New Way Chemistry for Hong Kong A-Level Book 41 1 Nitrogen and its Compound 43.1Introduction...

1New Way Chemistry for Hong Kong A-Level Book 4

1

Nitrogen and its Compound43.143.1 IntroductionIntroduction

43.243.2 Unreactive Nature of NitrogenUnreactive Nature of Nitrogen

43.343.3 Direct Combination of Nitrogen and Oxygen Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxidesleading to Formation of Nitrogen Oxides

43.443.4 AmmoniaAmmonia

43.543.5 Nitric(V) AcidNitric(V) Acid

43.643.6 Nitrates(V)Nitrates(V)

Chapter 43Chapter 43


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43.1 Introduction (SB p.110)

Nitrogen (first member of Group VA):

• Electronic configuration: 1s22s2

2p3

• Complete octet by forming diatomic molecules N N

• Non-metal, colourless and odourless gas

• Very low melting and boiling points

• Slightly soluble in water and does not support combustion

Covalent radius (nm) 0.074

Melting point (°C) –210

Boiling point (°C) –196

Bond enthalpy (kJ mol–

1)+944

First ionization enthalpy (kJ mol–1)

+1 400

Electron affinity (kJ mol–1)

+3

Electronegativity 3.0

Some information about nitrogen


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43.1 Introduction (SB p.110)

Nitrogen

• Mainly as free N2 molecules in the atmosphere (78%

by volume)

• Combine with other elements in the form of proteins in

all living things

• Liquid N2 is used as coolant

• Raw material for Haber process

(manufacture of ammonia)

• Ammonia is the major component

of nitrogenous fertilizers


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Nitrogen in gaseous state

• As diatomic molecules (N2) which are held by weak van der Waals’ forces

• 2 atoms are joined by extremely strong triple covalent bonds

• Bond enthalpy of the triple bond = +944 kJ mol–1

• Due to extremely strong covalent bonds and absence of bond polarity

Nitrogen molecule is very unreactive

43.2 Unreactive Nature of Nitrogen (SB p.111)


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Bond Bond

enthalpy (kJ mol–1)

Bond Bond

enthalpy (kJ mol–1)

N N

O = O

H – H

C – C

+944

+496

+436

+348

S – S

Cl – Cl

P – P

F – F

+264

+242

+172

+158

Bond enthalpies of some common covalent bonds


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• Reactions involving nitrogen usually have high activation en

ergies and unfavourable equilibrium constants

e.g. At 25°C

N2(g) + O2(g) 2NO(g) Kc = 4.5 10–31

• The presence of catalyst and high temperature and

pressure may be required for nitrogen to react

N2(g) + 3H2(g) 2NH3(g)400 – 500°C, 300 – 1000 atm

Fe as catalyst



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43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.112)

• N2 will not react at room temperature due to high bond

enthalpy

• At high temperature, N2 shows some reactions with

other elements

∵ sufficient energy to break N N triple bond

• At high temperature,

N2(g) + O2(g) 2NO(g)


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• The electric discharge in lightning provides sufficient e

nergy to break the N N triple bond and then react

with O2

N2(g) + O2(g) 2NO(g)

2NO(g) + O2(g) 2NO2(g)

lightning

colourless Reddish brown(poisonous)

colourless


9


• The above reactions are very important in nature

• The NO2 formed dissolves in rainwater to produce nit

ric(V) acid and nitric(III) acid

2NO2(g) + H2O(l) HNO3(aq) + HNO2(aq)


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• At high temperatures in car

engines, N2 & O2 react to form NO(g)

which emitted into air with

exhausted gas

• The NO formed will be oxidized to

NO2

• The NO2 absorbs sunlight and breaks down

into NO and O atom

NO2(g) NO(g) + O(g)

• These leads to formation of photochemical smog

sunlight


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• In laboratory, we use

the apparatus shown

on the right to convert

N2 into NO2

• When current is switched on, electric discharges occur in

the gap between the electrodes

• NO is formed and followed by NO2


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• Other than NO and NO2, N2 can form other oxides

e.g. 2 NO2 molecules (brown) can combine to form a

N2O4 molecule (yellow)

