10 the p-block elements
TRANSCRIPT
-
7/26/2019 10 the P-block Elements
1/33
10 The p-block Elements - 1
1100TThheepp--BBlloocckkEElleemmeennttss
Introduction
Groups IIIA to 0 constitute the p-block in the Periodic Table. The atoms of all the elements in
these groups have, besides two electrons in the s-orbitals, one or more valence electrons in
the p-orbitals. Unlike the s-block elements, which show similar chemistry and regular
variations in properties, elements of p-block, except the halogens, show more dissimilaritiesin their properties.
In general, p-block elements are mainly non-metals. Their first ionization enthalpies are high
because their atoms have large effective nuclear charge. Therefore, they have little tendency
to form positive ions and are more electronegative than s-block elements. When forming
compounds with metals, p-block elements tend to form negative ions by gaining electrons to
attain the octet configuration. However, they tend to form covalent bonds by sharing
electrons when combining with non-metals.
TThheeHHaallooggeennss
(I) General properties of the halogens
The elements fluorine, chlorine,
bromine, iodine and astatine are Group
VIIA elements. Astatine is radioactive.
Group VIIA elements are also called
halogens.
General properties of the halogens:
1. All the halogens are colored.
2. They have an outer electronic configuration of ns2np7.
3. As their outermost shell electrons are not effectively screened from nuclear
attraction, the force exerted on them is strong. Therefore, the halogens havecomparatively small atomic radii and high ionization enthalpies.
Chlorine
-
7/26/2019 10 the P-block Elements
2/33
10 The p-block Elements - 2
(II) Characteristic properties of the halogens
The halogens are very reactive. This is because they have seven valence electrons and
can complete the octet configuration either by gaining one electron or by sharing their
unpaired p electrons. They (except fluorine) can also expand their octet of electrons by
using the low-lying d-orbitals.
(i) Electronegativity
Electronegativity is a measure of the relative
tendency of an atom to attract a bonding
electron. Since the atoms of halogens are
relatively small with large effective nuclear
charge, they tend to attract one more electron
to complete the octet. Therefore, they have
very high electronegativity values. Their
electronegativities are the highest among the
elements in the same period and F is the most
electronegatice element. .
(ii) Electron affinity
The electron affinity is the amount of energy absorbed or released when one mole
of electrons is added to one mole of gaseous atoms or ions.
X(g) + eX
(g) HEA = ve
As halogen atoms are relatively small and they have high electronegativities, their
electron affinities are high as well. The value, HEA , gives an indication of how
easily a halogen atom forms a halide
ion by attracting an additional electron.
The electron affinities of halogens decrease from chlorine to astatine, with fluorine
breaking the trend. The decrease from chlorine to astatine is due to the increase in
atomic size and hence a smaller nuclear attraction for another electron. The low
electron affinity of fluorine is due to its small size, and hence large charge density.
The additional electron will experience a greater repulsion when approaching the
electron cloud of a fluorine atom.
-
7/26/2019 10 the P-block Elements
3/33
10 The p-block Elements - 3
(iii) Bonding and oxidation states
Electrons-in-boxes diagrams showing the various oxidation states of halogens.
Table. The various oxidation states of halogens in their compounds.
-
7/26/2019 10 the P-block Elements
4/33
10 The p-block Elements - 4
The outer electronic configuration of halogens is ns2np5, with one electron less
than that of a noble-gas configuration. As they are highly electronegative, the
halogen atoms can therefore accept an electron to form the halide ions, X. Hence,
all halogens have an oxidation state of 1 and form ionic bonds when combined
with metals. Besides, all halogen atoms can share their unpaired pelectrons with
other atoms to form covalent bonds. Depending on the electronegativity of the
atom which they combine with, halogens (except fluorine) can exhibit oxidation
state of +1.
Furthermore, with the exception of fluorine, all other halogens can expand their
octet of electrons by using the vacant, low-lying d-orbitals. Therefore, their
oxidation states range from 1 to +7 (excluding +2).
It should be noted that fluorine is always univalent. Since it does not have
low-lying d-orbitals and is the most electronegative element, it always has the
oxidation number 1 in its compounds.
(iv) Color
All halogens are colored. This is due to the absorption of visible light causing the
excitation of outer electrons to higher energy levels. The larger the atom, the less
the energy required for the excitation to occur. Therefore, small fluorine atoms
absorb high energy violet light for excitation and appear yellow; large iodine
atoms absorb low energy yellow light and appear violet.
Halogens exhibit different colors when dissolved in different solvents. As they are
molecular substances, halogens are not very soluble in water but are very soluble
in organic solvents such as 1,1,1trichloroethane.
Table. Colors of halogens.
ColorElement
color and state in water in 1,1,1-trichloroethane
F
Cl
Br
I
Pale yellow gas
Greenish yellow gas
Reddish brown liquid
Violet black solid
Pale yellow
Pale yellow
Yellow
Brown
Pale yellow
Yellow
Orange
Violet
(III) Variation in physical properties of the halogens
(i) Electronegativity: F > Cl > Br > I (see above notes)
(ii) Electron affinity: F < Cl > Br > I (see above notes)
(iii) Melting point and boiling point
In descending the group, there is a progressive increase in the sizes of the
halogen diatomic molecules. As a result, the van der Waals' forcesbetween
the molecules increase. So the melting points and boiling points of the
elements increase down the group. Fluorine and chlorine exist as gases,bromine as liquid and iodine as solid, at room temperature and pressure.
-
7/26/2019 10 the P-block Elements
5/33
10 The p-block Elements - 5
(IV) Variation in chemical properties of the halogens
Fluorine is the most reactive halogen. It reacts rapidly and forms stable compounds
with metals and non-metals. All other halogens react with metals to form metal
chlorides, the reactivity decreases down the halogen group, i.e. F2> C12> Br2> I2.
(i) Oxidizing power
All halogen are strong oxidizing agents. The halogens oxidize other substances,
themselves being reduced. The oxidizing power of halogens decreases in the
order of F2> C12> Br2> I2. The oxidizing ability depends on the electron affinity,
hydration enthalpy and enthalpy change of atomization, and can be illustrated by
a Born-Haber cycle of the reduction of halogens.
H = H
atom+ E.A. + Hhyd
Variation of melting points and boiling
points of halogens.
-
7/26/2019 10 the P-block Elements
6/33
10 The p-block Elements - 6
Halogens can oxidize both metals and non-metals.
Note: The extremely high oxidizing power of fluorine makes it one of the few
elements which can combine directly with a noble gas. Depending on the
reaction conditions and the amount of reagents present, xenon can form threedifferent fluorides, XeF2, XeF4and XeF6. For example, xenon combines with
fluorine around 500oC under high pressure to form xenon hexafluoride.
Xe(g) + F2(g) XeF2(s)
Xe(g) + 2F2(g) XeF4(s)
Xe(g) + 3F2(g) XeF6(s)
(ii) Reaction with sodium
All halogens combine directly with sodium to form sodium halides. The reactivitydecreases down the group from chlorine to iodine.
