2SO4
– Conc. H2SO4 (oil of vitriol) is a colourless, viscous liquid with a density of 1.84gcm-3
– It is very corrosive
– It turn blue litmus paper red
– Conc. H2SO4 has a great affinity for water. Hence, it dissolves readily in water.
– Conc. H2SO4 is hygroscopic
Chemical properties of H2SO4
- As an acid: the acid properties of dilute H2SO4 is shown in its reaction with metals, bases and trioxocarbonates (IV) e.g.
H2SO4(aq) + Zn(s) → ZnSO4(aq) + H2(g)
H2SO4(aq) + MgO(s) → MgSO4(aq) + H2O(l)
H2SO4(aq) + CuCO3(aq) → MgSO4(aq) + H2O(l) + CO2(g)
Note: H2SO4 is a dibasic acid. It can form both acid and normal salts.
- As an oxidizing agent: Hot conc. H2SO4 accept electrons from reducing agents. The oxidation number
of sulphur hence changes from +6 to +4 e.g.
Zn(s) + 2H2SO4(aq) → ZnSO4(aq) + 2H2O(l) + SO2(g)
C(s) + 2H2SO4(aq) → CO2(g) + 2H2O(l) + 2SO2(g)
H2S(s) + H2SO4(aq) → SO2(g) + 2H2O(l) + S(s)
- As a dehydrating agent: Conc. H2SO4 removes water molecules from compounds like sugar, ethanol, methanoic acid and ethanedioic acid. This removal of water from compounds is called dehydration reaction.
C12H22O11(s) – 11H2O(l) 12C
Sucrose conc. H2SO4 sugar charcoal
CH2OH-H2O(l) C2H2(g) CH2OH (l) conc. H2SO4
The dehydration reaction of H2SO4 accounts for:
– The corrosive action of the acid on cloth, paper, wood, skin etc.
– Removal of water of crystallization from hydrated salts to produce anhydrous salts e.g.
CuSO4.5H2O → 5H2O(l) + CuSO4
Blue conc.H2SO4 white
- Displacement of other acids from their salts: conc. H2SO4 displaces volatile acids from their salts e.g.
KCl(s) + H2SO4(aq) → KHSO4(aq) + HCl(g)
NaNO3(s) + H2SO4(aq) → NaHSO4(aq) + HNO3(aq)
The reactions depend solely on the boiling point of the acids. HCl and HNO3 distil over as gases.
Uses of H2SO4
- Production of fertilizer; Ca(HSO4)2 and (NH4)2SO4
- Production of pigments e.g titanium (IV) oxide used in paint and dyes
- For making cellulose film, natural and artificial fabrics and plastics.
- Purification of crude oil and manufacturing of artificial silk
- To clean or pickle metals before electroplating or enamelling
- Used as drying agent for many gases except alkaline gas (NH3) and reducing gases like H2S
- Used in the preparation of many important chemical compounds like HCl, HNO3 etc
TETRAOXOSULPHATES (VI)
They are the normal salts formed when all the H+ in H2SO4 are replaced by metallic or ammonium ion. Most of them are crystalline and soluble in water. The insoluble ones are CaSO4, PbSO4, BaSO4 and Hg2SO4
Laboratory preparation
They are prepared in the laboratory by any of the standard methods used for the preparation of salts depending on the position of the metallic radical in the electrochemical series and the nature of the Tetraoxosulphate (VI). The table below shows the methods:
Metal Method of preparation of SO4
K
Na
By the action of dil. H2SO4 on the oxides, hydroxides or the trioxocarbonates (IV) of the metal e.g.
H2SO4(aq) + CaO(s) CaSO4(aq) + H2O(l)
Ca H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H2O(l)
Mg
Al
Zn
Fe
By the action of dil. H2SO4 on the metals, their oxides or their trioxocarbonates (IV) aluminium requires the action of conc. H2SO4 e.g.
H2SO4(aq) + Mg(s) MgSO4(aq) + H2(g)
H2SO4(aq) + ZnO(s) ZnSO4(aq) + H2O(l)
H2SO4(aq) + FeCO3(s) FeSO4(aq) + CO2(g) + H2O(l)
Pb. By double decomposition e.g.
Pb(NO3)2(aq) + Na2SO4(aq) PbSO4(s) + 2NaNO3(aq)
Cu
Hg
Ag
By the action of hot. Conc. H2SO4 on the metals e.g.
Hg(l) + 2H2SO4(aq) HgSO4(aq) + 2H2O(l) + SO2(g)
2Ag(s) + 2H2SO4(aq) Ag2SO4(aq) + 2H2O(l) + SO2(g)
Properties of tetraoxosulphates(VI)
– Most of them are colourless and crystalline
– The soluble ones form hydrated salts/crystals
– The tetraoxosulphates (VI) of Na, K and Ca are very stable to heat while those of metals lower in
the electrochemical series tend to decompose on heating.
– When the hydrated tetraoxosulphate (VI) of Zn2+, Cu2+ and Fe2+ are heated, they lose their water of crystallization in stages. Further heating of the anhydrous salts brings about their decomposition into metallic oxides, sulphur (IV) oxides and trioxosulphates (IV).
Uses of tetraoxosulphates (VI)
– Na2SO4 and MgSO4 are used as laxatives
– (NH4)2SO4 is used as fertilizer
– For the production of SO2
– BaSO4 is used in the production of paints
– For the production of plaster of Paris PoP
Test for tetraoxosulphate (VI)
Put some of the unknown solution in a test tube and acidify it with dilute HCl. Then add a few drops of BaCl2 solution. A white precipitate of BaSO4 is formed if the unknown solution contains SO42-
The precipitate remains insoluble in an excess of the dilute acid.
Ba2+(aq) + SO42-(aq) → BaSO4(s)
Note: the acidification process is important to prevent the precipitation of other insoluble barium compounds like BaSO3 or BaCO3 which are soluble in the acidic solution.
Dilute HNO3 and BaNO3 can also be used instead of HCl and BaCl2.
Tutorial questions
– Briefly describe the contact process
– State three physical properties of H2SO4
– Describe any two chemical properties of H2SO4. Balance chemical equations important
– State three uses of H2SO4.
– State a laboratory test to distinguish between SO32- and SO42-.
– State the method of preparation of the tetraoxosulphates (VI) of (i) lead (ii) zinc