• NO2 & N2O4 exist in equilibrium in gas phase

2NO2(g) N2O4(g) H = –58 kJ mol–1brown yellow


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• The formation of N2O4 is exothermic

N2O4 predominantes at low

temperatures

NO2 predominantes at high

temperatures

the colour of mixture fades on

cooling, darkens on warming


NO2 (left) and N2O4 (right) predominate in hot water and ice water respectively


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Check Point 43-1 Check Point 43-1

(a) Draw the structures of the following compounds.

(i) Dinitrogen monoxide

(ii) Nitrogen monoxide

(iii) Dinitrogen trioxide

(iv) Nitrogen dioxide

(v) Dinitrogen tetraoxide

(vi) Dinitrogen pentaoxide Answer


(a) (i) Dinitrogen monoxide

(ii) Nitrogen monoxide


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(iii) Dinitrogen trioxide

(iv) Nitrogen dioxide


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(v) Dinitrogen tetraoxide

(vi) Dinitrogen pentaoxide


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(b) Arrange N2, O2 and F2 in an ascending order of reactivity. Explain the order briefly.

Answer(b) The ascending order of reactivity is: N2 < O2 < F2.

The reactivity of diatomic molecules depends on the bond enthalpy of covalent bonds. The bond enthalpy of N N is greater than that of O = O, which in turn is greater than that of F – F. Therefore, the breakage of N N bond requires the greatest amount of energy, whereas the breakage of F – F bond requires the least amount of energy.



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43.4 Ammonia (SB p.114)

Ammonia

• colourless, pungent gas

• polar molecules

• trigonal pyramidal shape with a lone pair of electrons on nitrogen

• extremely soluble in water and easy to condense to liquid due to hydrogen bonds

• good solvent for ionic compounds

• weakly alkaline

NH3(aq) + H2O(l) NH4+(aq) + OH–(aq)

Kb = 1.8 10–5 mol dm–3


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Ammonia

• one of the most fundamental raw materials for modern i

ndustries

• important source of fertilizers and 85% of ammonia i

s used to make nitrogenous fertilizers (e.g. (NH4)2SO4,

NH4NO3)

• making fibres and plastics (rayon, nylon)

• making nitric(V) acid (used to make fertilizers, dyes)

• making household cleaners

• making detergents


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Percentages of ammonia used in different industries


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Manufacture of Ammonia by the Haber ProcessManufacture of Ammonia by the Haber Process


• NH3 is manufactured industrially by t

he Haber Process, named after the G

erman chemist Fritz Haber

• The process involves direct combina

tion of N2 and H2 under special condi

tions

N2(g) + H2(g) 2NH3(g)

H = –92 kJ mol–1

Fritz Haber (1868 – 1934)


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Flow diagram for the Haber process


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Process of Haber Process:

• N2 is obtained from fractional distillation of liquid air

• H2 is obtained from methane, naphtha or mixture by ste

am reforming

CH4(g) + H2O(g) CO(g) + 3H2(g)

CH4(g) + air CO(g) + 2H2(g) + N2(g)

Ni900°C

Ni900°C



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• Mixture of CO & H2O is mixed with steam and passed over a heated catalyst

CO(g) + H2O(l) CO2(g) + H2(g)

The CO2 formed is dissolved in water under pressure

• The gases (N2 & H2) are purified before proceeding to the next stage

∵ Compounds of oxygen and sulphur will poison the catalyst



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• Purified N2 & H2 are mixed in ratio of 3 : 1 by volume

Compressed to 200 – 1000 atm and heated in the heat exchanger

Hot gaseous mixture is passed over iron in the catalytic chamber

Gases contain 10 – 15% of NH3 and unreacted N2

and H2 when leaving the chamber

The gases are cooled after passing through the heat exchanger

NH3 is liquefied under pressure and unreacted gas

es are recycled



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Physico-chemical principles:

• Synthesis of ammonia is an exothermic and reversible

reaction

N2(g) + H2(g) 2NH3(g) H = –92 kJ mol–1

• According to Le Chatelier’s principle,

(1) high pressure will increase the yield

(2) low temperature will increase the yield


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• Apart from increasing yield, the reaction rate should be

fast

Low temperatures would lower the rate of reaction

∴ optimum temperature is around 500°C which is high

enough for reaction to proceed quickly but low enough

to give satisfactory yield

• Catalyst is used to increase the reaction rate

poisoned by CO, CO2, H2S

Gases entering the catalytic chamber should have

high purity!!


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Chemical Properties of AmmoniaChemical Properties of Ammonia


As a base

• NH3 partly ionizes in water to give NH4+ and OH– ion

s

∴ NH3(aq) is alkaline

NH3(g) + H2O(l) NH4+(aq) + OH–(aq)

Kb = 1.8 10–5 mol

dm–3

∴ NH3(aq) is a weak base


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Reaction with Acids

• NH3 neutralizes acids to give ammonium salts

e.g. 2NH3(aq) + H2SO4(aq) (NH4)2SO4(aq)

NH3(aq) + HNO3(aq) NH4NO3(aq)

ammonium sulphate(VI)

ammonium nitrate(V)

Filter paper soaked with

NH3

Filter paper soaked with

HCl

• Formation of NH4Cl by reacting NH3 with HCl

NH3(aq) + HCl(aq) NH4Cl(s)


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Reaction with Metal Salts

• NH3 precipitates the hydroxide of many metals from solutio

ns of their salts

CaSO4(aq) + 2NH3(aq) +2H2O(l) Ca(OH)2(s) + (NH4)2SO4(aq)

ZnSO4(aq) + 2NH3(aq) +2H2O(l) Zn(OH)2(s) + (NH4)2SO4(aq)

Pb(NO3)2(aq) + 2NH3(aq) +2H2O(l) Pb(OH)2(s) + 2NH4NO3(aq)

CuSO4(aq) + 2NH3(aq) +2H2O(l) Cu(OH)2(s) + (NH4)2SO4(aq)

FeSO4(aq) + 2NH3(aq) +2H2O(l) Fe(OH)2(s) + (NH4)2SO4(aq)

Fe2(SO4)3(aq) + 6NH3(aq) +6H2O(l) 2Fe(OH)3(s) + 3(NH4)2SO4(aq)

white

white

white

blue

dirty green

reddish brown


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Pb(OH)2(s) Cu(OH)2(s) Fe(OH)2(s) Fe(OH)3(s)


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• Some metal hydroxides (e.g. Zn(OH)2 & Cu(OH)2) redis

solve in excess NH3 solution and form complex compound

s

Zn(OH)2(s) + 4NH3(aq) [Zn(NH3)4]2+(aq) + 2OH–(aq)

Cu(OH)2(s) + 4NH3(aq) [Cu(NH3)4]2+(aq) + 2OH–(aq)

colourless

deep blue

A solution containing Cu2+(aq)

Cu(OH)2(s)

[Cu(NH3)4]2+(aq)


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AgCl(s) dissolves

Addition of excess NH3(aq)

water

AgCl(s)

• Silver(I) ions also form a complex with ammonia

• AgCl is insoluble in water and acids, but dissolves in excess

NH3 forming soluble complex ion [Ag(NH3)2]+(aq)

AgCl(s) Ag+(aq) + Cl–(aq)

Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq)


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• NH3 does not burn in air or support combustion

• It burns in O2 with a yellow flame, forming N2 and water vapour


As a Reducing Agent

Laboratory set-up for oxidation of ammonia

Reaction with Oxygen

4NH3(g) + 3O2(g) 2N2(g) + 6H2O(g)


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• In the presence of catalyst (red hot spiral coil of platinum

at 800 – 900°C), NH3 is oxidized to NO by O2


4NH3(g) + 5O2(g)