2Na(s) + C12(g) 2NaCl(s) H = 411 kJ mol1
2Na(s) + Br2(g) 2NaBr(s) H = 360 kJ mol1
2Na(s) + I2(g) 2NaI(s) H = 288 kJ mol1
The enthalpy change of formation of sodium chloride is the highest among the
three sodium halides. This is due to the small size of the chloride ion and the
correspondingly high lattice enthalpy of sodium chloride.
As the sizes of halide ions increase down the halogen group, the lattice
enthalpies of sodium halides decrease. Therefore, their enthalpy changes offormation decrease.
(iii) Reaction with Iron(II) ions
Aqueous chlorine and bromine oxidize green iron(II) ions to yellowish brown
iron(III) ions.
2Fe2+(aq) + C12(aq) 2Fe3+(aq) + 2Cl(aq) E= +0.59 V
2Fe2+(aq) + Br2(aq) 2Fe3+(aq) + 2Br
(aq) E
= +0.30 V
However, iodine is a mild oxidizing agent. Its oxidizing power is not strong
enough to oxidize iron(II) ions.
Table. Enthalpy changes of formation of halides from the
corresponding halogens of standard states.
-
7/26/2019 10 the P-block Elements
7/33
10 The p-block Elements - 7
The spontaneity of a reaction can be worked out by adding the standard electrode
potentials of the two half-reactions concerned. If the overall standard electrode
potential is positive in value, the reaction is spontaneous. Try to verify that Cl2
and Br2 can react with Fe2+ spontaneously while there would be no reaction
between I2and Fe2+, using the following standard electrode potential data.
Also, would I(aq) (as from KI solution) react with (reduce) Fe3+(aq) ?
(iv) Reaction with phosphorus
All halogens react with red phosphorus to form phosphorus halides. As
phosphorus has low-lying vacant 3d orbitals, it is able to form molecules withmore than eight electrons in its outermost shell. The type(s) of product formed
depends on the oxidizing power of the halogens.
Table. Standard electrode potentials.
-
7/26/2019 10 the P-block Elements
8/33
10 The p-block Elements - 8
1. Fluorine Its oxidizing power is so great that it forces phosphorus (or other
elements it combines with) to exhibit the maximum oxidation state. Thus,
fluorine forms only PF5.
2P(s) + 5F2(g) 2PF5(s)
2. Chlorine Due to the strong oxidizing power, Cl2 forms PCl5 as the only
product.2P(s) + 5Cl2(g) 2PCl5(s)
3. Bromine and iodine - Bromine and iodine are very mild oxidizing agents and
forms tribromide and triiodide respectively.
2P(s) + 3Br2(1) 2PBr3(l)
2P(s) + 3I2(1) 2PI3(l)
(v) Reaction with water
Fluorine is the most powerful oxidizing agent and it oxidizes water readily to
form hydrogen fluoride and oxygen.
2F2(g) + 2H2O(1) 4HF(aq) + O2(g)
Chlorine is less reactive than fluorine. Chlorine reacts with water to form
hydrochloric acid and chloric(I) acid (hypochlorous acid). A mixture of
hydrochloric acid and chloric(I) acid is often called chlorine water.
In the reaction, the oxidation number of chlorine increases and decreases
simultaneously, i.e. chlorine undergoes oxidation and reduction at the same time.
This is an example of disproportionation.
Disproportionation is a reaction in which an element in the free state or in a
compound undergoes simultaneous reduction and oxidation.
When chlorine water is exposed to sunlight or high temperatures, oxygen isformed because the chlorate(I) ion is unstable and decomposes when exposed to
sunlight or high temperatures.
2OCl(aq) 2Cl
(aq) + O2(g)
The bleaching action of Cl2is due to the unstable OClwhich reacts with (oxidize)
colored dye to form colorless compounds.
C12(g) + H2O(1) HCl(aq) + HOCl(aq)0 +1
oxidation
reduction
C12(g) + H2O(1) 2H+(aq) + Cl
(aq) + OCl
(aq)
OCl(aq) + dye Cl(aq) + (dye+O)
colored compound colorless compound
-
7/26/2019 10 the P-block Elements
9/33
10 The p-block Elements - 9
Bromineis only slightly soluble in water and disproportionate in water to form.
hydrobromic acid and bromic(I) acid (hydrobromous acid).
Br2water also bleaches and the bleaching action is due to OBr(aq).
Iodineis only very slightly soluble in water and does not react with it. However,iodine is extremely soluble in potassium iodide solution and it exists as triiodide
ions in the solution.
I2(s) + KI(aq) KI3(aq) or I2(s) + I(aq) I3
(aq)
(vi) Reaction with alkalis
All halogens react with aqueous alkalis and undergo disproportionation. However,
they react differently under cold, hot, dilute and concentrated conditions. In general,
their reactivities decrease down the group.
Fluorinereacts with cold and very dilute (2%) sodium hydroxide solution.2F2(g) + 2NaOH(aq) 2NaF(aq) + OF2(g) + H2O(l)
With more concentrated sodium hydroxide solution, oxygen is formed:
2F2(g) + 4NaOH(aq) 4NaF(aq) + O2(g) + 2H2O(l)
Chlorinereacts with cold dilute sodium hydroxide solution to form sodium chloride
and sodium chlorate(I) (sodium hypochlorite).
C12(g) + 2NaOH(aq) NaCl(aq) + NaOCl(aq) + H2O(l)
With hot concentrated sodium hydroxide solution, chlorine forms sodiumchloride and sodium chlorate(V).
3C12(g) + 6NaOH(aq) 5NaCl(aq) + NaClO3(aq) + 3H2O(l)
Bromine and iodineundergo similar reactions with cold and dilute NaOH. However,
NaOBr and NaOI formed are not stable and disproportionate to bromide and
bromate(V), and iodide and iodate, respectively. The overall reaction would be:
(IV) Variation in properties of the compounds of the halogens
(i) Comparative study of the reactions of halide ions
(a) Reaction with halogens
The oxidizing power of halogens decreases down the group. Hence,
chlorine displaces bromide and iodide ions while bromine can only displace
iodide but not chloride ions.
C12(aq) + 2Br(aq) 2C1
(aq) + Br2(aq)
pale yellow colorless colorless yellow
C12(aq) + 2I(aq) 2Cl
(aq) + I2(aq)
Br2(g) + H2O(1) HBr(aq) + HOBr(aq)
3Br2(g) + 6NaOH(aq) 5NaBr(aq) + NaBrO3(aq) + 3H2O(l)
3I2(g) + 6NaOH(aq) 5NaI(aq) + NaIO3(aq) + 3H2O(l)
-
7/26/2019 10 the P-block Elements
10/33
10 The p-block Elements - 10
pale yellow colorless colorless brown
Br2(aq) + 2I(aq) 2Br
(aq) + I2(aq)
yellow colorless colorless brown
However, it is sometimes difficult to determine whether certain reactions have
taken place by examining the color changes only, especially when the reaction
involves the formation of Br2 and I2 as their colors are quite similar. Todetermine whether an aqueous solution contains bromine or iodine,
1,1,1-trichloroethane is added to the solution. Br2forms an orange red bottom
layer while I2forms a violet bottom layer.
