10 the P-block Elements

7/26/2019 10 the P-block Elements

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10 The p-block Elements - 1

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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.

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(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


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(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.


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(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.


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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.


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(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.


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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.


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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.


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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


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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)


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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)


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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.


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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)


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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


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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


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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.


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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


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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


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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.


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(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


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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


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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.


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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)


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(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.


24/33


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


25/33


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


26/33


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)


27/33


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.


28/33


(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.


29/33


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.


30/33


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.


31/33


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).


32/33


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


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

Documents

10 the P-block Elements