4NO(g) + 6H2O(g)

• This is called catalytic oxid

ation of ammonia

• Key reaction in the preparat

ion of HNO3

Pt

Laboratory set-up for catalytic oxidation of ammonia


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Reaction with Copper(II) Oxide

• When dry NH3 is passed over heated black CuO, NH3 i

s oxidized to N2 and H2O

• The CuO turns from black to reddish brown as it is r

educed to Cu

2NH3(g) + 3CuO(s) 3Cu(s) + N2(g) + 3H2O(g)

Laboratory set-up for oxidation of ammonia by copper(II) oxide


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(a) Write chemical equations to show how hydrogen is produced from

(i) the reaction of natural gas (mainly methane) with water;

(ii) the reaction of coal (mainly carbon) with water.

Answer


(a) (i) CH4(g) + H2O(g) CO(g) + 3H2(g)

(ii) C(s) + H2O(l) CO(g) + H2(g)


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Check Point 43-2 (cont’d) Check Point 43-2 (cont’d)

(b) Consider the following reversible reaction:

N2(g) + 3H2(g) 2NH3(g) H = –92 kJ mol–1

Discuss how each of the following factors affects the above equilibrium:

(i) increase in temperature

(ii) decrease in pressure

(iii) addition of a suitable catalyst Answer



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(b) (i) The forward reaction is exothermic. According to Le Chatelier’s principle, exothermic reactions are favoured at low temperatures. Therefore, an increase in temperature will favour the backward reaction, and thus decrease the yield of ammonia.

(ii) According to Le Chatelier’s principle, a high pressure will increase the yield of ammonia as the forward reaction is accompanied by a decrease of volume from four to two volumes of the gas. Therefore, a decrease in pressure will decrease the yield of ammonia.

(iii) Addition of a suitable catalyst will increase the rate of both forward and backward reactions to the same extent. As it does not change the position of the equilibrium, the yield of ammonia remains constant.



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(c) Ammonia reacts with oxygen in two different ways. Give equations for both of these reactions and explain how one of them is used industrially to produce nitric(V) acid.

Answer



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(c) In the absence of catalyst, ammonia burns to give molecular nitrogen and water vapour.

4NH3(g) + 3O2(g) 2N2(g) + 6H2O(g)

Industrially, in the presence of red hot platinum-rhodium at about 850°C, ammonia is catalytically oxidized to nitrogen monoxide.

4NH3(g) + 5O2(s) 4NO(g) + 6H2O(g)

The nitrogen monoxide formed then reacts with oxygen from the air to give nitrogen dioxide.

2NO(g) + O2(g) 2NO2(g)

The nitrogen dioxide reacts with excess air and water to produce aqueous nitric(V) acid.

4NO2(g) + O2(g) + 2H2O(l) 4HNO3(aq)

H2O(l) + 3NO2(g) 2HNO3(aq) + NO(g)

The NO(g) is recycled and subsequently combines with more oxygen and water to give more nitric(V) acid. Finally, the product is distilled to give concentrated nitric(V) acid (containing 68% HNO3).


Pt – Rh

850°C


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Nitric(V) acid

• a very strong acid

• turns yellow on storage as the formation

of dissolved NO2 from decomposition of s

ome acid

4HNO3(l) 4NO2(aq) + 2H2O(l) +O2(g)

• keep in brown bottles as light will speed

up decomposition

• used to make explosives, nylon, fertilizers

and dyestuff synthesis

43.5 Nitric(V) Acid (SB p.121)


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Manufacture of Nitric(V) Acid from the Catalytic Oxidation of Ammonia

Manufacture of Nitric(V) Acid from the Catalytic Oxidation of Ammonia


• Most of the ammonia formed is converted to nitric(V) ac

id by Ostward process

• Ostward process is divided into 3 stages:

1. Mixture of ammonia and excess air is passed over

Pt-Rh catalyst at around 700-800°C under low pres

sure

4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)Pt-Rh

850°C


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2. The NO formed then reacts with O2 to form NO2