(b) Reaction with concentrated sulphuric(VI) acid
1. Metal chloride
NaCl(s) + H2SO4(1) NaHSO4(s) + 2HCl(g)
(Conc. H2SO4acts as a non-volatile acid.)
2. Metal bromideNaBr(s) + H2SO4(1) NaHSO4(s) + 2HBr(g)
HBr further reacts with conc. H2SO4 to give sulphur dioxide and
bromine.
2HBr(g) + H2SO4(l) SO2(g) + Br2(g) + 2H2O(l)
(Conc. H2SO4acts as an oxidizing agent.)
3. Metal iodide
NaI(s) + H2SO4(1) NaHSO4(s) + 2HI(g)
HI further reacts with conc. H2SO4 to give hydrogen sulphide and iodine.
8HI(g) + H2SO4(l) H2S(g) + 4I2(g) + 4H2O(l)
(c) Reaction with concentrated phosphoric(V) acid
Phosphoric(V) acid is not a strong oxidizing agent. Hence, it reacts with
halides to form hydrogen halides which provides a general method for
preparing hydrogen halides in laboratory.
3NaCl(s) + H3PO4(l ) Na3PO4(s) + 3HCl(g)
3NaBr(s) + H3PO4(l ) Na3PO4(s) + 3HBr(g)3NaI(s) + H3PO4(l ) Na3PO4(s) + 3HI(g)
Unlike conc. H2SO4, conc. H3PO4 just acts as an acid (non-volatile), no
further reaction occurs.
Steamy fumes are formed as conc. H3PO4 reacts with the metal halide,
which forms dense white fumes with ammonia. This is a confirmatory test
for halides.
(d) Reaction with silver ion (from silver nitrate(V))
Halide ions react with Ag+to form precipitates of AgCl, AgBr nad AGI.
Ag+(aq) + Cl(aq) AgCl(s)
-
7/26/2019 10 the P-block Elements
11/33
10 The p-block Elements - 11
white precipitate
Ag+(aq) + Br(aq) AgBr(s)
pale yellow precipitate
Ag+(aq) + IAgI(s)
yellow precipitate
If excess aqueous ammonia is added to the precipitates, silver chloride
dissolves readily due to the formation of a soluble complex,diamminesilver(I) chloride.
AgCl(s) + 2NH3(aq) [Ag(NH3)2]Cl(aq)
Silver bromide is slightly soluble in aqueous ammonia and silver iodide
is insoluble.
When the precipitates are exposed to sunlight, silver chloride turns grey,
silver bromide turns yellowish grey, and silver iodide remains yellow.
This is due to photo-decomposition of the halides into their elements.
2AgCl(s) 2Ag(s) + C12(g)
2AgBr(s) 2Ag(s) + Br2(g)
Test for halide ions:
Excess dilute HNO3* is added to the sample, followed by addition of
AgNO3.
The sample contain Cl if a white precipitate which is soluble in ammonia
is formed.
The sample contain Br if a pale yellow precipitate which is slightly
soluble in ammonia is formed.
The sample contain I
if a yellow precipitate which is insoluble inammonia is formed.
*Note: Excess dilute nitric(V) acid must be added to the aqueous solutions of
halides and silver nitrate(V) to prevent the precipitation of other
insoluble silver compounds such as silver carbonate and silver
sulphate(IV).
2Ag+(aq) + CO32-(aq) Ag2CO3(s)
2Ag+(aq) + SO42-(aq) Ag2SO4(s)
(ii) Acidic properties of hydrogen halides
(a) Energetics of the hydrohalic acids
Pure hydrogen halides are predominantly covalent in nature. All hydrogen
halides are soluble in water and react with water to give an acidic solution
according to the general equation.
HX(g) + H2O(1) H3O+(aq) + X
(aq)
The steps involved are:
1. the breaking of the hydrogen-halogen bond;
2. the formation of the oxygen-hydrogen bond;3. the hydration of the hydrogen ion; and
4. the hydration of the halide ion.
-
7/26/2019 10 the P-block Elements
12/33
10 The p-block Elements - 12
When HX is HCI, HBr or HI, the energy liberated by the hydration of the
hydrogen ion (step 3) and halide ion (step 4) are greater than the amount of
energy required to break the hydrogen-halogen bond (step 1). All three
hydrogen halides are therefore very strong acids in water.
(b) Acid strengths
The acid strength of hydrogen halides decreases in the order:HI > HBr > HCl > HF
1. Hydrogen iodide is a very strong acidas the HI bond is weak due to
the large size of the iodine atom. The bond dissociation enthalpy of HI is
low. Hydrogen bromide and hydrogen chloride are quite strong acids as
well because they have small bond dissociation enthalpies.
2. HF is the weakest acid. It has very great bond dissociation enthalpy,
due to the small size of fluorine and the short HF bond length.
Moreover, the extensive hydrogen bonds among hydrogen fluoride
molecules make the dissociation process more difficult. The following
equilibrium lies essentially to the left. Hence, hydrogen fluoride is aweak acid in dilute aqueous solution.
Hydrogen
halide
Dissociation constant
Kc(mol dm3)
Degree of dissociation
in 0.1 M solution
HF
HCl
HBr
HI
7 104
1 107
1 109
1 1011
0.08
1.00
1.00
1.00
3. Concentrated HF is a strong acid.
In conc. HF solution, F reacts with undissociated HF molecules to form
HF2, hydrogen difluoride ion.
Removal of F(aq) shifts the following equilibrium to the right, thus
increasing the acid strength of HF. At a concentration of 5M to 15M, HF
behaves a strong acid.
The hydrogen difluoride ion is a resonance hybrid between:
and each of the resonance structures involves hydrogen bonding.
HF(l) + H2O(l) H3O+(aq) + F
(aq) Ka= 710
4mol dm3
F(aq) + HF(aq) HF2
(aq) K = 5.1 dm3mol
1
HF(l) + H2O(l) H3O+(aq) + F
(aq)
-
7/26/2019 10 the P-block Elements
13/33
10 The p-block Elements - 13
If a fluoride salt (e.g. KF) is dissolved in aqueous hydrofluoric acid, then
the ions present are K+and HF2. On evaporation, potassium hydrogen
difluoride (KHF2) is obtained.
Anhydrous HF can be obtained by heating the KHF2 solid by reversingthe above equilibrium.
4. Etching property of HF
Although hydrogen fluoride is a weak acid, it is very reactive. It will readily
etch glass. The glass object to be etched is coated with wax or a similar
acid-proof material. The pattern to be produced is cut through the wax layer
to expose the glass below. HF reacts with the silicate of the glass where
there is no protective coating.
CaSiO3(s) + 6HF(aq) CaF2(aq) + SiF4(aq) + 3H2O(1)
Note: Since HF attacks glass, hydrofluoric acid is stored in rubber or wax
bottles.
(V) Uses of halogens and halogen-containing compounds
Halogens are seldom used directly because of their high reactivity and toxicity. Halogen
compounds, however, are chemically very stable. They are used extensively in industry,
agriculture, medicine and households.
(i) Fluorine
Fluorine is used to make poly(tetrafluoroethene), which is commonly known
as teflon. It is used as a non-stick coating for frying pans.