2NO(g) + O2(g) NO2(g)

3. The NO2 reacts with excess air and water to give aqueo

us HNO3

4NO2(g) + O2(g) + 2H2O(l) 4HNO3(aq)



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Nitric(V) Acid as an Oxidizing AgentNitric(V) Acid as an Oxidizing Agent

• HNO3 is a strong oxidizing agent, especially when

concentrated

• NO3– acts as an electron acceptor when H+ ions are present

• HNO3 can be reduced to different nitrogen compounds

with different oxidation states, depending on

1. the conc. of HNO3

2. nature of substance being oxidized



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• If dilute or moderately concentrated HNO3 is reduce

d, NO will be formed

4HNO3(aq) + 3e– 3NO3 –(aq) + 2H2O(l) + NO(g)

or NO3–(aq) + 4H+(aq) + 3e– NO(g) + 2H2O(l)

• If concentrated HNO3 is reduced, NO2 will be formed

2HNO3(aq) + e– NO3 –(aq) + NO2(g) + H2O(l)

or NO3–(aq) + 2H+(aq) + e– NO2(g) + H2O(l)

• The electrons are supplied by the reducing agent in t

he reaction



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• Cu reacts with warm dilute HNO3 to give NO

3Cu(s) + 8HNO3(aq)

3Cu(NO3)2(aq) + 4H2O(l) + 2N

O(g)

• The NO formed reacts with atmospheric O2 to give

NO2

2NO(g) + O2(g) 2NO2(g)


Reaction with Copper


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• Conc. HNO3 (~14 M) reacts with Cu to give NO2 and a blue s

olution of Cu(NO3)2

Cu(s) + 4HNO3(aq) Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g)



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• Conc. HNO3 oxdizes green Fe2+ ions to brown Fe3+ io

ns while itself reduced to NO

3Fe2+(aq) + NO3–(aq) + 4H+(aq)

3Fe3+(aq) + NO(g) + 2H2

O(l)

• The NO formed reacts with atmospheric O2 to form N

O2

2NO(g) + O2(g) NO2(g)


Reaction with Iron(II) Ion


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• Hot concentrated HNO3 oxidizes sulphur to give sulph

uric(VI) acid and brown fumes of NO2

S(s) + 6HNO3(aq)

H2SO4(aq) + 6NO2(g) + 2H2

O(l)


Reaction with Sulphur


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Account for the following observation by giving a balanced equation.

(a) Nitrogen monoxide turns brown when exposed to air.

Answer


(a) Nitrogen monoxide reacts with atmospheric oxygen to give brown nitrogen dioxide gas.

2NO(g) + O2(g) 2NO2(g)


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(b) Nitric(V) acid turns yellowish brown on standing.Answer


(b) Nitric(V) acid turns yellowish brown on standing because of the dissolved nitrogen dioxide formed from the decomposition of some of the acid.

4HNO3(aq) 4NO2(aq) + 2H2O(l) + O2(g)


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(c) Silver dissolves in dilute nitric(V) acid, yielding a colourless gas.

Answer


(c) Dilute nitric(V) acid is reduced by silver to form colourless nitrogen monoxide gas.

3Ag(s) + 4HNO3(aq)

3Ag+(aq) + 3NO3–(aq) + 2H2O(l) + NO(g)


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(d) The nitrate of a metal ion decomposed on heat to give the metal.

Answer


(d) Both mercury nitrate(V) and silver nitrate(V) decompose on heating to give the corresponding metal.