Used to make Freon (difluorodichloromethane), a refrigerant gas and a
propellant for aerosols.
(ii) Fluoride
Sodium hexafluorosilicate, Na2SiF6, or sodium fluoride, NaF, is used to
fluoridate drinking water in Hong Kong to help decrease incidence of tooth
decay.
(iii) Chlorine
Chlorine is the raw material for the production of chloroethene, CHCl=CH2,
which can be polymerized to give poly(chloroethene) (polyvinyl chloride or
PVC). This polymer is widely used in electrical insulation, pipe making, etc.
Chlorine is used in the manufacture of industrial and domestic bleaches..
C12(g) + 2NaOH(aq) NaCl(aq) + NaOCl(aq) + H2O(1)
KF (aq) + HF(aq) KHF2
-
7/26/2019 10 the P-block Elements
14/33
10 The p-block Elements - 14
Used as a disinfectant and germicide, e.g. used in sterilizing water, sewage and
swimming pools.
(iv) Silver bromide
Silver bromide is coated on films for black and white photography. On exposure
to light, silver bromide decomposes to silver:
When the film is developed, the unexposed silver bromide is removed by some
chemicals, and the silver remains on the film as an opaque shadow.
(v) Iodine and iodide
Iodine dissolved in alcohol, water or potassium iodide is commonly called an
iodine tincture which is widely used as an antiseptic for cuts and wounds.
Iodine-131 is used in medical diagnosis to monitor and trace the flow of
thyroxine from the thyroid gland.
Iodide ions are added to table salt (sodium chloride) to prevent goiter.
2AgBr(s) 2Ag(s) + Br2(g)light
-
7/26/2019 10 the P-block Elements
15/33
10 The p-block Elements - 15
NNiittrrooggeennaannddiittssccoommppoouunnddss
(I) Nitrogen
General properties of nitrogen
(i) Nitrogen is the first member of Group VA in the Periodic Table. Its electronicconfiguration is 1s22s22p5. It is a colorless, odorless gas and is the major
component (78% by volume) of the atmosphere.
(ii) Nitrogen has very low melting (210oC) and boiling points (196oC) . It is
slightly less dense than air, slightly soluble in water and does not support
combustion.
(iii) Nitrogen can form a large number of inorganic compounds with reactive metals,
hydrogen and oxygen. It is also a major constituent of some organic compounds
such as amines, amino acids, amides, etc.
(iv) Unreactive nature of nitrogen
This is because its diatomic molecules are non-polar and the nitrogen atoms are
held together by very strong triple covalent bonds. The bond enthalpy of the triple
covalent bonds in nitrogen is 944 kJ mol1, which is much higher than the bond
enthalpies of other common bonds such as O=O, HH, CC, etc.
.
Table. Common compounds of nitrogen of different
oxidation numbers.
-
7/26/2019 10 the P-block Elements
16/33
10 The p-block Elements - 16
As a result, reactions involving nitrogen usually have high activation energies and
unfavorable equilibrium constants. For example, nitrogen and oxygen do not
combine to form nitrogen monoxide at 25oC, as the equilibrium constant of the
reaction is 4.5 1031.
Besides, a catalyst, high temperature and pressure may be required for nitrogen toreact. For example, the conditions for synthesis of ammonia from nitrogen and
hydrogen by the Haber Process are 500oC, 500 atm and using Fe as catalyst.
Reactions of nitrogen
(i) With metals
Nitrogen reacts with reactive metals such as lithium and magnesium when
heated to form metal nitrides.6Li(s) + N2(g) 2Li3N(s)
lithium nitride
3Mg(s) + N2(g) Mg3N2(s)
magnesium nitride
*Note: Burning magnesium in air produces both magnesium oxide and
magnesium nitride.
(ii) With oxygen
In nature, lightning causes nitrogen and oxygen to react to give nitrogen
monoxide, which immediately combines with oxygen in the air to give
nitrogen dioxide (a poisonous reddish brown gas with a pungent smell).
N2(g) + O2(g) 2NO(g)
2NO(g) + O2(g) 2NO2(g)
Nitrogen monoxide can also be formed from the reaction between nitrogen and
oxygen at high temperatures in motor car engines. The nitrogen monoxide
formed will be emitted to the air and further oxidized to nitrogen dioxide.
(iii) With hydrogen
In the presence of iron as the catalyst, nitrogen reacts with hydrogen to give
ammonia at high temperature and pressure..
(II) Ammonia
Ammonia is a colorless, pungent gas. It consists of polar trigonal pyramidal NH3molecules with a lone pair of electrons on the nitrogen atom. Because of the existence
of hydrogen bonds, gaseous ammonia is extremely soluble in water and is easily
condensed to liquid ammonia (with boiling point 33oC). Like water, liquid ammonia is
an excellent solvent for ionic compounds. Its aqueous solution is weakly alkaline.
N2(g) + O2(g) 2NO(g) Kc= 4.5 1031at 25oC
H = +180.5 kJ
N2(g) + 3H2(g) 2NH3(g)
NH3(aq) + H2O(l) 2NH4+(aq) + OH
(aq) Kb = 1.8 10
5
-
7/26/2019 10 the P-block Elements
17/33
10 The p-block Elements - 17
Manufacture of ammonia by the Haber Process
Ammonia is manufactured by direct combination of nitrogen and hydrogen in the
Haber Process.
Raw material: Nitrogen gas is obtained by the fractional distillation of liquefied air.Hydrogen is obtained from the reaction of naphtha or natural gas with
steam in a process called steam reforming. This involves passing the
gaseous hydrocarbon and superheated steam under a pressure of 10
atm over heated nickel catalyst at 700oC.
C5H12(g) + 5H2O(g) 5CO(g) + 11H2(g)
(from naphtha)
CH4(g) + H2O(g) CO(g) + 3H2(g)
The mixture of CO and H2 is further mixted with steam and passed
over a heated catalyst. The CO2produced is dissolved in water under
pressure.CO(g) + H2(g) + H2O(g) CO2(g) + 2H2(g)
Purification of reactants: Impurities would poison the catalyst for synthesis of
ammonia and therefore should be removed.
Process: Nitrogen and hydrogen are mixed in the ratio of 1 : 3 by volume and
react in the catalytic chamber at 500oC and 200 atm using finely
divided iron as catalyst. As the reaction is reversible, it does not go to
completion and the yield of ammonia is about 15%.
The hot gaseous mixture leaving the catalytic chamber containing
about 15% of ammonia is passed through the heat exchanger. This
heats up the incoming gaseous reactants and the ammonia is also
cooled.
The ammonia formed is removed from the gaseous mixture by
cooling and liquefied under pressure in the condenser. The unreacted
nitrogen and hydrogen are then recycled.
N2(g) + 3H2(g) 2NH3(g) H = 92 kJ mol1
-
7/26/2019 10 the P-block Elements
18/33
10 The p-block Elements - 18
Choice of reaction conditions - physicochemical principles
Equation for the reaction
(i) Effect of pressure on equilibrium
According to the equation, there are 4 moles of reactants at the left and 2 moles of
product at the right. According to Le Chatelier's principle, a high pressure will
shift the equilibrium to right and therefore increases the yield of the product.
However, with increasing pressure, the cost of industrial plant becomes more
expensive. The typical pressure used is between 200-1000 atmospheres,
depending on the scale of the plant.
N2(g) + 3H2(g) 2NH3(g) H = 92 kJ mol1
Change of % yield of ammonia with temperature and pressure.
-
7/26/2019 10 the P-block Elements
19/33
10 The p-block Elements - 19
(ii) Effect of temperature on equilibrium
Since the synthesis of ammonia is exothermic, a lower temperature will lead to a
higher yield of ammonia. However, a chemical reaction carried out at low
temperature is likely to be too slow. Decisions have to be made between high
yield in a long time and a lower yield in shorter time. In order to have the reaction
proceed quickly while maintaining an acceptable yield, the reaction is carried out
at about 500o
C, which is the optimum temperature for the reaction.
(iii) Effect of catalyst on equilibrium
Even with the above conditions of pressure and temperature, catalysts are
commonly used in chemical industries to speed up the reaction.
Finely divided iron is used in the Haber Process as it is effective and not easily
poisoned. It increases the rate of reaction and shortens the time to reach
equilibrium. However, it does not change the position of equilibrium and is
chemically unchanged at the end of the reaction.
Chemical properties of ammonia
(i) As a base
Ammonia is very soluble in water. However, only a little of the dissolved
ammonia gas reacts with water to form ammonium ions and hydroxide ions, by
accepting a hydrogen ion from water to its lone pain
The presence of hydroxide ions in water turns litmus blue, methyl orange yellowand phenolphthalein red. Ammonia behaves as a weak base in an aqueous
solution.
(ii) Reaction with acids
Ammonia solution neutralizes acids to give ammonium salts. For example,
2NH3(aq) + H2SO4(aq) (NH4)2SO4(aq)
ammonium sulphate(VI)
NH3(aq) + HNO3(aq) NH4NO3(aq)
ammonium nitrate(V)
When opened bottles of concentrated hydrochloric acid and concentrated
ammonia are brought near, white fumes of ammonium chloride is given out.
NH3(g) + HCl(g) NH4Cl(s)
(iii) Reaction with metal salts
Ammonia solution can precipitate insoluble metal hydroxides from solutions of
metal salts. The reaction is due to the OHions formed from ionization of NH3in
water. For example,
Ca2+(aq) + 2OH
(aq) Ca(OH)2(s) white precipitateMg2+(aq) + 2OH
(aq) Mg(OH)2(s) white precipitate
Al3+(aq) + 3OH(aq) Al(OH)3(s) white precipitate
NH3(aq) + H2O(l) 2NH4+(aq) + OH
(aq) Kb = 1.8 10
5
-
7/26/2019 10 the P-block Elements
20/33
10 The p-block Elements - 20
Zn2+(aq) + 2OH(aq) Zn(OH)2(s) white precipitate
Fe2+(aq) + 2OH(aq) Fe(OH)2(s) dirty green precipitate
Fe3+(aq) + 3OH(aq) Fe(OH)3(s) reddish brown precipitate
Pb2+(aq) + 2OH(aq) Pb(OH)2(s) white precipitate
Cu2+(aq) + 2OH(aq) Cu(OH)2(s) blue precipitate
Hydroxides of zinc and copper dissolve in excess ammonia solution to form
complex compounds.Zn(OH)2(s) + 4NH3(aq) Zn(NH3)4
+(aq) + 2OH(aq)
colorless solution
Cu(OH)2(s) + 4NH3(aq) Cu(NH3)42+(aq) + 2OH
(aq)
deep blue solution
Silver chloride is insoluble in water or acid. However, it dissolves in excess
ammonia solution. This is because ammonia shifts the equilibrium to the right by
taking up the Ag+ion in the form of the complex ion Ag(NH3)2+.
(iv) As a reducing agent
The oxidation number of N in ammonia is 3, the lowest possible value for N.
Therefore, it cannot be further reduced. It can be oxidized quite easily and is a
fairly strong reducing agent.
1. Reaction with oxygen
Ammonia burns in oxygen (but not in air) with a yellow flame, givingnitrogen and water.
4NH3(g) + 3O2(g) 2N2(g) + 6H2O(g)
Catalytic oxidation
Ammonia reacts with oxygen to form nitrogen(II) oxide in the presence red
hot platinum or copper which act as catalysts.
This is a key reaction in the industrial preparation of HNO3from ammonia.
In the laboratory, the reaction can be carried out using the following set-up.
AgCl(s) Ag+(aq) + Cl(aq)
Ag+(aq) + 2NH3(aq) Ag(NH3)2+(aq)
Oxidation of ammonia in
air enriched with oxygen.
4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)Pt
Heat
-
7/26/2019 10 the P-block Elements
21/33
10 The p-block Elements - 21
2. Reaction with copper(II) oxide
By passing ammonia over heated copper(II) oxide, ammonia is oxidized to
nitrogen and water.
2NH3(g) + 3CuO(s) 3Cu(s) + N2(g) + 3H2O(g)
(III) Nitric acid
Nitric(V) acid, HNO3, is a very strong acid. It turns yellow on storage because of the
dissolved nitrogen dioxide formed from the decomposition of some of the acid.
4HNO3(1) 4NO2(aq) + 2H2O(1) + O2(g)
As light will speed up this decomposition, HNO3acid is usually kept in a brown bottle
to avoid exposure to light.
Nitric(V) acid is commonly used in making dyes, explosives, nylon and fertilizers such
as ammonium nitrate.
Industrial preparation of nitric(V) acid
HNO3is manufactured from the catalytic oxidation of ammonia in the Ostwald process
(i) Catalytic oxidation of NH3to NO
Ammonia is oxidized to NO in the presence of red hot platinum-rhodium which
acts as a catalyst at about 850oC.
4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)
(ii) Oxidation of NO to NO2
The colorless nitrogen monoxide then reacts with oxygen to give nitrogen dioxide
which is a brown gas.
2NO(g) + O2(g) 2NO2(g)
(iii) Dissolving NO2in water in the presence of excess O2The resulting nitrogen dioxide is dissolved in water in the presence of oxygen,
giving nitric(V) acid.
Catalytic oxidation of
ammonia.
-
7/26/2019 10 the P-block Elements
22/33
10 The p-block Elements - 22
2NO2(g) + O2(g) + 2H2O(1) 4HNO3(aq)
The product is distilled to give concentrated nitric(V) acid, with its azeotrope
containing 68.5% HNO3by mass (15 M).
Note: An azeotrope is a liquid mixture of two or more substances that has the
same composition in the vapor and liquid state when distilled or partially
evaporated under a certain pressure.
Oxidizing properties of nitric(V) acid
Nitric(V) acid is a strong oxidizing agent. Half equations for
dilute or moderately concentrated nitric(V) acid:
NO3(aq) + 4H++ 3e
NO(g) + H2O(l)
concentrated nitric(V) acid:
NO3(aq) + 2H++ e
NO2(g) + H2O(l)
The electrons are supplied by reducing agents and HNO3acts as an electron acceptor.