Hg(NO3)2(s) Hg(s) + 2NO2(g) + O2(g)

2Ag(NO3)2(s) 2Ag(s) + 2NO2(g) + O2(g)


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43.6 Nitrates(V) (SB p.124)

• Metal nitrates(V) can be prepared by reacting very dilute nit

ric(V) acid with metals, metal oxides, hydroxides or carbon

ates

e.g. Mg(s) + 2HNO3(aq) Mg(NO3)2(aq) + H2(g)

CuO(s) + 2HNO3(aq) Cu(NO3)2(aq) + H2O(l)

NaOH(aq) + HNO3(aq) NaNO3(aq) + H2O(l)

Na2CO3(aq) + 2HNO3(aq)

2NaNO3(aq) + H2O(l) + CO2(g)

very dilute


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• Metal nitrates(V) can be prepared by reacting metals wit

h concentrated nitric(V) acid

e.g. Mg(s) + 4HNO3(aq)

Mg(NO3)2(aq) + 2H2O(l) + 2N

O2(g)

concentrated


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Action of Heat on Nitrates(V)Action of Heat on Nitrates(V)

When metal nitrates(V) in solid form are strongly heated, they decompose differently according to their thermal stability


Metal oxide, nitrogen dioxide and oxygen

2Ca(NO3)2(s) 2CaO(s) + 4NO2(g) + O2(g)

2Mg(NO3)2(s) 2MgO(s) + 4NO2(g) + O2(g)

4Al(NO3)3(s) 2Al2O3(s) + 12NO2(g) + 3O2(g)

2Zn(NO3)2(s) 2ZnO(s) + 4NO2(g) + O2(g)

2Fe(NO3)2(s) 2FeO(s) + 4NO2(g) + O2(g)

2Pb(NO3)2(s) 2PbO(s) + 4NO2(g) + O2(g)

2Cu(NO3)2(s) 2CuO(s) + 4NO2(g) + O2(g)

Calcium

Magnesium

Aluminium

Zinc

Iron

Lead

Copper

Metal nitrate(III), oxygen

2KNO3(s) 2KNO2(s) + O2(g)

2NaNO3(s) 2NaNO2(s) + O2(g)

PotassiumSodium

ProductReaction Nitrate of


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Dinitrogen oxide and water

NH4NO3(s) N2O(g) + 2H2O(l)Ammonium ion

Metal, nitrogen dioxide and oxygen

Hg(NO3)2(s) Hg(s) + 2NO2(g) + O2(g)

2AgNO3(s) 2Ag(s) + 2NO2(g) + O2(g)

Mercury(II)Silver

ProductReaction Nitrate of

Cont’d


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Brown Ring Test for Nitrate(V) IonsBrown Ring Test for Nitrate(V) Ions

• The brown ring test is used to detect nitrate(V) ions in

aqueous solutions


Procedure:

1. Mix a freshly prepared FeSO4 solution with a solution suspected of containing nitrate(V) ions in a test tube

2. Conc. H2SO4 is added carefully along the side tothe bottom of the test tube with the test tube tilted

Laboratory set-up for brown ring test


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• Formation of a brown ring confirms the presence of

nitrate(V) ions in the solution



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Reactions involved in the brown ring test:

• Nitrate(V) ions react with conc. H2SO4 to give HNO3

NO3–(aq) + H2SO4(l) HNO3(aq) + HSO4

–(aq)

• The nitric(V) acid oxidizes some FeSO4 to Fe2(SO4)3 and is itself reduced to NO

HNO3(aq) + 3Fe2+(aq) + 3H+(aq)

NO(g) + 3Fe3+(aq) + 2H2O(l)

• Finally, NO reacts with unreacted FeSO4 to form a brown complex

FeSO4(aq) + NO(g) FeSO4 • NO(aq)Brown complex


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Give the name of the ion responsible for the following observation.

(a) An ion produces a blue precipitate with ammonia solution. The blue precipitate redissolves in excess ammonia solution to give a clear deep blue solution.

Answer

(a) Copper(II) ion



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(b) An ion produces a dirty green precipitate with ammonia solution.

Answer

(b) Iron(II) ion



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(c) An ion gives a positive result in the brown ring test.

Answer

(c) Nitrate(V) ion



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The END

New Way Chemistry for Hong Kong A-Level Book 41 1 Nitrogen and its Compound 43.1Introduction...

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Transcript of New Way Chemistry for Hong Kong A-Level Book 41 1 Nitrogen and its Compound 43.1Introduction...