(i) Reaction with copper
With dilute HNO3
Copper reacts with dilute HNO3 to give nitrogen monoxide,
which reacts immediately with atmospheric oxygen to give a
brown gas of nitrogen dioxide.
3Cu(s) + 8HNO3(aq)
3Cu(NO3)2(aq) + 4H2O(1) + 2NO(g)[3Cu(s) + 8H++ 2NO3(aq) 3Cu2+(aq) + 2NO(g) + 4H2O(l)]
2NO(g) + O2(g) 2NO2(g)
With concentrated HNO3
Copper reacts with concentrated HNO3 (about 14 M) to give a brown gas of
nitrogen dioxide.
Cu(s) + 4HNO3(aq) Cu(NO3)2(aq) + 2H2O(1) + 2NO2(g)[Cu(s) + 4H++ 2NO3
(aq) Cu2+(aq) + 2NO2(g) + 2H2O(l)]
(ii) Reaction with iron(I1) ions
Dilute HNO3 can oxidize iron(II) compounds to iron(III) compounds. HNO3 is
reduced to nitrogen monoxide, which then turns to nitrogen dioxide in air.
3Fe2+(aq) + NO3(aq) + 4H+(aq) 3Fe3+(aq) + NO(g) + 2H2O(l)
pale green yellow / pale brown
2NO(g) + O2(g) 2NO2(g)
-
7/26/2019 10 the P-block Elements
23/33
10 The p-block Elements - 23
(iii) Reaction with sulphur
Sulphur reacts with hot concentrated HNO3to give H2SO4and nitrogen dioxide.
S(s) + 6HNO3(aq) H2SO4(aq) + 6NO2(g)+ 2H2O(1)
(IV) Nitrates(V)
Metal nitrates can be prepared by reactions of dilute nitric(V) acid with metal oxides,hydroxides or carbonates. For example,
CuO(s) + 2HNO3(aq) Cu(NO3)2(aq) + H2O(l)
NaOH(aq) + HNO3(aq) NaNO3(aq) + H2O(1)
Na2CO3(aq) + 2HNO3(aq) 2NaNO3(aq) + H2O(1) + CO2(g)
Mg(s) + 2HNO3(aq) Mg(NO3)2(aq) + H2(g)
very dilute
Action of heat on nitrates(V)
When strongly heated, solid metal nitrates(V) decompose differently according to their
thermal stabilities, which in turn depend on the positions of the metals in the reactivity
series.
Brown ring test for nitrate(V) ions
The presence of NO3 ions in a solution can be confirmed by the brown ring test. Fresh
iron(II) sulphate solution is mixed with a solution suspected of containing NO3ions in a
test tube. Concentrated sulphuric(VI) acid is then added carefully along the side to the
bottom of the test tube with the test tube tilted. After a while, abrown ring appears at the
junction of the two layers.
-
7/26/2019 10 the P-block Elements
24/33
10 The p-block Elements - 24
Equations for the reactions involved:
1. NO3ions react with concentrated H2SO4to give HNO3:
NO3(aq) + H2SO4(l) HNO3(aq) + HSO4
(aq)
2. HNO3oxidize iron(II) sulphate (FeSO4) to iron(III) sulphate (Fe2(SO4)3), itself is
reduced to nitrogen monoxide:
3Fe2+(aq) + HNO3(aq) + 3H+(aq) 3Fe3+(aq) + NO(g) + 2H2O(l)
3. Finally, nitrogen monoxide adds on to the excess iron(II) sulphate to form a brown
complex which forms the brown ring.
FeSO4(aq) + NO(g) FeSO4 NO(aq)
Freshly prepared
FeSO4(aq) and suspected
NO3(aq)
Concentrated H2SO4
Concentrated
H2SO4
Brown ring
Carrying out the brown ring test
-
7/26/2019 10 the P-block Elements
25/33
10 The p-block Elements - 25
SSuullpphhuurraannddiittssccoommppoouunnddss
(I) Sulphur
General Properties of Sulphur
Sulphur is the second member of Group VIA in the Periodic Table. It has an electronicconfiguration of 1s22s22p63s23p4. Sulphur can exist as different allotropes (allotropic
forms).
At room temperature (or temperatures up to 96oC), rhombic sulphur, with transparent
yellow crystals, is the stable form. It consists of eight sulphur atoms covalently bonded
together to form a crown structure ring. Rhombic sulphur is insoluble in water but
soluble in organic solvents.
Another allotrope of sulphur is monoclinic sulphur which is stable between 95.5oC
and 119oC. It exists as amber-yellow crystals and is also composed of S8molecules.
Sulphur atoms have six electrons in their outer shells. The atoms can accept two
electrons into their two singly occupied 3p orbitals to form the sulphide ion, S2.
Sulphur atoms can also form two, four or six covalent bonds, that is, it can also exist in
oxidation states +2, +4 and +6 as well as the 2 state in the sulphide ion. Sulphur atoms
form two covalent bonds by sharing the two electrons in the singly occupied 3p orbitals.
To form four covalent bonds, a sulphur atom promotes one of its paired 3p electrons to
a 3d orbital all of which are empty. Six covalent bonds can be formed by also
promoting a 3s electron into another 3d orbital.
rhombic sulphur monoclinic sulphur a S8
Electronic structure and
valency of sulphur
-
7/26/2019 10 the P-block Elements
26/33
10 The p-block Elements - 26
Table. Oxidation states of sulphur
Oxidation number Examples
+6 SO3, H2SO4, SO42
+4 SO2, H2SO3, SO32
+2 SCl2, S2O32
0 S8
2 H2S, S2
Shapes of sulphur compounds / ions:
Sulphur dioxide Sulphur trioxide Sulphuric(VI) acid
Sulphate(VI) ion Sulphur hexafluoride
Burning of sulphur
Sulphur burns with a dull blue flame in the presence of excess oxygen to form sulphur
dioxide gas which has a pungent choking smell. Traces of misty sulphur trioxide arealso found. The experiment is usually carried out in a fume cupboard, as sulphur
dioxide is a toxic gas.
S(s) + O2(g) SO2(g)
(II) Sulphur dioxide
Sulphur dioxide is a colorless gas with a characteristic pungent, choking smell. At room
temperature, it can be readily liquefied under pressure. It is very soluble in water and
forms an acidic solution. An aqueous solution of SO2 is called sulphuric(IV) acid or
sulphurous acid.
SO2(g) + H2OH+(aq) + HSO32
(aq)hydrogensulphate(IV) or hydrogensulphite
HSO32(aq)H+(aq) + SO3
2(aq)
-
7/26/2019 10 the P-block Elements
27/33
10 The p-block Elements - 27
Sulphur dioxide is known to be an acidic gaseous pollutant. It dissolves in raindrops,
forming acid rain. The acid rain causes damage to buildings made of concrete or
marble. This is because aqueous sulphur dioxide is an electrolyte that speeds up
corrosion. Besides, sulphur dioxide is also a highly irritating gas. It can cause damage
to the human respiratory system.
Oxidizing properties of sulphur dioxide
With strong reducing agents, sulphur dioxide acts as an oxidizing agent. In such cases,
the oxidation number of sulphur changes from +4 to 0.
SO2(g) + 4eS(s) + 2O2
(s)
or SO2(g) + 4H+(aq) + 4e
S(s) + 2H2O(l)`
(i) Reaction with Magnesium
Magnesium is a strong reducing agent. It
reacts with sulphur dioxide gas to give
sulphur and magnesium oxide.2Mg(s) +SO2(g) 2MgO(s) + S(s)
When a burning piece of magnesium is put
into a jar of sulphur dioxide, it continues to
burn, forming yellow specks of sulphur and
white magnesium oxide.
(ii) Reaction with hydrogen sulphide
Aqueous sulphur dioxide oxidizes hydrogen sulphide to give water and sulphur.
2H2S+ SO2(aq) 2H2O(1) + 3S(s)
Reducing properties of sulphur dioxide
Aqueous sulphur dioxide is a powerful reducing agent, acting as an electron donor.
SO2(g)+ 2H2O(1) SO42(aq) + 4H+(aq) + 2e
As SO32 ions are present in aqueous sulphur dioxide, the half equation can also be
written as
SO32(aq) + H2O(l) SO4
2(aq) + 2H+(aq) + 2e
Note that the oxidation number of sulphur increases from +4 to +6 when SO2or SO32
acts as reducing agent.
(i) Reaction with manganate(VII) ions(permanganate ion)
Manganate(VII) or permanganate ion, MnO4is a strong oxidizing agent. It reacts
with a reducing agent in an acidic medium to give Mn2+ion:
MnO4(aq) + 8H+(aq) + 5eMn2+(aq) + 4H2O(1)
2MnO4
(aq) + 5SO3
2(aq) + 6H+(aq) 2Mn2+(aq) + 5SO4
2(aq) + 4H2O(1)
purple colorless
The purple colour of MnO4(aq) is decolorized.
-
7/26/2019 10 the P-block Elements
28/33
10 The p-block Elements - 28
(ii) Reaction with dichromate(V1) ion
Dichromate(VI) ion, Cr2O72 is a strong oxidizing agent which reacts with a
reducing agent in an acidic medium to give chromium(III) ion.
Cr2O72(aq) + 14H+(aq) + 6e
2Cr3+(aq) + 7H2O(l)
Cr2O72
(aq) + 3SO32
(aq) + 8H+(aq) 2Cr3+(aq) + 3SO42
(aq) + 7H2O(l)orange green
(iii) Reaction with bromine
Bromine water, Br2(aq), is an oxidizing agent which reacts with a reducing agent
to give bromide ion.
Br2(aq) + 2e- 2Br(aq)
Br2(aq) + SO32(aq) + H2O(l) 2Br
(aq) + SO4
2(aq) + 2H+(aq)
yellowish brown (orange) colorless
(iv) Reaction with colored substance
Certain colored substances, such as dyes, are oxidizing agents. They can be
bleachedby moist sulphur dioxide, provided that the reduced form of the dye is
colorless.
dye(s) +SO32(aq) (dye O)(s) +SO4
2(aq)
colored colorless
Uses of sulphur dioxide
(i) Sulphur dioxide is a mild bleaching agent. It is commonly used to bleach
delicate materials such as paper, straw, silk and wool. Newspapers are bleached
by sulphur dioxide. When they are exposed in the air for a long time, they often
turn a pale yellow. This is because the oxygen in air reoxidize the reduced dye,
thus restoring the original color of raw paper.
(ii) Sulphur dioxide is commonly used to whiten some foodstuffs such as flour and
cheese. It is also used as a food preservative for fruit juices and jam.
(III) Sulphuric(VI) acid
Sulphuric(VI) acid is a corrosive, colorless, oily liquid. It is a strong dibasic acid. Pure
sulphuric(VI) acid boils and decomposes at 340oC, giving out fumes of sulphur trioxide
and steam.
H2SO4(l) SO3(g) + H2O(g)
The high boiling point and viscosity are due to the hydrogen bonding between the
hydrogen atom and oxygen atom of neighboring molecules.
-
7/26/2019 10 the P-block Elements
29/33
10 The p-block Elements - 29
Manufacture of sulphuric(VI) acid by the Contact Process
The Contact Process is an industrial method to manufacture sulphuric(VI) acid from
sulphur and oxygen/air via three stages.
(i) Preparation and purification of sulphur dioxide
Sulphur can be obtained naturally in elemental form in large undergrounddeposits which can be extracted by the Frasch process. Sulphur obtained is burnt
in air to give sulphur dioxide.
S(s) + O2(g) SO2(g)
Alternatively, sulphur dioxide can be obtained by roasting iron pyrite, FeS 2, or
galena, PbS , in oxygen/air.
4FeS2(s) + 11O2(g) 2Fe2O3(s) + 8SO2(g)
2PbS(s) + 3O2(g) 2PbO(s) + 2SO2(g)
The gases are then washed with water and dried by concentrated sulphuric(VI)
acid. This is necessary because impurities in sulphur dioxide and air may poisonthe catalyst used in the reaction.
Then, the purified sulphur dioxide and air are mixed and heated in a heat
exchanger by hot gases leaving the catalytic chamber.
(ii) Catalytic oxidation of sulphur dioxide to sulphur trioxide and choice for
optimum reacting conditions
Sulphur dioxide reacts with air to form sulphur trioxide.
2SO2(g) + O2(g)2SO3(g) H = 197 kJ mol1
However, the rate for the formation of sulphur trioxide is very slow. As the
reaction is exothermic and reversible, a low temperature favors the product side
of the equilibrium in the formation of sulphur trioxide but the time to reach the
equilibrium would be long. As a result, the manufacture of sulphur trioxide isusually carried out at 450oC.
The heat exchanger and catalytic chamber of
the Contact Process.
-
7/26/2019 10 the P-block Elements
30/33
10 The p-block Elements - 30
Besides, vanadium(V) oxide (V2O5) is used as a catalyst to increase the rate of
formation of sulphur trioxide. The common catalyst used in the past was platinum
which, however, is easily poisoned by arsenic compounds. Vanadium(V) oxide,
although has a lower efficiency than platinum, is cheaper and less susceptible to
poisoning. So, it is widely used today.
Furthermore, the volume occupied by the gaseous reactants (SO2 and O2) is
greater than that of the gaseous product (SO3). According to the Le Chateliersprinciple, high pressure will increase the yield. However, the provision of high
pressure is not economical. Therefore, the pressure under which the reaction takes
place is chosen to be one atmospheric pressure since the yield of sulphur trioxide
at this pressure is already 98%.
(iii) Conversion of sulphur trioxide to sulphuric(VI) acid
In the final stage, hot sulphur trioxide formed is sent back to the heat exchanger
to heat up the incoming sulphur dioxide and air. After it has cooled down, it is
dissolved in concentrated (98%) sulphuric(VI) acid in the absorption tower to
form oleum (fuming sulphuric acid).SO3(g) + H2SO4(1) H2S2O7(1)
Oleum is then added to the correct amount of water, forming concentrated
sulphuric(VI) acid of the required concentration.
H2S2O7(1) + H2O(l) 2H2SO4(l)
Note:Sulphur trioxide is not dissolved into water directly to form sulphuric(V1)
acid. This is because the reaction
SO3(g) + H2O(l) H2SO4(aq)
is highly exothermic and a mist of sulphuric(VI) acid will be formed instead
of a solution. It is difficult to collect the product and control its
concentration.
A flow diagram of manufacture of sulphuric acid by the Contact Process.
-
7/26/2019 10 the P-block Elements
31/33
10 The p-block Elements - 31
Chemical properties of sulphuric(VI) acid
(i) As a typical acid
Sulphuric(VI) acid is completely ionized in water:
H2SO4(l) + H2O(l) H3O+(aq) + HSO4
(aq)
hydrogensulphate(VI) ionHSO4
(aq) + H2O(l) H3O+(aq) + SO4
2(aq)
sulphate(VI) ion
Dilute sulphuric(VI) acid is a typical acid without any oxidizing property. The
following equations show the typical acidic properties of sulphuric(V1) acid.
Zn(s) + H2SO4(aq) ZnSO4(aq) + H2(g)
2NaOH(aq) + H2SO4Na2SO4(aq) + H2O(l)
2NH3(aq) + H2SO4(NH4)2SO4(aq)
CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l)
MgCO3(s) + H
2SO
4(aq) MgSO
4(aq) + CO
2(g) + H
2O(l)
NaHCO3(aq) + H2SO4(aq) Na2SO4(aq) + 2H2O(l) + 2CO2(g)
(ii) As an oxidizing agent
Concentrated sulphuric(VI) acid is a strong oxidizing agent, especially when it
is hot. It reacts with reducing agents to give a number of sulphur-containing
compounds. The oxidation numbers of sulphur usually decrease from +6 to +2
or +4.
1. Reaction with metals
Hot concentrated sulphuric(V1) acid reacts with all metals (except gold and
platinum):
Cu(s) + 2H2SO4(1) CuSO4(aq) + SO2(g) + 2H2O(l)
Zn(s) + 2H2SO4(1) ZnSO4(aq) + SO2(g) + 2H2O(l)
2. Reaction with non-metals
Some non-metals, such as carbon and sulphur, are oxidized by hot
concentrated sulphuric(VI) acid into their corresponding oxides.
C(s) + 2H2SO4(1) CO2(g) + 2SO2(g) + 2H2O(l)
S(s) + 2H2SO4(1) 3SO2(g) + 2H2O(l)
3. Reaction with hydrogen halides
Hot concentrated sulphuric(VI) acid oxidizes hydrogen bromide and
hydrogen iodide into bromine and iodine respectively.
2HBr(g) + H2SO4(1) Br2(g) + SO2(g) + 2H2O(1)
8HI(g) + H2SO4(l) 4I2(g) + H2S(g) + 4H2O(1)
However, concentrated H2SO4 cannot oxidize hydrogen fluoride (HF) orhydrogen chloride (HCl).
-
7/26/2019 10 the P-block Elements
32/33
10 The p-block Elements - 32
4. As a dehydrating agent
Concentrated sulphuric(VI) acid has a strong affinity for water and is
therefore a dehydrating agent.
Dehydrating hydrated salts
Remove water of crystallization from hydrated salts, for example,
Dehydrating organic compounds
Remove hydrogen and oxygen atoms in the ratio 2 : 1 to form H2O from
compounds that do not contain water molecules themselves, such as organic
compounds like alcohols or sucrose.
As concentrated sulphuric(VI) acid is capable of removing water or the
elements of water from other compounds, it acts as a dehydrating agent. It
can dehydrate wood, paper and cloth.
Note: Concentrated sulphuric(VI) acid can dehydrate skin which
contains protein and cause severe damage.
Test for sulphate(VI) Ions
The presence of sulphate(VI) ions in a solution can be confirmed by adding a solution
of barium chloride acidified with dilute nitric(V) acid to the suspected solution. If
SO42ions are present, a white precipitate will be formed.
Ba2+(aq) + SO42(aq) BaSO4(s)
Note: The barium chloride solution is acidified because carbonate ions and
sulphate(IV) ions also react with barium ions to give white precipitates.
Ba2+
(aq) + CO32
(aq) BaCO3(s)Ba2+(aq) + SO3
2(aq) BaSO3(s)
However, in the presence of acid, these precipitate will be dissolved again.
BaCO3(s) + 2HNO3(aq) Ba(NO3)2(aq) + H2O(l) + CO2(g)
BaSO3(s) + 2HNO3(aq) Ba(NO3)2(aq) + H2O(l) + SO2(g)
Uses of sulphuric(VI) acid
Sulphuric(VI) acid is one of the most important industrial chemicals. Many industries
require the use of sulphuric(VI) acid at some stage of the manufacturing process. For
instance, sulphuric(VI) acid is an important material used in the production of fertilizersand detergents, and many more.
CuSO45H2O(s) CuSO4(s) + 5H2O(l)
copper(II) sulphate-5-water white powder
blue crystals
conc. H2SO4
C2H5OH(l) CH2= CH2(g) + H2O(l)
ethanol ethene
conc. H2SO4
C12H22O11(s) 12C(s) + 11H2O(l)
sucrose
conc. H2SO4
-
7/26/2019 10 the P-block Elements
33/33
(i) Fertilizers
Calcium dihydrogenphosphate(V)
Calcium phosphate(V), which is found in
phosphate ores, is a water insoluble substance. It
can be converted to the more water-soluble
calcium dihydrogenphosphate(V) which is a good
phosphate fertilizer by reacting with concentratedH2SO4.
Ca3(PO4)2(s) + 2H2SO4(1)
Ca(H2PO4)2(s) + 2CaSO4(s)
Ammonium sulphate
Ammonia is a good nitrogenous fertilizer. However, it is highly soluble in water
and is likely to be washed away in heavy rain. It can be changed to ammonium
sulphate by reacting with concentrated H2SO4.
NH3(g) + H2SO4(1) (NH4)2SO4(aq)
(ii) Detergents
Long chain hydrocarbons are obtained either by cracking or by building up from
ethene units and benzene through addition reactions. For example,
phenyldodecane can be treated with concentrated H2SO4 to give a sulphonated
hydrocarbon. A sodium salt of sulphonated hydrocarbon which is a soapless
detergent or a synthetic detergent can be formed after neutralization with NaOH.
Synthetic detergents have several advantages over soapy detergents.
1. They do not form a scum with hard water or acidic solution.
2. They can be tailor-made to suit a particular cleaning purpose, e.g. washing
powder, shampoo and bath liquids.
(iii) Dyestuffs
Azo dyes are commonly used dyes which are made from nitrobenzene, which is
prepared from the reaction between benzene and a mixture of concentrated
sulphuric(VI) acid and concentrated nitric(V) acid.
(iv) Paints and pigments
A white pigment called titanyl sulphate, TiOSO4, is made by dissolving
titanium(IV) oxide, TiO4in hot concentrated sulphuric(VI) acid. Besides, barium
sulphate(VI) and calcium sulphate(VI) made from sulphuric acid are used as paint
additives.
Ba2+(aq) + SO42(aq) BaSO4(s)
2 2